User:Double sharp/Periodic table/Old periodic trends section

Valence electrons
Elements in the same column of the periodic table have analogous outer electron configurations. The number of valence electrons is thus constant down a column, with the exception of helium.

The valence subshells may be used for chemistry even when they are unoccupied in the atom.

This periodicity in valence configurations has significant consequences for the chemistries of the elements. As chemical reactions involve the valence electrons, elements with similar outer electron configurations may be expected to react similarly and form compounds with similar proportions of elements in them. Such elements are placed in the same group, and thus there tend to be clear similarities and trends in chemical behaviour as one proceeds down a group. As analogous configurations return at regular intervals, the properties of the elements thus exhibit periodic recurrences, hence the name of the periodic table and the periodic law. These periodic recurrences were noticed well before the underlying theory that explains them was developed.

For example, the valence of an element can be defined either as the number of hydrogen atoms that can combine with it to form a simple binary hydride, or as twice the number of oxygen atoms that can combine with it to form a simple binary oxide (that is, not a peroxide or a superoxide). The valences of the main-group elements are directly related to the group number: the hydrides in the main groups 1–2 and 13–17 follow the formulae MH, MH2, MH3, MH4, MH3, MH2, and finally MH. The highest oxides instead increase in valence, following the formulae M2O, MO, M2O3, MO2, M2O5, MO3, M2O7 (although there are many lower oxides as well: for example, phosphorus in group 15 forms two oxides, P2O3 and P2O5). The electron configuration suggests a ready explanation (the number of electrons available for bonding), although a full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone. Today the notion of valence has been extended by that of the oxidation state, which is the formal charge left on an element when all other elements in a compound have been removed as their ions.

Atomic radius
Atomic radii (the size of atoms) generally decrease going left to right along the main-group elements, because the nuclear charge increases but the outer electrons are still in the same shell. However, going down a column, the radii generally increase, because the outermost electrons are in higher shells that are thus further away from the nucleus.

In the transition elements, an inner shell is filling, but the size of the atom is still determined by the outer electrons. The increasing nuclear charge across the series and the increased number of inner electrons for shielding somewhat compensate each other, so the decrease in radius is smaller. The atoms coming immediately after a transition series is introduced are smaller than would have been expected. For example, this occurs in the 4p elements (right after the so-called scandide contraction of the 3d elements) and the 5d elements (right after the lanthanide contraction of the 4f elements).

Ionisation energy
The first ionisation energy of an atom is the energy required to remove an electron from it. This can be written as follows:


 * X (g) → X+ (g) + e−

This varies with the atomic radius: ionisation energy increases left to right and up to down, because electrons that are closer to the nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimised at the first element of each period – hydrogen and the alkali metals – and then generally rises until it reaches the noble gas at the right edge of the period.

For example, the ionisation energy of hydrogen is 13.6 eV. For helium this rises to 24.6 eV, as the helium nucleus has twice the charge of the hydrogen nucleus and the two valence electrons only incompletely shield each other from the nuclear charge. For lithium, the third electron is significantly farther from the nucleus and is shielded well from the nuclear charge by the filled inner shell. Therefore, it is much more easily removed: the ionisation energy of lithium is 5.39 eV. This rises for beryllium to 9.32 eV (although not so high as the inert gases). For boron, the new electron enters a 2p orbital, which is slightly higher in energy; therefore, the ionisation energy drops slightly to 8.30 eV. Carbon and nitrogen continue adding electrons to the 2p orbital and have higher ionisation energies of 11.3 eV and 14.5 eV respectively. With oxygen, electrons begin to pair up in the 2p shell, and so the increased interelectronic repulsion is sufficient to lower the ionisation energy slightly again to 13.6 eV, before the 2p shell is completed with fluorine (17.4 eV) and neon (21.6 eV). The trend in the rest of the periodic table may be rationalised similarly; in particular, because of the analogy of electronic structures, the trend from sodium to argon parallels that from lithium to neon.

In the transition series, the outer electrons are preferentially lost even though the inner orbitals are filling. For example, in the 3d series (scandium through zinc), the 4s electrons are lost first even though the 3d orbitals are being filled. The shielding effect of adding an extra 3d electron approximately compensates the rise in nuclear charge, and therefore the ionisation energies stay mostly constant, though there is a small increase (especially at the end of each transition series).

As metal atoms tend to lose electrons in chemical reactions, ionisation energy is generally correlated with chemical reactivity, although there are other factors involved as well.

The second and subsequent ionisation energies have analogous definitions, being the energies needed to remove subsequent electrons. For example, the second ionisation energy is the energy needed to remove a second electron from an already positive ion:
 * X+ (g) → X2+ (g) + e−

Each ionisation energy is higher than the one before, because of the stronger electrostatic attraction between the positive ion and its negative electrons. Periodicity may still be observed here, as ionisation energies show a large increase when they begin to correspond to taking electrons out of an inner shell. For example, the first three ionisation energies of beryllium (which has two outer electrons) are respectively 9.32 eV, 18.2 eV, and 154 eV.

Electron affinity
The opposite property to ionisation energy is the electron affinity, which is the energy released when adding an electron to the atom:


 * X (g) + e− → X− (g)

A passing electron will be more readily attracted to an atom if it feels the pull of the nucleus more strongly, and especially if there is an available partially filled outer orbital that can accommodate it. Therefore, electron affinity tends to increase down to up and left to right (except for the last column, the noble gases, which have a full shell and have no room for another electron). This gives the halogens in the next-to-last column the highest electron affinities.

Some atoms, like the noble gases, have no electron affinity: they cannot form stable gas-phase anions. If an electron is forced into them, the anion that forms will be metastable, and the electron soon autodetaches. This is sometimes expressed as a negative electron affinity. The noble gases, having high ionisation energies and no electron affinity, have little inclination towards gaining or losing electrons and are generally unreactive.

Some exceptions to the trends occur: oxygen and fluorine have lower electron affinities than their heavier homologues sulfur and chlorine, because they are small atoms and hence the newly added electron would experience significant repulsion from the already present ones. For the nonmetallic elements, electron affinity likewise somewhat correlates with reactivity, but not perfectly since other factors are involved. For example, fluorine has a lower electron affinity than chlorine, but is more reactive.

The second electron affinity would be the energy released when adding an electron to a singly-charged anion:


 * X− (g) + e− → X2− (g)

However, this process involves forcing a further electron into an already negatively charged anion, and no atoms favourably form these in the gas-phase. Such multiply charged anions, such as oxide and sulfide, do nevertheless exist in compounds.

Electronegativity
Another important property of elements is their electronegativity. Atoms form covalent bonds to each other by sharing electrons in pairs, creating an overlap of valence orbitals. The degree to which each atom attracts the shared electron pair depends on the atom's electronegativity – the tendency of an atom towards gaining or losing electrons. The more electronegative atom will tend to attract the electron pair more, and the less electronegative (or more electropositive) one will attract it less. Electronegativity is likewise dependent on how strongly the nucleus can attract the electron pair, and so it exhibits a similar variation to the other properties already discussed: electronegativity tends to fall going up to down, and rise going left to right. The alkali and alkaline earth metals are among the most electropositive elements, while the chalcogens, halogens, and noble gases are among the most electronegative ones.

Electronegativity is generally measured on the Pauling scale, on which the most electronegative atom (fluorine) is given electronegativity 4.0, and the least electronegative atom (caesium) is given electronegativity 0.79.

Electronegativity of an element also somewhat depends on the identity and number of the atoms it is bonded to, as well as how many electrons it has already lost (an atom becomes more electronegative when it has lost more electrons); these effects leave the general trend intact.

Bonding and metallicity
When two atoms of somewhat differing electronegativities form a covalent bond, the covalent bond is polarised: one atom holds on to the electron pair more tightly than the other, so electron density is greater around the more electronegative atom. If the difference in electronegativity is great enough, the end result is basically that the more electropositive atom surrenders its electron totally to the more electronegative one, creating an ionic bond. This can be approximated as being the result of electrostatic attraction between two ions, although in practice some charge transfer takes place between the nearby oppositely charged ions and the bonding always retains some partial covalent character.

A simple substance is a substance formed from atoms of one chemical element. The simple substances of the more electronegative atoms tend to form either discrete covalent molecules (like hydrogen or oxygen) or giant covalent structures (like carbon or silicon), with the noble gases simply staying as single atoms. The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations. The outer orbitals of one atom overlap to share electrons with all its neighbours, creating a giant structure of molecular orbitals extending all over the structure. (This likewise happens when the atoms are not all of the same kind, creating alloys.) This negatively charged "sea" pulls on all the ions and keeps them together in a metallic bond. Elements forming such bonds are often called metals.

The transition between the types of bonding is gradual. The following series gradually changes from more ionic to covalent:
 * NaF, MgF2, AlF3, SiF4, PF5, SF6, IF7, F2

and the following from more ionic to metallic:


 * NaCl, Na2O, Na2S, Na3P, Na3As, Na3Sb, Na3Bi, Na.

Due to the aforementioned trends, metals are generally found towards the left side of the periodic table, and nonmetals towards the right side. Metallicity tends to be correlated with electropositivity and the willingness to lose electrons, which increases right to left and up to down; therefore, the dividing line between metals and nonmetals is roughly diagonal. Elements near the borderline tend to have properties that are intermediate between those of metals and nonmetals, and may have some properties characteristic of both. They are often termed semi-metals or metalloids.

The following table considers the most stable allotropes at standard conditions. The elements coloured yellow form simple substances that are well-characterised by metallic bonding. Elements coloured light blue form giant covalent structures, whereas those coloured dark blue form small covalently bonded molecules that are held together by weaker van der Waals forces. The noble gases are coloured in violet: their molecules are single atoms and no covalent bonding occurs. Greyed-out cells are for elements which have not been prepared in sufficient quantities to experimentally tell.

Characteristically, metals are dense, as they can form closely-packed structures where the atoms are packed as efficiently as possible: in such a structure, each metal atom has twelve nearest neighbours. Some metals have structures distorted from this ideal (such as zinc and tin), and others have less efficiently packed structures; many have only eight nearest neighbours (such as the alkali metals, but also for example iron and tungsten), and near the borderline, antimony, bismuth, and polonium have only six. The structure of gallium is very unusual: the metal forms covalently bonded Ga2 units that are themselves linked by metallic bonding. This has similarities to the structure of iodine, a nonmetal.

A piece of metal is composed of various crystal grains, each of which shows the expected arrangement of atoms; where crystal grains touch, the arrangement is disrupted by misaligned atoms, but the atoms still touch and the metallic bond can pass through. As such, metals are malleable and ductile, as the atoms can move relative to each other without breaking the metallic bond. Putting stress on a metal simply moves the atoms and allows the layers of the structure to roll over one another; for a small stress, they will return to their original position after the stress is released (showing elasticity), but for a larger stress, the layers will roll over to a new position and permanently deform the metal. Heating a metal allows the atoms to rearrange themselves and decreases the number of grains; working it when cold increases this number. The more grains a metal has, the harder and more brittle it becomes.

Metals tend to have high melting and boiling points, due to the strength of the metallic bond (which has to be weakened to allow the metal to melt, and destroyed to allow it to boil). The alkali metals have lower melting and boiling points because they contribute few electrons to the bond, are inefficiently packed, and have relatively large atoms (which reduces the ion–electron attractions). Metals conduct electricity because their electrons are free to move in all three dimensions; similarly, they conduct heat, which is transferred by the electrons as extra kinetic energy (they move faster). These properties persist in the liquid state, as although the crystal structure is destroyed on melting, the atoms still touch and the metallic bond persists (though it is weakened). The characteristic lustre of metals, reflecting but not transmitting light, is also a consequence of the metallic bond.

The elements whose simple substances are not characterised by metallic bonding exhibit different properties. Those forming giant covalent crystals exhibit high melting and boiling points, as it takes a lot of energy to overcome the strong covalent bonds. Those forming discrete molecules are held together mostly by the intermolecular forces, which are more easily overcome; thus they tend to have lower melting and boiling points, and many are liquids or gases at room temperature. Nonmetals are often brittle when solid (as the atoms are held tightly in place), and are dull-looking and conduct electricity poorly as there are no mobile electrons; the orbitals that overlap to allow delocalisation are too high in energy to reach. However, especially near the borderline, this energy gap is small and electrons can readily cross it when thermally excited. Hence, many elements near the borderline are semiconductors, such as silicon and germanium. Arsenic goes further: it conducts electricity like a metal in its most stable form (but ceases to when its structure is destroyed), but each arsenic atom bonds most strongly to its three closest neighbours rather than delocalise its interactions to the extent of antimony and bismuth.

As such, it is common to designate a class of metalloids straddling the boundary between metals and nonmetals, as elements in that region are intermediate in both physical and chemical properties. However, no consensus exists in the literature for precisely which elements should be so designated. When such a category is used, boron, silicon, germanium, arsenic, antimony, and tellurium are usually included, but the majority of sources include other elements as well, without agreement on which extra elements should be added. For example, the periodic table used by the American Chemical Society includes polonium as a metalloid, but that used by the Royal Society of Chemistry does not, and that included in the Encyclopædia Britannica does not refer to metalloids or semi-metals at all.

Further manifestations of periodicity
There are some other relationships throughout the periodic table between elements that are not in the same group. The diagonal relationships Li–Mg, Be–Al, and B–Si occur due to a cancellation of trends; for example, electronegativity decreases down the table but increases to the right, and so moving one space down and to the right roughly compensates and gives rise to two elements with similar properties. For example, lithium is in some ways distinct from the heavier alkali metals and closer in properties to magnesium; it preferentially forms a simple oxide, burns in air to form a nitride, has a carbonate that decomposes on heating, and is considerably harder than the other alkali metals.

Just as similarities arise between elements in the same column of the periodic table, because of their analogous electron configurations, so there are also similarities between elements with similar but not quite as analogous electron configurations. For instance, the elements of group n and group n + 10 have the same number of valence electrons as n varies between 3 and 8, albeit in different valence orbitals. Some similarities can thus be found between the main groups and the transition metal groups. A similar secondary relationship exists between the early actinides and the early transition elements: thus uranium somewhat resembles chromium and tungsten in group 6, as all three have six valence electrons.

The first row of every block tends to show rather distinct properties from the other rows, because the first orbital of each type (1s, 2p, 3d, 4f, 5g, etc.) is significantly smaller than would be expected. This effect is responsible, among other things, for the general similarity of the lanthanides (because the 4f orbitals are small and involving them in bonding is difficult) and the extremely high electronegativities of the 2p elements. The degree of the anomaly is highest for the s-block, is moderate for the p-block, and is less pronounced for the d- and f-blocks.

With the exception of the s-block, the cores also change every two rows of the table; the 4p elements have an extra set of occupied 3d orbitals in the core that the 3p elements lack, and because atomic radii decrease across the 3d elements (the so-called scandide contraction), this cancels out the expected increase in size and metallicity of the 4p elements. A similar effect occurs for the 5d elements, impacted by the preceding 4f or lanthanide contraction. This creates an even-odd difference between the periods that is sometimes known as secondary periodicity, first described by the Russian chemist Evgeny Biron in 1915. Elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods differ in the opposite direction. Thus, many properties in the p-block show a zigzag rather than a smooth trend. For example, phosphorus and antimony (in odd periods of group 15) readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth (in even periods) prefer to stay at +3. The phenomenon was extended to the d-block by Chistyakov in 1968.

When atomic nuclei become highly charged, special relativity becomes needed to gauge the effect of the nucleus on the electron cloud. These relativistic effects result in heavy elements increasingly having differing properties compared to their lighter homologues in the periodic table. For example, relativistic effects explain why gold is golden and mercury is a liquid. These effects are expected to become very strong in the late seventh period. Although experiments cannot yet be conducted due to short half-lives, theoretical studies suggest that tennessine and oganesson do not behave chemically like the lighter halogens and noble gases respectively, but should be closer in behaviour to gallium and silicon respectively. Electron configurations and chemical properties are only clearly known till element 108 (hassium), so the chemical characterisation of the heaviest elements remains a topic of current research.

Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as melting points, boiling points, heats of fusion, heats of vaporisation, atomisation energy, and so on. Similar periodic variations appear for the compounds of the elements, which can be observed by comparing hydrides, oxides, sulfides, halides, and so on. Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include how oxidation states tend to vary in steps of 2 in the main-group elements, but in steps of 1 for the transition elements; the variation in the acidic and basic properties of the elements and their compounds; the stabilities of compounds and methods of isolating the elements; and trends in the stability of coordination complexes and the power of different ligands to donate electrons. Periodicity is and has been used very widely to predict the properties of unknown new elements and new compounds, and is central to modern chemistry.