Periodic table, main group elements

The main group elements of the periodic table are groups 1,2 and 13 through 18. Elements in these groups are collectively known as main group or representative elements. These groups contain the most naturally abundant elements, comprise 80% of the earth's crust and are the most important for life. Economically the most produced chemicals are main group elements or their compounds. It is in the main group elements that we most clearly see the trends in physical and chemical properties of the elements that chemists have used to understand the "stuff" things are made of.

Group I (Alkali Metals)

Group	1
Period
2	3  Li
3	11  Na
4	19  K
5	37  Rb
6	55  Cs
7	87  Fr

The alkali metals are the series of elements in Group 1 of the periodic table (excluding hydrogen in all but one rare circumstance): lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).

Properties

The alkali metals are silver-colored (caesium has a golden tinge), soft, low-density metals. These elements all have one valence electron which is easily lost to form an ion with a single positive charge. They have the lowest ionization energies in their respective periods. This makes them very reactive and they are the most active metals. Due to their activity they occur naturally in ionic compounds not in their elemental state.

The alkali metals react readily with halogens to form ionic salts, such as table salt, sodium chloride (NaCl). They are famous for their vigorous reactions with water to liberate hydrogen gas. These reactions also often liberate sufficient energy to ignite the hydrogen and can be quite dangerous. As we move down the group the reactions become increasingly violent. The reaction with water is as follows:

Alkali metal + water → Alkali metal hydroxide + hydrogen

With potassium as an example:

$2{K}_{(s)}+2{H_{2}O}_{(l)}\to 2{KOH}_{(aq)}+{H_{2}}_{(g)}$

The oxides, hydrides, and hydoxides of these metals are basic (alkaline). In particular the hydoxides resulting from the reaction with water are our most common laboratory bases (alkalis). It is from this character that they derive their group name.

Hydrogen also has a single valence electron and is usually placed at the top of Group 1, but it is not a metal (except under extreme circumstances as metallic hydrogen); rather it exists naturally as a diatomic gas. Hydrogen can form ions with a single positive charge, but removal of its single electron requires considerably more energy than removal of the outer electron from the alkali metals. Unlike the alkali metals hydrogen atoms can also gain an electron to form the negativly charged hydride ion. The hydride ion is an extremly stong base and does not usually occur except when combined with the alkali metals and some transition metals (i.e. the ionic sodium hydride, NaH). In compounds hydrogen most often forms covalent bonds.

Under extremely high pressure, such as is found at the core of Jupiter, hydrogen does become metallic and behaves like an alkali metal; see metallic hydrogen.

See also

• Science aid:Alkali metals A simple look at alkali metals
• Doc Brown, Alkali metals

Group 2

Group	2
Period
2	4 Be
3	12 Mg
4	20 Ca
5	38 Sr
6	56 Ba
7	88 Ra

The alkaline earth metals are the series of elements in Group 2 of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) (though radium is not always considered an alkaline earth due to its radioactivity).

Properties

The alkaline earth metals are silvery colored, soft, low-density metals, though are a bit harder than the alkali metals. These elements all have two valence electrons and tend to lose both to form ions with a two plus charge. Berylium is the least metallic element in the group and tends form covalent bonds in its compounds.

These metals are less active than the alkali metals, but are still fairly active. They react readily with halogens to form ionic salts, and can react slowly with water. Magnesium reacts only with steam and calcium with hot water. Beryllium is an exception: It does not react with water or steam, and its halides are covalent. The oxides are basic and dissolve in acids and the hydroxides are strong bases, though not as soluable as the alkali metal hydroxides.

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These were named alkaline earths because of their intermediate nature between the alkalis (oxides of the alkali metals) and the rare earths (oxides of rare earth metals). The classification of some apparently inert substances as 'earths' is millennia old. The earliest known system used by the ancient Greeks consisted of four elements, including earth. Later alchemists applied the term to any solid substance that did not melt and was not changed by fire. The realization that 'earths' were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them Substances simples salifiables terreuses, or salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths.

Group 13

Group	13
Period
2	5 B
3	13 Al
4	31 Ga
5	49 In
6	81 Tl
7	113 Uut

The Boron group is the series of elements in group 13 (IUPAC style) in the periodic table: boron (B), aluminium (Al), gallium (Ga), indium (In), thallium (Tl), and ununtrium (Uut) (unconfirmed).

Properties

In this group we begin to see the changeover toward non-metallic character. First appearing at the top of the group. Boron is a metalloid, it has characteristics intermediate between metals and non-metals, and the rest of the group are metals. These elements are characterized by having three valence electrons. The metals can loose all three electrons to form ions with a three plus charge, this exist in ionic compounds, but boron tends to form covalent bonds. The oxides of the metals dissolve in acids so may be considered basic, but aluminum oxide also dissolves in acids. It is amphoteric, that is it displays both acidic and basic characteristics. This is another indication of the changeover to non-metallic character. Aluminum is the third most abundant element in the earth's crust (7.4%).

Group 14

Group	14
Period
2	6 C
3	14 Si
4	32 Ge
5	50 Sn
6	82 Pb
7	114 Uuq

The carbon group is group 14 (IUPAC style) in the periodic table. In schools, it is often known as group 4.

Each element in this group has 4 electrons in its outer energy level. In most cases, the elements share their electrons. The tendency to lose electrons increases as the size of the atom increases, as it does with increasing atomic number. Carbon alone forms negative ions, in the form of carbide (C4-) ions. Silicon and germanium, both metalloids, each can form +4 ions.

The group consists of carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and ununquadium (Uuq).

Group 15

Group	15
Period
2	7 N
3	15 P
4	33 As
5	51 Sb
6	84 Bi
7	115 Uup

The Nitrogen group elements (a.k.a. group VA) are also known as IUPAC Group 15 (formerly Group V) of the periodic table.

This group has the defining characteristic that all the component elements have 5 electrons in their outermost shell, that is 2 electrons in the s subshell and 3 in the p subshell. They are therefore 3 electrons short of filling their outermost electron shell in their non-ionized state. The most important element of this group is Nitrogen (N), which in its diatomic form is the principal component of air.

Other members of the group include Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi) and ununpentium (UUp) (unconfirmed).

The collective name pnicogens (now also spelled pnictogens) is also sometimes used for elements of this group, with binary compounds being called pnictides: neither term is approved by IUPAC. Both spellings are said to derive from the Greek πνίγειν (pnigein), to choke or stifle, which is a property of nitrogen.

Group 16

Group	16
Period
2	8 O
3	16 S
4	34 Se
5	52 Te
6	84 Po
7	116 Uuh

The chalcogens (with the "ch" pronounced with a hard "c" as in "chemistry") are the name for the periodic table group 16 (old-style: VIB or VIA) in the periodic table. It is sometimes known as the oxygen family. It consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium (Te), the radioactive polonium (Po), and the synthetic ununhexium (Uuh). The compounds of the heavier chalcogens (particularly the sulfides, selenides, and tellurides) are collectively known as chalcogenides. Unless grouped with a heavier chalcogen, oxides are not considered chalcogenides.

The name is generally considered to mean "ore former" from the Greek chalcos "ore" and -gen "formation". [1]

Oxygen and sulfur are nonmetals, and polonium, selenium and tellurium are metalloid semiconductors (i.e., their electrical properties are between those of a metal and an insulator). Nevertheless, tellurium, as well as selenium, is often referred to as a metal when in elemental form.

Chalcogenides are quite common as minerals. For example, FeS₂ (pyrite) is an iron ore and AuTe₂ gave its name to the gold rush town of Telluride, Colorado in the United States.

The formal oxidation number of the chalcogen is generally -2 in a chalcogenide but other values (e.g. -1 in pyrite) can be attained.

The highest formal oxidation number +6 is found in sulfates, selenates and tellurates, e.g. in Na₂SeO₄ (sodium selenate). Modern chemical understanding based on quantum theory somewhat outdates the use of formal oxidation numbers in favour of a many-electron wavefunction approach allowing detailed computer simulation, though the concept, while flawed, is still useful in thought experiments.

See also

• Gold chalcogenides
• Periodic table

External links

• A Second Note on the Term "Chalcogen"

Group 17

This article discusses the group of chemical elements in the periodic table:for the light bulb, see the halogen lamp.

Group	17
Period
2	9 F
3	17 Cl
4	35 Br
5	53 I
6	85 At
7	117 Uus

The halogens are a chemical series. They are the elements in Group 17 (old-style: VII or VIIA) of the periodic table: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At) and the as yet undiscovered ununseptium (Uus). The term halogen was coined to mean elements which produce salt in union with a metal. It comes from 18th century scientific French nomenclature based on erring adaptations of Greek roots.

These elements are diatomic molecules in their natural form. They require one more electron to fill their outer electron shells, and so have a tendency to form a singly-charged negative ion. This negative ion is referred to as a halide ion; salts containing these ions are known as halides.

Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. Chlorine and iodine are both used as disinfectants for such things as drinking water, swimming pools, fresh wounds, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms, a process known as sterilization. Their reactive properties are also put to use in bleaching. Chlorine is the active ingredient of most fabric bleaches and is used in the production of most paper products.

Halide ions combined with single hydrogen atoms form the hydrohalic acids (i.e., HF, HCl, HBr, HI), a series of particularly strong acids. (HAt, or "hydrastatic acid", should also qualify, but it is not typically included in discussions of hydrohalic acid due to astatine's extreme instability toward alpha decay.)

They react with each other to form interhalogen compounds. Diatomic interhalogen compounds (BrF, ICl, ClF, etc.) bear strong superficial resemblance to the pure halogens.

Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen atoms; these are known as halogenated compounds or organic halides. Chlorine is by far the most abundant of the halogens, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in brain function by mediating the action of the inhibitory transmitter GABA and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of thyroid hormones such as thyroxine. On the other hand, neither fluorine nor bromine are believed to be really essential for humans, although small amounts of fluoride can make tooth enamel resistant to decay.

They show a number of trends when moving down the group - for instance, decreasing electronegativity and reactivity, increasing melting and boiling point.

Halogen	Atomic Mass (u)	Melting Point (K)	Boiling Point (K)	Electronegativity (Pauling)
Fluorine	18.998	53.53	85.03	3.98
Chlorine	35.453	171.6	239.11	3.16
Bromine	79.904	265.8	332.0	2.96
Iodine	126.904	386.85	457.4	2.66
Astatine	(210)	575	610 ?	2.2
Ununseptium	(291)*	*	*	*

* Ununseptium has not yet been discovered; values are either unknown if no value appears, or are estimates based on other similar chemicals.

See also

• pseudohalogen

Group 18

Group	18
Period
1	2 He
2	10 Ne
3	18 Ar
4	36 Kr
5	54 Xe
6	86 Rn
7	118 Uuo

The noble gases are the chemical elements in group 18 (old-style Group 0) of the periodic table. This chemical series contains helium, neon, argon, krypton, xenon, and radon.

Etymology

The noble gases were previously referred to as inert gases, but this term is not strictly accurate because several of them do take part in chemical reactions. Another older term was rare gases, although in fact argon forms a considerable part (0.93% by volume, 1.29% by mass) of the Earth's atmosphere.

The name 'noble gases' is an allusion to the similarly unreactive Noble metals, so called due to their preciousness, resistance to corrosion and long association with the aristocracy.

Chemistry

The general physical properties of Noble gases are:

They are all monatomic molecules and chemically inert (unreactive), except for Kr and Xe, which have shown some reactivity in the laboratory—see noble gas compounds.

The noble gases' lack of reactivity is due to their having a complete valence shell. They have little tendency to gain or lose electrons. The noble gases have high ionization energies and negligible electronegativities. The noble gases have low boiling points and are all gases at room temperature.

Because of their unreactivity, the noble gases were not discovered until 1868, when helium was detected spectrographically in the Sun. The isolation of helium on Earth had to wait until 1895. The noble gases have very weak inter-atomic forces of attraction, and consequently very low melting points and boiling points. This is why they are all gases under normal conditions, even those with larger atomic masses than many normally solid elements.

Ununoctium

No isotopes with 118 protons have yet been detected in nature or synthesized in the laboratory. In the meantime, the systematic name "ununoctium" is used to refer to this hypothetical element. If discovered, ununoctium is expected to be another noble gas, filling the empty space in the periodic table beneath radon. All its isotopes are likely to be radioactive with a very short half-life in the millisecond range.

Applications

One of the most commonly encountered uses of the noble gases in everyday life is in lighting. Argon is often used as a suitable safe and inert atmosphere for the inside of filament light bulbs. Some of the noble gases glow distinctive colours when used inside lighting tubes (neon lights). Helium, due to its unreactivity (compared to flammable hydrogen) and lightness, is often used in blimps and balloons.

Physical Properties

Property
Element Number	2	10	18	36	54	86	118
Element Name	Helium	Neon	Argon	Krypton	Xenon	Radon	Ununoctium
Normal Density	0.1786	0.9002	1.7818	3.708	5.851	9.97	N/A
Atomic Mass	4.00	20.2	39.9	82.92	130.2	222.4	N/A
Boiling Pt. (°C)	-268.83	-245.92	-185.81	-151.7	-106.6	-62	N/A
Melting Pt. (°C)	-272	-248.52	-189.6	-157	-111.5	-71	N/A

External links

• Ohio State University press release for uranium compounds with noble gases.
• Rare Gases - Neon, Krypton, Xenon Properties, Uses, Applications
• Argon Ar Properties, Uses, Applications
• Science Aid: Noble gases Simple look at the properties and uses of noble gases

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