acids and bases Flashcards
what is the bronsted-lowery definition for an acid
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)
what is the bronsted-lowery definition for a base
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)
what are conjugate acid-base pairs
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
how are acids related to bases in an equation
acid = proton + conjugate base
how are bases related to acids in eq
bases + proton = conjugate acid
examples of conjugate pairs
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-
give an example of acid-base eq
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
what is a conjugate pair
two species that are different from each other by a H+ ion
how to work out which is which in the equation - B-L acid or base
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
what does the strength of an acid refer to
pH not conc
what is a strong acid
low pH 1-3
high conc of H+ ions - which fully dissociate into ions
what is a weak acid
high pH 4-7
low conc of H+ ions - does not fully dissociate
what is the definition of pH
pH = -log[H+]
where [H+] is the concentration of hydrogen ions in mol dm–3
how do you find the conc of H+ ions
concentration of H+ of a solution can be calculated if the pH is known by rearranging the above equation to:
[H+] = 10-pH
what does log 10 mean
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.
to how many values is pH usually read
2.d.p
worked example of a pH calc
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
how are strong acids ionised in solution
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]
what are dibasic/diprotic acids
two replaceable protons and will react in a 1:2 ratio with bases
give an example of a dibasic acid
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
how to find log on calculator
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
how can water act as a base and acid
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).
how is equilibrium set up in water
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.
what is Kw
Kw is essentially just an equilibrium constant for the reactions shown.
it can be written as
Kw = [H+][OH-]
what is the value of kw at 298k
1.0x10-14 mol2dm-6
what is the ph of water at 298k
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.
how does kw change with temp
it varies
how does ph of pure water change in eq
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
how to work out the ph or pOH of mixed solution
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+))
what is a weak acid
A weak acid is an acid that partially (or incompletely) dissociates in aqueous solutions and donates a proton/H+
Eg. most organic acids (ethanoic acid), HCN (hydrocyanic acid), H2S (hydrogen sulfide) and H2CO3 (carbonic acid)
what is the eq constant for weak acids
[H+][A-]/[HA] which can be simplified to [H+]^2/[HA]
what does Ka mean
This constant is called the acid dissociation constant, Ka, and has the units mol dm-3
what does Ka indicate
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
how do you calculate the pH of weak acids
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])
calculate the pH of 0.100 mol of propanoic acid (pKa 4.87)
[H+] = sqaure root 10^-4.87 x 0.100
=1.16x10-3
pH =-log10 1.16x10-3 = 2.94
how to find the conc of a weak acid from pH
[H+] = 10-pH
[HA] = [H+]^2/Ka
how to calculate the pH of a mixture of weak acids and strong bases
-calculate moles of HA and OH-
-calculate moles excess HA and OH-
IF EXCESS HA / ACID
-calc [HA] and [A-]
- ICE
-then use E mols and divide by vol
-moles of new [H+] and [A-]
-calc [H+] by Ka x [HA]/[A-]
-log10H+
IF EXCESS OH-
-calc OH- mol over vol
- -log10 of conc
- then 14 - pOH
what is a buffer
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
what types of buffer can you get
acidic with ph less than 7
basic with ph greater than 7
how is an acidic buffer made
-mix of weak acid and one salt ie ethanote or ethanoic acid
-excess weak acid and strong alkali
how is a basic buffer made
-mix of weak alkali and salt
-mix of excess weak alkali and strong acid
how does ethanoic acid & sodium ethanoate act as a buffer
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
what happens when H+ ions are added to an acidic buffer
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
what happens when OH- ions are added to an acidic buffer
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
what happens when H+ ions are added to a basic buffer
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
what happens when OH- ions are added to basic buffer
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
how to calculate the pH of a buffer with a salt and weak acid
-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
how to find pH of buffer with weak acid and strong base
-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
how to find new pH of buffer when acid is added
-find start moles of HX and salt
-find moles of acid added
-HX mol + acid mol
-salt mol - acid mol
-[H+] = ka x salt mol/ new vol / HX mol /V
-then find pH
how to find new pH of buffer when alkali is added
-find moles of HX and other
-find moles of OH- added
-HX + OH- mole
-other mol - OH- mole
-find new conc of HA and A- - but add on extra vol on top of total - ka x other mol/vol / HX mol/v
-then find H+ conc
-then find pH
how do buffers control the pH of the blood
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
what is acidosis
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
what are indicators
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
how does the colour of an indicator change due to the pH
in a low pH = eq pushed left - colour 1 dominates
in a high pH = eq pushed right- colour 2 dominates
what colour does methyl orange give
- pH<3.2 - red
- pH>4.4 - yellow
- in between - orange
what colour does phenolpthalein give
-colourless pH<8.2
-bright pink pH>10.0
what is universal indicator
mixture of 4 indicators
-due to large range of colour it is unsuitable for titrations
what is necessary for indicators to work
colour change must be within the range of the pH change at the endpoint - colour change match the endpoint - vertical
strong acid - strong base indicators
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
weak acid - strong base indicators
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
strong acid - weak base indicators
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
weak acid - weak base indicators
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
what are pH curves and what are they used for
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
what does a pH curve look like
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
how are pH curves similar for acid and alkalis
they mirror one and other
-for acids - start at low pH
-for alkalis - start at high pH
Explain why the volume of sodium hydroxide solution added between each pH measurement is smaller as the end point of the titration is approached
To avoid missing the endpoint
Suggest why the pH probe is washed with distilled water between each of the calibration measurements.
To wash off any residual solution which could interfere with the reading
Explain why [H2O] is not shown in the Kw expression.
H2O is almost constant
Calculate the pH of the solution when half of the acid has reacted.
pH at half neutralisation = pKa
-log10Ka
what type of titration is it if the equivilance point is less than 7
strong acid -weak base
what type of titration is it if the equivilance point is above 7
weak acid - strong base
what type of titration could it be if the equivilance point at ph 7
strong acid - strong base
weak acid -weak base - there should be no curve for this one
if given a graph and asked to find pH or Ka what do u do
-find vol at half equivilance - find ph 7 and the vol of this then divide this by 2
-find ph on graph at this vol
-then for ka - 10-pH
if given a question asking calc the ph after an addition of hydroxide/ water - including KX
-find moles HX
-find new moles HX by subtracting moles of added OH or water to moles
-find KX moles by adding KX moles to OH/ water moles
-then find H+ - Ka x HX/KX
-then -log10H+
what formula can u use if the conc of both substances are the same to find the pH including pKa
pH= pKa + log(salt conc/acid conc)
how do u find new moles of a buffer solution when add an acid
-old moles HX + mol added
-old mol salt - mol added
