Inorganic Compounds and Nomenclature — IMAT Formulas and Traps

Oxidation Numbers, Common Ions and Writing Formulas

Rules for assigning oxidation numbers

Sum of oxidation numbers

\sum (\text{number of atoms} \times \text{oxidation number}) = \text{overall charge}

overall charge0 for a compound, the ion's charge for an ion

Writing a correct formula by balancing charges

Charge balance

a \times (\text{cation charge}) + b \times (\text{anion charge}) = 0

$a, b$ smallest whole numbers of cations and anions in the formula

Common ions and their charges

Name	Formula	Charge
Ammonium	$\mathrm{NH_4^+}$	1+
Hydroxide	$\mathrm{OH^-}$	1−
Nitrate / nitrite	$\mathrm{NO_3^-}$ / $\mathrm{NO_2^-}$	1−
Hydrogen carbonate	$\mathrm{HCO_3^-}$	1−
Hydrogen sulfate	$\mathrm{HSO_4^-}$	1−
Ethanoate (acetate)	$\mathrm{CH_3COO^-}$	1−
Permanganate	$\mathrm{MnO_4^-}$	1−
Cyanide	$\mathrm{CN^-}$	1−
Carbonate	$\mathrm{CO_3^{2-}}$	2−
Sulfate / sulfite	$\mathrm{SO_4^{2-}}$ / $\mathrm{SO_3^{2-}}$	2−
Chromate / dichromate	$\mathrm{CrO_4^{2-}}$ / $\mathrm{Cr_2O_7^{2-}}$	2−
Phosphate	$\mathrm{PO_4^{3-}}$	3−

Common traps

An element in its free state is always zero

In

\mathrm{Br_2}

,

\mathrm{O_2}

or

\mathrm{P_4}

, each atom is 0, however many atoms the molecule has.

\mathrm{Br_2}

is one molecule of the element bromine: not a compound, not two separate atoms, and not ions.

Hydrogen is negative in metal hydrides

In NaH and

\mathrm{CaH_2}

hydrogen is bonded to a less electronegative metal, so it is

-1

(the hydride ion,

\mathrm{H^-}

). Using

+1

there makes the metal's oxidation number come out impossible.

Carbonate is 2−, hydrogen carbonate is 1−

Each added H takes away one negative charge. So a Group 1 metal forms

\mathrm{M_2CO_3}

but

\mathrm{MHCO_3}

; a formula like

\mathrm{NaCO_3}

leaves the carbonate charge unbalanced.

Brackets matter: Ca(OH)₂ is not CaOH₂

Without brackets the 2 applies only to the H. Whenever a formula needs more than one polyatomic ion, the whole ion goes in brackets with the number outside:

\mathrm{Ca(OH)_2}

,

\mathrm{(NH_4)_2SO_4}

.

Oxides, Hydroxides, Acids, Hydrides and Salts

Learn this subtopic in the notes

Oxides with water: hydroxides and oxoacids

Oxides and water

\text{metal oxide} + \mathrm{H_2O} \rightarrow \text{hydroxide} \qquad \text{non-metal oxide} + \mathrm{H_2O} \rightarrow \text{oxoacid}

Basic, acidic, amphoteric and neutral oxides

Type	Formed by	Behaviour	Examples
Basic	Metals, especially Groups 1 and 2	Neutralises acids; with water gives a hydroxide	$\mathrm{Na_2O}$ , $\mathrm{K_2O}$ , CaO, MgO, CuO
Acidic	Non-metals	Neutralises bases; with water gives an oxoacid	$\mathrm{CO_2}$ , $\mathrm{SO_2}$ , $\mathrm{SO_3}$ , $\mathrm{NO_2}$ , $\mathrm{P_4O_{10}}$ , $\mathrm{SiO_2}$
Amphoteric	Metals near the metal and non-metal border	Reacts with both acids and bases	$\mathrm{Al_2O_3}$ , ZnO, PbO
Neutral	A few non-metals	Reacts with neither acids nor bases	CO, NO, $\mathrm{N_2O}$

Silicon dioxide is acidic although it does not dissolve in water: it reacts with hot concentrated alkali.

Binary acids, oxoacids, hydroxides and hydrides

Class	Made of	Examples
Binary acid (hydracid)	H + one non-metal	HCl, HBr, HF, $\mathrm{H_2S}$
Oxoacid	H + non-metal + O	$\mathrm{H_2SO_4}$ , $\mathrm{HNO_3}$ , HClO, $\mathrm{H_3PO_4}$ , $\mathrm{H_2CO_3}$
Hydroxide	Metal ion + $\mathrm{OH^-}$	NaOH, $\mathrm{Ca(OH)_2}$ , $\mathrm{Al(OH)_3}$
Metal hydride	Metal + H (H is $-1$ )	NaH, $\mathrm{CaH_2}$ , LiH
Non-metal hydride	Non-metal + H (H is $+1$ )	$\mathrm{CH_4}$ , $\mathrm{NH_3}$ , $\mathrm{H_2O}$

Salts: normal, acid, basic and hydrated

Type	Contains	Examples
Normal salt	No replaceable H left	NaCl, $\mathrm{K_2SO_4}$ , $\mathrm{CaCO_3}$ , $\mathrm{CH_3COONa}$
Acid salt	Anion still carries acidic H	$\mathrm{NaHCO_3}$ , $\mathrm{NaHSO_4}$ , $\mathrm{KH_2PO_4}$
Basic salt	$\mathrm{OH^-}$ as well as another anion	Mg(OH)Cl, $\mathrm{Cu_2(OH)_2CO_3}$
Hydrated salt	Water of crystallisation	$\mathrm{CuSO_4 \cdot 5H_2O}$

Common traps

Not every non-metal oxide is acidic

CO, NO and

\mathrm{N_2O}

are neutral: they do not form acids in water. Most other non-metal oxides are acidic, including

\mathrm{NO_2}

, which reacts with water to give nitric acid.

The acid keeps the oxidation number of its oxide

\mathrm{SO_2}

gives sulfurous acid,

\mathrm{H_2SO_3}

(S

+4

), not sulfuric acid. Only

\mathrm{SO_3}

gives

\mathrm{H_2SO_4}

(S

+6

). Match the oxidation numbers before choosing the acid.

An acid is not an oxide, and a hydroxide is not an acid

HClO contains H, Cl and O, so it is an oxoacid, not an acidic oxide (which has no H at all, like

\mathrm{Cl_2O}

). A metal hydroxide such as NaOH is a base. Classify by the elements present before thinking about behaviour.

Hydrogen in a formula does not make an acid salt

Only acidic H counts. Sodium ethanoate,

\mathrm{CH_3COONa}

, and ammonium chloride contain H atoms, but none of them is a leftover acidic H of the parent acid, so they are normal salts.

IUPAC and Traditional Names of Inorganic Compounds

Learn this subtopic in the notes

From name to formula and from formula to name

Charge on each metal ion

\text{charge on each cation} = \frac{\text{number of anions} \times |\text{anion charge}|}{\text{number of cations}}

Naming metal compounds and binary covalent compounds

Formula	Stock (IUPAC) name	Traditional name
$\mathrm{FeCl_2}$	iron(II) chloride	ferrous chloride
$\mathrm{FeCl_3}$	iron(III) chloride	ferric chloride
$\mathrm{Cu_2O}$	copper(I) oxide	cuprous oxide
CuO	copper(II) oxide	cupric oxide
$\mathrm{SnCl_4}$	tin(IV) chloride	stannic chloride
PbO	lead(II) oxide	plumbous oxide
$\mathrm{Hg_2Cl_2}$	mercury(I) chloride	mercurous chloride

Oxoacids and oxoanions: hypo-, -ous, -ic, per- and -ite, -ate

Acid	Acid name	Anion	Anion name	Oxidation number
HClO	hypochlorous acid	$\mathrm{ClO^-}$	hypochlorite	Cl +1
$\mathrm{HClO_2}$	chlorous acid	$\mathrm{ClO_2^-}$	chlorite	Cl +3
$\mathrm{HClO_3}$	chloric acid	$\mathrm{ClO_3^-}$	chlorate	Cl +5
$\mathrm{HClO_4}$	perchloric acid	$\mathrm{ClO_4^-}$	perchlorate	Cl +7
$\mathrm{H_2SO_3}$	sulfurous acid	$\mathrm{SO_3^{2-}}$	sulfite	S +4
$\mathrm{H_2SO_4}$	sulfuric acid	$\mathrm{SO_4^{2-}}$	sulfate	S +6
$\mathrm{HNO_2}$	nitrous acid	$\mathrm{NO_2^-}$	nitrite	N +3
$\mathrm{HNO_3}$	nitric acid	$\mathrm{NO_3^-}$	nitrate	N +5
$\mathrm{H_3PO_4}$	phosphoric acid	$\mathrm{PO_4^{3-}}$	phosphate	P +5
$\mathrm{H_2CO_3}$	carbonic acid	$\mathrm{CO_3^{2-}}$	carbonate	C +4
HCl	hydrochloric acid	$\mathrm{Cl^-}$	chloride	Cl −1

Common traps

-ous and -ic on a metal give its charge, not the anion

Ferrous means

\mathrm{Fe^{2+}}

and ferric means

\mathrm{Fe^{3+}}

, whatever the anion. Do not confuse these endings with -ite and -ate, which belong to the anion. Ferrous sulfate is

\mathrm{FeSO_4}

; iron(II) sulfite is

\mathrm{FeSO_3}

.

-ite is the lower oxidation state, not the higher

Sulfite (S

+4

) has one O fewer than sulfate (S

+6

); nitrite (N

+3

) one fewer than nitrate (N

+5

). Swapping the two endings gives a formula with the wrong number of oxygen atoms.

Work out the metal's charge from the anion first

In

\mathrm{Cu_2O}

the oxide is

-2

, shared by two copper ions, so each is

+1

: copper(I) oxide, cuprous oxide. The subscript 2 is not the charge. Always divide the anion charge by the number of metal atoms.

Inorganic Compounds and Nomenclature formulas

Oxidation Numbers, Common Ions and Writing Formulas

Rules for assigning oxidation numbers

Writing a correct formula by balancing charges

Common ions and their charges

Common traps

Oxides, Hydroxides, Acids, Hydrides and Salts

Oxides with water: hydroxides and oxoacids

Basic, acidic, amphoteric and neutral oxides

Binary acids, oxoacids, hydroxides and hydrides

Salts: normal, acid, basic and hydrated

Common traps

IUPAC and Traditional Names of Inorganic Compounds

From name to formula and from formula to name

Naming metal compounds and binary covalent compounds

Oxoacids and oxoanions: hypo-, -ous, -ic, per- and -ite, -ate

Common traps

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