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Chemistry > Topic 8 & 18 - Acids and bases > Flashcards

Topic 8 & 18 - Acids and bases Flashcards

1
Q

Define Bronsted-Lowry acid and base

A

A Bronsted-Lowry acid donates a proton (BADP)

A Bronsted-Lowry base accepts a proton (BBAP)

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2
Q

Define Lewis acid and base

A

A Lewis acid accepts (a pair of) electrons (LAAE)

A Lewis base donates (a pair of) electrons (LBDE)

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3
Q

What are the requirements for a substance to be a Bronsted-Lowry acid or base?

A

An acid must be able to dissociate and release H+

A base must be able to accept H+ (must have a lone pair of electrons)

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4
Q

Deduce the conjugate acid and base pairs of a Bronsted-Lowry base and acid

A
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5
Q

What can we say about the relationship between Bronsted-Lowry and Lewis acids?

A

All Bronsted-Lowry acids are Lewis acids but not vice versa

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6
Q

What are alkalis?

A

Bases that dissolve in water to form OH-1

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7
Q

List four types of bases that are not hydroxides

A
  1. Metal oxides
  2. Ammonia
  3. (soluble) Carbonates
  4. Hydrogencarbonates
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8
Q

What is a universal indicator?

A

A mixture of several indicators that can be used to identify acids and bases on the whole pH scale

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9
Q

Define salt

A

The compound formed when the hydrogen of an acid is replaced by a metal or another positive ion

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10
Q

What are spectator ions in acid-base reactions?

A

Species that do not change during the reaction (can be cancelled out)

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11
Q

What is the reaction between acids and metals? Give three examples

A

Acid + metal → salt + hydrogen

2 HCl(aq) + Zn(s) → ZnCl2(aq) + H2(g)

H2SO4(aq) + Fe(s) → FeSO4(aq) + H2(g)

2 CH3COOH(aq) + Mg(s) → Mg(CH3COOH)2(aq) + H2(g)

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12
Q

What is the reaction between acids and bases? Give three examples

A

Acid + base → salt + water

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

HNO3(aq) + NH4OH(aq) → NH4NO3(aq) + H2O(l)

2 CH3COOH(aq) + CuO(s) → Cu(CH3COO)2(aq) + H2O(l)

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13
Q

What is the reaction between acids and carbonates? Give three examples

A

Acid + carbonate → salt + water + carbon dioxide

2 HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(l) + CO2(g)

H2SO4(aq) + Na2CO3(s) → Na2SO4(aq) + H2O(l) + CO2(g)

CH3COOH(aq) + KHCO3(s) → KCH3COO(aq) + H2O(l) + CO2(g)

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14
Q

Define effervescence

A

A reaction that gives off gas and produces visible bubbles (acid + metal, acid + carbonate)

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15
Q

Distinguish between strong and weak acids and bases

A

Strong acids and bases dissociate almost completely in solution.

Weak acids and bases only partially dissociate in solution.

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16
Q

What are three examples of strong acids?

A
  1. Hydrochloric acid, HCl
  2. Nitric acid, HNO3
  3. Sulfuric acid, H2SO4
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17
Q

What are three examples of weak acids?

A
  1. Ethanoic acid, CH3COOH
  2. Carbonic acid, H2CO3
  3. Phosphoric acid, H3PO4
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18
Q

What are four examples of strong bases?

A
  1. Lithium hydroxide, LiOH
  2. Sodium hydroxide, NaOH
  3. Potassium hydroxide, KOH
  4. Barium hydroxide, Ba(OH)2
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19
Q

What are two examples of weak bases?

A
  1. Ammonia, NH3
  2. Ethylamine, C2H5NH2

(other amines)

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20
Q

When are comparisons between the strength of acids and bases valid?

A

When the concentration and the temperature of the solutions are the same

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21
Q

List and describe three ways of distinguishing between strong and weak acids and bases

A
  1. Electrical conductivity
    - Depends on the concentration of mobile ions
    - Strong show higher conductivity than weak
    - Can be measured by a conductivity meter or probe
  2. Rate of reaction
    - Strong will produce H+ ions at a faster rate than weak
  3. pH
    - The higher the H+ concentration, the lower the pH
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22
Q

What is the relationship between the concentration of H+ ions and pH?

A

pH = –log[H+]

[H+] = 10–pH

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23
Q

What is the change in [H+] when the pH changes by one unit?

A

A 10-fold change

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24
Q

What are values of pH for solutions that are acidic, neutral, and alkaline?

A

**Acidic **= pH < 7

**Neutral **= pH = 7

Alkaline (basic) = pH > 7

25
Q

What is the expression for the ionic product constant of water (Kw)?

A

Kw= [H+][OH–]

26
Q

Kw is dependent on?

A

Temperature

27
Q

What is the relationship between the concentrations of H+ and OH–?

A

Inversely proportional

28
Q

What is the pOH scale used for?

A

Describing the OH– concentration of solutions

29
Q

What is the relationship between pOH and [OH–]?

A

pOH = –log[OH–]

[OH–] = 10–pOH

30
Q

What is the relationship between pH and pOH at 25°C?

A

pH + pOH = 14

31
Q

What is the relationship between pKw and Kw?

A

pKw = –log(Kw)

Kw = 10–pKw

32
Q

What is the relationship between pH, pOH, and pKw?

A

pH + pOH = pKw

33
Q

What are dissociation constants used for?

A

Expressing the strength of acids and bases

34
Q

Define the acid dissociation constant

A

The greater the value of Ka, the greater the dissociation, and so the stronger the acid

35
Q

Define the base dissociation constant

A

The greater the value of Kb​, the greater the dissociation, and so the stronger the base

36
Q

What is the relationship between Ka, Kb, and Kw?

A

Ka X Kb = Kw

37
Q

What are the relationships between Ka and pKa and between Kb and pKb?

A

pKa = –logKa pKb = –logKb

Ka = 10–pKa Kb = 10–pKb

The relationship between Ka and pKa and between Kb and pKbis inverse

→ The larger the pKa, the weaker the acid

→ The larger the pKb, the weaker the base

38
Q

What is the relationship between pKa, pKb, and pKw?

A

pKa + pKb = pKw = 14

39
Q

What is a buffer solution?

A

A solution that resist the changes in pH on the addition of small amounts of acid or alkali

40
Q

What are the two types of buffer solution?

A
  1. Acidic buffers that maintain the pH below 7
  2. Basic buffers that maintain the pH above 7
41
Q

How is an acidic buffer made? Give an example

A

By mixing an aqueous solution of a weak acid with a solution of its salt of a strong alkali

CH3COOH(aq) + NaCH3COO(aq)

The mixture will contain high concentration of CH3COOH and CH3COO– (acid and conjugate base) that are ready to react with OH– and H+

42
Q

Describe the reaction of CH3COOH + NaCH3COO buffer solution with acid

A

H+ combines with the base CH3COO– to for CH3COOH, decreasing the acidity

43
Q

Describe the reaction of CH3COOH + NaCH3COO buffer solution with base

A

OH–combines with the acid CH3COOH to form CH3COO–and H2O2, removing the OH– ions.

44
Q

How is a basic buffer made? Give an example

A

By combining a weak base with its salt of a strong acid

NH3(aq) + NH4Cl(aq)

Basic buffers work exactly the same way as acidic except the other way round

45
Q

How are pH and pOH of a buffer solution determined?

46
Q

How does dilution influence buffers?

A

Does not change the pH but lowers its buffering capacity

47
Q

What factors can influence buffers?

A

Temperature and dilution

48
Q

What happens to the pH in anion hydrolysis?

A

It increases

49
Q

What happens to the pH in cation hydrolysis?

A

It decreases

50
Q

Outline the possible hydrolysis combinations and the nature of their salts

A
  1. Strong acid + strong base = neutral salt
  2. Weak acid + strong base = basic salt
  3. Strong acid + weak base = acidic salt
51
Q

What is the equivalence point?

A

The point when stoichiometrically equivalent amounts of acid and base have been reacted together. At this point the solution contains only salt and water.

52
Q

Give an example and draw the titration curve of the reaction of a strong acid and strong base

A

HCl + NaOH → NaCl + H2O

53
Q

Give an example and draw the titration curve of the reaction of a weak acid and strong base

A

CH3COOH + NaOH ⇔ NaCH3COO + H2O

54
Q

Give an example and draw the titration curve of the reaction of a strong acid and weak base

A

HCl(aq) + NH3(aq) ⇔ NH4Cl(aq)

55
Q

Give an example and draw the titration curve of the reaction of a weak acid and weak base

A

CH3COOH(aq) + NH3(aq) ⇔ CH3COONH4(aq)

56
Q

Describe the action of an acid-base indicator

A

HIn(aq) ⇔ H+(aq) + In–(aq)

  • Weak acids or weak bases in which the undissociated and dissociated forms have different colours
  • HIn has colour A and In– has colour B
  • H+ concentration determines to which side the equilibrium shifts and thus changes colour
57
Q

When do indicators change colour?

A

When the pH is equal to their pKa (the end point)

Ka = [H+] or pKa = pH

58
Q

What can indicators be used for in titrations?

A

To identify the equivalence point

Chemistry (13 decks)

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