"Acids and Bases: Comprehensive Study Guide with Practice Questions"
Description:
This document provides an in-depth exploration of Acids and Bases, designed to simplify the topic and prepare students for exams. It includes:
Detailed Summaries: Clear explanations of key concepts such as pH s...
Acids and bases are two broad classes of compounds with great importance in both chemistry
and biochemistry. In industry, acids and bases are used in various reactions to form fertilizers
for agriculture, plastics, paints, dyes, production of fabrics, paper, and cleaning agents as well
as to purify petroleum products.
Acids and bases are also common in our everyday lives. They have a sour taste and many sour-
tasting foods are acidic. Vinegar, for example, is diluted acetic acid (normal household vinegar is
a 3% solution of acetic acid). Other familiar foods with sour flavors get their tartness from acids:
oranges and lemons contain citric acid, wine contains tartaric acid, and aspirin contains
acetylsalicylic acid.
Bases have a bitter flavour and feel slippery because they are soapy in nature, which is why
they are used in cleaning detergents. Sodium hydroxide, a strong base, can dissolve grease and
protein, and is used in oven cleaners, products for unclogging drains, and in hair-removal
lotions.
The table below shows a list of common acids and bases.
Common acids
NAME FORMULA
Hydrochloric acid HCl
Sulphuric acid H2 SO4
Phosphoric acid H3 PO4
Carbonic acid H2 CO3
Ethanoic acid (acetic acid) CH3 COOH
Nitric acid HNO3
Oxalic acid H2 C2 O4
1
,Common bases
NAME FORMULA
Sodium hydroxide NaOH
Potassium hydroxide KOH
Magnesium hydroxide Mg(OH)2
Calcium hydroxide Ca(OH)2
Sodium carbonate Na2 CO3
Ammonia NH3
Ammonium hydroxide NH3 OH
Definition of acids and bases
Arrhenius theory
• An acid is a substance that produces hydrogen ions (H + ) or hydronium ions
(H3 O+ ) when it dissolves in water. In other words, an acid increases the concentration
of H + ions in an aqueous solution. This protonation of water yields the hydronium ion
(H3 O+ ) i.e. H + + H2 O → H3 O+ . H3 O+ is used as a shorthand for H + because it is a
bare proton H + does not exist as a free species in aqueous solution. Examples of an
Arrhenius acid are:
H2 O
HCl → H + + Cl−
H2 SO4 + H2 O → HSO4 − + H3 O+
• A base is a substance that produces hydroxide ions (OH − ) when it dissolves in water. In
other words, a base increases the concentration of OH − ions in an aqueous solution. For
example:
H2 O
NaOH → Na+ + OH −
H2 O
KOH → K + + OH −
2
, Lowry-Brønsted theory
• An acid is a proton donor i.e.
acid → proton + ion
For example:
HCl → H + + Cl−
• A base is a proton acceptor i.e.
base + proton → ion
For example:
NH3 + H + → NH4 +
Acids and bases must always react in pairs. This is because if a compound is to behave as an
acid, donating its proton, then there must necessarily be a base present to accept that proton.
In order to decide which substance is an acid and which is a base, we need to look at what
happens to each reactant as shown below:
Example 1:
HCl(aq) + H2 O(l) → Cl− (aq) + H3 O+ (aq)
The HCl acts as an acid because it loses a proton to become Cl− and H2 O acts as a base because
it accepted a proton to become H3 O+ .
Example 2:
NH3 (g) + H2 O(l) → NH4 + (aq) + OH −
H2 O acts as an acid because it loses a proton to become OH − . NH3 acts as an base because it
gains a proton to become NH4 + .
In the above examples, water acts as either an acid (example 1) or a base (example 2).
A substance that can act as either an acid or a base is called an ampholyte or an amphoteric
substance.
Conjugate acid base pairs
The general scheme for a Brønsted-Lowry acid/base reaction can be visualised in the form:
acid + base ⇌ conjugate base + conjugate acid
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