Nomenclature of inorganic acids and bases

• investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases (ACSCH067) 

Nomenclature – Naming Acids

Acid names apply to the following two different groups of acids:

• binary acids do not contain oxygen (particularly hydrohalic acids).
• oxoacids (oxyacids) are inorganic compounds made up of oxygen.

 

Hydrohalic acids

• Hydrohalic acids are aqueous solutions of binary inorganic compounds in which hydrogen, H, is combined with a halogen (Group 17) element.

 

Molecular formular	Prefix	+	Modified name of element	+ acid =	“acid” name
HF	Hydro	+	Fluorine + ic	+ acid =	Hydrofluoric acid
HCl	Hydro	+	Chlorine + ic	+ acid =	Hydrochloric acid
HBr	Hydro	+	Bromine + ic	+ acid =	Hydrobromic acid
HI	Hydro	+	Iodine + ic	+ acid =	Hydroiodic acid

 

Oxyacids

• Oxoacids (or oxyacids) are inorganic compounds made up of oxygen (O), hydrogen (H) and one other element (E) called the central atom or central element.
• Examples of molecular formula and their corresponding possible structures showing the general relative positions of hydrogen (H), oxygen (O) and the central element (E) are shown below:

• Oxoacids are named with the name of the central element first using a modified ending (suffix) to indicate the relative amount of oxygen present, followed by the word "acid"

 

Non-halogenic oxyacids

• The "ic" suffix indicates more oxygen is present in the compound than for the "ous" suffix. The table below includes compounds containing oxygen and hydrogen and one other element that is not a halogen (Group 17) element.

 

Table: naming nomenclature of oxyacids where the central atom is not a halogen.

Central element in oxyacid	Most oxygen (highest oxidation state)	Least oxygen (lowest oxidation state)
Nitrogen	Nitric acid (HNO3)	Nitrous acid (HNO2)
Phosphorus	Phosphoric acid (H3PO4)	Phosphorous acid H3PO3
Sulfur	Sulfuric acid (H2SO4)	Sulfurous acid (H2SO3)

 

Halogenic oxyacids

• perhalic acid has the most oxygen of all with the general molecular formula HXO₄
• halic acid has less oxygen than perhalic acid and has the general molecular formula HXO₃
• halous acid has less oxygen than halic acid has the general molecular formula HXO₂.
• hypohalous acid has the least oxygen of all and has the general molecular formula HXO

 

Table: naming nomenclature of oxyacids where the central atom is a halogen.

Central element in oxyacid	More oxygen (highest oxidation state)		Less oxygen (lowest oxidation state)
Chlorine	Perchloric acid (HClO4)	Chloric acid (HClO3)	Chlorous acid (HClO2)	Hypochlorous acid (HClO)
Bromine	Perbromic acid (HBrO4)	Bromic acid (HBrO3)	Bromous acid (HBrO2)	Hypobromous acid (HBrO)

Common Acids and Bases

Examples of common acids and bases

 

Table: common examples of strong and weak acids

Strong acid		Weak acid
Molecular formula	Name	Molecular formula	Name
HClO4	Perchloric acid	H3PO4	Phosphoric acid
HI	Iodic acid	HF	Hydrofluoric acid
HBr	Hydrobromic acid	CH3COOH	Ethanoic acid (acetic acid)
H2SO4	Sulfuric acid	CH2OOH	Methanoic aicd
HCl	Hydrochloric acid	C6H8O7	Citric acid
HNO3	Nitric acid	C2H2O4	Oxalic acid

 

Table: common examples of strong and weak bases

Strong base		Weak base
Molecular formula	Name	Molecular formula	Name
NaOH	Sodium hydroxide	NH3	Ammonia
KOH	Potassium hydroxide	NaHCO3	Sodium bicarbonate
Ba(OH)2	Barium hydroxide	CH3NH2	Methylamine
Ca(OH)2	Calcium hydroxide	(CH3CH2)2NH	Diethylamine

 

 

Acids ranked by their dissociation constant, K_a, in water

Strong acids (complete ionisation) are positioned above hydronium ion (Ka > 55) and weak acids are positioned below.   Strongest acid has the greatest Kaat a given temperature. This will correspond to the lowest pKa.   Ka provides an indication of the tendency of an acid to lost its proton(s). For example, carbonic acid has lower tendency to lose its protons compared with acetic acid.

 

Deprotonation/ionisation of acids

• Not all hydrogen atoms (protons) of an acidic species can be deprotonated or donated away.
• Some acidic species may be able to donate more than one proton per acid molecule.
  • Monoprotic – an acid molecule can only donate one proton. E.g. HCl, HNO₃, CH₃COOH (acetic acid)
  • Diprotic – an acid molecule can donate up to two protons. E.g. H₂SO₄
  • Triprotic – an acid molecule can donate up to three protons. E.g. H₃PO₄

 

If there are more than one proton that could be donated, each has a different K_a value. This means diprotic and triprotic acids typically have multiple K_a and pK_a values. Consequently, the strength of acidic nature of these protons varies.

 

• The number of ionisable protons does not indicate acid strength. A triprotic acid (e.g. H₃PO₄) can have lower K_a values than a monoprotic acid e.g. HCl.

 

Phosphoric acid(H3PO4)	Sulfuric acid (H2SO4)	Hydrochloric acid (HCl)	Nitric acid (HNO3)
Triprotic	Diprotic	Monoprotic	Monoprotic
Ka1 7.1 ´ 10–3 Ka2 6.3 ´ 10–8 Ka3 4.5 ´ 10–13	Ka1 Large Ka2 1.2 ´ 10–2	Ka1 Large	Ka1 Large

 

Acidic Hydrogens

Davy’s hydrogen theory did not account of the acidic nature of hydrogens in hydrogen-containing molecules. Not all hydrogens can be ionised or deprotonated. Hydrogen atoms that can be deprotonated is referred to as the acidic hydrogen.   Evidently, the pKa scale can be as low as –8 and as high as 50. Technically, all hydrogens can be deprotonated but ones with extraordinarily high pKavalues are effectively not acidic because it is very unlikely for these hydrogens to be deprotonated.   Polarity   When all other factors are kept constant, acids become stronger as the X–H bond becomes more polar. The second-row nonmetal hydrides, for example, become more acidic as the difference between the electronegativity of the X and H atoms increases. HF is the strongest of these four acids, and CH4 (methane) is one of the weakest Brønsted-Lowry acids known.   When these compounds act as an acid, a H–X bond is broken to form H+ and X­– ions. The more polar this bond, the easier it is to form these ions. Thus, the more polar the bond, the stronger the acid.

Atomic Radius of the X Atom   At first glance, we might expect that HF, HCl, HBr, and HI would become weaker acids as we go down this column of the periodic table because the X-H bond becomes less polar. Experimentally, we find the opposite trend. These acids actually become stronger as we go down this column.   Acids become stronger as the X-H bond becomes weaker, and bonds generally become weaker as the atoms get larger as shown in the figure to the right.