M6-S2 Nomenclature of inorganic acids and bases

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"
Structure of inorganic oxyacids

 

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, Ka, 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 Ka value. This means diprotic and triprotic acids typically have multiple Ka and pKa 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 Ka 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.