Compounds

A substance is a large (macroscopic) collection of molecules. An element is a special case when all the molecules consist of one type of atom; otherwise it is a compound.

Molecules are only stable within a certain range of temperature and pressure and depending on what other molecules are present; otherwise they react into other types.

Generally speaking, there are three broad classes of compounds, depending on whether the atoms are prevalently electropositive, electronegative, or mixed. Electronegativity is how much an atom attracts electrons towards itself in a covalent bond.

Periodic table of electronegativity

Exceptions: Hydrogen is in a class of its own: it forms a metal, covalent bonds and both positive and negative ions. Beryllium is often classified as electropositive but mostly forms covalent bonds.

The compounds shown in these pages are only the simplest, there are, of course, innumerably more complicated molecules.

The Simplest Molecules

The simplest molecules are the atoms of the 'noble' elements He, Ne, Ar,...; they have very low melting and boiling points (mp, bp), and density; He is a superfluid at 2.2 K with no heat capacity.

The next-simplest combinations of the molecule-forming atoms are:

Note about the colors: O is not red, N blue..., atoms only "feel" each other's electric fields; but colors allow us to recognize the atoms.

The strongest molecules of all are CO and nitrogen N₂, with its triple bond, hence it is very inert, even at 3000K. The next strongest simple molecules are HF, CO₂, and H₂O.

• Hydrogen H₂ has the lightest molecule of all, hence has very low melting and boiling points.
• Hydrogen fluoride H-F has a bond that is almost ionic but being so small it doesn't separate so easily (pK=3.2).
• Fluorine F₂ has a very weak and easily broken bond. Only a few fluorine compounds are stable; most are reactive, even at low temperatures: FCN, HOF, F₂CO, NHF₂ .... produce HF, often explosively or when water is added (X-F + H₂O → X-OH + HF); for example,
  • F₂ + H₂O → 2HF + O,
  • 2F₂ + CH₄ → 4HF + C,
  • FNO + H₂O → HF + HNO₂
• Oxygen difluoride has very weak bonds; with a little extra energy, it decomposes into O₂ and F₂. Similarly, NF₃ has weak bonds, liable to lose its fluorine atoms.
• Water has strong intermolecular forces because of its polarity (hydrogen bonding) and geometry. It has the highest boiling point among the simple molecules and it forms aggregates of about 40 molecules fluctuating at 10 GHz when a cold liquid. One in 500 million molecules ionizes as H₃O⁺ and OH^-, each 'hydrated' by four other water molecules (pK=15.7).
• Oxygen O₂ has a double bond and is easily polarized. The two outer electrons in the bond have the same spin, making the molecule magnetic. Ozone O₃ is a less stable form of oxygen; the extra O is held weakly by a coordinate bond.
• Ammonia NH₃ oscillates (by flipping) at 10GHz. It is not as stable as the other molecules: in water, it converts to NH₄⁺ + OH^- (pK=4.8).
• NO₂ has fairly weak N-O bonds and is not very stable since it has an unfilled orbital; at higher temperatures, it tends to decompose as 2NO₂ → N₂O₄ → N₂O₂ + O₂ → 2NO + O₂ → N₂ + O₂. Other possible compounds, such as N₂O and N₂O₃ are less stable. N₂O₅ often breaks up as NO₂ + NO₃
• Carbon dioxide CO₂ has strong double polar bonds. Carbon monoxide CO is even more stable than N₂, but reverts to CO₂ if given the chance. OCCO splits into 2CO immediately.

The elements just below C, N, O, F in the periodic table have similar chemical properties:

They form similar molecules, but because of their bigger size and longer bonds, they are less stable and more likely to join up into larger molecules (with higher mp ...) instead of forming double or triple bonds.

The non-polar molecules CCl₄ (or SiCl₄), NCl₃ (or PCl₃), and OCl₂ (or SCl₂) are respectively similar to CF₄, NF₃, and OF₂, but they are heavier (so have a higher m.p., b.p.).

Like CF₄, the carbon-fluorides are very stable, non-polar and unreactive; but bonds with the larger Br or I are very weak.

Silane SiH₄ is similar to methane CH₄, being non-polar, but weaker and more reactive.

NCl₃ is very explosive: 2NCl₃ → N₂ + 3Cl₂, NCl₃ + 3H₂O → NH₃ + 3ClOH. PCl₃ is slightly more stable but still reacts with H₂O and O₂; it can accept extra chlorine atoms to form PCl₅ (but reverts to PCl₃ when hot).

The polar molecules PH₃, HCl, and H₂S are similar to, but less polar than, NH₃, HF, and H₂O; they have weaker bonds, have a lower m.p./b.p., and are more reactive.

The presence of more orbitals in Si, P, S allows for novel combinations and larger molecules amongst themselves.

Hydrochloric acid HCl dissociates more easily than HF, so it is a stronger acid: HCl + H₂O → H₃O⁺ + Cl^-.

Reactivity of Molecules

In general, molecules are reactive depending on the strength of their bonds and how easily they lose or gain electrons. Weak molecules like F₂ or strong molecules that are charged or polar and lose or gain electrons easily, are liable to react. Reactions occur when molecules collide with each other in the correct orientation and with enough energy ("activation energy"). If only the fastest or precisely oriented molecules react, or their concentration is low, then the reaction rate is slower but still occurs. Molecules in solids are obviously less likely to react than liquids and gases because they only collide with their immediate neighbors, especially if covalent.

At low temperatures, most molecules are stable because they do not have enough energy to react; but as the temperature increases, reactions start to occur and only the more stable molecules survive.

Every molecule decomposes or reacts at high enough temperatures. Beyond 4000 K all bonds are broken; in a flame (1500—3000 K) or UV light, bonds are broken and reform all the time. How do the atoms recombine? Most probably as some of the small stable molecules shown previously, first N₂, HF, CO, and metal chlorides/oxides, then H₂O, HCl, and CO₂, finally any excess atoms form an element (O₂, H₂, Cl₂, C, metals ...)

Reactions occur only if they result in an increase in entropy. This is mostly accompanied by a release of energy — exothermic; more rare are endothermic reactions that absorb surrounding heat (yet the entropy still increases).

A reaction that gives energy to its molecules (exothermic) has a runaway effect. The extra heat may dissipate, but if enough heat is generated, it raises the temperature substantially. If the reaction occurs at the surface and depends on an inward diffusion of molecules to continue, the result is a flame (e.g., a solid/liquid reacting at its surface), otherwise it produces an explosion as gas is produced in a small volume.

Molecules with a net electric field, such as ions or polar molecules, are more likely to react with molecules of opposite charges that they attract. The most stable compounds are either non-polar small molecules, water, and ionic/covalent solids. Any molecule is as strong as its weakest bond.

Stability also depends on the type of surrounding molecules: In the Earth's atmosphere, stability with respect to oxygen and water is what's important.

If there is a chance that a molecule can dissociate, it will do so, at very low concentrations, because once the two parts separate there is a very low chance of rejoining, especially if the molecule can escape 'forever' (e.g. as a gas in an open container). At higher concentrations an equilibrium is reached between the joined and decomposed forms. The ratio of the forms is measured by a logarithmic scale called pK. A readily dissociated molecule has a low or negative pK value.

At high pressure, large molecules in solids are preferred; at low pressure, small gas molecules that form in a reaction escape, even if the reaction absorbs energy (endothermic).

Reducers and Oxidizers

Molecules can be broadly classified as reducers or oxidizers, depending on whether they have high or low-energy molecular orbitals available and thus liable to give or take electrons.

Acids are solutions with a predominance of H⁺ (i.e., H₃O⁺ in water; a 'naked' proton is very reactive and immediately becomes part of a covalent bond).

Alkalis or bases are solutions with a preponderance of OH^-, e.g., NaOH.

In general, reducers react with oxidizers, especially if they are far apart in the following list. Molecules lose electrons/H to others lower in the list:

So metals are mostly chemically unstable (although they may need heat to activate them): their conduction electrons readily transfer to H⁺ (acids/water), fluorine/chlorine, or burn in oxygen,
e.g., magnesium metal in an acid gives off hydrogen: Mg + 2H⁺ → Mg²⁺ + H₂
calcium metal slowly reacts in water: Ca + H₂O → Ca(OH)₂ + H₂
sodium burns in oxygen or chlorine: 4Na + O₂ → 4Na⁺ + 2O^2-
ammonia burns in pure oxygen: 4NH₃ + 3O₂ → 2N₂ + 6H₂O
nitric oxide annihilates ozone: NO + O₃ → NO₂ + O₂
nitric acid is reduced to nitrogen dioxide by hydrogen sulfide: H₂S + 2HNO₃ → S + 2NO₂ + 2H₂O
lead sulfide is oxidized to sulfate by hydrogen peroxide: PbS + 4H₂O₂ → PbSO₄ + 4H₂O
The presence of oxidizers and reducers in the same molecule is usually explosive: e.g., NH₄NO₂ → N₂ + 2H₂O

A reaction that ought to occur but doesn't is usually because only the surface reacts forming a layer that protects it.

Metals

When electropositive atoms have no nearby electronegative atoms to give their electrons to, they form metals.

Alkali metal atoms (Na, K) are relatively large, so have a low interatomic attraction and form large soft bcc lattices, hence low mp/bp, density, heat of fusion; their outer electron is very energetic, so reactive and very metallic.

Alkaline metals (Mg, Ca) lose two electrons per atom but less easily than alkalis.

Transition metals: generally have high mp, are dense, hard, ductile, and insoluble in water; are less reactive; their conduction band characteristics are complicated, so their conduction properties vary; although they have two outer s-shell electrons, they have d-shell electrons in reserve for covalent bonding.
Fe, Co, Ni are ferromagnetic, Cr is anti-ferromagnetic;
W, Ta,... have covalent bonding in addition to being metallic, hence their high mp and strength.
Cu, Ag, Au, Pt are much less reactive;
Zn, Cd, Hg have an almost empty conduction band, hence are weakly metallic with a low mp;
Sn, Pb, Bi have an almost empty conduction band; they can form covalent bonds, e.g., PbCl₄.

Mixtures of electropositive atoms form metal alloys: crystals if their atoms are the same size, or amorphous solids if not. Often, they are stronger but more brittle than one-element metals.

Complexes

When transition element atoms lose their outer two s-electrons, they expose their inner shell of d-orbitals, allowing them to make coordinate bonds with up to 6 polar/ionic molecules such as H₂O (hydrated), NH₃, CN^-, NO, or larger, and form complexes. The bonds are weak, and easily dissociate with heat. Charged complexes are limited up to 4±.

Be has no d-orbitals but is still hydrated by 4 water molecules. Larger metal atoms, such as Pt, cannot form strong coordinate bonds, so they act as good reactive catalysts.

Large Covalent Solids

When atoms combine together by large rigid networks of covalent bonds, the result is a large covalent solid which is extremely strong and stable. Large covalent solids are exceptionally strong, hard, and durable due to the presence of a vast number of covalent bonds. They typically have high melting points, making them resistant to heat and pressure and often do not conduct heat and electricity well.

Carbon's four outer electrons allow it to repeatedly form covalent bonds with other carbon atoms to form a covalent solid; there is a wide variety of carbon forms:

Pure silicon, and SiO₂ (silica), SiC (carborundum), C₃N₄, B₄C, BN also form diamond-like covalent crystals, and almost as hard.

Metalloids

Metalloids are a group of elements in between the metals and nonmetals. They can form covalent bonds but can also lose electrons like metals; their properties range from conductors with a semi-metallic look and fairly high m.p., to semi-conductors, to insulators.

Ge, Si form unreactive covalent solids which are semiconducting if doped with other atoms.
As, Sb, Bi have an almost empty conduction band.
III-V alloys form semiconducting zincblende crystals; II-VI alloys form semiconducting crystals.

Simple Ionic Compounds

Electronegative atoms can accept electrons until all their outer orbitals are filled up.

Halides

Metals react readily with fluorine and chlorine to produce ionic solids:

Fluorides F^-: ionic, e.g., cryolite Na₃AlF₆, or part of complexes.

Chlorides Cl^-: ionic, mostly cubic solids, e.g., with Na⁺ (halite, "common salt"), K⁺ (sylvite) especially mixed as (K⁺,Mg²⁺) (carnallite), and with Ca²⁺ (fluorite); but LiCl forms covalent bonds.

Aluminium chloride is covalent as a gas Al₂Cl₆ or AlCl₃ but is ionic in water. Ionic solids with complexes like (SnCl₆)^2- are also possible.

Hydrides H^-

Hydrogen usually loses its electron, but it can also accept an electron, H^-, from strongly electropositive atoms like Na, Ca, or Mg to form ionic solids.

However its hold on its extra electron is tenuous, and it easily gives it off, acting as a strong reducer, especially with acids, water, or oxygen:

• X⁺H^- + H⁺ → X⁺ + H₂
• X⁺H^- + H₂O → X⁺ + H₂ + OH^-
• 2X⁺H^- + O₂ → X⁺₂O^2- + H₂O

Hydrides of Al, Be, B are covalent: e.g., BeH₂ is a polymer; boron hydrides are in the shape of octahedra, dodecahedra etc. (borane B₂H₆, up to B₆H₁₀, or more, are mostly unstable.) BH₄ borohydride

Oxides O^2-

Electrons transfer readily from metals to O₂ to give the oxide, at least on their surface. Common oxides are cubic ionic lattices, e.g., with Li⁺, Na⁺, Mg²⁺ (periclase), Ca²⁺ (lime), Fe²⁺ possibly with Ti⁴⁺ (ilmenite), as well as Ni²⁺, Co²⁺, Mn²⁺, Cu⁺ (cuprite).

Na⁺ and K⁺ also form oxides of O₂^2-, or O₂^-.

Fe³⁺ (hematite) has a different ionic crystal structure; the same as the oxide of Al³⁺ (corundum) which is a very hard covalent solid.

Ti⁴⁺ (rutile), Sn⁴⁺ (cassiterite), Mn⁴⁺ (pyrolusite), have a different lattice; also Pb⁴⁺Cr⁴⁺ (crocoite).

Most oxides are insoluble in water (except Na₂O) because of the strong ionic attraction. But they react with acids, and in solution, these anions are often very reactive.

O^2- + H₃O⁺ → OH^- + H₂O
OH^- + H₃O⁺ → 2H₂O

Hydroxides OH^-

Hydroxides are less tightly held than the oxides, so they are slightly more soluble because of their lesser charge; examples include Na⁺ (caustic soda), Ca²⁺, Mg²⁺ (brucite); most are hydrated complexes, e.g., Cu²⁺(H₂O)₄(OH^-)₂ (insoluble when uncharged); some are mixed with other anions, e.g., FeO OH (limonite), Cu₂(OH)₃Cl (atacamite). Some metals absorb OH^- to become complexes, e.g., Al(OH)₄^-, Zn(OH)₄^-, Sn(OH)₆^-

Heating solid hydroxides gives back the oxides: 2OH^- → H₂O + O^2-

Hydroxides react with acids: OH^- + H₃O⁺ → 2H₂O.

Al(OH)₃ (bauxite) forms crystals made up of sheets; it reacts with both acids and OH^-: Al(OH)₃ + OH^- → Al(OH)₄^-

Sulfides S^2-

Like O^2-, sulfides form mostly cubic ionic solids, e.g. with Zn²⁺ (sphalerite/wurtzite), Mo⁴⁺ (molybdenite), Cd²⁺ ('cadmium yellow'), Pb²⁺ (galena; semiconducts), sulfosalts (Cu, Ag, Pb)(As, Sb, Bi) such as (Cu,As)²⁺ (enargite), Hg²⁺ (cinnabar).

Thiols SH^- are like the hydroxides; heating gives the sulfides.

Disulfides S₂^2-, e.g., with Fe²⁺ (pyrite), (Ni,Co,Fe)²⁺ (ullmannite)

Sulfides react with acids to form thiols, then hydrogen sulfide: S^2- + H₃O⁺ → HS^- + H₂O
HS^- + H₂O → H₂S + H₂O
Sulfides react slowly with water or oxygen to give oxides (or sulfates).
S^2- + H₂O → O^2- + H₂S
S^2- + 2O₂ → SO₂ + 2O^-

Nitrides N^3-

Nitride anions join mostly with highly charged Mg²⁺, Ca²⁺; but AlN is a hard ionic/covalent solid.

Nitrides react with acids (and water) to give ammonia NH₃.

Phosphides are similar to nitrides.

Nitrites NO₂^-

Nitrates NO₃^-

On heating, nitrates give nitrite and O₂ but it can also explode giving off N₂, O₂, and oxides, e.g., ammonium nitrate: NH₄NO₃ → N₂ + 2H₂O + O.

Nitrates form ionic solids e.g., with Na⁺ (soda nitre), K⁺ (nitre), Ca²⁺.

NO₃^- is often robbed of one of its oxygens, by sulfur, phosphorus, or carbon, releasing NO₂.

• S+6NO₃^- → SO₄^2- + 6NO₂ + 2O^2-; 2O^2- + 4H₃O⁺ → 6H₂O
• 4NO₃^- + C + H₃O⁺ → CO₃^2- + 4NO₂ + OH^- + H₂O

Cyanide

Carbides C^4-

Carbide anions are partly covalent, e.g. with Al³⁺; like the oxides and nitrides, they react with acids and water to give methane CH₄.

With Na⁺ and Ca²⁺, the carbide ion is C₂^2-, or even C₃^4-, as with Li⁺ and Mg²⁺ in order to lessen the charge concentration.

Carbonates CO₃^2-

CO + OH^- → HCO₂^-
Carbonates form ionic solids with Na⁺ (soda), K⁺ (potash), Ca²⁺ (calcite), Mg²⁺ (magnesite), (Ca,Mg)²⁺ (dolomite), CuCO₃.Cu(OH)₂ (malachite), ZnCO₃ (calamine).

Carbonates are liable to give out carbon dioxide, e.g., CO₃^2- + SO₂ → SO₃^2- + CO₂
They react with acids: CO₃^2- + 2H₃O⁺ → CO₂ + 3H₂O

Cyanates OCN^-

Cyanic acid O-C≡N-H is unstable in water, giving off CO₂ + NH₃.

Chlorates

Perchlorate is stable because of its symmetry. But in all the chlorates, the O are held weakly by Cl, so they are highly reactive oxidizers (e.g., give out O₂), even explosive with a reducer. Chlorine in water produces hypochlorite: Cl₂ + H₂O → ClOH + HCl. Which anion predominates depends on the thermodynamic conditions, including concentrations of other molecules such as H⁺.

Sulfites SO₃^-

Sulfites can be oxidized to sulfates: 2SO₃^2- + O₂ → 2SO₄^2-

Sulfates SO₄^2-

Sulfur trioxide eagerly absorbs water to become the more stable symmetric sulfate molecule.

Sulfates form ionic solids, often hydrated, with mixtures (polyhalite) of Na⁺ (mirabilite), Mg²⁺ (epsomite), and Ca²⁺ (gypsum), as well as Al³⁺, Fe²⁺ (melanterite), Cu²⁺ (chalcanthite, antlerite, brochantite), Pb²⁺ (anglesite), Ba²⁺ (barite), KAl₃(SO₄)₂(OH)₆ (alunite).

Sulfates are stable, unreactive, and mostly insoluble; but they can be reduced to sulfites;

Pyrosulfate S₂O₇^2- is a powerful oxidizer, readily giving SO₃ and sulfate.

SO₃ solidifies as a polymer ...-SO₂-O-SO₂-O-....

Phosphites HPO₃^2-

Phosphites are liable to give off PH₃ and phosphate, but is more stable than PCl₃. They can be oxidized to phosphate: 4HPO₃^2- + HSO₄^- → 4HPO₄^2- + HS^-.

Hypophosphite H₂PO₂^-

Phosphates PO₃^3-

Phosphates can form insoluble ionic solids, often with OH^-, e.g., Al₃(PO₄)₂(OH)₃ (wavellite), Ca₅(PO₄)₃(OH,Cl) (apatite), CuAl₆(PO₄)₄(OH)₈ (turquoise). They are able to either accept or release H⁺ depending on the pH.

Phosphoranes PO₅^5- are reactive oxidizers.

Arsenates AsO₄^3- are similar to phosphates, but heavier.

Others Si(OH)₂O^-

Silicic acid

Borate BO₃^3-
Borinic acid accepts an OH^- from water.

Aluminate AlO₃^3- is similar to borate.

Tungstenate WO₄^2- with, for example, Ca²⁺ (scheelite), (Fe,Mn)²⁺ (wolframite), and Molybdates MoO₄^2- with, for example, Pb²⁺ (wulfenite), are similar.

Chromates CrO₄^2-, manganates MnO₄^2-, and permanganates MnO₄^- are reactive, giving off the oxide + H₂O.

When given the chance, these anions join up so that they reduce the charge concentration: XO₄^2- + 2H⁺ → X₂O₇^2- + H₂O ... or structures as complex as the silicates.

Salts

Salts are ionic compounds that do not react with water.

The various salts have different solubilities at room temperature, depending on the balance between lattice energy and hydration energy. In general, smaller highly-charged ions have tightly-held lattices but are also more hydrated, e.g., MgO is tighter than Mg(OH)₂, Ca²⁺ is more hydrated than Na⁺ which is more hydrated than the larger K⁺.

Mixing two salts together, or with an acid, causes precipitation when a cation and an anion form an insoluble solid. For example,

Na₂CO₃ + Cu(NO₃)₂ → CuCO₃ (precipitates) + 2NaNO₃
(What is happening is simply Cu²⁺ + CO₃^2- → CuCO₃ (solid)

Such distinct changes are used to test whether certain cations or anions are present in a solution.