Acids & Bases

Question 1

What is the Lewis definition of acids and bases?

Answer 1

Acids are electron pair acceptors, bases are electron pair donors.

Question 2

What is the Brønsted-Lowry definition of acids and bases?

Answer 2

Acids donate protons (H⁺), bases accept protons.

Question 3

How are Lewis and Brønsted-Lowry definitions related?

Answer 3

Brønsted-Lowry acids and bases are a subset of Lewis acids and bases.

Question 4

Give examples of Lewis acids.

Answer 4

H⁺, BF₃, SiF₄, Al(OH)₃, Fe²⁺.

Question 5

Give examples of Lewis bases.

Answer 5

NH₃, OH⁻, H₂O, CN⁻, Cl⁻.

Question 6

What structural features help identify Lewis acids?

Answer 6

Presence of empty orbitals or partial positive charges.

Question 7

What structural features help identify Lewis bases?

Answer 7

Lone electron pairs or negative charges.

Question 8

In HCl + NaOH, identify each species using both acid-base theories.

Answer 8

H⁺ = B-L acid & Lewis acid, OH⁻ = B-L base & Lewis base.

Question 9

Why is CH₃CH₃ not a Brønsted-Lowry acid?

Answer 9

It lacks a labile proton (non-polar C-H bond).

Question 10

Examples of Brønsted-Lowry acids?

Answer 10

HCOOH, CH₃COOH, HCN, H₂CO₃.

Question 11

Examples of Brønsted-Lowry bases?

Answer 11

NaOH, NH₃.

Question 12

What affects acid strength?

Answer 12

Bond polarity and stability of the conjugate base.

Question 13

Why can’t H⁺ exist freely in water?

Answer 13

It is always solvated (e.g. H₃O⁺).

Question 14

Common hydrated proton species?

Answer 14

H₃O⁺, H₅O₂⁺, H₇O₃⁺.

Question 15

What is a conjugate acid-base pair?

Answer 15

Two species that differ by a proton (HA ⇌ H⁺ + A⁻).

Question 16

What direction does equilibrium favor in acid-base reactions?

Answer 16

Toward the weaker acid-base pair.

Question 17

Write the autoionization of water.

Answer 17

H₂O + H₂O ⇌ H₃O⁺ + OH⁻.

Question 18

What does amphiprotic mean?

Answer 18

A substance that can act as both an acid and a base.

Question 19

What is Kw at 25°C?

Answer 19

1.0 × 10⁻¹⁴.

Question 20

In pure water, what are [H⁺] and [OH⁻]?

Answer 20

Each is 1.0 × 10⁻⁷ M.

Question 21

How does temperature affect Kw?

Answer 21

Increases with temperature.

Question 22

What happens to [OH⁻] if [H⁺] increases?

Answer 22

It decreases (inverse relationship via Kw).

Question 23

How is pH calculated?

Answer 23

pH = –log[H⁺].

Question 24

How is pOH calculated?

Answer 24

pOH = –log[OH⁻].

Question 25

What is the relationship between pH and pOH?

Answer 25

pH + pOH = 14.

Question 26

How to calculate pH of a weak acid?

Answer 26

Use ICE table and Ka expression: [H₃O⁺] = √(Ka·[HA]).

Question 27

What is the 5% rule?

Answer 27

If x < 5% of initial concentration, approximation is valid.

Question 28

How to calculate pH of a weak base?

Answer 28

Use ICE table and Kb expression: [OH⁻] = √(Kb·[B]).

Question 29

How to find pH from weak base [OH⁻]?

Answer 29

pOH = –log[OH⁻], then pH = 14 – pOH.

Question 30

What is the Ka–Kb relationship?

Answer 30

Ka·Kb = Kw or pKa + pKb = 14.

Question 31

What is a buffer solution?

Answer 31

A solution that resists pH changes by neutralising added H⁺ or OH⁻.

Question 32

What happens when OH⁻ is added to a CH₃COOH/CH₃COO⁻ buffer?

Answer 32

CH₃COOH + OH⁻ ⇌ CH₃COO⁻ + H₂O

Question 33

What happens when H⁺ is added to a CH₃COOH/CH₃COO⁻ buffer?

Answer 33

CH₃COO⁻ + H⁺ ⇌ CH₃COOH

Question 34

What is the buffer capacity range?

Answer 34

pKa ± 1

Question 35

When is buffer capacity at its maximum?

Answer 35

When pH = pKa and [A⁻] = [HA]

Question 36

What are two methods of preparing a buffer?

Answer 36

(1) Mix weak acid with conjugate base, (2) Partially neutralise weak acid with strong base

Question 37

How is buffer pH calculated?

Answer 37

(1) ICE table & Ka expression, (2) Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA])

Question 38

What does a speciation diagram show?

Answer 38

Relative amounts of HA and A⁻ vs pH

Question 39

What is the equivalence point?

Answer 39

The point where [acid] = [base]

Question 40

Which types of acid-base titrations exist?

Answer 40

Strong acid/strong base, weak acid/strong base, weak base/strong acid

Question 41

How does concentration affect titration curves?

Answer 41

More concentrated = sharper pH change

Question 42

General equation for indicator equilibrium?

Answer 42

HIn + H₂O ⇌ In⁻ + H₃O⁺

Question 43

What causes colour change in an indicator?

Answer 43

A shift in equilibrium due to pH changes

Question 44

Why are indicators used in small amounts?

Answer 44

To avoid affecting pH of the solution

Question 45

Methyl Orange colour range and pKa?

Answer 45

pKa = 3.46. Red ⇌ Orange ⇌ Yellow

Question 46

How is acid/base strength measured?

Answer 46

By Ka/Kb or pKa/pKb values

Question 47

How does electronegativity affect acid strength?

Answer 47

Higher EN → holds e- tighter → better at stabilising charge on conj. base → stronger acid

Question 48

How does atomic size affect acid strength?

Answer 48

large atoms form weaker bond w H -> H-A bond breaks easier = able to donate H+ easier -> stronger acid

Question 49

How does molecular charge affect acid strength?

Answer 49

Anions = neg charge -> less eager to lose p+ (H+) -> weaker acid Cations = v/v

Question 50

How does hybridisation affect acid strength?

Answer 50

hybridisation changes how tightly atom holds onto H's p+ (H+) - more s-character in hybrid orbital → e- held closer to nucleus → more EN → stronger acid (sp > sp² > sp³)

Question 51

How does resonance affect acid strength?

Answer 51

Stabilises conjugate base → stronger acid

Question 52

Effect of substituents on acid strength?

Answer 52

Electron-withdrawing groups → stronger acids Electron-donating groups → weaker acids They work by affecting how stable the molecule is after losing a proton (H⁺).EWG stabilise base → stronger acid; EDG destabilise base → weaker acid

Question 53

Define solubility and Ksp

Answer 53

Solubility = amount dissolved; Ksp = product of ion concentrations at equilibrium

Question 54

What is the Ksp of CaF₂ with [Ca²⁺]=2.4×10⁻⁴, [F⁻]=4.8×10⁻⁴?

Answer 54

Ksp = (2.4×10⁻⁴)(4.8×10⁻⁴)² = 5.5×10⁻¹¹

Question 55

Why is HgS considered insoluble?

Answer 55

Very low Ksp (1.6×10⁻⁵²)

Question 56

What is the common ion effect?

Answer 56

Presence of shared ion reduces solubility

Question 57

Why is K₂SO₄ added to BaSO₄ in medical imaging?

Answer 57

Reduces Ba²⁺ solubility via common ion effect → safer ingestion

Question 58

What does it mean if Q < Ksp?

Answer 58

More solid can dissolve

Question 59

What does it mean if Q > Ksp?

Answer 59

Precipitation occurs

Question 60

Example: Q = 1.4×10⁻¹⁰, Ksp = 1.8×10⁻¹⁰. What happens?

Answer 60

Q < Ksp → More solid dissolves

Question 61

What determines competition between complexation and protonation?

Answer 61

Relative concentrations and pH