Topic 8: Acids & Bases Flashcards
Arrhenius (ionic) theory definitions of acids, bases, and alkalis
Acid: produces H+ when dissolved
Base: produces OH- when dissolved
Alkali: soluble base
Limitation of Arrhenius (ionic) theory definition
the rxn between NH3 and HCl gas can’t be explained, as NH3 doesn’t contain OH-
NH3 (g) + HCl (g) -> NH4Cl (s)
Bronsted-Lowry theory
Acid: proton (H+) donor
Base: proton (H+) acceptor
In an aq soln, a proton can be represented as either hydrogen (H+) or hydronium (H3O+)
in what conditions will H3O+ form?
when a water molecule forms a coordinate bond with a proton
monoprotic
- type of acid
- donates 1 proton
eg. HCl
diprotic
- type of acid
- donates 2 protons
eg. H2SO4
triprotic
- type of acid
- donates 3 protons
eg. H3PO4
What can be concluded in a reversible rxn involving an acid/base? Give an example.
- 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
conjugate
- 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)
conjugate acid-base pair
- conjugate acids and bases will differ from one another by a single proton
- they are called conjugate acid-base pairs
amphiprotic species
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)
zwitter ion
acts as an acid in the presence of a strong base by donating a proton, and vice versa for strong acids
eg. H2O
requirements for BL acid
must be able to dissociate and release H+
requirements for BL base
must be able to accept H+ (have lone e- pair)
requirements for BL amphiprotic
must possess both a lone e- pair and a H+ ion
difference between amphiprotic and amphoteric
- 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
types of bases
- metal oxides/hydroxides
- ammonia
- soluble carbonates
- hydrogen carbonates
why doesn’t HNO3 release H?
because of its oxidising properties
ACID + METAL → ? (specific to metals more reactive than Cu)
acid + metal → salt + H2
ACID + BASE → ?
acid + base → salt + H2O
ACID + METAL OXIDE → ?
acid + metal oxide → salt + H2O
ACID + AMMONIA → ?
acid + ammonia → salt
ACID + CARBONATE → ?
acid + carbonate → salt + H2O + CO2
ΔH(neut)
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
acid turns phenolphthalein…
colourless
acid turns methyl orange…
red
acid turns litmus paper…
red
acid tastes…
sour
bases taste…
bitter
bases turn litmus paper…
blue
bases turn methyl orange…
yellow
bases turn phenolphthalein…
pink
pH scale
- stands for power of Hydrogen
- the negative log of the conc of H3O+ or H2
- expressed in moles/litre
pH 0 turns UI…
red
pH 4 turns UI…
orange
pH 7 turns UI…
green
pH 10 turns UI…
blue
pH 14 turns UI…
purple
examples of strong acids
- HCl
- HNO3
- H2SO4
examples of strong bases
- NaOH
- KOH
- Ba(OH)2
examples of weak acids
- CH3COOH
- H2CO3 (dissociates to form CO2)
examples of weak bases
- NH3
- C2H5NH2
experiments to distinguish between strong and weak acids/bases
- 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
differences between strong and weak acids/bases
- 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
difference between strong and weak bases
- strong bases are good proton acceptors
- weak bases are poor proton acceptors
effect of dilution on pH
- if diluted 10x, [H+] conc will be 10% of the original
- thus pH increases
effect of volume on pH
- no effect!
- pH is a measure of concentration
relationship between [H+] and [OH-]
in aqueous state:
[H+] + [OH-] = 10^-14
acid deposition
the process by which acid-forming pollutants are deposited on the Earth’s surface
acid rain
- most prevalent form of acid deposition
- caused by acidic oxides reacting with and dissolving in water in the air
sulphur oxides
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)
nitrogen oxides
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
how acid deposition occurs
- SO2 and NOx are carried into the atmosphere
- oxidation and hydrolysis occur, forming H2SO4 HNO3 etc
- acid clouds form
effect of acid deposition on plants
- 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
effect of acid deposition on lakes and rivers
- 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
effect of acid deposition on buildings
- many buildings are made of CaCO3 (limestone)
- they can be eroded by H2SO4 in acid rain
effect of acid deposition on human health
- 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
pre-combustion method of reducing SO2 emission
- 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)
HDS (hydrodesulphurisation)
- 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
post-combustion method of reducing SO2 emissions
flue gas desulphurisation
- this involves adding CaO and CaCO3 to flue gas in the smoke released by power stations
flue gas desulphurisation
CaO (s) + SO2 (g) -> CaSO3 (s)
CaCO3 (s) + SO2 (g) -> CaSO3 (s) + CO2
how to counteract acid rain in lakes
- adding CaO or CaOH (lime)
- this neutralizes acidity, precipitates Al from solution, and increases Ca2+ conc
characteristics of Lewis acid-base reactions
- always results in coordinate covalent bond
- because all electrons come from the base
Lewis definition of acids
- lone pair acceptors
- thus they are electrophilic
- they use their LUMO (lowest unoccupied molecular orbital) when bonding with a base
species that can function as Lewis acids
- 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
electrophile
- 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
nucleophile
- 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
autoionisation of water
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
is ionisation of water significant?
it only has a real effect on pH in extremely dilute acids/bases
weak acids in environment
- sulfurous acid, H2SO3
- nitrous acid, HNO2
- carbonic acid, H2CO3
effect of H2SO3 on environment
- leaching of soil
- corrosion of limestone buildings/statues/marble
- harms/kills plants
effect of HNO2 on environment
- leaching of soil
- corrodes marble/limestone buildings/statues
- harms/kills plants
effect of H2CO3 on environment
- acidification of lakes
- corrosion of marble/limestone buildings/marble
why would a very acidic aq solution contain OH- ions?
- in v acidic solns, [H+] increases and [OH-] decreases but there are still some present
[OH-] = Kw / [H+] therefore [OH-] cannot be zero.
what is the effect of increasing temp on the equilibrium constant of water dissociation equation?
- 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
pKw
- the negative log of Kw
- pKw = 14 if condition is room temp
pH + pOH = pKw
pH buffer
resistant to changes in pH on the addition of small amounts of acids/bases
making pH buffers
- the weak acid and strong base (2:1 ratio)
- conjugate pairs (1:1 ratio)
- weak acid and its salt (1:1 ratio)
factors affecting the pH of the buffer
- the pKa/pKb of its constituents
- the ratio of initial concentrations of the constituents used in preparation
factors affecting buffers
- dilution
- temperature
factors affecting buffers: dilution
- 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
factors affecting buffers: temperature
- temperature affects equilibrium constants
- so it affects buffer pH
reducing NOx emissions
- 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
other ways of reducing acid deposition
using CaO and Ca(OH)2 to neutralize the acid
assumptions made when making pH buffers
- 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
