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A2 Chemistry > 5C - Acids & pH > Flashcards

5C - Acids & pH Flashcards

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1
Q

Define:

Acid

Base

Alkali

Salt

A

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

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2
Q

Define:

Conjugate acid and conjugate base

A

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

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3
Q

Define + give an example of each:

Monobasic acid

Dibasic acid

Tribasic acid

A

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

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4
Q

Define:

Brønsted - Lowry acid

Brønsted - Lowry base

A

Brønsted - Lowry acid = Proton donor

Brønsted - Lowry base = Proton acceptor

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5
Q

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+?

A
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6
Q

Equation:

pH = ?

[H+] = ?

A

pH = -log( [H+] )

[H+] = 10-pH

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7
Q

Working out pH from a strong acid

Assumptions?

A

HA(aq) → H+(aq) + A-(aq)

[HA] = Conc of Acid

[H+] = [HA]

Assumptions:

  • All of HA dissociates
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8
Q

Question:

Calculate the pH of 0.43moldm-3 of HCl(aq)

A

HCl is a strong acid ∴

[HA] = [H+]

Assumptions:

  • No HA dissociates

[H+] = 0.43moldm-3

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

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9
Q

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 H2SO4(aq)

What is particular about the ionic equation?

A

Construct ionic equations

NOTE:

The ionic equations of both reactions are the same

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10
Q

Define spectator ion

A

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

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11
Q

What is the Acid dissociation constant (Ka)?

Units = ?

Assumptions?

A

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

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12
Q

Equation:

PKa = ?

Ka from PKa = ?

A

PKa = -log(Ka)

Ka = 10-Pka

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13
Q

Strong alkalis in solution = ?

Assumptions?

A

[OH-] = [Alkali]

Assumption:

  • Alkali fully dissociates
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14
Q

Equation:

Ionic product of water (Kw) = ?

Units = ?

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

A

Kw = [OH-] [H+]

At 25°C, Kw = 1 x 10-14

Units = mol2dm-6

OR

pOH = 14 - pH because pH + pOH = 14

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15
Q

Question:

A

As no’ of carbons increase, acid strength decreases

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16
Q

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

A

Yes - sometimes

17
Q

Define a buffer solution

A

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

18
Q

State 2 different type of buffers that can be made

A

Acid buffer

Alkali Buffer

19
Q

How to form an ACID BUFFER

Assumptions?

A

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
20
Q

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

A

Ethanoic acid + sodium ethanoate

21
Q

How do ACID BUFFERS work to minimise pH changes?

A

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
22
Q

How do ALKALINE BUFFERS work to minimise pH changes?

Use an example to explain this

A

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
23
Q

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

A

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
24
Q

Question:

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

A
25
Q

Question:

Use the pOH method for this question

A
26
Q

Buffer question

A
27
Q

Define

Equivalence point

End point

A

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

28
Q

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

A

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

29
Q

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
A
30
Q

Explain the purpose and function how blood pH is controlled

A

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
31
Q

At what pH is blood mainted at and why?

A

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

32
Q

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?

A

A2 Chemistry (17 decks)

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