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Chemistry IB HL (2016+) > Topic 8: Acids & Bases > Flashcards

Topic 8: Acids & Bases Flashcards

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1
Q

Arrhenius (ionic) theory definitions of acids, bases, and alkalis

A

Acid: produces H+ when dissolved
Base: produces OH- when dissolved
Alkali: soluble base

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2
Q

Limitation of Arrhenius (ionic) theory definition

A

the rxn between NH3 and HCl gas can’t be explained, as NH3 doesn’t contain OH-

NH3 (g) + HCl (g) -> NH4Cl (s)

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3
Q

Bronsted-Lowry theory

A

Acid: proton (H+) donor

Base: proton (H+) acceptor

In an aq soln, a proton can be represented as either hydrogen (H+) or hydronium (H3O+)

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4
Q

in what conditions will H3O+ form?

A

when a water molecule forms a coordinate bond with a proton

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5
Q

monoprotic

A
  • type of acid
  • donates 1 proton
    eg. HCl
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6
Q

diprotic

A
  • type of acid
  • donates 2 protons
    eg. H2SO4
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7
Q

triprotic

A
  • type of acid
  • donates 3 protons
    eg. H3PO4
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8
Q

What can be concluded in a reversible rxn involving an acid/base? Give an example.

A
  • the acid/base is weak as they don’t fully dissociate

e.g. CH3COOH (aq) + H2O (l) → CH3COO- (aq) + H3O+ (aq)
CH3COOH: BL acid
H3O+: conjugate acid

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9
Q

conjugate

A
  • species that remains after the acid has lost a proton (in forward rxn)
  • will act as a species opposite the original species in backwards rxn (e.g. if it was acid in forward, its conjugate will be a base in backward rxn)
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10
Q

conjugate acid-base pair

A
  • conjugate acids and bases will differ from one another by a single proton
  • they are called conjugate acid-base pairs
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11
Q

amphiprotic species

A

species that can act as either BL acid or BL base depending on the rxn

e.g. HCO3 - (aq) + H2O (l) → CO3 2- (aq) + H3O+ (l)
HCO3 - (aq) + H2O (l) → H2CO3 (aq) + OH- (aq)

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12
Q

zwitter ion

A

acts as an acid in the presence of a strong base by donating a proton, and vice versa for strong acids
eg. H2O

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13
Q

requirements for BL acid

A

must be able to dissociate and release H+

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14
Q

requirements for BL base

A

must be able to accept H+ (have lone e- pair)

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15
Q

requirements for BL amphiprotic

A

must possess both a lone e- pair and a H+ ion

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16
Q

difference between amphiprotic and amphoteric

A
  • amphiprotic is specifically related to BL theory (where emphasis is on proton transfer)
  • amphoteric has a broader meaning, describing a substance that can act as both acid and base (even in rxns that don’t involve proton transfer)
  • all amphiprotic substances are amphoteric but the opposite can’t be said
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17
Q

types of bases

A
  • metal oxides/hydroxides
  • ammonia
  • soluble carbonates
  • hydrogen carbonates
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18
Q

why doesn’t HNO3 release H?

A

because of its oxidising properties

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19
Q

ACID + METAL → ? (specific to metals more reactive than Cu)

A

acid + metal → salt + H2

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20
Q

ACID + BASE → ?

A

acid + base → salt + H2O

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21
Q

ACID + METAL OXIDE → ?

A

acid + metal oxide → salt + H2O

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22
Q

ACID + AMMONIA → ?

A

acid + ammonia → salt

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23
Q

ACID + CARBONATE → ?

A

acid + carbonate → salt + H2O + CO2

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24
Q

ΔH(neut)

A

enthalpy change occurring when an acid and base react together to form 1 mol of water

for all strong acids & bases, enthalpy change is very similar: ΔH = -57 kJ/mol

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25
Q

acid turns phenolphthalein…

A

colourless

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26
Q

acid turns methyl orange…

A

red

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27
Q

acid turns litmus paper…

A

red

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28
Q

acid tastes…

A

sour

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29
Q

bases taste…

A

bitter

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30
Q

bases turn litmus paper…

A

blue

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31
Q

bases turn methyl orange…

A

yellow

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32
Q

bases turn phenolphthalein…

A

pink

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33
Q

pH scale

A
  • stands for power of Hydrogen
  • the negative log of the conc of H3O+ or H2
  • expressed in moles/litre
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34
Q

pH 0 turns UI…

A

red

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35
Q

pH 4 turns UI…

A

orange

36
Q

pH 7 turns UI…

A

green

37
Q

pH 10 turns UI…

A

blue

38
Q

pH 14 turns UI…

A

purple

39
Q

examples of strong acids

A
  • HCl
  • HNO3
  • H2SO4
40
Q

examples of strong bases

A
  • NaOH
  • KOH
  • Ba(OH)2
41
Q

examples of weak acids

A
  • CH3COOH

- H2CO3 (dissociates to form CO2)

42
Q

examples of weak bases

A
  • NH3

- C2H5NH2

43
Q

experiments to distinguish between strong and weak acids/bases

A
  • pH measurement
  • conductivity measurements: strong acids/bases conduct better
  • reaction rates (acids only): H conc is much greater in strong acids, so their rxns with metal compounds are much faster
44
Q

differences between strong and weak acids/bases

A
  • electrical conductivity
  • extent of dissociation
  • rate of reaction
  • pH
  • ability to accept/donate protons
  • strong acids/bases form weak/neutral conjugates
  • weak acids/bases form strong conjugates
45
Q

difference between strong and weak bases

A
  • strong bases are good proton acceptors

- weak bases are poor proton acceptors

46
Q

effect of dilution on pH

A
  • if diluted 10x, [H+] conc will be 10% of the original

- thus pH increases

47
Q

effect of volume on pH

A
  • no effect!

- pH is a measure of concentration

48
Q

relationship between [H+] and [OH-]

A

in aqueous state:

[H+] + [OH-] = 10^-14

49
Q

acid deposition

A

the process by which acid-forming pollutants are deposited on the Earth’s surface

50
Q

acid rain

A
  • most prevalent form of acid deposition

- caused by acidic oxides reacting with and dissolving in water in the air

51
Q

sulphur oxides

A

natural source: volcanoes
industrial source: combustion of fossil fuels, smelting of sulfide ores

S (s) + O2 (g) → SO2 (g)

SO2 is further oxidized in the presence of sunlight:
SO2 (g) + 1/2 O2(g) → SO3 (g)

the oxides can react with water in the air to form sulfurous/sulfuric acid
SO2 (g) + H2O (l) → H2SO3 (aq)
SO3 (g) + H2O (l) → H2SO4 (aq)

52
Q

nitrogen oxides

A

natural source: lightning, bacterial action
industrial source: internal combustion engine, jet engines

N2 (g) + O2(g) → 2NO (g)

NO can be further oxidized:
N2 (g) + 2O2 (g) → 2NO2 (g)

Then NO either reacts with H2O:
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)

Or reacts with O2 in the presence of H2O:
4NO2 (g) + O2 (g) + H2O (l) → 4HNO3

53
Q

how acid deposition occurs

A
  • SO2 and NOx are carried into the atmosphere
  • oxidation and hydrolysis occur, forming H2SO4 HNO3 etc
  • acid clouds form
54
Q

effect of acid deposition on plants

A
  • increased acidity in soil leaches important nutrients (e.g. Ca2+, Mg2+, K+)
  • ↓ Mg = ↓ levels of chlorophyll = ↓ ability to photosynthesize
  • Al3+ leaching into soil water damages roots and prevents uptake
  • leading to stunted growth, yellowing, and thinning tree tops
55
Q

effect of acid deposition on lakes and rivers

A
  • aquatic life is highly sensitive to pH
  • below pH 6 many fish decline
  • below pH 5 many microscopic species disappear
  • below pH 4 the lake is effectively dead
  • Al3+ ions interfere with fish gills and reduce oxygen-carrying ability
  • nitrates in acid rain also causes eutrophification
56
Q

effect of acid deposition on buildings

A
  • many buildings are made of CaCO3 (limestone)

- they can be eroded by H2SO4 in acid rain

57
Q

effect of acid deposition on human health

A
  • acid rain irritates mucous membranes
  • thus increasing the risk of respiratory illness
  • high levels of Al in drinking water may be linked to Alzheimer’s
58
Q

pre-combustion method of reducing SO2 emission

A
  • sulphur present as metal sulphides in coal can be removed by crushing the coal and washing it with water
  • metal sulphides have high density, so they sink and separate from coal
  • HDS (hydrodesulphurisation)
59
Q

HDS (hydrodesulphurisation)

A
  • catalytic removal of sulphur from refined petroleum products
  • by reacting it with hydrogen to produce H2S
  • then can be captured and converted to pure sulphur
60
Q

post-combustion method of reducing SO2 emissions

A

flue gas desulphurisation

- this involves adding CaO and CaCO3 to flue gas in the smoke released by power stations

61
Q

flue gas desulphurisation

A

CaO (s) + SO2 (g) -> CaSO3 (s)

CaCO3 (s) + SO2 (g) -> CaSO3 (s) + CO2

62
Q

how to counteract acid rain in lakes

A
  • adding CaO or CaOH (lime)

- this neutralizes acidity, precipitates Al from solution, and increases Ca2+ conc

63
Q

characteristics of Lewis acid-base reactions

A
  • always results in coordinate covalent bond

- because all electrons come from the base

64
Q

Lewis definition of acids

A
  • lone pair acceptors
  • thus they are electrophilic
  • they use their LUMO (lowest unoccupied molecular orbital) when bonding with a base
65
Q

species that can function as Lewis acids

A
  • all cations
  • atoms/ions/molecules with incomplete octet or can accept more than 8 e-s
  • molecules with multiple bonds between atoms of differing electronegativity
66
Q

electrophile

A
  • likes electrons
  • electron-deficient species
  • generally has positive or partially positive charge (dipole) or an incomplete octet
  • accepts a lone pair from another reactant to form a covalent bond
  • basically a Lewis acid
67
Q

nucleophile

A
  • likes nucleus
  • e-rich species
  • has negative or partially negative charge (dipole)
  • donates a lone pair to form a new covalent bond in a rxn
  • basically a Lewis base
68
Q

autoionisation of water

A

when 2 H2O molecules turn into a hydronium and hydroxide ion

  • the hydronium ion takes the proton from the H of the other H2O molecule
  • the hydroxide ion got its proton (H atom) stolen but retains the 2 e-s from the bond
69
Q

is ionisation of water significant?

A

it only has a real effect on pH in extremely dilute acids/bases

70
Q

weak acids in environment

A
  • sulfurous acid, H2SO3
  • nitrous acid, HNO2
  • carbonic acid, H2CO3
71
Q

effect of H2SO3 on environment

A
  • leaching of soil
  • corrosion of limestone buildings/statues/marble
  • harms/kills plants
72
Q

effect of HNO2 on environment

A
  • leaching of soil
  • corrodes marble/limestone buildings/statues
  • harms/kills plants
73
Q

effect of H2CO3 on environment

A
  • acidification of lakes

- corrosion of marble/limestone buildings/marble

74
Q

why would a very acidic aq solution contain OH- ions?

A
  • in v acidic solns, [H+] increases and [OH-] decreases but there are still some present

[OH-] = Kw / [H+] therefore [OH-] cannot be zero.

75
Q

what is the effect of increasing temp on the equilibrium constant of water dissociation equation?

A
  • remember that dissociation of water is endothermic
  • forward rxn favoured
  • Kw increases as [OH-] and [H+] increase
  • endothermic favoured as it will use up some of the heat supplied
76
Q

pKw

A
  • the negative log of Kw
  • pKw = 14 if condition is room temp

pH + pOH = pKw

77
Q

pH buffer

A

resistant to changes in pH on the addition of small amounts of acids/bases

78
Q

making pH buffers

A
  • the weak acid and strong base (2:1 ratio)
  • conjugate pairs (1:1 ratio)
  • weak acid and its salt (1:1 ratio)
79
Q

factors affecting the pH of the buffer

A
  • the pKa/pKb of its constituents

- the ratio of initial concentrations of the constituents used in preparation

80
Q

factors affecting buffers

A
  • dilution

- temperature

81
Q

factors affecting buffers: dilution

A
  • Ka and Kb are equilibrium constants so they are unaffected by dilution
  • so dilution does not change pH, only concentration
  • HOWEVER it does decrease buffer capacity (the amount of acid/base the buffer can absorb without changing pH)
  • because buffering capacity depends on molar concentration of its constituents
82
Q

factors affecting buffers: temperature

A
  • temperature affects equilibrium constants

- so it affects buffer pH

83
Q

reducing NOx emissions

A
  • using catalytic converters in vehicles: hot gases are mixed with air and passed over a platinum/palladium catalyst
    2CO + 2NO –> 2CO2 + N2
  • lower temp combustion: combusting at lower temps reduces the amount of NO formed
84
Q

other ways of reducing acid deposition

A

using CaO and Ca(OH)2 to neutralize the acid

85
Q

assumptions made when making pH buffers

A
  • weak acid/base’s dissociation is considered negligible

- (when using the salt of the weak acid/base) the salt is considered to be fully dissociated

Chemistry IB HL (2016+) (13 decks)

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