1.1 Nature and Definitions of Acids and Bases
1.1.1 – General Properties of Acids and Bases
Acids: sour taste, dissolve metals, neutralize bases, turn blue litmus red.
Bases: bitter, slippery, neutralize acids, turn red litmus blue.
Two main theories of acid and bases to be discussed
1.1.2 – Arrhenius and Brønsted–Lowry Definitions
Arrhenius definition:
- Acid: substances that produce H+ ions in solution
- Base: substances that produce OH− ions in solution
1.1.3 – Hydronium Ion Hydrogen ions (H+)
Hydronium ion (H3O+): H+ ions are so reactive that they bond with water molecules, forming H3O+.
Brønsted–Lowry definition:
- Acid: proton donor (H+ donor).
- Base: proton acceptor (contains lone pair or can bond with H+).
Example: HCl + H2O → H3O+ + Cl− Here: HCl is acid (donates H+), H2O is base (accepts H+).
1.1.4 – Relation Between H⁺ and H₃O⁺ (Simplified)
In water, H⁺ doesn’t float around by itself.
It attaches to a water molecule (H₂O) and becomes H₃O⁺ (hydronium ion).
That’s why chemists often write H⁺ and H₃O⁺ as if they mean the same thing.
1.1.5 – Conjugate Acid–Base Pairs Definition:
A conjugate acid–base pair consists of two substances related by the gain or loss of a proton. Where once H is added to a base it turn into conjugate acid.
1.1.6 – Amphoteric Substances
Amphoteric = substances that can act as either acid or base depending on the reaction. Because the both have transferable H and atom with lone pair e.
What they turn into depends on what they are reacting with.
Example: Water (H2O)
1.1.7 – Types of Acids
Oxyacids
Have hydrogen, oxygen, and another element.
Example: HNO₃, H₂SO₄.
Organic acids
Contain a –COOH group (carboxylic group).
Example: CH₃COOH (acetic acid).
Hydrohalic acids
Made of hydrogen + halogen.
Example: HCl, HF.
Polyprotic acids
Can give away more than one proton (H⁺).
Diprotic (2 H⁺): H₂SO₄
Triprotic (3 H⁺): H₃PO₄
1.1.9 – Strong and Weak Acids
Strong Acids
Electrolyte: Strong (they ionize 100%, so conduct electricity very well).
Conjugate Base: Very weak (almost no ability to accept H⁺ back).
Example: Cl⁻ from HCl is so weak it’s neutral.
Equilibrium Position: Lies far to the right → almost all HA becomes H⁺ + A⁻.
Examples: HCl, HNO₃, H₂SO₄.
Weak Acids
Electrolyte: Weak (they ionize only a little, so conduct electricity poorly).
Conjugate Base: Relatively strong (it can accept H⁺ back more easily).
Example: CH₃COO⁻ from acetic acid can act as a base.
Equilibrium Position: Lies to the left → most HA stays undissociated, only a little becomes H⁺ + A⁻.
Examples: HF, CH₃COOH.
1.1.10 – Comparison of Strong and Weak Acids
Strong acids: single arrow reaction (→), Ka >> 1, strong electrolytes.
Weak acids: equilibrium reaction (reversible ), Ka << 1, weak electrolytes.
1.1.11 – Equilibrium in Acid Reactions
Equilibrium position shifts toward products for strong acids (full ionization).
For weak acids, equilibrium favors reactants (limited ionization).
1.1.12 – Significance of Ka
Ka = equilibrium constant for acid ionization, measures acid strength when reacts with H20
Large Ka (≫1):
Acid ionizes almost completely.
Strong acid.
Very weak conjugate base.
Equilibrium lies far to the right.
Small Ka (≪1):
Acid ionizes only slightly.
Weak acid.
Conjugate base is stronger.
Equilibrium lies to the left.
1.1.13 – Equilibrium Expression
For weak acids: Ka = [H3O+][A−] / [HA]. Only weak acids have measurable Ka values because strong acids ionize completely.
1.1.14 – Competition for Protons
Different bases in a solution compete to accept protons. The stronger base (with higher proton affinity) will dominate proton capture.
1.1.15 – Acid–Conjugate Base Relation
The stronger the acid, the weaker its conjugate base. The weaker the acid, the stronger its conjugate base.
1.1.16 – Ionic Attraction and Acid Strength
Acid strength depends on:
- Bond polarity: more polar H–X bond → stronger acid.
- Bond strength: weaker H–X bond → stronger acid.
Example: HI is stronger than HF because H–I bond is weaker and easier to break.