5C - Acids & pH

Question 1

Define:

Acid

Base

Alkali

Salt

Answer 1

Acid = Proton donor in aqueous solution

Base = Proton acceptor in aqueous solution OR a compound that neutralises an acid to form a salt

Alkali = A type of base that dissolves in water forming OH^- ions

Salt = The product of reaction in which the H+ ions of an acid is replaced by metal or ammonium ions

Question 2

Define:

Conjugate acid and conjugate base

Answer 2

Conjugate acid = A species that releases a proton to form a conjugate base

Conjugate base = A species that accepts a proton to form a conjugate acid

Question 3

Define + give an example of each:

Monobasic acid

Dibasic acid

Tribasic acid

Answer 3

Monobasic acids = Reacts with one base as it has only 1 proton - HCl

Dibasic acids = Reacts with 2 bases as it has 2 protons - H₂SO₄

Tribasic acids = Reacts with 3 bases as it has 3 protons - H₃PO₄

Question 4

Define:

Brønsted - Lowry acid

Brønsted - Lowry base

Answer 4

Brønsted - Lowry acid = Proton donor

Brønsted - Lowry base = Proton acceptor

Question 5

Identify the acid base pairs:

Conjugate acid base paris method:

• HCl + NH₃ → NH₄⁺ + Cl^-

Acid 1 / Base 1 method:

• H₂O + HCl → H₃O⁺ + Cl^-

What ion is H₃O⁺?

Answer 5

Question 6

Equation:

pH = ?

[H⁺] = ?

Answer 6

pH = -log( [H⁺] )

[H⁺] = 10^-pH

Question 7

Working out pH from a strong acid

Assumptions?

Answer 7

HA_(aq) → H⁺_(aq) + A^-_(aq)

[HA] = Conc of Acid

[H⁺] = [HA]

Assumptions:

• All of HA dissociates

Question 8

Question:

Calculate the pH of 0.43moldm^-3 of HCl_(aq)

Answer 8

HCl is a strong acid ∴

[HA] = [H⁺]

Assumptions:

• No HA dissociates

[H⁺] = 0.43moldm^-3

pH = log(0.43) = 0.376 (3 sig figs)

Question 9

How do you show the role of H⁺ in the reactions of acids?

Question:

Show the role of H⁺ ions in the rold of the reaction between:

1. Mg_(s) with HCl_(aq)
2. Mg_(s) with H₂SO_4(aq)

What is particular about the ionic equation?

Answer 9

Construct ionic equations

NOTE:

The ionic equations of both reactions are the same

Question 10

Define spectator ion

Answer 10

Spectator ion = Ions that are present but don’t take part in a chemical reaction

Question 11

What is the Acid dissociation constant (Ka)?

Units = ?

Assumptions?

Answer 11

Weak acids partially dissociate: HA_(aq) ⇌ H⁺_(aq) + A^-_(aq)

Ka = [H⁺] [A^-] / [HA]

Assumptions:

• Contribution of [H⁺] from H2O is negligable
• [H⁺] = [A^-]
• No acid dissociate ∴ [HA]_equil = [HA]_initial

∴ Ka = [H⁺]² / [HA]

[H⁺] = √(Ka x [HA])

Units = moldm^-3

Question 12

Equation:

PKa = ?

Ka from PKa = ?

Answer 12

PKa = -log(Ka)

Ka = 10^-Pka

Question 13

Strong alkalis in solution = ?

Assumptions?

Answer 13

[OH^-] = [Alkali]

Assumption:

• Alkali fully dissociates

Question 14

Equation:

Ionic product of water (K_w) = ?

Units = ?

What is another way to work out pH from [OH^-]?

Answer 14

Kw = [OH-] [H+]

At 25°C, Kw = 1 x 10^-14

Units = mol²dm^-6

OR

pOH = 14 - pH because pH + pOH = 14

Question 15

Question:

Answer 15

As no’ of carbons increase, acid strength decreases

Question 16

Is it possible to calculate Ka without making the assumption that no acid dissociates so: [HA]_equil = [HA]_initial

Answer 16

Yes - sometimes

Question 17

Define a buffer solution

Answer 17

Buffer solution = A solution that minimises the change in the pH by small additions of H⁺ or OH^-

Question 18

State 2 different type of buffers that can be made

Answer 18

Acid buffer

Alkali Buffer

Question 19

How to form an ACID BUFFER

Assumptions?

Answer 19

WEAK ACID + SALT OF WEAK ACID → ACID BUFFER

HA_(aq) ⇌ H⁺_(aq) + A^-_(aq)

NaA → Na⁺ + A^- (Na is just used as an example)

Ka = [HA] [A^-] / [HA] ∴ [H⁺] = Ka x [HA] / [A^-]

Assumptions:

• [H⁺] from H₂O is negligable
• [A^-] from acid is negligable, all [A^-] comes from salt
• [HA]_equil = [HA]_initial i.e. no acid dissociates

Question 20

Give an example of a mixture which would form an acid buffer solution in water

Answer 20

Ethanoic acid + sodium ethanoate

Question 21

How do ACID BUFFERS work to minimise pH changes?

Answer 21

HA ⇌ H⁺ + A^- (**Acid partially dissociates)

NaA → Na⁺ + A^- (Salt fully dissociates) - Note that Na is just an example of the salt

ADD H⁺:

• [H⁺] increases, equil shifts to left
• More A^- combining with H⁺ to form HA to do so
• Buffer action maintained until SALT used up

ADD OH^-:

• [OH^-] increases, H⁺ + OH^- → H₂O
• H⁺ decreases compared to OH^- ratio ∴ equil shifts to right
• More HA dissociates to produce H⁺
• Buffer action maintained until ACID used up

Question 22

How do ALKALINE BUFFERS work to minimise pH changes?

Use an example to explain this

Answer 22

Alkali used in this example = Ammonia

NH₃ + H₂O ⇌ NH₄⁺ + OH^-

NH₄Cl → NH₄⁺ + Cl^-

ADD H⁺:

• H⁺ + OH^- → H₂O
• [H⁺] increases ∴ reduces the [OH^-] ratio ∴ equil shifts to right
• NH₃ will react with more H₂O to produce more OH^-
• Buffer action maintained until BASE used up

ADD OH^-:

• [OH^-] increases ∴ equil shifts to left
• More NH₄⁺ will react with excess OH^-
• Buffer action maintained until SALT used up

Question 23

What are the 2 methods to calculate the pH of a buffer solution?

Answer 23

METHOD 1 - Acid dissociation constant

• Ka = [HA] [A^-] / [HA] ∴ [H⁺] = Ka x [HA] / [A^-]
• pH = -log( [H⁺] )

METHOD 2 - PKa form

• PKa = pH - log( [A^-] / [HA] )
• To find PKa neutralise half of the acid ∴ [A^-] and [HA] will cancel out so pKa = pH

Question 24

Question:

Calculate the pH of 0.0280 moldm^-3 NaOH_(aq) using the Kw method

Answer 24

Question 25

Question: Use the pOH method for this question

Answer 25

Question 26

Buffer question

Answer 26

Question 27

Define ## Footnote **Equivalence point** **End point**

Answer 27

**Equivalence point** = The point at which you have equal moles of acid and alkali in a chemical reaction **End point** = The point at which the indicator changes colour

Question 28

How do you pick a suitable indicator for a titration reaction?

Answer 28

Each indicator covers a certain range of pH values. This range must cover the vertical section of the titration curve

Question 29

What are the 4 titration curves that you need to know? * Shape of graph + axis * Estimate start and end point * Equivalence point \> or \< 7 * Where buffer action is

Answer 29

Question 30

Explain the purpose and function how blood pH is controlled

Answer 30

_Carbonic acid - hydrogencarbonate buffer system:_ Maintains blood pH between **7.35-7.45** (needed for blood plasma to function properly) **H₂CO₃ ⇌ H⁺ + HCO₃^-** **H⁺ + OH^- → H₂O** _ADD H⁺:_ * [H⁺] increases, equil shifts to left * More H⁺ reacts with HCO₃^- to produce more H₂CO₃ * Buffer action maintained until SALT of HCO₃^- used up _ADD OH^-:_ * [OH^-] increases, equil shifts to favour right * More H⁺ reacts with OH^- to produce more H₂O * More H₂CO₃ dissociates to provide more H⁺ * Buffer action mainted until H₂CO₃ used up

Question 31

At what pH is blood mainted at and why?

Answer 31

Maintains blood pH between **7.35-7.45** (needed for blood plasma to function properly)

Question 32

_Question:_ The pKa, for carbonic acid - hydrogencarbonate equilibrium is 6.1 at body temperature. What is the ratio of HCO₃^- / H₂CO₃ in healthy blood at a pH of 7.40?

Answer 32