I completed the A-Level in one year teaching most of the course to myself, earning an A overall. This is a summary of the textbook chapter with all details you will need for the exam, I found this most useful as it streamlines the information and keeps everything I needed in one place. Comparing ma...
Acids, Bases and Buffers
21.1 Defining an acid: 1. Calculate the concentration of ions at equilibrium
2. The [H+] = [A-]and so you can rearrange to have
Brønsted-Lowry Acid: a substance that donates protons Ka = [H+]2 / [HA]
(H+ ion) pKa: a measurement that shows how strong a weak acid
Brønsted-Lowry Base: a substance that can accept a is, the smaller the value of pKa the stronger the acid.
proton (H+ ion) pKa = - log10Ka
Water as an acid and base, HCl can donate a proton to
water and so acts as a base. 21.4 Acid-base titrations:
HCl + H2O ——> H2NO3+ + HSO4-
Water can donate an electron to ammonia so acts as an Titrations are used to find the concentration of a solution
acid. by gradually adding a second solution of a known
H2O + NH3 ——-> OH- + NH4+ molarity and volume.
This shows water to be slightly ionised in a reversible 1. Acid of a known concentration is added via a burette
reaction to a measured amount of basic solution
H2O(l) <——-> H+(aq) + OH-(aq), writing the 2. Add the acid slowly, swirling the mixture so it mixes
expression for Kc is different because the [H2O] is 3. Repeat until the indicator shows there’s been a colour
constant, changing the expression. change indicating neutralisation
4. Repeat until you have concordant results, then
Ionic product of water (Kw): the constant of water calculate the average amount of acid used for the
concentration, 1.0 x 10-14 mol2dm-6 titration for your calculation
Kw = [H+(aq)][OH-(aq)] Or you can measure the pH with a pH meter, calibrated by
In pure water there is equal concentrations of OH- and placing the probe in a buffer solution of a known pH
H+ions
Titration Curves-
21.2 The pH scale: This shape is
when base is
The acidity of a solution depends on the concentration added from the
of H+ ions, this is then measured on the pH scale. burette into an
• the smaller the pH value, the greater the [H+] acidic solution in
• A difference of 1 on the pH scale gives a difference the conical flask.
tenfold on the [H+] The equivalence
• A pH < 7 is acidic , pH > 7 is alkaline point will
determine the type
pH = -log10[H+(aq)] of indicator you
use as it must
Calculating [H+] from pH, you can use antilog to have a colour
find the [H+], you can substitute this value into the change for that
equation for Kw to find the [OH-]. pH.
Strong acids: acids that dissociate completely Equivalence point: when sufficient amount of base/acid has
been added to the solution to just neutralise the solution.
21.3 Weak acids and bases:
Calculating concentration-
1. Write the balanced symbol equation for the reaction
Strong bases: bases that completely dissociate in
2. From the results of your titration find the equivalence
aqueous solutions
point (how much acid / base needs to be added)
Weak acids: acids that only partially dissociate in
3. Calculate the number of moles of the substance you
aqueous solutions, forming an equilibrium
know the concentration and volume of.
HA <———> H+ + A-
4. Using the balanced equation calculate the number of
Acid dissociation constant (Ka): the equilibrium
moles of the other substance
constant for a weak acid. The larger Ka the more the
5. Rearrange the formula to calculate the concentration.
equilibrium is to the right-acid is more dissociated and
is stronger.
To find the pH of a weak acid you cannot assume it 21.5 Choice of indicators:
dissociates fully and so must calculate [H+] using Ka
End point: the volume of base/acid added when the
indicator just changes colour
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