topic 8/18 - acids and bases Flashcards
give the Bronsted-Lowry definition of an acid and a base
acid - proton/H+ donor
base - proton/H+ acceptor
define amphiprotic species
can act as both Bronsted Lowry acids and bases
define a conjugate acid-base pair
a pair of species differing by a single proton
give 2 ways of representing a proton in aqueous solution
H+
H3O+
give an example of an amphiprotic substance
water
h2o -> h+ + OH-
h2o + h+ -> H3O+
define an amphoteric substance
a more general term that refers to substances that can act as an acid or a base
give the difference between an amphoteric and amphiprotic substance
all amphiprotic substances are amphoteric but not all amphoteric substances are amphiprotic
define a Lewis acid and base
a substance that can accept an electron pair
a substance that can donate an electron pair
give 5 properties of acids
- taste sour
- pH < 7
- litmus is red
- phenolphthalein is colourless
- methyl orange is red
give 5 properties of bases
- taste bitter
- pH > 7.0
- litmus is blue
- phenolphthalein is pink
- methyl orange is yellow
give the equation for the reaction of an acid with:
- metal
- base
- metal carbonate/metal hydrogencarbonate
acid + metal -> salt + hydrogen
acid + base -> salt + water
acid + metal carbonate/metal hydrogen carbonate -> salt + carbon dioxide + water
salt and water are produced in…
exothermic neutralisation reactions
give the two equations for pH
pH = -log[H+] or -log[H3O+]
[H+] = 10^-pH
a change in one pH unit represents
a ten fold change in the hydrogen ion concentration
pH 1-6
acid
pH 7
neutral
pH 8-14
alkaline
give the equilibrium expression for the auto-ionisation of water
H2O (l) <-> H+ (aq) + OH- (aq)
Kc expression for ionisation of water
[H+][OH-]/[H2O]
Kw =
[H+][OH-] = 10^-14
define a strong acid
an effective proton donor that is assumed to completely dissociate in water
3 examples of strong acids
HCl, H2SO4, HNO3
define a weak acid
a poor proton donor that dissociates only partially in water.
describe the dissociation of an acid
a reversible reaction that reaches an equilibrium where only a small proportion of the acid molecules have dissociated
compare the conjugate base of a weak acid to that of a strong acid
the conjugate base of weak acid is stronger (has a higher affinity for a proton) than does the conjugate base of a strong acid
define a strong base
an effective proton acceptor that completely dissociates in water
give an alternative explanation of strong/weak bases
A strong base is a base that ionizes completely in an aqueous solution. A weak base is a base that ionizes only slightly in an aqueous solution.
why does a metal hydroxide not act as a bronsted-lowry base
because it does not have the capacity to accept a proton.
when does this change and why?
in solution the hydroxide ion acts as a base, accepting a proton:
OH- + H3O+ -> 2H2O
ammonia is a
weak base
ammonia + water
NH3 (aq) + h2O (l) <-> NH4+ (aq) + OH- (aq)
effect of strength of acid/base on electrical conductivity
in a stronger acid/base, concentration of ions is greater, making it more conductive
effect of strength of acid/base on pH
acid - in a strong acid, there is a higher concentration of H+ ions, so lower pH
base - in a strong base, there is a higher concentration of OH- ions, so higher pH
effect of strength of acid/base on rate of reaction
stronger acid/base means higher concentration of H+/OH- ions, leading to a higher rate
why is the enthalpy change of neutralisation for a strong acid almost identical to that for a weak acid?
the neutralisation reaction removes ionised species from the dissociation reaction, so driving the reaction to completion
state the effects of acid deposition on the earth
Vegetation:
The acid (H+ ions) can displace metal ions from the soil that are consequently washed away (calcium, magnesium, potassium ions).
Mg2+ ions needed to produce chlorophyll, so photosynthesis is stopped
Acid rain also causes aluminium ions to dissolve from rocks, which damages plant roots and limits water growth.
This can cause stunted growth and thinning/yellowing of leaves on trees.
Lakes and rivers:
Insect, larvae, fish and invertebrates cannot survive below pH 5.2.
No life will survive below pH 4.0.
Acid rain can also dissolve minerals like Al3+ from rocks, which can damage aquatic life (aluminium damages fish gills).
Buildings:
Limestone and marble are eroded by acid rain and dissolve away.
This exposes a fresh surface to react with more acid.
CaCO3 (s) + H2SO4 (aq)CaSO4 (s)+H2O (l) + CO2 (g)
Human health
Acids irritate mucous membranes and cause respiratory illnesses such as asthma and bronchitis
Acidic water can dissolve heavy metal compounds releasing poisonous ions such as Cu2+, Pb2+ and Al3+ which may be linked to Alzheimer’s.
typical pH value of rainwater
5.6
why is rain naturally acidic?
due to the presence of dissolved carbon dioxide, which forms weak carbonic acid, H2CO3
equations for rain becoming acidic due to carbon dioxide
CO2 + H2O <-> H2CO3
H2CO3 <-> H+ + HCO3-
HCO3- <-> H+ + CO32-
acid deposition due to sulphur
Fossil fuels are often contaminated with small amounts of sulfur impurities:
- When these contaminated fossil fuels are combusted, the sulfur in the fuels get oxidised to sulfur dioxide
- S (s) + O2 (g) → SO2 (g)
- Sulfur dioxide may be further oxidised to sulfur trioxide
- 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)
The sulfur dioxide and sulfur trioxide then dissolve in rainwater droplets to form sulfurous acid and sulfuric acid
- SO2(g) + H2O (l) → H2SO3 (aq)
- SO3 (g) + H2O (l) → H2SO4 (aq)
dissociation H2SO3
H2SO3 (aq) + H2O (l) <-> HSO3- (aq) + H3O+ (aq)
acid deposition due to nitrogen
The temperature in an internal combustion engine can reach over 2000 °C
Here, nitrogen and oxygen, which at normal temperatures don’t react, combine to form nitrogen monoxide:
N2 (g)+ O2 (g) ⇌ 2NO (g)
Nitrogen monoxide reacts further forming nitrogen dioxide:
2NO (g) + O2 (g) ⇌ 2NO2 (g)
Nitrogen dioxide gas reacts with rain water to form a mixture of nitrous and nitric acids, which contribute to acid rain:
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)
what are the main causes of acid deposition?
increased industrialisation and economic development in many parts of the world have led to rapidly increasing emissions of the nitrogen and sulfur oxides that cause acid rain
pre-combustion methods to reduce sulfur emissions
Pre-combustion:
Hydrodesulfurization (hydrotreating)
Heating crude oil fractions with hydrogen in the presence of a catalyst.
This converts the sulphur to hydrogen sulphide (H2S), which can be removed from the reaction mixture by bubbling it through an alkaline solution.
The H2S can be subsequently converted back into sulphur and sold to other companies to make sulfuric acid.
post-combustion methods to reduce sulfur emissions
Post-combustion:
Commonly used in coal-fired power stations
Involves passing the exhaust gases from the furnace through a vessel where SO2 can react with alkalis/bases like calcium oxide, carbonate, or hydroxide.
CaCO3 (s) + SO2 (g) → CaSO3 (s) + CO2 (g)
reaction of calcium carbonate with acid rain
CaCO3 (s) + H2SO4 -> CaSO4 + CO2 + H2O
lewis acid
electron pair acceptor
lewis base
electron pair donor
when a Lewis base reacts with a Lewis acid, what is formed?
a coordinate bond
nucleophile
Lewis base
electrophile
Lewis acid
acid dissociation constant for weak acids
HA + H2O <-> A- + H3O+
Ka = [A-][H3O+]/[HA]
base dissociation constant for weak bases
B + H2O <-> BH+ + OH-
Kb = [BH+][OH-]/[B]
Kw =
Ka x Kb
the stronger the acid
- the larger the Ka
- the weaker the conjugate base
- the smaller the Kb of the conjugate base
the stronger the base
- the larger the Kb
- the weaker the conjugate acid
- the smaller the Ka of the conjugate acid
is ionisation of water exothermic or endothermic?
endothermic
effect of temperature on Kw
- a rise in temperature will result in the forward reaction being favoured
- thus increasing concentration of hydroxide and hydrogen ions
- increase in magnitude of Kw
- decrease in pH
pKa
-log10Ka
pKb
-log10Kb
Ka
10^-pKa
Kb
10^-pKb
calculate the pH of a solution of 0.08 mol/dm3 methanoic acid, for which pKa = 3.75 at 298K
p401
3 examples of strong acids
hydrochloric acid, Hal
nitric acid, HNO3
sulphuric acid, H2SO4
3 examples of strong bases
lithium hydroxide, LiOH
sodium hydroxide, NaOH
potassium hydroxide, KOH
barium hydroxide, Ba(OH)2
3 examples of weak acids
ethanoic acid, CH3COOH
carbonic acid, H2CO3
phosphoric acid, H3PO4
3 examples of weak bases
ammonia, NH3
ethylamine, C2H5NH2
what are buffer solutions?
a solution that resists changes in pH when small amounts of acid or alkali are added to itt
state the two types of buffer solutions
acidic buffer solutions
basic buffer solutions
acidic buffer solutions (giving a pH below 7)
made from a mixture of weak acid and the salt of its conjugate base
eg. ethnic acid, CH3COOH (aq) and sodium ethanoate, CH3COONa (aq), a salt of the conjugate base
basic buffer solutions (giving a pH above 7)
made from a mixture of a weak base and the salt of its conjugate acid
eg. ammonia, NH3 (aq), a weak base, and ammonium chloride, NH4Cl, a salt of the conjugate acid
describe how acidic buffer solutions work using CH3COOH as an example
CH3COOH, a weak acid, dissociates only partially:
- equilibrium 1: CH3COOH (aq) <-> CH3COO- (aq) + H+ (aq)
CH3COONa, the salt of the conjugate base, is soluble in water so it dissociates completely:
- reaction 2: CH3COONa (aq) -> CH3COO- (aq) + Na+ (aq)
the buffer solution will contain a high concentration of CH3COO- ions, produced in reaction 2, meaning that the position of equilibrium 1 lies very far to the left, resulting in a high concentration of CH3COOH acid, undissociated, and a relatively low conc of H+ ions
- adding H+ ions:
almost all of the ions will react with the reservoir of CH3COO- ions and equilibrium 1 shifts to the left. the conc of H+ ions hardly changes and, therefore, pH remains almost constant - adding OH- ions;
almost all of the OH- ions react with the low concentration of H+ ions present to form water. equilibrium 1 then shifts to the right, replacing the H+ ions and therefore pH remains almost constant.
describe how basic buffer solutions work using NH3 as an example
NH3 + H2O <-> NH4+ + OH-
NH4Cl -> NH4+ + Cl-
when acid (H+) is added, it will react with OH- to form water, causing the equilibrium to shift to the right.
when alkali (OH-) is added, it will react with NH4+ and cause the equilibrium to shift to the left.
how can buffer solutions be made?
acid:
if a weak acid, such as ethnic acid, is partially neutralised by a strong base, such as sodium hydroxide:
NaOH + CH3COOH -> CH3COONa + H2O
all of the alkali reacts but some acid remains so the solution contains the weak acid, CH3COOH and its salt CH3COONa, so is a buffer solution.
base:
when a weak base (NH3) is partially neutralised by a strong acid (HCl)
NH3 + HCl -> NH4Cl
NH3 is excess, no HCl left, so solution contains NH3 and NH4+ (buffer)
how do u calculate the pH of a buffer solution
using the Henderson hasselbalch equation
when [HA]=[A-],
pH = pKa
describe the carbonic acid/hydrogen carbonate buffer
the pH of blood must stay between 7.35 and 7.45; controlled by a mixture of buffers.
H2CO3 <-> H+ + HCO3-
(H2CO3 <-> CO2.H2O)
at high pH values the hydrogen carbonate anion can lose the remaining hydrogen as a proton to give a carbonate buffer:
HCO3- <-> CO32- + H+
4 types of titration pH curve
- strong acid/base
- weak acid/strong base
- weak base/strong base
- weak base/weak acid
strong acid + strong baseale
draw (p405)
- starting point on the pH axis is an important feature of a pH curve as it is an indication of the relative strength of the acid. The strong acid gives an initial pH reading of =1
- there is a gradual rise in the pH as the titration approaches the equivalence point
- the sharp rise in the pH at the equivalence point (pH=7) is described as the point of inflection of the curve
- once there is no remaining acid to be neutralised, the curve flattens and finishes at a high pH reflecting the strong base
eg HCl + NaOH -> NaCl + H2O
weak acid + strong base
- weak acid gives an initial pH reading =3
- the initial rise is steep, as a strong base is being added to a weak acid so neutralisation is rapid
- as the weak acid begins to be neutralised the strong conjugate base sodium ethanoate is formed, creating a buffer that resists change in pH. ethanoic acid is in equilibrium with the ethanoate ion:
CH3COOH + H2O <-> CH3COO- + H3O+ - the continued addition of base to the solution uses up H+ ions, hence the forward reaction is favoured (=gradual change in pH)
- half equivalence is the stage of the titration at which half of the amount of weak acid has been neutralised (pKa=pH)
- there is a sharp rise in pH at equivalence point (pH>7), as a result of salt hydrolysis.
- with no remaining acid to be neutralised, the curve flattens out.
eg CH3COOH + NaOH -> CH3COONa + H2O
weak base + strong acid
HCl + NH3 -> NH4Cl
- weak base gives an initial pH reading =11
- as the weak base begins to be neutralised, the ammonium ion NH4+ (conjugate acid) is created resulting in a buffer that resists change in pH:
NH3 + H2O <-> NH4+ + OH- - at the half equivalence point half of the amount of weak base has been neutralised. At this point, pOH=pKb
- there is a gradual fall in the pH due to the buffering effect as the titration approaches the equivalence point
- the pH falls sharply at the equivalence point (pH<7). the equivalence point is the result of salt hydrolysis.
- with no remaining base to be neutralised, the curve flattens and ends at a low pH due to the presence of excess strong acid.
weak base + weak acid
- weak base gives an initial reading =11.0
- the change in pH throughout the titration is very gradual
- the point of inflection in the pH curve is not as steep as previous pH curves. the point of equivalence is difficult to determine, so this kind of titration has little/no practical use
- with no remaining base to be neutralised, the curve flattens and ends at a pH that indicates the presence of a weak acid
NH3 + CH3COOH -> CH3COONH4
define indicators and give an equation to show them
usually weak acids (represented by HIn) in which the acid (H+) and its conjugate base (In-) are different colours
HIn <-> H+ + In-
colour a -> colour b
when added to acid the equilibrium is pushed to the left and colour A is seen. when added to alkali the OH- combines with H+ so the equilibrium is pushed to the right and colour B is seen.
Ka =
[H+][In-]/[HIn]
when is the midpoint of the colour changed observed?
when [HIn] = [In-]
at this point
[H+] = Ka
pH = pKa
what does the choice of indicator for an acid-base titration depend on?
the relative strengths of the acid and base and therefore the pH of the equivalence point. The midpoint of an indicator’s colour change must correspond to the equivalence point of the titration.
strong acid + strong base
phenol red
pKa = 7.9
pH range: 6.8-8.4
acidic colour: yellow
alkaline colour: red
strong acid + weak base
methyl orange
pKa = 3.7
pH range: 3.1-4.4
acidic colour: red
alkaline colour: yellow
weak acid + strong base
phenolphthalein
pKa = 9.6
pH range: 8.3-10.0
acidic colour: colourless
alkaline colour: pink
indicator with a weak base
BOH (aq) <-> B+ + OH-
colour a -> colour b
