This chapter explores how to write electronic structures for atoms and ions using s, p, and d notation. It supposes that we know about simple atomic orbitals - at least as far as the way they are named, and their relative energies. If we want to look at the electronic structures of easy monatomic ions (such as Cl-, Ca2+ and Cr3+), we require knowing the electronic structure of the atom and the number of electrons gained/lost in forming the ion.
Elements in the periodic table are arranged in periods (rows) and groups (columns). Each of the seven periods is packed sequentially through atomic number. Groups comprise elements having the similar electron configuration in their outer shell, which consequences in group elements sharing similar chemical properties. The electrons in the outer shell are termed valence electrons. Valence electrons find out the properties and chemical reactivity of the element and contribute in chemical bonding. The Roman numerals establish above each group specify the usual number of valence electrons. There are 2 sets of groups. Groups 1 & 2 and 13 - 18 (formerly Group A) elements are the representative elements that have s or p sublevels as their outer orbitals. Group's 3 - 12 (formerly Group B) elements are the non-representative elements, which have partly filled d sublevels (the conversion elements) or moderately filled f sublevels (the lanthanide series and the actinide series). The Roman numeral and letter descriptions provide the electron configuration for the valence electrons (for instance the valence electron configuration of a Group 15 (formerly group VA) element will be s2p3 through 5 valence electrons).
Even in the initial studies of chemistry, it became evident that definite elemental substances were extremely much like other substances in their physical and chemical properties. For instance, common alkali metals (Na and K) were approximately indistinguishable to early chemists. Both metals are alike in appearance and undergo reactions through the similar reagents.
As more and more elemental substances were recognized, more and more similarities between the new elements and previously recognized elements were noticed. Chemists began to wonder why similarities existed. In the modern periodic table, elements are organized in order in order of rising atomic number, and the properties of the elements drop into a entirely regular order.
An element is described through its atomic number (the number of protons in the nucleus of its atoms), but its chemical reactivity is found out through the number of electrons in its outer shells, a property fundamental to the organization of the periodic table of the elements. In the periodic table, elements through the similar number of outermost electrons fall into the equivalent group.
Structures of Monatomic Ions: The electrons in the outermost shell (the ones by the highest value of n) are the most energetic, and are the ones that are exposed to other atoms. This shell is recognized as the valence shell. The inner, core electrons (inner shell) don't generally play a role in chemical bonding.
Elements through similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the 'extra' s1electron is the one that is lost:
The next shell down is now the outermost shell, which is now full - meaning there is very little tendency to gain or lose more electrons. The ion's electron configuration is the same as the nearest noble gas - the ion is said to be isoelectronic by the nearest noble gas. Atoms 'prefer' to have a filled outermost shell since this builds them more automatically stable.
The Group IV and V metals can lose either the electrons from the p subshell, or from both the s and p subshells, thus attaining a pseudo-noble gas configuration.
The Groups 14 - 17 non-metals gain electrons until their valence shells are full (8 electrons). Group 14 C 1s
The Group 18 noble gases already possess a full outer shell, so they have no tendency to form ions.
Properties of the Representative Elements
The major 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 1 (alkali metals)
The alkali metals are the series of elements in Group 1 of the periodic table (excluding hydrogen in all but one rare circumstance). The series consists of the elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).
The alkali metals are silver-colored (caesium has a golden tinge), soft, low-density metals. Such elements every have one valence electron that is simply lost to form an ion through a single positive charge. They have the lowest ionisation energies in their respective periods. This builds them extremely reactive and they are the most active metals. Due to their activity they take place naturally in ionic compounds not in their elemental state.
The alkali metals react readily through halogens to form ionic salts, these as table salt, sodium chloride (NaCl). They are well-known for their vigorous reactions through water to release hydrogen gas. Such reactions as well often release enough energy to ignite the hydrogen and can be fairly dangerous. As we shift down the group the reactions become increasingly violent. The reaction through water is as follows:
Alkali metal + water → Alkali metal hydroxide + hydrogen
Through potassium as an instance: 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
The oxides, hydrides, and hydroxides of these metals are basic (alkaline). In particular, the hydroxides resulting from the reaction through water are our most general laboratory bases (alkalis). It is from this character that they obtain their group name.
Hydrogen as well has a single valence electron and is generally positioned at the top of Group 1, but it isn't a metal (except under extreme circumstances as metallic hydrogen); rather it exists naturally as a diatomic gas. Hydrogen can form ions through a single positive charge, but elimination of its single electron needs considerably more energy than elimination of the outer electron from the alkali metals. Unlike the alkali metals, hydrogen atoms can as well increase an electron to form the negatively charged hydride ion. The hydride ion is a very strong base and doesn't generally take place except when combined through the alkali metals and several transition metals (that is the ionic sodium hydride, NaH). In compounds, hydrogen most often structures covalent bonds.
Under very high pressure, these as it is found at the core of Jupiter, hydrogen does become metallic and behaves as an alkali metal
Group 2 (alkaline earth metals)
The alkaline earth metals are the series of elements in Group 2 of the periodic table. The series consists of the components beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) (though radium isn't always considered an alkaline on earth due to its radioactivity).
The alkaline earth metals are silvery colored, soft, low-density metals, though are a bit harder than the alkali metals. Such elements all have 2 valence electrons and tend to lose both to form ions by a 2 plus charge. Beryllium is the least metallic element in the group and tends to form covalent bonds in its compounds.
Such metals are less active than the alkali metals, but are still fairly active. They react readily through halogens to form ionic salts, and can react slowly by water. Magnesium reacts only by steam and calcium by hot water. Beryllium is an exception: it doesn't react by water or steam, and its halides are covalent. The oxides are basic and melt in acids and the hydroxides are strong bases, although not as soluble 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. Such were called alkaline earths since of their middle nature between the alkalis (oxides of the alkali metals) and the rare earths (oxides of rare earth metals). The classification of several apparently inert materials as 'earths' is millennia old. The initial known system utilized via the Greeks consisted of 4 elements, including earth. Later alchemists applied the term to any solid substance that didn't melt and was not changed by fire. The realization that 'earths' were not elements but compounds is attributed to the chemist Anoine Lavoisier. In his Traité Élémentaire de Chimie ('Elements of Chemistry') of the year 1789 he termed them Substances simples salifiables terreuses, or salt-forming earth elements. Later, he proposed that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In the year 1808, acting on Lavoisier's idea, Humphrey Davy became the 1st to attain examples of the metals through electrolysis of their molten earths.
Group 13 (boron group)
The Boron group is the series of components in group 13 (formerly group III) in the periodic table. It consists of the elements boron (B), aluminium (Al), gallium (Ga), indium (In), thallium (Tl), and ununtrium (Uut) (unconfirmed).
In this group, we start to see modify over toward non-metallic character. 1st 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 groups are metals. Such elements are characterized by having 3 valence electrons. The metals can lose all 3 electrons to form ions by a 3 plus charge in ionic compounds, but boron tends to form covalent bonds. The oxides of such metals dissolve in acids so may be considered basic, but aluminum oxide as well dissolves in bases. It is amphoteric; that is, it shows both acidic and basic traits.
This is an additional suggestion of the changeover to non-metallic character. Aluminum is the 3rd most abundant element in the earth's crust (7.4%), and is extensively utilized in packaging materials. Aluminum is an active metal, but the stable oxide shapes a shielding coating over the metal making opposed to corrosion.
Group 14 (carbon group)
The carbon group is the series of elements in group 14 ([formerly group IV) in the periodic table. It consists of the elements carbon (C), silicon (Si), and germanium (Ge), tin (Sn), lead (Pb), and ununquadium (Uuq).
This group has blended kinds of element; by the non-metal carbon, 2 metalloids, and 2 metals. The general characteristic is 4 valence electrons. The 2 metals, tin and lead, are fairly unreactive metals and both can shape ions through a 2 plus or a four plus charge in ionic compounds. Carbon forms 4 covalent bonds in compounds rather than form monatomic ions. In the elemental state, it has several forms, the most known of which are graphite and diamond. Carbon is the basis of organic chemistry and of biological molecules. Life depends on carbon. One oxide of carbon, carbon dioxide (CO2), dissolves in water to give a weakly acidic solution. Acidic oxides are characteristic of non-metals. Silicon in several respects is alike to carbon in that it forms 4 covalent bonds, but it doesn't form the broad range of compounds.
Silicon is the 2nd most abundant element in the earth's crust (25.7%) and we are surrounded through silicon having materials: bricks, pottery, porcelain, lubricants, sealants, computer chips, and solar cells. The easiest oxide, silicon dioxide (SO2) or silica is a component of many rocks and minerals.
Group 15 (nitrogen group)
The nitrogen group is the series of elements in group 15 (formerly Group V) of the periodic table. It consists of the elements Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi) and Ununpentium (UUp) (unverified). 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 (pnigein), to choke or stifle, which is a property of nitrogen.
Such elements all have 5 valence electrons. Nitrogen and phosphorous are non-metals. They can grow 3 electrons to form fairly unstable ions with a three minus charge, the nitride and phosphide ions. In compounds, they more often form covalent bonds. Although they aren't in the top ten most general elements in the earth's crust, they are extremely significant elements. Nitrogen, as a diatomic molecule is the chief constituent of air and both elements are necessary for life. Nitrogen comprises about 3% of the weight of the human body and phosphorous about 1.2%. Commercially, such elements are significant for fertilizers. Arsenic and Antimony are metalloids, and bismuth is the only metal in the group. Bismuth can lose 3 electrons to form an ion through a 3 plus charge. Bismuth is as well the heaviest entirely stable element that doesn't decay radioactively to other simpler elements.
Group 16 (Chalcogens)
The chalcogens (by the 'ch' pronounced through a hard "c" as in "chemistry") are the name for the periodic table Group 16 (formerly Group VIb or VIa) in the periodic table. It is sometimes known as the oxygen family. Elements in this group include oxygen (O), sulfur (S), selenium (Se), tellurium (Te), the radioactive polonium (Po), and the synthetic ununhexium (Uuh). The compounds of the heavier chalcogens (chiefly the sulphides, selenides, and tellurides) are collectively recognized as chalcogenides. Unless grouped through a heavier chalcogen, oxides aren't considered chalcogenides.
This group has 6 valence electrons. Oxygen and sulphur are non-metals; their elemental shape is molecular, and they can increase 2 electrons to form ions through a 2 minus charge. Oxygen is through far the most abundant element in the earth's crust (49.5%), and is present in approximately all. It existents elementally in the air as a diatomic molecule, is part of water and many huge minerals, and is needed for life. Sulphur has almost certainly the most allotropes of any element, although the most general and stable form is the yellow crystals of S8 molecules. Although selenium is lumped through the non-metals, and can form selenides similar to oxides and sulphides, its elemental state is that of a metalloid semiconductor as is tellurium and polonium. In their elemental state, they are often termed to as metals. Oxygen can join through sulphur, selenium and tellurium to form polyatomic ion oxo-anions. Oxygen is more electronegative than such elements (S, Se and Te), so they suppose a positive oxidation number in such ions; example is SO42-. The name chalcogen is commonly considered to mean "ore former" from the Greek chalcos 'ore' and -gen 'formation'. Chalcogenides are quite general as minerals. For instance, FeS2 (pyrite) is an iron ore and AuTe2 gave its name to the gold rush town of Telluride, Colorado in the United States.
Group 17 (halogens)
The halogens are the elements in Group 17 (formerly Group VII or VIIa) of the periodic table. They are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At) and the as yet undiscovered ununseptium (Uus).
Such elements all have 7 valence electrons. This group is the 1st one to consist of completely non-metals. They exist as diatomic molecules in their natural state and have a progressive variation of physical properties (see Table below). Fluorine and chlorine exist as gases at room temperature, bromine as a liquid, and iodine as a solid. They require one more electron to fill their outer electron shells, and so have a tendency to gain one electron to shape singly-charged negative ions. Such negative ions are termed to as halide ions, and salts containing these ions are recognized as halides.
Halogens are extremely reactive, and as such can be damaging or lethal to biological organisms in sufficient quantities. Fluorine is the most reactive and the reactivity declines as we go down the group. Chlorine and iodine are both utilized as disinfectants. In their elemental state, the halogens are oxidizing agents and are utilized in bleaches. Chlorine is the active ingredient of most fabric bleaches and is being utilized in the production of most paper products. The oxides and hydrides, like those of most non-metals, of the halogens are acidic. Halide ions joined through single hydrogen atoms to form the hydrohalic acids (that is HF, HCl, HBr, HI), a series of mainly strong acids. (HAt, or 'hydrastatic acid', should as well qualify, but it isn't classically comprised in discussions of hydrohalic acid due to astatine's great instability toward radioactive alpha decay.) They can react through each other to form inter-halogen compounds, and can combine through oxygen in polyatomic oxoanions (e.g., SO42-) Diatomic inter-halogen compounds (BrF, ICl, ClF, etc.) bear strong superficial resemblance to the pure halogens.
Many synthetic organic compounds, and a few natural ones, enclose halogen atoms; such are known as halogenated compounds or organic halides. Chlorine is through far the most abundant of the halogens, and the only one required in comparatively huge amounts (as chloride ions) via human beings. For instance, chloride ions play a key role in brain function during mediating the action of the inhibitory transmitter Gamma-aminobutyric acid (GABA). Chloride ions are as well employed through the body to create stomach acid. Iodine is required in trace amounts for the production of thyroid hormones these as thyroxin. On the other hand, neither fluorine nor bromine is believed to be really necessary for humans, although small amounts of fluoride can create tooth enamel opposed to decay.
The term halogen was coined to mean elements which create salt in union through a metal. It comes from 18th century scientific French nomenclature based on erring adaptations of Greek roots.
Table: Trends in Melting Point, Boiling Point, and Electronegativity of Halogens
Group 18 (noble gases)
The noble gases are the chemical elements in group 18 (formerly group VIII) of the periodic table. They are helium, neon, argon, krypton, xenon, and radon. They are sometimes termed inert gases or rare gases. The name 'noble gases' is a reference to the likewise unreactive noble metals, so termed due to their preciousness, opposition to corrosion and long association through the aristocracy.
The noble gasses are all non-metals and are characterized through having entirely filled shells of electrons. In common this makes them extremely unreactive chemically since it is hard to add or remove electrons. Physically they exist as monatomic gases at room temperature; even those with larger atomic masses (see Table). This is since they have extremely weak inter-atomic forces of attraction, and consequently extremely low melting points and boiling points. Krypton and Xenon are the only noble gasses that doesn't form any compounds at all. Such elements can do this because they have the potential to shape a developed octet through accepting electrons in a vacant d subshell.
Since of their unreactivity, the noble gases were not determined until in the year 1868, whenever helium was noticed spectrographically in the Sun. The isolation of helium on Earth had to wait until in the year 1895. The noble gasses are generally encountered in helium balloons (safer than flammable hydrogen) and lighting. Several of the noble gases glow distinctive colors when utilized inside discharge tubes (neon lights), and Argon is frequently employed inside filament light bulbs.
Table: Trends in Melting Point, Boiling Point and Density of Noble Gases
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