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Chemistry - A-level > acids and bases > Flashcards

acids and bases Flashcards

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1
Q

what is the bronsted-lowery definition for an acid

A

an acid is a substance which can behave as a proton donor - H+
hydrogen chloride (HCl) is a Brønsted acid as it can lose a proton to form a hydrogen (H+) and chloride (Cl-) ion
HCl (aq) → H+ (aq) + Cl- (aq)

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2
Q

what is the bronsted-lowery definition for a base

A

a base is a substance which can behave as a proton acceptor - has a lone pair

hydroxide (OH-) ion is a Brønsted base as it can accept a proton to form water
OH- (aq) + H+ (aq) → H2O (l)

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3
Q

what are conjugate acid-base pairs

A

most acid-base reactions are reversible
the acid which gives up a proton can accept a proton and thus behave as a base

the species formed when a base accepts a proton can give up a proton and behave ad an acid

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4
Q

how are acids related to bases in eq

A

acid = proton + conjugate base

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5
Q

how are bases related to acids in eq

A

bases + proton = conjugate acid

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6
Q

examples of conjugate pairs

A

Identify the acid-base conjugate pairs in the following reactions:

HCO3- (aq) + H2O (l) ⇌CO32- (aq) + H3O+ (aq)
HCO3- (aq) + H3O+(aq) ⇌ CO2 (g) + H2O (l) + H2O (l)
H2SO4 (aq) + HNO3 (aq) ⇌ HSO4- (aq) + NO2+ (aq) + H2O (l)
HSO4- (aq) + OH- (aq) ⇌ SO42- (aq) + H2O (l)
Answers

The pairs in the order acid/base are:

HCO3- and CO32- ; H3O+ and H2O
H3O+ and H2O ; (CO2 + H2O) and HCO3-
H2SO4 and HSO4- ; (NO2+ + H2O) and HNO3
HSO4- and SO42- ; H2O and OH-

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7
Q

give an example of acid-base eq

A

For example, ethanoic acid (CH3COOH) is a weak acid that partially dissociates in solution
When equilibrium is established there are CH3COOH, H2O, CH3COO- and H3O+ ions present in the solution
The species that can donate a proton are acids and the species that can accept a proton are bases
CH3COOH (aq) + H2O (l) ⇌ CH3COO- (aq) + H3O+ (aq)

                                acid                   base            conjugate base      conjugate acid  The reactant CH3COOH is linked to the product CH3COO- by the transfer of a proton from the acid (CH3COOH) to the base (CH3COO-) Similarly, the H2O molecule is linked to H3O+ ion by the transfer of a proton These pairs are therefore called conjugate acid-base pairs
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8
Q

what is a conjugate pair

A

two species that are different from each other by a H+ ion

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9
Q

how to work out which is which in the equation

A

the B-L base is the reactant that has gained a hydrogen on the product side - thus the conjugate acid - now has a positive charge

the B-L acid is the one that loses the hydrogen on the product side - thus the conjugate base

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10
Q

what does the strength of an acid refer to

A

pH not conc

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11
Q

what is a strong acid

A

low pH 1-3
high conc of H+ ions - which fully dissociate into ions

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12
Q

what is a weak acid

A

high pH 4-7
low conc of H+ ions - does not fully dissociate

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13
Q

what is the definition of pH

A

pH = -log[H+]

where [H+] is the concentration of hydrogen ions in mol dm–3

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14
Q

how do you find the conc of H+ ions

A

concentration of H+ of a solution can be calculated if the pH is known by rearranging the above equation to:
[H+] = 10-pH

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15
Q

what does log 10 mean

A

pH scale is a logarithmic scale with base 10
This means that each value is 10 times the value below it. For example, pH 5 is 10 times more acidic than pH 6.

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16
Q

to how many values is pH usually read

A

2.d.p

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17
Q

worked example of a pH calc

A

Question 1: Find the pH when the hydrogen concentration is 1.60 x 10-4 mol dm-3
Answer 1:

The pH of the solution is:

pH = -log[H+]

  = -log 1.6 x 10-4

  = 3.80

Question 2: Find the hydrogen concentration when the pH is 3.10

Answer 2:
The hydrogen concentration can be calculated by rearranging the equation for pH

pH = -log[H+]

[H+] = 10-pH

   = 10-3.10

   = 7.94 x 10-4 mol dm-3
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18
Q

how are strong acids ionised in solution

A

Strong acids are completely ionised in solution
HA (aq) → H+ (aq) + A- (aq)

Therefore, the concentration of hydrogen ions, H+, is equal to the concentration of acid, HA
The number of hydrogen ions formed from the ionisation of water is very small relative to the [H+] due to ionisation of the strong acid and can therefore be neglected
The total [H+] is therefore the same as the [HA]

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19
Q

what are dibasic/diprotic acids

A

two replaceable protons and will react in a 1:2 ratio with bases

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20
Q

give an example of a dibasic acid

A

Sulfuric acid is an example
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O (l)

You might think that being a strong acid it is fully ionised so the concentration of the hydrogen is double the concentration of the acid
This would mean that 0.1 mol dm-3 would be 0.2 mol dm-3 in [H+] and have a pH of 0.69
However, measurements of the pH of 0.1 mol dm-3 sulfuric acid show that it is actually about pH 0.98, which indicates it is not fully ionised
The ionisation of sulfuric acid occurs in two steps
H2SO4 → HSO4- + H+

HSO4- ⇌ SO42- + H+

Although the first step is thought to be fully ionised, the second step is suppressed by the abundance of hydrogen ions from the first step creating an equilibrium
The result is that the hydrogen ion concentration is less than double the acid concentration

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21
Q

how to find log on calculator

A

Make sure you know how to use the antilog (base 10) feature on your calculator. On most calculators it is the 10x button, but on other models it could be LOG-1, ALOG or even a two-button sequence such as INV + LOG

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22
Q

how can water act as a base and acid

A

Water molecules can function as both acids and bases. One water molecule (acting as a base) can accept a hydrogen ion from a second one (acting as an acid).

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23
Q

how is equilibrium set up in water

A

However, the hydroxonium ion is a very strong acid, and the hydroxide ion is a very strong base. As fast as they are formed, they react to produce water again.

The net effect is that an equilibrium is set up.
2H2O = H3O+ + OH-

At any one time, there are incredibly small numbers of hydroxonium ions and hydroxide ions present. Further down this page, we shall calculate the concentration of hydroxonium ions present in pure water. It turns out to be 1.00 x 10-7 mol dm-3 at room temperature.

You may well find this equilibrium written in a simplified form:
H2O = H+ + OH-

This is OK provided you remember that H+(aq) actually refers to a hydroxonium ion.

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24
Q

what is Kw

A

Kw is essentially just an equilibrium constant for the reactions shown.

it can be written as
Kw = [H+][OH-]

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25
Q

what is the value of kw at 298k

A

1.0x10-14 mol2dm-6

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26
Q

what is the ph of water at 298k

A

7

That means that you can replace the [OH-] term in the Kw expression by another [H+].

[H+]2 = 1.00 x 10-14

Taking the square root of each side gives:

[H+] = 1.00 x 10-7 mol dm-3

Converting that into pH:

pH = - log10 [H+]

pH = 7

That’s where the familiar value of 7 comes from.

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27
Q

how does kw change with temp

A

it varies

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28
Q

how does ph of pure water change in eq

A

The formation of hydrogen ions (hydroxonium ions) and hydroxide ions from water is an endothermic process. Using the simpler version of the equilibrium:

The forward reaction absorbs heat.

According to Le Chatelier’s Principle, if you make a change to the conditions of a reaction in dynamic equilibrium, the position of equilibrium moves to counter the change you have made.

According to Le Chatelier, if you increase the temperature of the water, the equilibrium will move to lower the temperature again. It will do that by absorbing the extra heat.

That means that the forward reaction will be favoured, and more hydrogen ions and hydroxide ions will be formed. The effect of that is to increase the value of Kw as temperature increases.

  • Kw increases so H+ also increase when temp increases so shift right opppose change in conc / increase in temp
29
Q

how to work out the ph or pOH of mixed solution

A

calculate the moles of H+
calculate moles of OH-
calculate moles excess of H+ and OH-
calculate excess [H+] or [OH-] depending on which is larger
calculate ph = -log10 (H+)
or 14 - (-log10(h+))

30
Q

what is a weak acid

A

A weak acid is an acid that partially (or incompletely) dissociates in aqueous solutions
Eg. most organic acids (ethanoic acid), HCN (hydrocyanic acid), H2S (hydrogen sulfide) and H2CO3 (carbonic acid)

31
Q

what is the eq constant for weak acids

A

[H+][A-]/[HA] which can be simplified to [H+]^2/[HA]

32
Q

what does Ka mean

A

This constant is called the acid dissociation constant, Ka, and has the units mol dm-3

33
Q

what does Ka indicate

A

The value of Ka indicates the extent of dissociation
The higher the value of Ka the more dissociated the acid and the stronger it is
The lower the value of Ka the weaker the acid

34
Q

how do you calculate the pH of weak acids

A

This means you can simplify and re-arrange the expression to
Ka x [HA] = [H+]2

[H+]2 = Ka x [HA]

Taking the square roots of each side
[H+] = √(Ka x [HA])

Then take the negative logs
pH = -log[H+] = -log√(Ka x [HA])

35
Q

calculate the pH of 0.100 mol of propanoic acid (pKa 4.87)

A

[H+] = sqaure root 10^-4.87 x 0.100
=1.16x10-3
pH =-log10 1.16x10-3 = 2.94

36
Q

how to find the conc of a weak acid from pH

A

[H+] = 10-pH
[HA] = [H+]^2/Ka

37
Q

how to calculate the pH of a mixture of weak acids and strong bases

A

-calculate moles of HA and OH-
-calculate moles excess HA and OH-
IF EXCESS HA / ACID
-calc [HA] and [A-]
-use Ka to find [H+] by doing ICE
-then use E mols and divide by vol
-calc pH by 10^-Kax[HA]/[A-]
-log10

IF EXCESS OH-
-calc OH- mol over vol
- -log10 of conc
- then 14 - pOH

38
Q

what is a buffer

A

solution which resists changes in pH when small amounts of acids or alkalis are added
A buffer solution is used to keep the pH almost constant

39
Q

what types of buffer can you get

A

acidic with ph less than 7
basic with ph greater than 7

40
Q

how is an acidic buffer made

A

-mix of weak acid and one salt ie ethanote or ethanoic acid
-excess weak acid and strong alkali

41
Q

how is a basic buffer made

A

-mix of weak alkali and salt
-mix of excess weak alkali and strong acid

42
Q

how does ethanoic acid & sodium ethanoate act as a buffer

A

A common buffer solution is an aqueous mixture of ethanoic acid and sodium ethanoate
Ethanoic acid is a weak acid and partially ionises in solution to form a relatively low concentration of ethanoate ions

Sodium ethanoate is a salt which fully ionises in solution

There are reserve supplies of the acid (CH3COOH) and its conjugate base (CH3COO-)
The buffer solution contains relatively high concentrations of CH3COOH (due to partial ionisation of ethanoic acid) and CH3COO- (due to full ionisation of sodium ethanoate)
In the buffer solution, the ethanoic acid is in equilibrium with hydrogen and ethanoate ions

43
Q

what happens when H+ ions are added to an acidic buffer

A

The equilibrium position shifts to the left as H+ ions react with CH3COO- ions to form more CH3COOH until equilibrium is re-established
As there is a large reserve supply of CH3COO- the concentration of CH3COO- in solution doesn’t change much as it reacts with the added H+ ions
As there is a large reserve supply of CH3COOH the concentration of CH3COOH in solution doesn’t change much as CH3COOH is formed from the reaction of CH3COO- with H+
As a result, the pH remains reasonably constant

44
Q

what happens when OH- ions are added to an acidic buffer

A

the OH- reacts with H+ to form water
OH- (aq) + H+ (aq) → H2O (l)

The H+ concentration decreases
The equilibrium position shifts to the right and more CH3COOH molecules ionise to form more H+ and CH3COO- until equilibrium is re-established
CH3COOH (aq) → H+ (aq) + CH3COO- (aq)

As there is a large reserve supply of CH3COOH the concentration of CH3COOH in solution doesn’t change much when CH3COOH dissociates to form more H+ ions
As there is a large reserve supply of CH3COO- the concentration of CH3COO- in solution doesn’t change much
As a result, the pH remains reasonably constant

45
Q

what happens when H+ ions are added to a basic buffer

A

eq shift right
excess H+ removed when reacts with OH- so some of NH3 reacts to replace OH-
some of NH3 falls
but remains roughly constant

46
Q

what happens when OH- ions are added to basic buffer

A

eq shift left
added OH- removed by creating with NH4+ to form NH3
NH3 rises slightly and NH4+ falls but both are bigger than OH- so remain roughly the same

47
Q

how to calculate the pH of a buffer with a salt and weak acid

A

-find moles of both HA and salt
-find [A-] by doing mol x vol
-then find H+ conc - ka x [HA]/[A-]
-then find pH -log10

48
Q

how to find pH of buffer with weak acid and strong base

A

-find moles of HA and OH-
-then find before and after moles by doing ICE via the one in XS
-then find left over [HA] and [A-] - and divide by total volume
-the find H+
-then find pH

49
Q

how to find new pH of buffer when acid is added

A

-find mol of H+ added
-the find before and after reaction moles by misusing H+ from the A- and using the HA and A- from previous calc
-find new conc of HA and A- - but add on extra vol on top of total
-then sub into ka to find H+
-then work out pH

50
Q

how to find new pH of buffer when alkali is added

A

-find moles of OH- added
-then find before and after moles by minus OH- from HA and using the before values from previous question
-find new conc of HA and A- - but add on extra vol on top of total
-then find H+ conc
-then find pH

51
Q

how do buffers control the pH of the blood

A

In humans, HCO3- ions act as a buffer to keep the blood pH between 7.35 and 7.45
Body cells produce CO2 during aerobic respiration
This CO2 will combine with water in blood to form a solution containing H+ ions
CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3- (aq)

This equilibrium between CO2 and HCO3- is extremely important
If the concentration of H+ ions is not regulated, the blood pH would drop and cause ‘acidosis’

The equilibrium position shifts to the left until equilibrium is restored
CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3- (aq)

This reduces the concentration of H+ and keeps the pH of the blood constant
If there is a decrease in H+ ions
The equilibrium position shifts to the right until equilibrium is restored
CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3- (aq)

This increases the concentration of H+ and keeps the pH of the blood constant

52
Q

what is acidosis

A

Acidosis refers to a condition in which there is too much acid in the body fluids such as blood
This could cause body malfunctioning and eventually lead to coma
If there is an increase in H+ ions

53
Q

what are indicators

A

weak acids which have a different colour to their conjugate base
HIn (aq) + H2O (l) ⇌ H3O+ (aq) + In- (aq)

colour 1 colour 2

HIn and its conjugate base In- are different colours
it is an example of equilibrium

54
Q

how does the colour of an indicator change due to the pH

A

in a low pH = eq pushed left - colour 1 dominates
in a high pH = eq pushed right- colour 2 dominates

55
Q

what colour does methyl orange give

A
  • pH<3.2 - red
  • pH>4.4 - yellow
  • in between - orange
56
Q

what colour does phenolpthalein give

A

-colourless pH<8.2
-bright pink pH>10.0

57
Q

what is universal indicator

A

mixture of 4 indicators
-due to large range of colour it is unsuitable for titrations

58
Q

what is necessary for indicators to work

A

colour change must be within the range of the pH change at the endpoint - colour change match the endpoint - vertical

59
Q

strong acid - strong base indicators

A

In strong acid - strong base titrations, the pH changes from 4 to 10 at the end-point so a suitable indicator must change colour within this range
Methyl red and phenolphthalein are suitable indicators for these titrations
Methyl orange is not ideal but it shows a significant enough colour change at the end point so is widely used

60
Q

weak acid - strong base indicators

A

In weak acid - strong base titrations, the pH changes from 7 to 10 at the end-point so a suitable indicator must change colour within this range
Phenolphthalein is the only suitable indicator for weak acid - strong base titrations that is widely available

61
Q

strong acid - weak base indicators

A

In strong acid - weak base titrations, the pH changes from 4 to 7 at the end-point so a suitable indicator must change colour within this range
Methyl red is the most suitable indicator for these titrations
However methyl orange is often used since it shows a significant enough colour change at the end-point and is more widely available than methyl red

62
Q

weak acid - weak base indicators

A

In weak acid -weak alkali titrations, there is no sudden pH change at the end-point and thus there are no suitable indicators for these titrations
The end-points of these titrations cannot be easily determined

63
Q

what are pH curves and what are they used for

A

A pH curve is a graph showing how the pH of a solution changes as the acid (or base) is added
The result is characteristically shaped graph which can yield useful information about how the particular acid and alkali react together with stoichiometric information

64
Q

what does a pH curve look like

A

All pH curves show an s-shape curve and the midpoint of the inflection is called the equivalence or stochiometric point
From the curves, you can
Determine the pH of the acid by looking where the curve starts on the y-axis
Find the pH at the equivalence point
Find volume of base at the equivalence point
Obtain the range of pH at the vertical section of the curve

65
Q

how are pH curves similar for acid and alkalis

A

they mirror one and other
-for acids - start at low pH
-for alkalis - start at high pH

66
Q

Explain why the volume of sodium hydroxide solution added between each pH measurement is smaller as the end point of the titration is approached

A

To avoid missing the endpoint

67
Q

Suggest why the pH probe is washed with distilled water between each of the calibration measurements.

A

To wash off any residual solution which could interfere with the reading

68
Q

Explain why [H2O] is not shown in the Kw expression.

A

H2O is almost constant

69
Q

Calculate the pH of the solution when half of the acid has reacted.

A

pH at half neutralisation = pKa
-log10Ka

Chemistry - A-level (8 decks)

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