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IB Chemistry HL > UNIT 8 - Acids And Bases > Flashcards

UNIT 8 - Acids And Bases Flashcards

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1
Q

State the equation of pKw

A

pH + pOH

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2
Q

State how is SO2 most usually produced and its environmental consequences.

A
  • Produced from the burning of fossil fuels especially coal and heavy oil in power plants used to generate electricity.
  • Dissolved in water to form sulphurous acid or can be oxidized to suffer trioxide which then dissolves in water to form sulphuric acid.
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3
Q

State the ways of reduction of SO2 emissions.

A
  • pre-combustion methods - reducing the super present in coal or oil before combustion.
  • post-combustion methods - flue-gas desulfurisation
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4
Q

State how is NO most usually produced and its environmental consequences.

A
  • mainly from internal combustion engines where the burning of fuel releases heat energy causing nitrogen and oxygen from the air to combine.
  • nitrogen dioxide is formed from the oxidation of nitrogen monoxide and NO2 will dissolve in water to form a mixture of nitrous acid and nitric acid.
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5
Q

Define hydrodesulfuristation.

A

Catalytic process of removing soulful from refined petroleum products by reacting it with hydrogen sulphide, a highly toxic gas which is captured and later converted into elemental soulful for use in the manufacture of sulphuric acid.

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6
Q

Define flue-gas desulphurisation.

A

Removal of up to 90% of SO2 from flue gas in the smoke of coal-fired power station before it is released into the atmosphere through reacting wet slurry of CaO and CaCO3 with SO2 to form the neutral product calcium sulphate CaSO4.

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7
Q

Describe the ways of reducing NO emissions.

A
  • Catalytic converters in vehicles - controlling of the exhaust gases with the use of catalytic converters where the hot gases are mixed with air and passed over a platinum or palladium based catalyst which converts them into relatively harmless products.
  • Lower temperature combustion - the formation of nitrogen monoxide at lower temperature.
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8
Q

State the factors determining the pH of a buffer.

A
  • pKa / pKb of its acid/base
  • ratio of its initial concentrations of acid and salt/base and salt.
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9
Q

List the 3 main types of reaction by which acids react to form salts.

A
  1. Acid + reactive metal -> salt + hydrogen
  2. Acid + base -> salt + water
  3. Acid + carbonate -> salt + water + carbon dioxide
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10
Q

Define alkalis.

A

Bases that dissolve in water to form the hydroxide ion OH- (subgroup of bases).

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11
Q

Define leaching.

A

Important minerals for the soil such as Mg2+ , Ca2+ and K+ become soluble and so wash away before they can be absorbed by plants.

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12
Q

State what are the pKa / pKb values dependant on and its consequence.

A

Temperature and thus they need to be quotes at a specific temperature.

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13
Q

List the only soluble hydroxides.

A
  • NH4OH
  • LiOH
  • NaOH
  • KOH
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14
Q

State the pH at half equivalence point.

A

pH = pKa
pOH = pKb

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15
Q

List the only soluble hydrogencarbonates.

A
  • NaHCO3
  • KHCO3
  • Ca(HCO3)2
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16
Q

State the effect of temperature on the value of Kw and equilibrium.

A
  • it will shift to the right
  • increase in the Kw value (thus decrease in pH and a reduction in temperature)
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17
Q

Define Ka and list an expression for it.

A

Acid dissociation constant/acid ionization constant is a direct measure of the strength of an acid as it depends on the position of the equilibrium of acid dissociation.

Ka = ([H3O+][A-]) /[HA]

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18
Q

State the condition necessary for the indicator to be effective in singling the equivalence point of a titration.

A

End-point coincides with the pH are equivalence point.

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19
Q

List the weak acids.

A
  • CH3COOH
  • H2CO3
  • H2PO4
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20
Q

State the formulas for pKa and pKb.

A

pKa = -log 10 Ka
pKb = -log 10 kb

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21
Q

State the formula for Kw.

A

Kw = [H+][OH-]

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22
Q

List the only soluble carbonates.

A
  • (NH4)2CO3
  • Na2CO3
  • K2CO3
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23
Q

Describe the impact of acid deposition on human health.

A
  • Asthma
  • Bronchitis
  • Emphysema
  • Irritation of the eyes
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24
Q

State the Brønsted-Lowry-Lowry theory.

A
  • acid is a proton (H+) donor
  • base is a proton (H+) acceptor
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25
Q

State the main components of acid rain.

A
  • H2SO3
  • H2SO4
  • HN02
  • HNO3
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26
Q

Distinguish through definition strong and weak acids and bases.

A
  • strong acids and bases ionize almost completely in a solution
  • weak acids and bases ionize only partially in a solution (the reaction is reversible)
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27
Q

State where does the point of inflection occur on the pH curve and how is it related to the equivalence point.

A
  • at equivalence a big jump in pH occurs
  • equivalence point is half-way up this jump
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28
Q

State the formula for concentration of protons based on the pH.

A

[H+] = 10 -pH

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29
Q

Define the buffer region of the pH curve.

A

Region relatively resistant to change in pH on the addition of small amounts of base.

30
Q

Define eutrophication.

A

Over-fertilization of bodies of water caused by nitrates present in acid rain and causing algal booms leading to oxygen depletion.

31
Q

State the composition of a buffer solution and reason for it.

A
  • mixing of a weak acid with a solution of its salt of a strong alkali.
  • mixture contains relatively high concentrations of both acid and its conjugate base ready to react with added OH- and H+ in neutralization reactions.
32
Q

State the formula for Kw and pKw for any conjugate acid-base pair.

A

pKa + pKb = pKw
Ka x Kb = Kw

33
Q

State the consequence and effect on pH of the reaction of the salt of a weak acid and a strong base.

A

Anion hydrolysis because the acid is so weak the conjugate base is strong enough to cause hydrolysis release of OH- causes the pH of the solution to increase.

34
Q

State the mechanism of working of a buffer upon addition of an acid or base.

A
  • addition of acid: will combine with the base and thus remove most of the added H+.
  • addition of base: will combine with the acid and thus remove most of the OH-.

Thus the added H+ and OH- are used in the reactions and do not persist in the solution so no change in pH.

35
Q

State the relationship between pKa / pKb and Ka / Kb.

A

Strong acid -> weak acid = pKa increases
Weak acid -> strong acid = Ka increases

Strong base -> weak base = pKb increases
Weak base -> strong base = Kb increases

36
Q

List the weak bases

A
  • NH3
  • C2H5NH2
37
Q

Define acid deposition.

A

All processes by which acidic components as precipitates or gases leave the atmosphere.

38
Q

Define an amphiprotic substance.

A

One which can act as both a proton donor and a proton acceptor.

39
Q

Define the end-point of the indicator.

A

pH at which it changes colour which occurs at the pH equal to its pKa.

40
Q

State where do ligands lie within the Lewis acid/base defintion.

A

Donors of lone pairs thus Lewis bases.

41
Q

List the strong acids.

A
  • HCl
  • HNO3
  • H2SO4
42
Q

Describe the impact of acid deposition on plant life.

A
  • Slower growth
  • Injury or death of plants
  • leaching
  • release of toxic substances such as Al3+ causing damage to plant roots
  • blocking of stomata
43
Q

State the examples of acid-base indicators with their corresponding colours in acids and bases.

A
  • litmus -> acid: pink -> alkali: blue
  • methyl orange -> acid: red -> alkali: yellow
  • phenolphtalein -> acid: colourless -> alkali: pink
44
Q

State where do nucleophiles and electrophiles lie within the Lewis acid/base definition.

A

nucleophile - Lewis base
electrophile - Lewis acid

45
Q

Define an equivalence point in titration.

A

The acid and base exactly neutralize each other.

46
Q

Define a neutralization reaction, state its common ionic equation and thermic effect.

A

reaction in which an acid and base react together to form salt and water

H+(aq) + OH-(aq) -> H2O(l)

exothermic reaction

47
Q

State the relationship between Ka and temperature.

A

Higher value for Ka for a particular temperature -> greater dissociation -> stronger acid.

48
Q

Define acid rain.

A

Rain water, whcih is naturally acidic due to dissolved CO2 with pH < 5.6.

49
Q

State the name for the H3O+ ion.

A
  • hydroxonium ion
  • oxonium ion
  • hydronium ion
50
Q

State the name of Kw and reason for it.

A
  • Ionic product constant of water.
  • concentration of water can be
    considered constant because so little of it ionizes and thus can be combined with Kc to produce a modified equilibrium constant - Kw.
51
Q

State the effect of dillution on the buffer.

A
  • no change in pH
  • lowers its buffering capacity
52
Q

List 2 main types of acid deposition.

A
  • wet acid deposition: rain, snow, sleet, hail, fog
  • dry acid deposition: acidifying particles, gases falling to ground as dust and smoke, later dissolve in water to form acids.
53
Q

State the Lewis acid base theory.

A

Lewis acid -> lone pair acceptor
Lewis base -> lone pair donor

54
Q

State the relationship between the Bronsted-Lowry acids and Lewis acids.

A

All Bronsted-Lowry acids are Lewis acids but the Lewis acids will also include molecules with incomplete valence shell.

55
Q

Describe the impact of acid deposition on building materials.

A

Marble and limestone are forms of CaCO3 which will react with both sulfur dioxide and sulfuric acid forming CaSO4 calcium sulfate washes out of the limestone or flakes off and causes expansions and stress in the stonework because of its greater molar volume.

This leads to erosion of structures

56
Q

State the utility of using pKa and pKb.

A

Useful for weak acids and basis where the extent of dissociation is small as the pKa and pKb values will be usually positive and have no units.

57
Q

Define the conjugate acid-base pairs present in Bronsted-Lowry theory.

A

A consequential formation of a base after a conjugate acid accepts a proton.

! must differ by just one proton.

58
Q

State the pH formula.

A

pH = -log 10[H+]

59
Q

Define Kb and list an expression for it.

A

Base dissociation constant

Kb=([BH+][OH-])/[B]

60
Q

Define half-equivalence point on the pH graph.

A

Represents where exactly half of the acid has been neutralised by the base and converted into salt while the other half remains unreacted.

61
Q

State the principle behind acid-base indicators.

A

Change colour irreversibly because of the H+ ion concentration in the solution which occurs because they are weak acids/bases whose conjugates have different colours.

62
Q

Describe the impact of acid deposition on water.

A

Aluminium ions interfering with the ability of the fish’s gills to take in oxygen.

Eutrophication

Lowering the pH of lakes and thus making them unable to support life.

63
Q

Define a buffer solution.

A

Solution resistant to changes in pH on the addition of small amounts of acid or alkali.

64
Q

State the effect of a reaction of a salt of a strong acid and a strong base.

A

No hydrolysis as the conjugate acid and base are weak and thus there is virtuallly no hydrolysis of ions and the pH is close to neutral.

65
Q

State if there is an effect of temperature change on the buffer.

A

of course as it affects the Ka and Kb values
and thus operations on buffers need to be conducted with stabilized temperature.

66
Q

Describe the impact of acid deposition on metals.

A

Components of dry deposition (sulfur dioxide) and acid rain (sulfuric acid) cause corrosion of the metals.
Acid rain is able to react with and remove the protective oxide layer on the surface of the metals such as aluminium causing significant damage to metallic structures such as bridges, rail road tracks and vehicles.

67
Q

List the strong bases.

A
  • LiOH
  • NaOH
  • KOH
  • Ba(OH)2
68
Q

State the consequence and effect on pH of the reaction of the salt of a strong acid and a weak base.

A
  • cation hydrolysis
  • release of H+ causes the pH of the solution to decrease.
69
Q

State with reason the bond formed in Lewis acid-base reactions.

A

Covalent coordinate bond because both electrons come from the base.

70
Q

Define an indicator.

A

Substances (usually weak acid or bases) that change colour reversibly according to the pH of the solution.

IB Chemistry HL (13 decks)

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