the+elements



=The Element Hydrogen=  1 H  Hydrogen 1.00794

**What's in a name?** From the Greek words **hydro** and **genes**, which together mean "water forming." **Say what?** Hydrogen is pronounced as **HI-dreh-jen**. **History and Uses:** Scientists had been producing hydrogen for years before it was recognized as an element. Written records indicate that Robert Boyle produced hydrogen gas as early as 1671 while experimenting with [|iron] and acids. Hydrogen was first recognized as a distinct element by Henry Cavendish in 1766. Composed of a single [|proton] and a single [|electron], hydrogen is the simplest and [|most abundant element in the universe]. It is estimated that 90% of the visible universe is composed of hydrogen. Hydrogen is the raw fuel that most stars 'burn' to produce energy. The same process, known as fusion, is being studied as a possible power source for use on earth. The sun's supply of hydrogen is expected to last another 5 billion years. Hydrogen is a commercially important element. Large amounts of hydrogen are combined with [|nitrogen] from the air to produce ammonia (NH 3 ) through a process called the Haber process. Hydrogen is also added to fats and oils, such as peanut oil, through a process called hydrogenation. Liquid hydrogen is used in the study of superconductors and, when combined with liquid [|oxygen], makes an excellent rocket fuel. Hydrogen combines with other elements to form numerous compounds. Some of the common ones are: water (H 2 O), ammonia (NH 3 ), methane (CH 4 ), table sugar (C 12 H 22 O 11 ), hydrogen peroxide (H 2 O 2 ) and hydrochloric acid (HCl). Hydrogen has three common [|isotopes]. The simplest isotope, called protium, is just ordinary hydrogen. The second, a stable isotope called [|deuterium], was discovered in 1932. The third isotope, [|tritium], was discovered in 1934. **Estimated Crustal Abundance:** 1.40×10 3 milligrams per kilogram **Estimated Oceanic Abundance:** 1.08×10 5 milligrams per liter **Number of Stable Isotopes:** 2 ( [|View all isotope data] )http://education.jlab.org/itselemental/iso001.htmlhttp://education.jlab.org/itselemental/iso001.htmlhttp://education.jlab.org/itselemental/iso001.html **Ionization Energy:** 13.598 eV **Oxidation States:** +1, -1 Helium  2 He Helium 4.002602 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** For the Greek god of the sun, [|**Helios**]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Helium is pronounced as **HEE-lee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Helium, the second [|most abundant element in the universe], was discovered on the sun before it was found on the earth. Pierre-Jules-César Janssen, a French astronomer, noticed a yellow line in the sun's spectrum while studying a total solar eclipse in 1868. Sir Norman Lockyer, an English astronomer, realized that this line, with a wavelength of 587.49 nanometers, could not be produced by any element known at the time. It was hypothesized that a new element on the sun was responsible for this mysterious yellow emission. This unknown element was named helium by Lockyer. <span style="font-family: 'times new roman',serif; font-size: 16px;">The hunt to find helium on earth ended in 1895. Sir William Ramsay, a Scottish chemist, conducted an experiment with a mineral containing [|uranium] called clevite. He exposed the clevite to mineral acids and collected the gases that were produced. He then sent a sample of these gases to two scientists, Lockyer and Sir William Crookes, who were able to identify the helium within it. Two Swedish chemists, Nils Langlet and Per Theodor Cleve, independently found helium in clevite at about the same time as Ramsay. <span style="font-family: 'times new roman',serif; font-size: 16px;">Helium makes up about 0.0005% of the [|earth's atmosphere]. This trace amount of helium is not gravitationally bound to the earth and is constantly lost to space. The earth's atmospheric helium is replaced by the decay of radioactive elements in the earth's crust. [|Alpha decay], one type of radioactive decay, produces particles called [|alpha particles]. An alpha particle can become a helium atom once it captures two electrons from its surroundings. This newly formed helium can eventually work its way to the atmosphere through cracks in the crust. <span style="font-family: 'times new roman',serif; font-size: 16px;">Helium is commercially recovered from natural gas deposits, mostly from Texas, Oklahoma and Kansas. Helium gas is used to inflate blimps, scientific balloons and party balloons. It is used as an inert shield for arc welding, to pressurize the fuel tanks of liquid fueled rockets and in supersonic windtunnels. Helium is combined with [|oxygen] to create a [|nitrogen] free atmosphere for deep sea divers so that they will not suffer from a condition known as nitrogen narcosis. Liquid helium is an important cryogenic material and is used to study superconductivity and to create superconductive magnets. The Department of Energy's [|Jefferson Lab] uses large amounts of liquid helium to operate its superconductive [|electron accelerator]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Helium is an inert gas and does not easily combine with other elements. There are no known compounds that contain helium, although attempts are being made to produce helium diflouride (HeF 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 8×10 -3 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 7×10 -6 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2
 * Atomic Number:** 1
 * Atomic Weight:** 1.00794
 * Melting Point:** 13.81 K (-259.34°C or -434.81°F)
 * Boiling Point:** 20.28 K (-252.87°C or -423.17°F)
 * Density:** 0.00008988 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 1 **Group Number:** 1 **Group Name:** none
 * Atomic Number:** 2
 * Atomic Weight:** 4.002602
 * Melting Point:** 0.95 K (-272.2°C or -458.0°F)
 * Boiling Point:** 4.22 K (-268.93°C or -452.07°F)
 * Density:** 0.0001785 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 1 **Group Number:** 18 **Group Name:** Noble Gas

<span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 24.587 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** 0 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :1s2** = Lithium = <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 3 Li  Lithium 6.941

<span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for stone, **lithos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Lithium is pronounced as **LITH-ee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Lithium was discovered in the mineral petalite (LiAl(Si 2 O 5 ) 2 ) by Johann August Arfvedson in 1817. It was first isolated by William Thomas Brande and Sir Humphrey Davy through the electrolysis of lithium oxide (Li 2 O). Today, larger amounts of the metal are obtained through the electrolysis of lithium chloride (LiCl). Lithium is not found free in nature and makes up only 0.0007% of the earth's crust. <span style="font-family: 'times new roman',serif; font-size: 16px;">Many uses have been found for lithium and its compounds. Lithium has the highest specific heat of any solid element and is used in heat transfer applications. It is used to make special glasses and ceramics, including the Mount Palomar telescope's 200 inch mirror. Lithium is the lightest known metal and can be alloyed with [|aluminium], [|copper] , [|manganese] , and [|cadmium] to make strong, lightweight metals for aircraft. Lithium hydroxide (LiOH) is used to remove carbon dioxide from the atmosphere of spacecraft. Lithium stearate (LiC 18 H 35 O 2 ) is used as a general purpose and high temperature lubricant. Lithium carbonate (Li 2 CO 3 ) is used as a drug to treat manic depression disorder. <span style="font-family: 'times new roman',serif; font-size: 16px;">Lithium reacts with water, but not as violently as [|sodium]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.8×10 -1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 ( [|View all isotope data] ) <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 5.392 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 1s 2 2s 1 =Beryllium= <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 4 Be  Beryllium 9.012182
 * Atomic Number:** 3
 * Atomic Weight:** 6.941
 * Melting Point:** 453.65 K (180.50°C or 356.90°F)
 * Boiling Point:** 1615 K (1342°C or 2448°F)
 * Density:** 0.534 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 2 **Group Number:** 1 **Group Name:** Alkali Metal
 * [|Electron Shell Configuration] :**

<span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word **beryl**, a type of mineral. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Beryllium is pronounced as **beh-RIL-ee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although emeralds and beryl were known to ancient civilizations, they were first recognized as the same mineral (Be 3 Al 2 (SiO 3 ) 6 ) by Abbé Haüy in 1798. Later that year, Louis-Nicholas Vauquelin, a French chemist, discovered that an unknown element was present in emeralds and beryl. Attempts to isolate the new element finally succeeded in 1828 when two chemists, Friedrich Wölhler of Germany and A. Bussy of France, independently produced beryllium by reducing beryllium chloride (BeCl 2 ) with [|potassium] in a [|platinum] crucible. Today, beryllium is primarily obtained from the minerals beryl (Be 3 Al 2 (SiO 3 ) 6 ) and bertrandite (4BeO·2SiO 2 ·H 2 O) through a chemical process or through the electrolysis of a mixture of molten beryllium chloride (BeCl 2 ) and sodium chloride (NaCl). <span style="font-family: 'times new roman',serif; font-size: 16px;">Beryllium is relatively transparent to X-rays and is used to make windows for X-ray tubes. When exposed to [|alpha particles], such as those emitted by [|radium] or [|polonium] , beryllium emits [|neutrons] and is used as a neutron source. Beryllium is also used as a moderator in nuclear reactors. <span style="font-family: 'times new roman',serif; font-size: 16px;">Beryllium is alloyed with [|copper] (2% beryllium, 98% copper) to form a wear resistant material, known as beryllium bronze, used in gyroscopes and other devices where wear resistance is important. Beryllium is alloyed with [|nickel] (2% beryllium, 98% nickel) to make springs, spot-welding electrodes and non-sparking tools. Other beryllium alloys are used in the windshield, brake disks and other structural components of the space shuttle. <span style="font-family: 'times new roman',serif; font-size: 16px;">Beryllium oxide (BeO), a compound of beryllium, is used in the nuclear industry and in ceramics. <span style="font-family: 'times new roman',serif; font-size: 16px;">Beryllium was once known as glucinum, which means sweet, since beryllium and many of its compounds have a sugary taste. Unfortunately for the chemists that discovered this particular property, beryllium and many of its compounds are poisonous and should never be tasted or ingested. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.8 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 5.6×10 -6 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1
 * Atomic Number:** 4
 * Atomic Weight:** 9.012182
 * Melting Point:** 1560 K (1287°C or 2349°F)
 * Boiling Point:** 2744 K (2471°C or 4480°F)
 * Density:** 1.85 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 2 **Group Number:** 2 **Group Name:** Alkaline Earth Metal

<span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 9.323 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 1s 2 2s 2
 * [|Electron Shell Configuration] :**

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 5 B  Boron 10.811 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Arabic word **Buraq** and the Persian word **Burah**, which are both words for the material "borax." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Boron is pronounced as **BO-ron**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Boron was discovered by Joseph-Louis Gay-Lussac and Louis-Jaques Thénard, French chemists, and independently by Sir Humphry Davy, an English chemist, in 1808. They all isolated boron by combining boric acid (H 3 BO 3 ) with [|potassium]. Today, boron is obtained by heating borax (Na 2 B 4 O 7 ·10H 2 O) with [|carbon], although other methods are used if high-purity boron is required. <span style="font-family: 'times new roman',serif; font-size: 16px;">Boron is used in pyrotechnics and flares to produce a green color. Boron has also been used in some rockets as an ignition source. Boron-10, one of the naturally occurring [|isotopes] of boron, is a good absorber of [|neutrons] and is used in the control rods of nuclear reactors, as a radiation shield and as a neutron detector. Boron filaments are used in the aerospace industry because of their high-strength and lightweight.
 * Atomic Number:** 5
 * Atomic Weight:** 10.811
 * Melting Point:** 2348 K (2075°C or 3767°F)
 * Boiling Point:** 4273 K (4000°C or 7232°F)
 * Density:** 2.37 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Semi-metal
 * Period Number:** 2 **Group Number:** 13 **Group Name:** none

<span style="font-family: 'times new roman',serif; font-size: 16px;">Boron forms several commercially important compounds. The most important boron compound is sodium borate pentahydrate (Na 2 B 4 O 7 ·5H 2 O). Large amounts of this compound are used in the manufacture of fiberglass insulation and sodium perborate bleach. The second most important compound is boric acid (H 3 BO 3 ), which is used to manufacture textile fiberglass and is used in cellulose insulation as a flame retardant. Sodium borate decahydrate (Na 2 B 4 O 7 ·10H 2 O), better known as borax, is the third most important boron compound. Borax is used in laundry products and as a mild antiseptic. Borax is also a key ingredient in a substance known as [|Oobleck], a strange material 6th grade students experiment with while participating in [|Jefferson Lab's BEAMS program]. Other boron compounds are used to make borosilicate glasses, enamels for covering steel and as a potential medicine for treating arthritis. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 4.44 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 8.298 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 1s 2 2s 2 2p 1
 * [|Electron Shell Configuration] :**

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 6 C Carbon 12.0107 **Atomic Number:** 6 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for charcoal, **carbo**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Carbon is pronounced as **KAR-ben**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Carbon, the sixth [|most abundant element in the universe], has been known since ancient times. Carbon is most commonly obtained from coal deposits, although it usually must be processed into a form suitable for commercial use. Three naturally occurring allotropes of carbon are known to exist: amorphous, graphite and diamond. <span style="font-family: 'times new roman',serif; font-size: 16px;">Amorphous carbon is formed when a material containing carbon is burned without enough [|oxygen] for it to burn completely. This black soot, also known as lampblack, gas black, channel black or carbon black, is used to make inks, paints and rubber products. It can also be pressed into shapes and is used to form the cores of most dry cell batteries, among other things. <span style="font-family: 'times new roman',serif; font-size: 16px;">Graphite, one of the softest materials known, is a form of carbon that is primarily used as a lubricant. Although it does occur naturally, most commercial graphite is produced by treating petroleum coke, a black tar residue remaining after the refinement of crude oil, in an oxygen-free oven. Naturally occurring graphite occurs in two forms, alpha and beta. These two forms have identical physical properties but different crystal structures. All artificially produced graphite is of the alpha type. In addition to its use as a lubricant, graphite, in a form known as coke, is used in large amounts in the production of steel. Coke is made by heating soft coal in an oven without allowing oxygen to mix with it. Although commonly called [|lead], the black material used in pencils is actually graphite. <span style="font-family: 'times new roman',serif; font-size: 16px;">Diamond, the third naturally occurring form of carbon, is one of the hardest substances known. Although naturally occurring diamond is typically used for jewelry, most commercial quality diamonds are artificially produced. These small diamonds are made by squeezing graphite under high temperatures and pressures for several days or weeks and are primarily used to make things like diamond tipped saw blades. Although they posses very different physical properties, graphite and diamond differ only in their crystal structure. <span style="font-family: 'times new roman',serif; font-size: 16px;">A fourth allotrope of carbon, known as white carbon, was produced in 1969. It is a transparent material that can split a single beam of light into two beams, a property known as birefringence. Very little is known about this form of carbon. <span style="font-family: 'times new roman',serif; font-size: 16px;">Large molecules consisting only of carbon, known as buckminsterfullerenes, or buckyballs, have recently been discovered and are currently the subject of much scientific interest. A single buckyball consists of 60 or 70 carbon atoms (C 60 or C 70 ) linked together in a structure that looks like a soccer ball. They can trap other atoms within their framework, appear to be capable of withstanding great pressures and have magnetic and superconductive properties. <span style="font-family: 'times new roman',serif; font-size: 16px;">Carbon-14, a radioactive [|isotope] of carbon with a [|half-life] of 5,730 years, is used to find the age of formerly living things through a process known as radiocarbon dating. The theory behind carbon dating is fairly simple. Scientists know that a small amount of naturally occurring carbon is carbon-14. Although carbon-14 decays into [|nitrogen] -14 through [|beta decay], the amount of carbon-14 in the environment remains constant because new carbon-14 is always being created in the upper atmosphere by cosmic rays. Living things tend to ingest materials that contain carbon, so the percentage of carbon-14 within living things is the same as the percentage of carbon-14 in the environment. Once an organism dies, it no longer ingests much of anything. The carbon-14 within that organism is no longer replaced and the percentage of carbon-14 begins to decrease as it decays. By measuring the percentage of carbon-14 in the remains of an organism, and by assuming that the natural abundance of carbon-14 has remained constant over time, scientists can estimate when that organism died. For example, if the concentration of carbon-14 in the remains of an organism is half of the natural concentration of carbon-14, a scientist would estimate that the organism died about 5,730 years ago, the half-life of carbon-14. <span style="font-family: 'times new roman',serif; font-size: 16px;">There are nearly ten million known carbon compounds and an entire branch of chemistry, known as organic chemistry, is devoted to their study. Many carbon compounds are essential for life as we know it. Some of the most common carbon compounds are: carbon dioxide (CO 2 ), carbon monoxide (CO), carbon disulfide (CS 2 ), chloroform (CHCl 3 ), carbon tetrachloride (CCl 4 ), methane (CH 4 ), ethylene (C 2 H 4 ), acetylene (C 2 H 2 ), benzene (C 6 H 6 ), ethyl alcohol (C 2 H 5 OH) and acetic acid (CH 3 COOH). <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.00×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2.8×10 1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 11.260 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +4, +2, -4 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 2
 * Atomic Weight:** 12.0107
 * Melting Point:** 3823 K (3550°C or 6422°F)
 * Boiling Point:** 4098 K (3825°C or 6917°F)
 * Density:** 2.2670 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Non-metal
 * Period Number:** 2 **Group Number:** 14 **Group Name:** none

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 7 N Nitrogen 14.0067 **Atomic Number:** 7 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek words **nitron** and **genes**, which together mean "saltpetre forming." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Nitrogen is pronounced as **NYE-treh-gen**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Nitrogen was discovered by the Scottish physician Daniel Rutherford in 1772. It is the fifth [|most abundant element in the universe] and makes up about 78% of the [|earth's atmosphere], which contains an estimated 4,000 trillion tons of the gas. Nitrogen is obtained from liquefied air through a process known as fractional distillation. <span style="font-family: 'times new roman',serif; font-size: 16px;">The largest use of nitrogen is for the production of ammonia (NH 3 ). Large amounts of nitrogen are combined with [|hydrogen] to produce ammonia in a method known as the Haber process. Large amounts of ammonia are then used to create fertilizers, explosives and, through a process known as the Ostwald process, nitric acid (HNO 3 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Nitrogen gas is largely inert and is used as a protective shield in the semiconductor industry and during certain types of welding and soldering operations. Oil companies use high pressure nitrogen to help force crude oil to the surface. [|Liquid nitrogen] is an inexpensive cryogenic liquid used for refrigeration, preservation of biological samples and for low temperature scientific experimentation. [|Jefferson Lab's Frostbite Theater] features videos of many basic liquid nitrogen experiments, such as this one: <span style="font-family: 'times new roman',serif; font-size: 16px;">#video <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.9×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 5×10 -1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 14.534 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +4, +3, +2, +1, -1, -2, -3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 3
 * Atomic Weight:** 14.0067
 * Melting Point:** 63.15 K (-210.00°C or -346.00°F)
 * Boiling Point:** 77.36 K (-195.79°C or -320.44°F)
 * Density:** 0.0012506 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 2 **Group Number:** 15 **Group Name:** Pnictogen

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 8 O Oxygen 15.9994 **Atomic Number:** 8 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the greek words **oxys** and **genes**, which together mean "acid forming." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Oxygen is pronounced as **OK-si-jen**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Oxygen had been produced by several chemists prior to its discovery in 1774, but they failed to recognize it as a distinct element. Joseph Priestley and Carl Wilhelm Scheele both independently discovered oxygen, but Priestly is usually given credit for the discovery. They were both able to produce oxygen by heating mercuric oxide (HgO). Priestley called the gas produced in his experiments 'dephlogisticated air' and Scheele called his 'fire air'. The name oxygen was created by Antoine Lavoisier who incorrectly believed that oxygen was necessary to form all acids. <span style="font-family: 'times new roman',serif; font-size: 16px;">Oxygen is the third [|most abundant element in the universe] and makes up nearly 21% of the [|earth's atmosphere]. Oxygen accounts for nearly half of the mass of the [|earth's crust], two thirds of the mass of the human body and nine tenths of the mass of water. Large amounts of oxygen can be extracted from liquefied air through a process known as fractional distillation. Oxygen can also be produced through the electrolysis of water or by heating potassium chlorate (KClO 3 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Oxygen is a highly reactive element and is capable of combining with most other elements. It is required by most living organisms and for most forms of combustion. Impurities in molten pig [|iron] are burned away with streams of high pressure oxygen to produce steel. Oxygen can also be combined with acetylene (C 2 H 2 ) to produce an extremely hot flame used for welding. Liquid oxygen, when combined with liquid [|hydrogen], makes an excellent rocket fuel. Ozone (O 3 ) forms a thin, protective layer around the earth that shields the surface from the sun's ultraviolet radiation. Oxygen is also a component of hundreds of thousands of organic compounds. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 4.61×10 5 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 8.57×10 5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 13.618 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** -2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 4
 * Atomic Weight:** 15.9994
 * Melting Point:** 54.36 K (-218.79°C or -361.82°F)
 * Boiling Point:** 90.20 K (-182.95°C or -297.31°F)
 * Density:** 0.001429 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 2 **Group Number:** 16 **Group Name:** Chalcogen

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 9 F Fluorine 18.9984032 **Atomic Number:** 9 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin and French words for flow, **fluere**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Fluorine is pronounced as **FLU-eh-reen or as FLU-eh-rin**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Fluorine is the most reactive of all elements and no chemical substance is capable of freeing fluorine from any of its compounds. For this reason, fluorine does not occur free in nature and was extremely difficult for scientists to isolate. The first recorded use of a fluorine compound dates to around 1670 to a set of instructions for etching glass that called for Bohemian emerald (CaF 2 ). Chemists attempted to identify the material that was capable of etching glass and George Gore was able to produce a small amount of fluorine through an electrolytic process in 1869. Unknown to Gore, fluorine gas explosively combines with [|hydrogen] gas. That is exactly what happened in Gore's experiment when the fluorine gas that formed on one electrode combined with the hydrogen gas that formed on the other electrode. Ferdinand Frederic Henri Moissan, a French chemist, was the first to successfully isolate fluorine in 1886. He did this through the electrolysis of potassium fluoride (KF) and hydrofluoric acid (HF). He also completely isolated the fluorine gas from the hydrogen gas and he built his electrolysis device completely from [|platinum]. His work was so impressive that he was awarded the Nobel Prize for chemistry in 1906. Today, fluorine is still produced through the electrolysis of potassium fluoride and hydrofluoric acid as well as through the electrolysis of molten potassium acid fluoride (KHF 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Fluorine is added to city water supplies in the proportion of about one part per million to help prevent tooth decay. Sodium fluoride (NaF), stannous(II) fluoride (SnF 2 ) and sodium monofluorophosphate (Na 2 PO 3 F) are all fluorine compounds added to toothpaste, also to help prevent tooth decay. Hydrofluoric acid (HF) is used to etch glass, including most of the glass used in light bulbs. Uranium hexafluoride (UF 6 ) is used to separate [|isotopes] of [|uranium]. Crystals of calcium fluoride (CaF 2 ), also known as fluorite and fluorspar, are used to make lenses to focus infrared light. Fluorine joins with [|carbon] to form a class of compounds known as fluorocarbons. Some of these compounds, such as dichlorodifluoromethane (CF 2 Cl 2 ), were widely used in air conditioning and refrigeration systems and in aerosol spray cans, but have been phased out due to the damage they were causing to the earth's ozone layer. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 5.85×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 17.423 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** -1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p
 * Atomic Weight:** 18.9984032
 * Melting Point:** 53.53 K (-219.62°C or -363.32°F)
 * Boiling Point:** 85.03 K (-188.12°C or -306.62°F)
 * Density:** 0.001696 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 2 **Group Number:** 17 **Group Name:** Halogen

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 10 Ne Neon 20.1797 **Atomic Number:** 10 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for new, **neos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Neon is pronounced as **NEE-on**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Neon was discovered by Sir William Ramsay, a Scottish chemist, and Morris M. Travers, an English chemist, shortly after their discovery of the element [|krypton] in 1898. Like krypton, neon was discovered through the study of liquefied air. Although neon is the fourth [|most abundant element in the universe], only 0.0018% of the [|earth's atmosphere] is neon. <span style="font-family: 'times new roman',serif; font-size: 16px;">The largest use for neon gas is in advertising signs. Neon is also used to make high voltage indicators and is combined with [|helium] to make helium-neon lasers. Liquid neon is used as a cryogenic refrigerant. Neon is highly inert and forms no known compounds, although there is some evidence that it could form a compound with [|flourine]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 5×10 -3 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.2×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 21.565 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** 0 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6
 * Atomic Weight:** 20.1797
 * Melting Point:** 24.56 K (-248.59°C or -415.46°F)
 * Boiling Point:** 27.07 K (-246.08°C or -410.94°F)
 * Density:** 0.0008999 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 2 **Group Number:** 18 **Group Name:** Noble Gas

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 11 Na Sodium 22.98976928 **Atomic Number:** 11 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the English word **soda** and from the Medieval Latin word **sodanum**, which means "headache remedy." Sodium's chemical symbol comes from the Latin word for sodium carbonate, **natrium**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Sodium is pronounced as **SO-dee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although sodium is the sixth [|most abundant element on earth] and comprises about 2.6% of the earth's crust, it is a very reactive element and is never found free in nature. Pure sodium was first isolated by Sir Humphry Davy in 1807 through the electrolysis of caustic soda (NaOH). Since sodium can ignite on contact with water, it must be stored in a moisture free environment. <span style="font-family: 'times new roman',serif; font-size: 16px;">Sodium is used in the production of [|titanium], sodamide, sodium cyanide, sodium peroxide, and sodium hydride. Liquid sodium has been used as a coolant for nuclear reactors. Sodium vapor is used in streetlights and produces a brilliant yellow light. <span style="font-family: 'times new roman',serif; font-size: 16px;">Sodium also forms many useful compounds. Some of the most common are: table salt (NaCl), soda ash (Na 2 CO 3 ), baking soda (NaHCO 3 ), caustic soda (NaOH), Chile saltpeter (NaNO 3 ) and borax (Na 2 B 4 O 7 ·10H 2 O). <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.36×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.08×10 4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 5.139 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 1
 * Atomic Weight:** 22.98976928
 * Melting Point:** 370.95 K (97.80°C or 208.04°F)
 * Boiling Point:** 1156 K (883°C or 1621°F)
 * Density:** 0.97 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 3 **Group Number:** 1 **Group Name:** Alkali Metal

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 12 Mg Magnesium 24.3050 **Atomic Number:** 12 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** For Magnesia, a district in the region of Thessaly, Greece. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Magnesium is pronounced as **mag-NEE-zhi-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although it is the eighth [|most abundant element in the universe] and the seventh [|most abundant element in the earth's crust], magnesium is never found free in nature. Magnesium was first isolated by Sir Humphry Davy, an English chemist, through the electrolysis of a mixture of magnesium oxide (MgO) and mercuric oxide (HgO) in 1808. Today, magnesium can be extracted from the minerals dolomite (CaCO 3 ·MgCO 3 ) and carnallite (KCl·MgCl 2 ·6H 2 O), but is most often obtained from seawater. Every cubic kilometer of seawater contains about 1.3 billion kilograms of magnesium (12 billion pounds per cubic mile). <span style="font-family: 'times new roman',serif; font-size: 16px;">Magnesium burns with a brilliant white light and is used in pyrotechnics, flares and photographic flashbulbs. Magnesium is the lightest metal that can be used to build things, although its use as a structural material is limited since it burns at relatively low temperatures. Magnesium is frequently alloyed with [|aluminum], which makes aluminum easier to roll, extrude and weld. Magnesium-aluminum alloys are used where strong, lightweight materials are required, such as in airplanes, missiles and rockets. Cameras, horseshoes, baseball catchers' masks and snowshoes are other items that are made from magnesium alloys. <span style="font-family: 'times new roman',serif; font-size: 16px;">Magnesium oxide (MgO), also known as magnesia, is the second [|most abundant compound in the earth's crust]. Magnesium oxide is used in some antacids, in making crucibles and insulating materials, in refining some metals from their ores and in some types of cements. When combined with water (H 2 O), magnesia forms magnesium hydroxide (Mg(OH) 2 ), better known as milk of magnesia, which is commonly used as an antacid and as a laxative. <span style="font-family: 'times new roman',serif; font-size: 16px;">Hydrated magnesium sulphate (MgSO 4 ·7H 2 O), better known as Epsom salt, was discovered in 1618 by a farmer in Epsom, England, when his cows refused to drink the water from a certain mineral well. He tasted the water and found that it tasted very bitter. He also noticed that it helped heal scratches and rashes on his skin. Epsom salt is still used today to treat minor skin abrasions. <span style="font-family: 'times new roman',serif; font-size: 16px;">Other magnesium compounds include magnesium carbonate (MgCO 3 ) and magnesium fluoride (MgF 2 ). Magnesium carbonate is used to make some types of paints and inks and is added to table salt to prevent caking. A thin film of magnesium fluoride is applied to optical lenses to help reduce glare and reflections. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.33×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.29×10 3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.646 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 13 Al Aluminum 26.9815386 **Atomic Number:** 13 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for alum, **alumen**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Aluminum is pronounced as **ah-LOO-men-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although aluminum is the [|most abundant metal in the earth's crust], it is never found free in nature. All of the earth's aluminum has combined with other elements to form compounds. Two of the most common compounds are alum, such as potassium aluminum sulfate (KAl(SO 4 ) 2 ·12H 2 O), and aluminum oxide (Al 2 O 3 ). About 8.2% of the earth's crust is composed of aluminum. <span style="font-family: 'times new roman',serif; font-size: 16px;">Scientists suspected than an unknown metal existed in alum as early as 1787, but they did not have a way to extract it until 1825. Hans Christian Oersted, a Danish chemist, was the first to produce tiny amounts of aluminum. Two years later, Friedrich Wöhler, a German chemist, developed a different way to obtain aluminum. By 1845, he was able to produce samples large enough to determine some of aluminum's basic properties. Wöhler's method was improved in 1854 by Henri Étienne Sainte-Claire Deville, a French chemist. Deville's process allowed for the commercial production of aluminum. As a result, the price of aluminum dropped from around $1200 per kilogram in 1852 to around $40 per kilogram in 1859. Unfortunately, aluminum remained too expensive to be widely used. <span style="font-family: 'times new roman',serif; font-size: 16px;">Two important developments in the 1880s greatly increased the availability of aluminum. The first was the invention of a new process for obtaining aluminum from aluminum oxide. Charles Martin Hall, an American chemist, and Paul L. T. Héroult, a French chemist, each invented this process independently in 1886. The second was the invention of a new process that could cheaply obtain aluminum oxide from bauxite. Bauxite is an ore that contains a large amount of aluminum hydroxide (Al 2 O 3 ·3H 2 O), along with other compounds. Karl Joseph Bayer, an Austrian chemist, developed this process in 1888. The Hall-Héroult and Bayer processes are still used today to produce nearly all of the world's aluminum. <span style="font-family: 'times new roman',serif; font-size: 16px;">With an easy way to extract aluminum from aluminum oxide and an easy way to extract large amounts of aluminum oxide from bauxite, the era of inexpensive aluminum had begun. In 1888, Hall formed the Pittsburgh Reduction Company, which is now known as the Aluminum Company of America, or Alcoa. When it opened, his company could produce about 25 kilograms of aluminum a day. By 1909, his company was producing about 41,000 kilograms of aluminum a day. As a result of this huge increase of supply, the price of aluminum fell rapidly to about $0.60 per kilogram. <span style="font-family: 'times new roman',serif; font-size: 16px;">Today, aluminum and aluminum alloys are used in a wide variety of products: cans, foils and kitchen utensils, as well as parts of airplanes, rockets and other items that require a strong, light material. Although it doesn't conduct electricity as well as [|copper], it is used in electrical transmission lines because of its light weight. It can be deposited on the surface of glass to make mirrors, where a thin layer of aluminum oxide quickly forms that acts as a protective coating. Aluminum oxide is also used to make synthetic rubies and sapphires for lasers. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 8.23×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2×10 -3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 5.986 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 14 Si Silicon 28.0855 **Atomic Number:** 14 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for flint, **silex**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Silicon is pronounced as **SIL-ee-ken**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Silicon was discovered by Jöns Jacob Berzelius, a Swedish chemist, in 1824 by heating chips of [|potassium] in a silica container and then carefully washing away the residual by-products. Silicon is the seventh [|most abundant element in the universe] and the second [|most abundant element in the earth's crust]. Today, silicon is produced by heating sand (SiO 2 ) with [|carbon] to temperatures approaching 2200°C. <span style="font-family: 'times new roman',serif; font-size: 16px;">Two allotropes of silicon exist at room temperature: amorphous and crystalline. Amorphous appears as a brown powder while crystalline silicon has a metallic luster and a grayish color. Single crystals of crystalline silicon can be grown with a process known as the Czochralski process. These crystals, when doped with elements such as [|boron], [|gallium] , [|germanium] , [|phosphorus] or [|arsenic] , are used in the manufacture of solid-state electronic devices, such as transistors, solar cells, rectifiers and microchips. <span style="font-family: 'times new roman',serif; font-size: 16px;">Silicon dioxide (SiO 2 ), silicon's most common compound, is the [|most abundant compound in the earth's crust]. It commonly takes the form of ordinary sand, but also exists as quartz, rock crystal, amethyst, agate, flint, jasper and opal. Silicon dioxide is extensively used in the manufacture of glass and bricks. Silica gel, a colloidal form of silicon dioxide, easily absorbs moisture and is used as a desiccant. <span style="font-family: 'times new roman',serif; font-size: 16px;">Silicon forms other useful compounds. Silicon carbide (SiC) is nearly as hard as diamond and is used as an abrasive. Sodium silicate (Na 2 SiO 3 ), also known as water glass, is used in the production of soaps, adhesives and as an egg preservative. Silicon tetrachloride (SiCl 4 ) is used to create smoke screens. Silicon is also an important ingredient in silicone, a class of material that is used for such things as lubricants, polishing agents, electrical insulators and medical implants. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.82×10 5 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2.2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 8.152 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +4, +2, -4 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 2
 * Atomic Weight:** 24.3050
 * Melting Point:** 923 K (650°C or 1202°F)
 * Boiling Point:** 1363 K (1090°C or 1994°F)
 * Density:** 1.74 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 3 **Group Number:** 2 **Group Name:** Alkaline Earth Metal
 * Atomic Weight:** 26.9815386
 * Melting Point:** 933.437 K (660.323°C or 1220.581°F)
 * Boiling Point:** 2792 K (2519°C or 4566°F)
 * Density:** 2.70 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 3 **Group Number:** 13 **Group Name:** none
 * Atomic Weight:** 28.0855
 * Melting Point:** 1687 K (1414°C or 2577°F)
 * Boiling Point:** 3538 K (3265°C or 5909°F)
 * Density:** 2.3296 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Semi-metal
 * Period Number:** 3 **Group Number:** 14 **Group Name:** none

<span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 15 P Phosphorus 30.973762 **Atomic Number:** 15 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for light bearing, **phosphoros**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Phosphorus is pronounced as **FOS-fer-es**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">In what is perhaps the most disgusting method of discovering an element, phosphorus was first isolated in 1669 by Hennig Brand, a German physician and alchemist, by boiling, filtering and otherwise processing as many as 60 buckets of urine. Thankfully, phosphorus is now primarily obtained from phosphate rock (Ca 3 (PO 4 ) 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Phosphorus has three main allotropes: white, red and black. White phosphorus is poisonous and can spontaneously ignite when it comes in contact with air. For this reason, white phosphorus must be stored under water and is usually used to produce phosphorus compounds. Red phosphorus is formed by heating white phosphorus to 250°C (482°F) or by exposing white phosphorus to sunlight. Red phosphorus is not poisonous and is not as dangerous as white phosphorus, although frictional heating is enough to change it back to white phosphorus. Red phosphorus is used in safety matches, fireworks, smoke bombs and pesticides. Black phosphorus is also formed by heating white phosphorus, but a [|mercury] [|catalyst] and a seed crystal of black phosphorus are required. Black phosphorus is the least reactive form of phosphorus and has no significant commercial uses. <span style="font-family: 'times new roman',serif; font-size: 16px;">Phosphoric acid (H 3 PO 4 ) is used in soft drinks and to create many phosphate compounds, such as triple superphosphate fertilizer (Ca(H 2 PO 4 ) 2 ·H 2 O). Trisodium phosphate (Na 3 PO 4 ) is used as a cleaning agent and as a water softener. Calcium phosphate (Ca 3 (PO 4 ) 2 ) is used to make china and in the production of baking powder. Some phosphorus compounds glow in the dark or emit light in response to absorbing radiation and are used in fluorescent light bulbs and television sets. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.05×10 3 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 6×10 -2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 10.487 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +3, -3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 16 S Sulfur 32.065 **Atomic Number:** 16 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Sanskrit word **sulvere** and the Latin word **sulphurium**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Sulfur is pronounced as **SUL-fer**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Sulfur, the tenth [|most abundant element in the universe], has been known since ancient times. Sometime around 1777, Antoine Lavoisier convinced the rest of the scientific community that sulfur was an element. Sulfur is a component of many common minerals, such as galena (PbS), gypsum (CaSO 4 ·2(H 2 O), pyrite (FeS 2 ), sphalerite (ZnS or FeS), cinnabar (HgS), stibnite (Sb 2 S 3 ), epsomite (MgSO 4 ·7(H 2 O)), celestite (SrSO 4 ) and barite (BaSO 4 ). Nearly 25% of the sulfur produced today is recovered from petroleum refining operations and as a byproduct of extracting other materials from sulfur containing ores. The majority of the sulfur produced today is obtained from underground deposits, usually found in conjunction with salt deposits, with a process known as the Frasch process. <span style="font-family: 'times new roman',serif; font-size: 16px;">Sulfur is a pale yellow, odorless and brittle material. It displays three allotropic forms: orthorhombic, monoclinic and amorphous. The orthorhombic form is the most stable form of sulfur. Monoclinic sulfur exists between the temperatures of 96°C and 119°C and reverts back to the orthorhombic form when cooled. Amorphous sulfur is formed when molten sulfur is quickly cooled. Amorphous sulfur is soft and elastic and eventually reverts back to the orthorhombic form. <span style="font-family: 'times new roman',serif; font-size: 16px;">Most of the sulfur that is produced is used in the manufacture of sulfuric acid (H 2 SO 4 ). Large amounts of sulfuric acid, nearly 40 million tons, are used each year to make fertilizers, lead-acid batteries, and in many industrial processes. Smaller amounts of sulfur are used to vulcanize natural rubbers, as an insecticide (the Greek poet Homer mentioned "pest-averting sulphur" nearly 2,800 years ago!), in the manufacture of gunpowder and as a dying agent. <span style="font-family: 'times new roman',serif; font-size: 16px;">In addition to sulfuric acid, sulfur forms other interesting compounds. Hydrogen sulfide (H 2 S) is a gas that smells like rotten eggs. Sulfur dioxide (SO 2 ), formed by burning sulfur in air, is used as a bleaching agent, solvent, disinfectant and as a refrigerant. When combined with water (H 2 O), sulfur dioxide forms sulfurous acid (H 2 SO 3 ), a weak acid that is a major component of acid rain. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 3.50×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 9.05×10 2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:4** <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 10.360 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +6, +4, -2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 4 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 17 Cl Chlorine 35.453 **Atomic Number:** 17 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for greenish yellow, **chloros**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Chlorine is pronounced as **KLOR-een or as KLOR-in**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Since it combines directly with nearly every element, chlorine is never found free in nature. Chlorine was first produced by Carl Wilhelm Scheele, a Swedish chemist, when he combined the mineral pyrolusite (MnO 2 ) with hydrochloric acid (HCl) in 1774. Although Scheele thought the gas produced in his experiment contained [|oxygen], Sir Humphry Davy proved in 1810 that it was actually a distinct element. Today, most chlorine is produced through the electrolysis of aqueous sodium chloride (NaCl). <span style="font-family: 'times new roman',serif; font-size: 16px;">Chlorine is commonly used as an antiseptic and is used to make drinking water safe and to treat swimming pools. Large amounts of chlorine are used in many industrial processes, such as in the production of paper products, plastics, dyes, textiles, medicines, antiseptics, insecticides, solvents and paints. <span style="font-family: 'times new roman',serif; font-size: 16px;">Two of the most familiar chlorine compounds are sodium chloride (NaCl) and hydrogen chloride (HCl). Sodium chloride, commonly known as table salt, is used to season food and in some industrial processes. Hydrogen chloride, when mixed with water (H 2 O), forms hydrochloric acid, a strong and commercially important acid. Other chlorine compounds include: chloroform (CHCl 3 ), carbon tetrachloride (CCl 4 ), potassium chloride (KCl), lithium chloride (LiCl), magnesium chloride (MgCl 2 ) and chlorine dioxide (ClO 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Chlorine is a very dangerous material. Liquid chlorine burns the skin and gaseous chlorine irritates the mucus membranes. Concentrations of the gas as low as 3.5 parts per million can be detected by smell while concentrations of 1000 parts per million can be fatal after a few deep breaths. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.45×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.94×10 4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 12.968 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +7, +5, +1, -1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 5 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 18 Ar Argon 39.948 **Atomic Number:** 18 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for inactive, **argos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Argon is pronounced as **AR-gon**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Argon was discovered by Sir William Ramsay, a Scottish chemist, and Lord Rayleigh, an English chemist, in 1894. Argon makes up 0.93% of the [|earth's atmosphere], making it the third most abundant gas. Argon is obtained from the air as a byproduct of the production of [|oxygen] and [|nitrogen]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Argon is frequently used when an inert atmosphere is needed. It is used to fill incandescent and fluorescent light bulbs to prevent oxygen from corroding the hot filament. Argon is also used to form inert atmospheres for arc welding, growing semiconductor crystals and processes that require shielding from other atmospheric gases. <span style="font-family: 'times new roman',serif; font-size: 16px;">Once thought to be completely inert, argon is known to form at least one compound. The synthesis of argon fluorohydride (HArF) was [|reported] by Leonid Khriachtchev, Mika Pettersson, Nino Runeberg, Jan Lundell and Markku Räsänen in August of 2000. Stable only at very low temperatures, argon fluorohydride begins to decompose once it warms above -246°C (-411°F). Because of this limitation, argon fluorohydride has no uses outside of basic scientific research. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 3.5 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 4.5×10 -1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 15.760 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** 0 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 19 K Potassium 39.0983 **Atomic Number:** 19 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the English word **potash**. Potassium's chemical symbol comes from the Latin word for alkali,**kalium**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Potassium is pronounced as **poh-TASS-ee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although potassium is the eighth [|most abundant element on earth] and comprises about 2.1% of the earth's crust, it is a very reactive element and is never found free in nature. Metallic potassium was first isolated by Sir Humphry Davy in 1807 through the electrolysis of molten caustic potash (KOH). A few months after discovering potassium, Davy used the same method to isolate [|sodium]. Potassium can be obtained from the minerals sylvite (KCl), carnallite (KCl·MgCl 2 ·6H 2 O), langbeinite (K 2 Mg 2 (SO 4 ) 3 ) and polyhalite (K 2 Ca 2 Mg(SO 4 ) 4 ·2H 2 O). These minerals are often found in ancient lake and sea beds. Caustic potash, another important source of potassium, is primarily mined in Germany, New Mexico, California and Utah. <span style="font-family: 'times new roman',serif; font-size: 16px;">Pure potassium is a soft, waxy metal that can be easily cut with a knife. It reacts with [|oxygen] to form potassium superoxide (KO 2 ) and with water to form potassium hydroxide (KOH), [|hydrogen] gas and heat. Enough heat is produced to ignite the hydrogen gas. To prevent it from reacting with the oxygen and water in the air, samples of metallic potassium are usually stored submerged in mineral oil. <span style="font-family: 'times new roman',serif; font-size: 16px;">Potassium forms an alloy with sodium (NaK) that is used as a heat transfer medium in some types of nuclear reactors. <span style="font-family: 'times new roman',serif; font-size: 16px;">Potassium forms many important compounds. Potassium chloride (KCl) is the most common potassium compound. It is used in fertilizers, as a salt substitute and to produce other chemicals. Potassium hydroxide (KOH) is used to make soaps, detergents and drain cleaners. Potassium carbonate (KHCO 3 ), also known as pearl ash, is used to make some types of glass and soaps and is obtained commercially as a byproduct of the production of ammonia. Potassium superoxide (KO 2 ) can create oxygen from water vapor (H 2 O) and carbon dioxide (CO 2 ) through the following reaction: 2KO 2 + H 2 O + 2CO 2 => 2KHCO 3 + O 2. It is used in respiratory equipment and is produced by burning potassium metal in dry air. Potassium nitrate (KNO 3 ), also known as saltpeter or nitre, is used in fertilizers, match heads and pyrotechnics. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.09×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 3.99×10 2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 4.341 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 20 Ca Calcium 40.078 **Atomic Number:** 20 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for lime, **calx**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Calcium is pronounced as **KAL-see-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although calcium is the fifth [|most abundant element in the earth's crust], it is never found free in nature since it easily forms compounds by reacting with [|oxygen] and water. Metallic calcium was first isolated by Sir Humphry Davy in 1808 through the electrolysis of a mixture of lime (CaO) and mercuric oxide (HgO). Today, metallic calcium is obtained by displacing calcium atoms in lime with atoms of [|aluminum] in hot, low-pressure containers. About 4.2% of the [|earth's crust] is composed of calcium. <span style="font-family: 'times new roman',serif; font-size: 16px;">Due to its high reactivity with common materials, there is very little demand for metallic calcium. It is used in some chemical processes to refine [|thorium], [|uranium] and [|zirconium]. Calcium is also used to remove [|oxygen], [|sulfur] and [|carbon] from certain alloys. Calcium can be alloyed with [|aluminum], [|beryllium] , [|copper] , [|lead] and [|magnesium]. Calcium is also used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes. <span style="font-family: 'times new roman',serif; font-size: 16px;">Calcium carbonate (CaCO 3 ) is one of the common compounds of calcium. It is heated to form quicklime (CaO) which is then added to water (H 2 O). This forms another material known as slaked lime (Ca(OH) 2 ) which is an inexpensive base material used throughout the chemical industry. Chalk, marble and limestone are all forms of calcium carbonate. Calcium carbonate is used to make white paint, cleaning powder, toothpaste and stomach antacids, among other things. Other common compounds of calcium include: calcium sulfate (CaSO 4 ), also known as gypsum, which is used to make dry wall and plaster of Paris, calcium nitrate (Ca(NO 3 ) 2 ), a naturally occurring fertilizer and calcium phosphate (Ca 3 (PO 4 ) 2 ), the main material found in bones and teeth. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 4.15×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 4.12×10 2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.113 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 21 Sc Scandium 44.955912 **Atomic Number:** 21 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for Scandinavia. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Scandium is pronounced as **SKAN-dee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Scandium was discovered by Lars Fredrik Nilson, a Swedish chemist, in 1879 while attempting to produce a sample of pure ytterbia from 10 kilograms of the mineral euxenite ((Y, Ca, Er, La, Ce, U, Th)(Nb, Ta, Ti) 2 O 6 ). Scandium can be obtained from the minerals thortveitite ((Sc, Y) 2 Si 2 O 7 ), bazzite (Be 3 (Sc, Al) 2 Si 6 O 18 ) and wiikite, but is usually obtained as a byproduct of refining [|uranium]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Metallic scandium was first produced in 1937 and the first pound (0.45 kilograms) of pure scandium was produced in 1960. Scandium is a soft, light metal that might have applications in the aerospace industry. With a cost of $270 per gram ($122,500 per pound), scandium is too expensive for widespread use. <span style="font-family: 'times new roman',serif; font-size: 16px;">Alloys of scandium and [|aluminum] are used in some kinds of athletic equipment, such as aluminum baseball bats, bicycle frames and lacrosse sticks. It is expected that scandium-aluminum alloys will be important in the manufacture of fuel cells. <span style="font-family: 'times new roman',serif; font-size: 16px;">Scientists have only studied a few compounds of scandium. About 20 kilograms (44 pounds) of scandium oxide (Sc 2 O 3 ), also known as scandia, are used each year in the United States in the production of high intensity lights. Scandium iodide (ScI 3 ) is added to [|mercury] vapor lamps so that they will emit light that closely resembles sunlight. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.2×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 6×10 -7 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.561 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 22 Ti Titanium 47.867 **Atomic Number:** 22 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word [|**Titans**], the mythological "first sons of the Earth." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Titanium is pronounced as **tie-TAY-nee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Titanium was discovered in 1791 by the Reverend William Gregor, an English pastor. Pure titanium was first produced by Matthew A. Hunter, an American metallurgist, in 1910. Titanium is the ninth [|most abundant element in the earth's crust] and is primarily found in the minerals rutile (TiO 2 ), ilmenite (FeTiO 3 ) and sphene (CaTiSiO 5 ). Titanium makes up about 0.57% of the earth's crust. <span style="font-family: 'times new roman',serif; font-size: 16px;">Titanium is a strong, light metal. It is as strong as steel and twice as strong as [|aluminum], but is 45% lighter than steel and only 60% heavier than aluminum. Titanium is not easily corroded by sea water and is used in propeller shafts, rigging and other parts of boats that are exposed to sea water. Titanium and titanium alloys are used in airplanes, missiles and rockets where strength, low weight and resistance to high temperatures are important. Since titanium does not react within the human body, it is used to create artificial hips, pins for setting bones and for other biological implants. Unfortunately, the high cost of titanium has limited its widespread use. <span style="font-family: 'times new roman',serif; font-size: 16px;">Titanium oxide (TiO 2 ) is used as a pigment to create white paint and accounts for the largest use of the element. Pure titanium oxide is relatively clear and is used to create titania, an artificial gemstone. Titanium tetrachloride (TiCl 4 ), another titanium compound, has been used to make smoke screens. <span style="font-family: 'times new roman',serif; font-size: 16px;">A final bit of titanium trivia -- titanium is the only element that will burn in an atmosphere of pure [|nitrogen]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 5.65×10 3 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1×10 -3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 5 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.828 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +4, +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 <span style="background-color: #ffffff; font-family: 'times new roman',serif; font-size: 16px;">4s <span style="background-color: #ffffff; font-family: 'times new roman',serif; font-size: 12px;">2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 23 V Vanadium 50.9415 **Atomic Number:** 23 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for the Scandinavian goddess [|**Vanadis**]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Vanadium is pronounced as **veh-NAY-dee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Vanadium was discovered by Andrés Manuel del Rio, a Mexican chemist, in 1801. Rio sent samples of vanadium ore and a letter describing his methods to the Institute de France in Paris, France, for analysis and confirmation. Unfortunately for Rio, his letter was lost in a shipwreck and the Institute only received his samples, which contained a brief note describing how much this new element, which Rio had named erythronium, resembled [|chromium]. Rio withdrew his claim when he received a letter from Paris disputing his discovery. Vanadium was rediscovered by Nils Gabriel Sefstrôm, a Swedish chemist, in 1830 while analyzing samples of [|iron] from a mine in Sweden. Vanadium was isolated by Sir Henry Enfield Roscoe, an English chemist, in 1867 by combining vanadium trichloride (VCl 3 ) with [|hydrogen] gas (H 2 ). Today, vanadium is primarily obtained from the minerals vanadinite (Pb 5 (VO) 3 Cl) and carnotite (K 2 (UO 2 ) 2 VO 4 ·1-3H 2 O) by heating crushed ore in the presence of [|carbon] and [|chlorine] to produce vanadium trichloride. The vanadium trichloride is then heated with [|magnesium] in an [|argon] atmosphere. <span style="font-family: 'times new roman',serif; font-size: 16px;">Vanadium is corrosion resistant and is sometimes used to make special tubes and pipes for the chemical industry. Vanadium also does not easily absorb [|neutrons] and has some applications in the nuclear power industry. A thin layer of vanadium is used to bond [|titanium] to steel. <span style="font-family: 'times new roman',serif; font-size: 16px;">Nearly 80% of the vanadium produced is used to make ferrovanadium or as an additive to steel. Ferrovanadium is a strong, shock resistant and corrosion resistant alloy of iron containing between 1% and 6% vanadium. Ferrovanadium and vanadium-steel alloys are used to make such things as axles, crankshafts and gears for cars, parts of jet engines, springs and cutting tools. <span style="font-family: 'times new roman',serif; font-size: 16px;">Vanadium pentoxide (V 2 O 5 ) is perhaps vanadium's most useful compound. It is used as a mordant, a material which permanently fixes dyes to fabrics. Vanadium pentoxide is also used as a [|catalyst] in certain chemical reactions and in the manufacture of ceramics. Vanadium pentoxide can also be mixed with [|gallium] to form superconductive magnets. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.20×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2.5×10 -3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.746 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +4, +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 24 Cr Chromium 51.9961 **Atomic Number:** 24 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for color, **chroma**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Chromium is pronounced as **KROH-mee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Chromium was discovered by Louis-Nicholas Vauquelin while experimenting with a material known as Siberian red lead, also known as the mineral crocoite (PbCrO 4 ), in 1797. He produced chromium oxide (CrO 3 ) by mixing crocoite with hydrochloric acid (HCl). Although he believed a method for isolating chromium didn't yet exist, Vauquelin was pleasantly surprised in 1798 to discover that he was able to obtain metallic chromium by simply heating chromium oxide in a charcoal oven. Today, chromium is primarily obtained by heating the mineral chromite (FeCr 2 O 4 ) in the presence of [|aluminum] or [|silicon]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Chromium is a blue-white metal that is hard, brittle and very corrosion resistant. Chromium can be polished to form a very shiny surface and is often plated to other metals to form a protective and attractive covering. Chromium is added to steel to harden it and to form stainless steel, a steel alloy that contains at least 10% chromium. Other chromium-steel alloys are used to make armor plate, safes, ball bearings and cutting tools. <span style="font-family: 'times new roman',serif; font-size: 16px;">Chromium forms many colorful compounds that have industrial uses. Lead chromate (PbCrO 4 ), also known as chrome yellow, has been used as a yellow pigment in paints. Chromic oxide (Cr 2 O 3 ), also known as chrome green, is the ninth [|most abundant compound in the earth's crust] and is a widely used green pigment. Rubies and emeralds also owe their colors to chromium compounds. Potassium dichromate (K 2 Cr 2 O 7 ) is used in the tanning of leather while other chromium compounds are used as mordants, materials which permanently fix dyes to fabrics. Chromium compounds are also used to anodize aluminum, a process which coats aluminum with a thick, protective layer of oxide. Chromite, chromium's primary ore, is used to make molds for the firing of bricks because of its high melting point, moderate thermal expansion and stable crystal structure. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.02×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 3×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.767 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +6, +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 <span style="background-color: #ffffff; font-family: 'times new roman',serif; font-size: 16px;">4s <span style="background-color: #ffffff; font-family: 'times new roman',serif; font-size: 12px;">1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 25 Mn Manganese 54.938045 **Atomic Number:** 25 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for magnet, **magnes**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Manganese is pronounced as **MAN-ge-nees**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Proposed to be an element by Carl Wilhelm Scheele in 1774, manganese was discovered by Johan Gottlieb Gahn, a Swedish chemist, by heating the mineral pyrolusite (MnO 2 ) in the presence of charcoal later that year. Today, most manganese is still obtained from pyrolusite, although it is usually burned in a furnace with powdered [|aluminum] or is treated with sulfuric acid (H 2 SO 4 ) to form manganese sulfate (MnSO 4 ), which is then electrolyzed. <span style="font-family: 'times new roman',serif; font-size: 16px;">Nearly 90% of all of the manganese produced each year is used in the production of steel. Manganese is added to molten steel to remove [|oxygen] and [|sulfur] and is alloyed with steel to make it easier to form and work with and to increase steel's strength and resistance to impact. Railroad tracks, for example, are made with steel that contains as much as 1.2% manganese. Manganese is also used to give glass an amethyst color and is responsible for the color of amethyst gemstones. <span style="font-family: 'times new roman',serif; font-size: 16px;">Manganese dioxide (MnO 2 ), the most common compound of manganese, makes up about 0.14% of [|the Earth's crust]. It is used in dry cell batteries to prevent the formation of [|hydrogen], to remove the green color in glass that is caused by the presence of [|iron] contaminants, and as a drying agent in black paints. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 9.50×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.434 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +7, +4, +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 26 Fe Iron 55.845 **Atomic Number:** 26 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Anglo-Saxon word **iron**. Iron's chemical symbol comes from the Latin word for iron, **ferrum**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Iron is pronounced as **EYE-ern**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Archaeological evidence suggests that people have been using iron for at least 5000 years. Iron is the cheapest and one of the most abundant of all metals, comprising nearly [|5.6% of the earth's crust] and nearly all of the earth's core. Iron is primarily obtained from the minerals hematite (Fe 2 O 3 ) and magnetite (Fe 3 O 4 ). The minerals taconite, limonite (FeO(OH)·nH 2 O) and siderite (FeCO 3 ) are other important sources. <span style="font-family: 'times new roman',serif; font-size: 16px;">Huge amounts of iron are used to make steel, an alloy of iron and [|carbon]. Steel typically contains between 0.3% and 1.5% carbon, depending on the desired characteristics. The addition of other elements can give steel other useful properties. Small amounts of [|chromium] improves durability and prevents rust (stainless steel); [|nickel] increases durability and resistance to heat and acids; [|manganese] increases strength and resistance to wear; [|molybdenum] increases strength and resistance to heat; [|tungsten] retains hardness at high temperatures; and [|vanadium] increases strength and springiness. Steel is used to make paper clips, skyscrapers and everything in between. <span style="font-family: 'times new roman',serif; font-size: 16px;">In addition to helping build the world around us, iron helps keep plants and animals alive. Iron plays a role in the creation of chlorophyll in plants and is an essential part of hemoglobin, the substance that carries oxygen within red blood cells. Iron sulfate (FeSO 4 ) is used to treat the blood disease anemia. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 5.63×10 4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2×10 -3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 4 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.902 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 27 Co Cobalt 58.933195 **Atomic Number:** 27 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the German word for goblin or evil spirit, **kobald** and the Greek word for mine, **cobalos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Cobalt is pronounced as **KO-bolt**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Cobalt was discovered by Georg Brandt, a Swedish chemist, in 1739. Brandt was attempting to prove that the ability of certain minerals to color glass blue was due to an unknown element and not to [|bismuth], as was commonly believed at the time. Cobalt's primary ores are cobaltite (CoAsS) and erythrite (Co 3 (AsO 4 ) 2 ). Cobalt is usually recovered as a byproduct of mining and refining [|nickel], [|silver] , [|lead] , [|copper] and [|iron]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Although cobalt is used in electroplating to give objects an attractive surface that resists oxidation, it is more widely used to form alloys. Alnico, an alloy consisting of [|aluminum], nickel and cobalt is used to make powerful permanent magnets. Stellite alloys, which contain cobalt, chromium and tungsten, are used to make high-speed and high temperature cutting tools and dyes. Cobalt is also used to make alloys for jet engines and gas turbines, magnetic steels and some types of stainless steels. <span style="font-family: 'times new roman',serif; font-size: 16px;">Cobalt-60, a radioactive [|isotope] of cobalt, is an important source of gamma rays and is used to treat some forms of cancer and as a medical tracer. Cobalt-60 has a [|half-life] of 5.27 years and decays into nickel-60 through [|beta decay]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Cobalt compounds have been used for centuries to color porcelain, glass, pottery, tile and enamel. Some of these compounds are known as: cobalt blue, ceruleum, new blue, smalt, cobalt yellow and cobalt green. In addition to being used as a dye, cobalt is also important to human nutrition as it is an essential part of vitamin B 12. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.5×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.881 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 28 Ni Nickel 58.6934 **Atomic Number:** 28 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the German word **Nickel**, which means "Old Nick," a name for the devil. Also from the German word for the mineral niccolite, **kupfernickel**, which means "Old Nick's copper." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Nickel is pronounced as **NIK-'l**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Nickel was discovered by the Swedish chemist Axel Fredrik Cronstedt in the mineral niccolite (NiAs) in 1751. Today, most nickel is obtained from the mineral pentlandite (NiS·2FeS). Most of the world's supply of nickel is mined in the Sudbury region of Ontario, Canada. It is believed that this large deposit of nickel ore is a result of an ancient meteor impact. <span style="font-family: 'times new roman',serif; font-size: 16px;">Nickel is a hard, corrosion resistant metal. It can be electroplated onto other metals to form a protective coating. Finely divided nickel is used as a [|catalyst] for the hydrogenation of vegetable oils. Adding nickel to glass gives it a green color. A single kilogram of nickel can be drawn into 300 kilometers of wire. Nickel is also used to manufacture some types of coins and batteries. <span style="font-family: 'times new roman',serif; font-size: 16px;">Nickel is alloyed with other metals to improve their strength and resistance to corrosion. Nickel is alloyed with steel to make armor plate, vaults and machine parts. It is alloyed with copper to make pipes that are used in desalination plants. Very powerful permanent magnets, known as Alnico magnets, can be made from an alloy of [|aluminum], nickel, [|cobalt] and [|iron]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 8.4×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 5.6×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 5 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.640 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 8 4s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 29 Cu Copper 63.546 **Atomic Number:** 29 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word **cuprum**, which means "from the island of Cyprus." <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Copper is pronounced as **KOP-er**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Archaeological evidence suggests that people have been using copper for at least 11,000 years. Relatively easy to mine and refine, people discovered methods for extracting copper from its ores at least 7,000 years ago. The Roman Empire obtained most of its copper from the island of Cyprus, which is where copper's name originated. Today, copper is primarily obtained from the ores cuprite (CuO 2 ), tenorite (CuO), malachite (CuO 3 ·Cu(OH) 2 ), chalcocite (Cu 2 S), covellite (CuS) and bornite (Cu 6 FeS 4 ). Large deposits of copper ore are located in the United States, Chile, Zambia, Zaire, Peru and Canada. <span style="font-family: 'times new roman',serif; font-size: 16px;">Used in large amounts by the electrical industry in the form of wire, copper is second only to [|silver] in electrical conductance. Since it resists corrosion from the air, moisture and seawater, copper has been widely used in coins. Although once made nearly entirely from copper, American pennies are now made from [|zinc] that has been coated with copper. Copper is also used to make water pipes and jewelry, as well as other items. <span style="font-family: 'times new roman',serif; font-size: 16px;">Pure copper is usually too soft for most uses. People first learned about 5,000 years ago that copper can be strengthened if it is mixed with other metals. The two most familiar alloys of copper are bronze and brass. Bronze, the first alloy created by people, is a mix of copper that contains as much as 25% [|tin]. Early people used bronze to make tools, weaponry, containers and ornamental items. Brass, a mix of copper that contains between 5% and 45% zinc, was first used about 2,500 years ago. The Romans were the first to make extensive use of brass, using it to make such things as coins, kettles and ornamental objects. Today, brass is also used in some musical instruments, screws and other hardware that must resist corrosion. <span style="font-family: 'times new roman',serif; font-size: 16px;">Hydrated copper sulfate (CuSO 4 ·H 2 O), also known as blue vitriol, is the best known copper compound. It is used as an agricultural poison, as an algicide in water purification and as a blue pigment for inks. Cuperic chloride (CuCl 2 ), another copper compound, is used to fix dyes to fabrics. Cuprous chloride (CuCl) is a poisonous white powder that is chiefly used to absorb carbon dioxide (CO 2 ). Copper cyanide (CuCN) is commonly used in electroplating. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 6.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2.5×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.726 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2, +1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 30 Zn Zinc 65.38 **Atomic Number:** 30 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the German word **zink**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Zinc is pronounced as **ZINK**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although zinc compounds have been used for at least 2,500 years in the production of brass, zinc wasn't recognized as a distinct element until much later. Metallic zinc was first produced in India sometime in the 1400s by heating the mineral calamine (ZnCO 3 ) with wool. Zinc was rediscovered by Andreas Sigismund Marggraf in 1746 by heating calamine with charcoal. Today, most zinc is produced through the electrolysis of aqueous zinc sulfate (ZnSO 4 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Roughly one third of all metallic zinc produced today is used in a process known as galvanization. During galvanization, an object that is subject to corrosion, such as an [|iron] nail, is given a protective coating of zinc. The zinc can be applied to an object by dipping it in a pool of molten zinc, but it is most often applied through an electroplating process. Sacrificial zinc anodes are used in cathodic protection systems to protect exposed iron from corrosion. Metallic zinc is also used to make dry cell batteries, roof cladding and die castings. <span style="font-family: 'times new roman',serif; font-size: 16px;">Zinc is used to make many useful alloys. Brass, an alloy of zinc that contains between 55% and 95% [|copper], is probably the best known zinc alloy. Brass was first used about 2,500 years ago and was widely used by the ancient Romans, who used it to make such things as coins, kettles and decorative items. Brass is still used today, particularly in musical instruments, screws and other hardware that must resist corrosion. Zinc is alloyed with [|lead] and [|tin] to make solder, a metal with a relatively low melting point used to join electrical components, pipes and other metallic items. Prestal ®, an alloy containing 78% zinc and 22% [|aluminum] , is a strange material that is nearly as strong as steel but is molded as easily as plastic. Nickel silver, typewriter metal, spring brass and German silver are other common zinc alloys. <span style="font-family: 'times new roman',serif; font-size: 16px;">Zinc oxide (ZnO), a common zinc compound, forms when metallic zinc is exposed to the air and forms a protective coating that protects the rest of the metal. Zinc oxide is used in paints, some rubber products, cosmetics, pharmaceuticals, plastics, printing inks, soap and batteries, among other things. Zinc sulfide (ZnS), another zinc compound, glows when it is exposed to ultraviolet light, X-rays or [|electrons] and is used to make luminous watch dials, television screens and fluorescent light bulbs. Zinc chloride (ZnCl 2 ) is another zinc compound that is used to protect wood from decay and insects. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 7.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 4.9×10 -3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 3 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 9.394 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 31 Ga Gallium 69.723 **Atomic Number:** 31 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for France, **Gallia**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Gallium is pronounced as **GAL-ee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">First proposed to exist by Dmitri Mendeleyev in 1871 based on gaps in his newly created Periodic Table of Elements, gallium was discovered spectroscopically by the French chemist Paul-Émile Lecoq de Boisbaudran in 1875. Later that same year, Lecoq was able to obtain pure gallium through the electrolysis of a solution of gallium hydroxide (Ga(OH) 3 ) in potassium hydroxide (KOH). Trace amounts of gallium are found in diaspore, sphalerite, germanite and bauxite as well as in the byproducts of burning coal. <span style="font-family: 'times new roman',serif; font-size: 16px;">Gallium melts near room temperature and has one of the largest liquid ranges of any metal, so it has found use in high temperature thermometers. Gallium easily forms alloys with most metals and has been used to create low melting alloys. Gallium is used as a doping material for semiconductors and has been used to produce solid-state items like transistors and light emitting diodes. Gallium arsenide (GaAs) can produce laser light directly from electricity. Large amounts of gallium trichloride (GaCl 3 ) have been gathered to build the [|Gallium Neutrino Observatory], an observatory located in Italy built to study particles called neutrinos which are produced inside the sun during the process of nuclear fusion. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.9×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 3×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 5.999 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 32 Ge Germanium 72.64 **Atomic Number:** 32 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for the country of Germany. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Germanium is pronounced as **jer-MAY-ni-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">First proposed to exist by Dmitri Mendeleyev in 1871 based on gaps in his newly created Periodic Table of Elements, germanium was discovered by the German chemist Clemens Winkler in the mineral argyrodite (Ag 8 GeS 6 ) in 1886. Today, germanium is primarily obtained from the smelting of [|zinc] ores and from the byproducts of burning certain types of coal. <span style="font-family: 'times new roman',serif; font-size: 16px;">The largest use of germanium is in the semiconductor industry. When doped with small amounts of [|arsenic], [|gallium] , [|indium] , [|antimony] or [|phosphorus] , germanium is used to make transistors for use in electronic devices. Germanium is also used to create alloys and as a phosphor in fluorescent lamps. Both germanium and germanium oxide (GeO) are transparent to infrared radiation and are used in infrared optical instruments and infrared detectors. Some germanium compounds seem to be effective in killing some types of bacteria and are currently being studied for use in chemotherapy. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.5 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 5×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 5 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.900 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +4, +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 33 As Arsenic 74.92160 **Atomic Number:** 33 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word **arsenicum**, the Greek word **arsenikon** and the Arabic word **Az-zernikh**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Arsenic is pronounced as **AR-s'n-ik**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Although arsenic compounds were mined by the early Chinese, Greek and Egyptian civilizations, it is believed that arsenic itself was first identified by Albertus Magnus, a German alchemist, in 1250. Arsenic occurs free in nature, but is most often found in the minerals arsenopyrite (FeAsS), realgar (AsS) and orpiment (As 2 S 3 ). Today, most commercial arsenic is obtained by heating arsenopyrite. <span style="font-family: 'times new roman',serif; font-size: 16px;">Arsenic and its compounds are poisonous. They have been used to make rat poison and some insecticides. Small amounts of arsenic are added to [|germanium] to make transistors. [|Gallium] arsenide (GaAs) can produce laser light directly from electricity. <span style="font-family: 'times new roman',serif; font-size: 16px;">If you were paying careful attention to the physical data listed above, you may have noticed that arsenic's boiling point is lower than its melting point. This occurs because these two temperatures are measured at different atmospheric pressures. When heated at standard atmospheric pressure, arsenic changes directly from a solid to a gas, or sublimates, at a temperature of 887 K. In order to form liquid arsenic, the atmospheric pressure must be increased. At 28 times standard atmospheric pressure, arsenic melts at a temperature of 1090 K. If it were also measured at a pressure of 28 atmospheres, arsenic's boiling point would be higher than its melting point, as you would expect. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.8 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 3.7-3 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 9.815 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +3, -3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 34 Se Selenium 78.96 **Atomic Number:** 34 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for the moon, **selene**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Selenium is pronounced as **si-LEE-nee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Selenium was discovered by Jöns Jacob Berzelius, a Swedish chemist, in 1817 after analyzing an impurity that was contaminating the sulfuric acid (H 2 SO 4 ) being produced at a particular factory in Sweden. Originally believing the material was [|tellurium], Berzelius eventually realized that it was actually a previously unknown element. Selenium occurs in minerals such as eucairite (CuAgSe), crooksite (CuThSe) and clausthalite (PbSe), but these minerals are too rare to use as a major source of selenium. Today, most selenium is obtained as a byproduct of refining [|copper]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Selenium's resistance to the flow of electricity is greatly affected by the amount of light shining on it. The brighter the light, the better selenium conducts electricity. This property has made selenium useful in devices that respond to the intensity of light, such as electric eyes, photo cells, light meters for cameras and copiers. Selenium can also produce electricity directly from sunlight and is used in solar cells. Selenium is also a semiconductor and is used in some types of solid-state electronics as well as in rectifiers, devices which convert alternating current electricity into direct current electricity. In addition to its use in electrical devices, selenium is also used to make a ruby-red color in glasses and enamels, as a photographic toner and as an additive to stainless steel. <span style="font-family: 'times new roman',serif; font-size: 16px;">Selenium forms few inorganic compounds, none of which are commercially important. They include selenious acid (H 2 SeO 3 ), selenium dichloride (SeCl 2 ) and selenium oxychloride (SeOCl 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 5×10 -2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 6 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 9.752 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +6, +4, -2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 35 Br Bromine 79.904 **Atomic Number:** 35 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for stench, **bromos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Bromine is pronounced as **BRO-meen**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">The only nonmetallic element that is a liquid at normal room temperatures, bromine was produced by Carl Löwig, a young chemistry student, the summer before starting his freshman year at Heidelberg. When he showed his professor, Leopold Gmelin, the red, smelly liquid he had produced, Gmelin realized that this was an unknown substance and encouraged Löwig to produce more of it so they could study it in detail. Unfortunately, winter exams and the holidays delayed Löwig's work long enough for another chemist, Antoine-Jérôme Balard, to publish a paper in 1826 describing the new element. Balard was credited with the discovery and named it after the greek word for stench, bromos. Today, bromine is primarily obtained by treating brines from wells in Michigan and Arkansas with [|chlorine]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Elemental bromine is a hazardous material. It causes severe burns when it comes in contact with the skin and its vapor irritates the eyes, nose and throat. Most of the bromine produced in the United States was used in the manufacture of ethylene dibromide(C 2 H 4 Br 2 ), a chemical added to leaded gasolines that prevented the accumulation of [|lead] compounds within the engine. With the discontinuation of leaded gasolines in favor of unleaded gasolines, the demand for bromine has been greatly reduced. Silver bromide (AgBr), a chemical used in photography, now accounts for the largest use of bromine. Other bromine compounds are used in fumigants, in flameproofing agents and in some compounds used to purify water. Tyrian purple, an expensive purple dye known to ancient civilizations, was produced from an organic bromine compound secreted from a sea mussel known as the murex. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 6.73×10 1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 2 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 11.814 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +1, -1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 36 Kr Krypton 83.798 **Atomic Number:** 36 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for hidden, **kryptos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Krypton is pronounced as **KRIP-ton**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Krypton was discovered on May 30, 1898 by Sir William Ramsay, a Scottish chemist, and Morris M. Travers, an English chemist, while studying liquefied air. Small amounts of liquid krypton remained behind after the more volatile components of liquid air had boiled away. The [|earth's atmosphere] is about 0.0001% krypton. <span style="font-family: 'times new roman',serif; font-size: 16px;">The high cost of obtaining krypton from the air has limited its practical applications. Krypton is used in some types of photographic flashes used in high speed photography. Some fluorescent light bulbs are filled with a mixture of krypton and [|argon] gases. Krypton gas is also combined with other gases to make luminous signs that glow with a greenish-yellow light. In 1960, the length of the meter was defined in terms of the orange-red spectral line of krypton-86, an [|isotope] of krypton. <span style="font-family: 'times new roman',serif; font-size: 16px;">Once thought to be completely inert, krypton is known to form a few compounds. Krypton difluoride (KrF 2 ) is the easiest krypton compound to make and gram amounts of it have been produced. <span style="font-family: 'times new roman',serif; font-size: 16px;">For those that are curious, [|pictures of krypton gas and krypton plasma] can be found in the [|Questions and Answers] section of this site. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1×10 -4 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 2.1×10 -4 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 5 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 14.000 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** 0 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 37 Rb Rubidium 85.4678 **Atomic Number:** 37 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Latin word for deepest red, **rubidus**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Rubidium is pronounced as **roo-BID-ee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Rubidium was discovered by the German chemists Robert Bunsen and Gustav Kirchhoff in 1861 while analyzing samples of the mineral lepidolite (KLi 2 Al(Al, Si) 3 O 10 (F, OH) 2 ) with a device called a spectroscope. The sample produced a set of deep red spectral lines they had never seen before. Bunsen was eventually able to isolate samples of rubidium metal. Today, most rubidium is obtained as a byproduct of refining [|lithium]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Rubidium is used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes. It is also used in the manufacture of photocells and in special glasses. Since it is easily ionized, it might be used as a propellant in ion engines on spacecraft. Recent discoveries of large deposits of rubidium suggest that its usefulness will increase as its properties become better understood. <span style="font-family: 'times new roman',serif; font-size: 16px;">Rubidium forms a large number of compounds, although none of them has any significant commercial application. Some of the common rubidium compounds are: rubidium chloride (RbCl), rubidium monoxide (Rb 2 O) and rubidium copper sulfate Rb 2 SO 4 ·CuSO 4 ·6H 2 0). A compound of rubidium, [|silver] and [|iodine], RbAg 4 I 5 , has interesting electrical characteristics and might be useful in thin film batteries. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 9.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.2×10 -1 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 4.177 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 38 Sr Strontium 87.62 **Atomic Number:** 38 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for the town of Strontian, Scotland. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Strontium is pronounced as **STRON-she-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Strontium was discovered by Adair Crawford, an Irish chemist, in 1790 while studying the mineral witherite (BaCO 3 ). When he mixed witherite with hydrochloric acid (HCl) he did not get the results he expected. He assumed that his sample of witherite was contaminated with an unknown mineral, a mineral he named strontianite (SrCO 3 ). Strontium was first isolated by Sir Humphry Davy, an English chemist, in 1808 through the electrolysis of a mixture of strontium chloride (SrCl 2 ) and mercuric oxide (HgO). Today, strontium is obtained from two of its most common ores, celestite (SrSO 4 ) and strontianite (SrCO 3 ), by treating them with hydrochloric acid, forming strontium chloride. The strontium chloride, usually mixed with potassium chloride (KCl), is then melted and electrolyzed, forming strontium and [|chlorine] gas (Cl 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">Most of the strontium produced today is used in the manufacture of color television picture tubes. It is also used to refine [|zinc] and is combined with [|iron] to make magnets. <span style="font-family: 'times new roman',serif; font-size: 16px;">Two strontium compounds, strontium carbonate (SrCO 3 ) and strontium nitrate (Sr(NO 3 ) 2 ), burn with a bright, red flame and are used in fireworks and signal flares. Strontium carbonate is also used to make certain kinds of glass and is the base material for making most other strontium compounds. <span style="font-family: 'times new roman',serif; font-size: 16px;">Strontium-90, a radioactive [|isotope] of strontium, is a common product of nuclear explosions. It has a [|half-life] of about 28.8 years and decays into [|yttrium] -90 through [|beta decay]. Strontium-90 is especially deadly since it has a relatively long half-life, is strongly radioactive and is absorbed by the body, where it accumulates in the skeletal system. The radiation affects the production of new blood cells, which eventually leads to death. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 3.70×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 7.9 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 4 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 5.695 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 39 Y Yttrium 88.90585 **Atomic Number:** 39 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for the village of Ytterby, Sweden. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Yttrium is pronounced as **IT-ree-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Yttrium was discovered by Johan Gadolin, a Finnish chemist, while analyzing the composition of the mineral gadolinite ((Ce, La, Nd, Y) 2 FeBe 2 Si 2 O 10 ) in 1789. Gadolinite, which was named for Johan Gadolin, was discovered several years earlier in a quarry near the town of Ytterby, Sweden. Today, yttrium is primarily obtained through an ion exchange process from monazite sand ((Ce, La, Th, Nd, Y)PO 4 ), a material rich in rare earth elements. <span style="font-family: 'times new roman',serif; font-size: 16px;">Although metallic yttrium is not widely used, several of its compounds are. Yttrium oxide (Y 2 O 3 ) and yttrium orthovanadate (YVO 4 ) are both combined with [|europium] to produce the red phosphor used in color televisions. Garnets made from yttrium and [|iron] (Y 3 Fe 5 O 12 ) are used as microwave filters in microwave communications equipment. Garnets made from yttrium and [|aluminum] (Y 3 Al 5 O 12 ) are used in jewelry as simulated diamond. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 3.3×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1.3×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.217 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 1 5s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 40 Zr Zirconium 91.224 **Atomic Number:** 40 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Persian word for gold-like, **zargun**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Zirconium is pronounced as **zer-KO-nee-em**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Zirconium was discovered by Martin Heinrich Klaproth, a German chemist, while analyzing the composition of the mineral jargon (ZrSiO 4 ) in 1789. Zirconium was isolated by Jöns Jacob Berzelius, a Swedish chemist, in 1824 and finally prepared in a pure form in 1914. Obtaining pure zirconium is very difficult because it is chemically similar to [|hafnium], an element which is always found mixed with deposits of zirconium. Today, most zirconium is obtained from the minerals zircon (ZrSiO 4 ) and baddeleyite (ZrO 2 ) through a process known as the Kroll Process. <span style="font-family: 'times new roman',serif; font-size: 16px;">Zirconium is a corrosion resistant metal that is used in high performance pumps and valves. Since it also does not easily absorb [|neutrons], zirconium is widely used in nuclear reactors. The nuclear power industry uses nearly 90% of the zirconium produced each year, which must be nearly free of hafnium. Zirconium is also used as an alloying agent in steel, to make some types of surgical equipment and as a getter, a material that combines with and removes trace gases from vacuum tubes. <span style="font-family: 'times new roman',serif; font-size: 16px;">Zircon (ZrSiO 4 ) is a zirconium compound that can take many different forms, the most popular of which is a clear, transparent gemstone that can be cut to look like diamond and is frequently used in jewelry. Zirconium dioxide (ZrO 2 ) can withstand very high temperatures and is used to make crucibles and to line the walls of high temperature furnaces. Zirconium carbonate (3ZrO 2 ·CO 2 ·H 2 O) is used in lotions to treat poison ivy. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.65×10 2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 3×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 4 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.634 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +4 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 2 5s 2 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 41 Nb Niobium 92.90638 **Atomic Number:** 41 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** Named for the Greek mythological figure [|**Niobe**]. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Niobium is pronounced as **ni-OH-bee-um**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">The story of niobium's discovery is a bit confusing. The first governor of Connecticut, John Winthrop the Younger, discovered a new mineral around 1734. He named the mineral columbite ((Fe, Mn, Mg)(Nb, Ta) 2 O 6 ) and sent a sample of it to the British Museum in London, England. The columbite sat in the museum's mineral collection for years until it was analyzed by Charles Hatchett in 1801. Hatchett could tell that there was an unknown element in the columbite, but he was not able to isolate it. He named the new element columbium. <span style="font-family: 'times new roman',serif; font-size: 16px;">The fate of columbium took a drastic turn in 1809 when William Hyde Wollaston, an English chemist and physicist, compared the minerals columbite and tantalite ((Fe, Mn)(Ta, Nb) 2 O 6 ) and declared that columbium was actually the element [|tantalum]. This confusion arose because tantalum and niobium are similar metals, are always found together and are very difficult to isolate. <span style="font-family: 'times new roman',serif; font-size: 16px;">Niobium was rediscovered and renamed by Heinrich Rose in 1844 when he produced two new acids, niobic acid and pelopic acid, from samples of columbite and tantalite. These acids are very similar to each other and it took another twenty-two years and a Swiss chemist named Jean Charles Galissard de Marignac to prove that these were two distinct chemicals produced from two different elements. Metallic niobium was finally isolated by the Swedish chemist Christian Wilhelm Blomstrand in 1864. Today, niobium is primarily obtained from the minerals columbite and pyrochlore ((Ca, Na) 2 Nb 2 O 6 (O, OH, F)). <span style="font-family: 'times new roman',serif; font-size: 16px;">Niobium is used as an alloying agent and for jewelry, but perhaps its most interesting applications are in the field of superconductivity. Superconductive wire can be made from an alloy of niobium and [|titanium] which can then be used to make superconductive magnets. Other alloys of niobium, such as those with [|tin] and [|aluminum], are superconductive as well. Pure niobium is itself a superconductor when it is cooled below 9.25 K (-442.75°F). [|Superconductive niobium cavities] are at the heart of a machine built at the [|Thomas Jefferson National Accelerator Facility]. This machine, called an [|electron accelerator], is used by scientists to study the [|quark] structure of matter. The accelerator's 338 niobium cavities are bathed in liquid [|helium] and accelerate [|electrons] to nearly the speed of light. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 2.0×10 1 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1×10 -5 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 1 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 6.759 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +5, +3 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 4 5s 1 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;"> 42 Mo Molybdenum 95.96 **Atomic Number:** 42 <span style="font-family: 'times new roman',serif; font-size: 16px;">**What's in a name?** From the Greek word for lead, **molybdos**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**Say what?** Molybdenum is pronounced as **meh-LIB-deh-nem**. <span style="font-family: 'times new roman',serif; font-size: 16px;">**History and Uses:** <span style="font-family: 'times new roman',serif; font-size: 16px;">Molybdenum was discovered by Carl Welhelm Scheele, a Swedish chemist, in 1778 in a mineral known as molybdenite (MoS 2 ) which had been confused as a [|lead] compound. Molybdenum was isolated by Peter Jacob Hjelm in 1781. Today, most molybdenum is obtained from molybdenite, wulfenite (PbMoO 4 ) and powellite (CaMoO 4 ). These ores typically occur in conjunction with ores of [|tin] and [|tungsten]. Molybdenum is also obtained as a byproduct of mining and processing tungsten and [|copper]. <span style="font-family: 'times new roman',serif; font-size: 16px;">Molybdenum has a high melting point and is used to make the electrodes of electrically heated glass furnaces. Some electrical filaments are also made from molybdenum. The metal is used to make some missile and aircraft parts and is used in the nuclear power industry. Molybdenum is also used as a [|catalyst] in the refining of petroleum. <span style="font-family: 'times new roman',serif; font-size: 16px;">Molybdenum is primarily used as an alloying agent in steel. When added to steel in concentrations between 0.25% and 8%, molybdenum forms ultra-high strength steels that can withstand pressures up to 300,000 pounds per square inch. Molybdenum also improves the strength of steel at high temperatures. When alloyed with [|nickel], molybdenum forms heat and corrosion resistant materials used in the chemical industry. <span style="font-family: 'times new roman',serif; font-size: 16px;">Molybdenum disulphide (MoS 2 ), one of molybdenum's compounds, is used as a high temperature lubricant. Molybdenum trioxide (MoO 3 ), another molybdenum compound, is used to adhere enamels to metals. Other molybdenum compounds include: molybdic acid (H 2 MoO 4 ), molybdenum hexafluoride (MoF 6 ) and molybdenum phosphide (MoP 2 ). <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Crustal Abundance:** 1.2 milligrams per kilogram <span style="font-family: 'times new roman',serif; font-size: 16px;">**Estimated Oceanic Abundance:** 1×10 -2 milligrams per liter <span style="font-family: 'times new roman',serif; font-size: 16px;">**Number of Stable Isotopes:** 6 <span style="font-family: 'times new roman',serif; font-size: 16px;">**Ionization Energy:** 7.092 eV <span style="font-family: 'times new roman',serif; font-size: 16px;">**Oxidation States:** +6 <span style="display: block; font-family: 'times new roman',serif; font-size: 16px;">** [|Electron Shell Configuration] :** 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 5 5s 1
 * Atomic Weight:** 30.973762
 * Melting Point:** 317.30 K (44.15°C or 111.47°F)
 * Boiling Point:** 553.65 K (280.5°C or 536.9°F)
 * Density:** 1.82 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Non-metal
 * Period Number:** 3 **Group Number:** 15 **Group Name:** Pnictogen
 * Atomic Weight:** 32.065
 * Melting Point:** 388.36 K (115.21°C or 239.38°F)
 * Boiling Point:** 717.75 K (444.60°C or 832.28°F)
 * Density:** 2.067 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Non-metal
 * Period Number:** 3 **Group Number:** 16 **Group Name:** Chalcogen
 * Atomic Weight:** 35.453
 * Melting Point:** 171.65 K (-101.5°C or -150.7°F)
 * Boiling Point:** 239.11 K (-34.04°C or -29.27°F)
 * Density:** 0.003214 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 3 **Group Number:** 17 **Group Name:** Halogen
 * Atomic Weight:** 39.948
 * Melting Point:** 83.80 K (-189.35°C or -308.83°F)
 * Boiling Point:** 87.30 K (-185.85°C or -302.53°F)
 * Density:** 0.0017837 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 3 **Group Number:** 18 **Group Name:** Noble Gas
 * Atomic Weight:** 39.0983
 * Melting Point:** 336.53 K (63.38°C or 146.08°F)
 * Boiling Point:** 1032 K (759°C or 1398°F)
 * Density:** 0.89 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 1 **Group Name:** Alkali Metal
 * Atomic Weight:** 40.078
 * Melting Point:** 1115 K (842°C or 1548°F)
 * Boiling Point:** 1757 K (1484°C or 2703°F)
 * Density:** 1.54 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 2 **Group Name:** Alkaline Earth Metal
 * Atomic Weight:** 44.955912
 * Melting Point:** 1814 K (1541°C or 2806°F)
 * Boiling Point:** 3109 K (2836°C or 5137°F)
 * Density:** 2.99 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 3 **Group Name:** none
 * Atomic Weight:** 47.867
 * Melting Point:** 1941 K (1668°C or 3034°F)
 * Boiling Point:** 3560 K (3287°C or 5949°F)
 * Density:** 4.5 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 4 **Group Name:** none
 * Atomic Weight:** 50.9415
 * Melting Point:** 2183 K (1910°C or 3470°F)
 * Boiling Point:** 3680 K (3407°C or 6165°F)
 * Density:** 6.0 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 5 **Group Name:** none
 * Atomic Weight:** 51.9961
 * Melting Point:** 2180 K (1907°C or 3465°F)
 * Boiling Point:** 2944 K (2671°C or 4840°F)
 * Density:** 7.15 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 6 **Group Name:** none
 * Atomic Weight:** 54.938045
 * Melting Point:** 1519 K (1246°C or 2275°F)
 * Boiling Point:** 2334 K (2061°C or 3742°F)
 * Density:** 7.3 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 7 **Group Name:** none
 * Atomic Weight:** 55.845
 * Melting Point:** 1811 K (1538°C or 2800°F)
 * Boiling Point:** 3134 K (2861°C or 5182°F)
 * Density:** 7.874 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 8 **Group Name:** none
 * Atomic Weight:** 58.933195
 * Melting Point:** 1768 K (1495°C or 2723°F)
 * Boiling Point:** 3200 K (2927°C or 5301°F)
 * Density:** 8.86 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 9 **Group Name:** none
 * Atomic Weight:** 58.6934
 * Melting Point:** 1728 K (1455°C or 2651°F)
 * Boiling Point:** 3186 K (2913°C or 5275°F)
 * Density:** 8.912 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 10 **Group Name:** none
 * Atomic Weight:** 63.546
 * Melting Point:** 1357.77 K (1084.62°C or 1984.32°F)
 * Boiling Point:** 2835 K (2562°C or 4644°F)
 * Density:** 8.933 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 11 **Group Name:** none
 * Atomic Weight:** 65.38
 * Melting Point:** 692.68 K (419.53°C or 787.15°F)
 * Boiling Point:** 1180 K (907°C or 1665°F)
 * Density:** 7.134 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 12 **Group Name:** none
 * Atomic Weight:** 69.723
 * Melting Point:** 302.91 K (29.76°C or 85.57°F)
 * Boiling Point:** 2477 K (2204°C or 3999°F)
 * Density:** 5.91 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 4 **Group Number:** 13 **Group Name:** none
 * Atomic Weight:** 72.64
 * Melting Point:** 1211.40 K (938.25°C or 1720.85°F)
 * Boiling Point:** 3106 K (2833°C or 5131°F)
 * Density:** 5.323 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Semi-metal
 * Period Number:** 4 **Group Number:** 14 **Group Name:** none
 * Atomic Weight:** 74.92160
 * Melting Point:** 1090 K (817°C or 1503°F)
 * Boiling Point:** 887 K (614°C or 1137°F)
 * Density:** 5.776 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Semi-metal
 * Period Number:** 4 **Group Number:** 15 **Group Name:** Pnictogen
 * Atomic Weight:** 78.96
 * Melting Point:** 493.65 K (220.5°C or 428.9°F)
 * Boiling Point:** 958 K (685°C or 1265°F)
 * Density:** 4.809 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Non-metal
 * Period Number:** 4 **Group Number:** 16 **Group Name:** Chalcogen
 * Atomic Weight:** 79.904
 * Melting Point:** 265.95 K (-7.2°C or 19.0°F)
 * Boiling Point:** 331.95 K (58.8°C or 137.8°F)
 * Density:** 3.11 grams per cubic centimeter
 * Phase at Room Temperature:** Liquid
 * Element Classification:** Non-metal
 * Period Number:** 4 **Group Number:** 17 **Group Name:** Halogen
 * Atomic Weight:** 83.798
 * Melting Point:** 115.79 K (-157.36°C or -251.25°F)
 * Boiling Point:** 119.93 K (-153.22°C or -243.80°F)
 * Density:** 0.003733 grams per cubic centimeter
 * Phase at Room Temperature:** Gas
 * Element Classification:** Non-metal
 * Period Number:** 4 **Group Number:** 18 **Group Name:** Noble Gas
 * Atomic Weight:** 85.4678
 * Melting Point:** 312.46 K (39.31°C or 102.76°F)
 * Boiling Point:** 961 K (688°C or 1270°F)
 * Density:** 1.53 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 1 **Group Name:** Alkali Metal
 * Atomic Weight:** 87.62
 * Melting Point:** 1050 K (777°C or 1431°F)
 * Boiling Point:** 1655 K (1382°C or 2520°F)
 * Density:** 2.64 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 2 **Group Name:** Alkaline Earth Metal
 * Atomic Weight:** 88.90585
 * Melting Point:** 1795 K (1522°C or 2772°F)
 * Boiling Point:** 3618 K (3345°C or 6053°F)
 * Density:** 4.47 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 3 **Group Name:** none
 * Atomic Weight:** 91.224
 * Melting Point:** 2128 K (1855°C or 3371°F)
 * Boiling Point:** 4682 K (4409°C or 7968°F)
 * Density:** 6.52 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 4 **Group Name:** none
 * Atomic Weight:** 92.90638
 * Melting Point:** 2750 K (2477°C or 4491°F)
 * Boiling Point:** 5017 K (4744°C or 8571°F)
 * Density:** 8.57 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 5 **Group Name:** none
 * Atomic Weight:** 95.96
 * Melting Point:** 2896 K (2623°C or 4753°F)
 * Boiling Point:** 4912 K (4639°C or 8382°F)
 * Density:** 10.2 grams per cubic centimeter
 * Phase at Room Temperature:** Solid
 * Element Classification:** Metal
 * Period Number:** 5 **Group Number:** 6 **Group Name:** none