Acids & Bases Flashcards
Reacting acids and bases
1) Acid + Metal → Salt + Hydrogen
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
2) Acid + Metal Oxide → Salt + Water
Li2O(s) + 2HCl(aq) → 2LiCl(aq) + H2O
3) Acid + Metal Hydroxide → Salt + Water
Ca(OH)2(aq) + 2HCl(aq) → CaCl2(aq) +2H2O
4) Acid + Metal Carbonate → Salt + CO2 + Water
CaCO3(g) + 2CH3COOH(aq) → (CH3COO)2Ca(aq) + CO2 + H2O
5) Acid + Ammonia → Ammonium salt
NH3(g) + HCl (aq) → NH4Cl
Strong/ weak acid examples
Strong:
HCl = Cl-
H2SO4= SO4 2-
HNO3= NO3-
H3PO4 = PO4 3-
Weak:
CH3COOH= COO-
Any other acid
Strong/ weak base examples
Strong:
Metal oxides (MgO, Na2O)
Metal hydroxide ( NaOH, Ba(OH)2 )
Weak:
NH3
Carbonate (Na2CO3) = CO3 2-
What is an acid?
An ACID is a proton donor
HCl(aq) → H+(aq) + Cl-(aq)
What is a base?
A BASE is a proton acceptor
-OH(aq) + H+(aq) → H2O
*H2O can act as an acid and a base
Dilutions
pH expressions
pH = -Log10[H+]
[H+] = 10^-pH
- The smaller the pH, the greater the concentration of H+ ions.
- A difference of 1 on the pH scale means a 10x difference in [H+]
Strong acids and calculations
A STRONG acid FULLY DISSOCIATES
HCl(aq) → H+(aq) + Cl-(aq)
Calculations:
-Log[HA] monoprotic
-Log2[HA] diprotic (H2SO4)
-Log3[HA] triprotic (H3PO4)
Explain why chloroethanoic acid is a weaker acid than ethanoic acid?
Chloroethanoic acid is weaker as it is more polar.
Ethanoic acid is stronger, less polar due to the positive inductive effect
Weak acids and calculations
A WEAK acid only PARTIALLY DISSOCIATES
CH3COOH(aq) ⇌ CH3COO-(aq) + H+(aq)
Ka = [H+][A-]
[HA]
[H+] = √Ka[HA]
-Log[H+]
How to differentiate if strong or weak acid using a single pH mesurement?
strong acid pH=1
weak acid pH= greater than 1
pKa values
pKa = -Log10Ka
Ka = 10^-pKa
Pure water and calculations
H2O(l) ⇌ H+(aq) + -OH(aq) ∆H = +57.3kJmol-1
very slightly dissociates
always neutral because [H+] = [OH-] at all times
Kw = [H+][-OH]
[H+] = √Kw
-log[H+]
-water conc is constant as it very slightly dissociates
-as temp ↑, equilibrium shifts to the right in endo direction to oppose ↑ in temp
-[H+] increases
Strong alkali and calculations
Strong bases fully dissociate
MOH → M+ + OH-
Kw = [H+][-OH]
[H+] = Kw / [nMOH]
-Log[H+]
Buffer Solutions
A BUFFER is a solution which can resist changes in pH when a small amount of acid or base is added
An ACIDIC BUFFER is made of a weak acid and a soluble salt of that acid. It maintains a pH below 7.
A BASIC BUFFER is made of a weak BASE and a soluble salt of that BASE. It maintains a pH above 7.
Calculating the pH of buffer
Ka = [H+][A-]
[HA]
[H+] = Ka[HA] / [A-] (A- = conc of salt)
-log[H+]
Half neutralisation
Equal number of moles of acid and salt
Ka= [H+]
pKa = pH
Calculating the pH change of a buffer when a small amount of acid is added
When acid is added, the equilibrium shifts to the left to oppose increase in [H+]
H+ + CH3COO- → CH3COOH
Calculating the pH change of a buffer when a small amount of alkali is added
How does a buffer resist pH changes when a small amount of alkali is added?
-OH + H+ → H2O
CH3COOH → H+ + CH3COO-
equilibrium shifts right
2 ways to conduct pH titrations
1) normal titration
2) Using a pH probe:
- Add 25cm^3acid into 250ml conical flask
- Measure the pH of the acid solution and record.
- Add 1cm3 of the base solution
- Stir the mixture
- Measure the pH and record.
- Repeat the process until the base is in excess.
- Add base in smaller increments near the end point
It is important to calibrate the pH meter as after storage it may not give an accurate reading. To calibrate it you place the pH meter in a solution of known pH and then adjust the meter accordingly.
pH titration curves
Acid base titration calculations
at the equivalence point:
nNaOH= CxV = 25/1000 x 0.1
= 0.0025
nNaOH= nAcid
nAcid= 0.0025
[Acid] = 0.0025/ vol
Equivalence point definition
when exactly enough acid has been added to neutralise the base
End Point definition
the exact volume of acid or base which needs to be added to cause an indicator to change colour
Indicators
A SUITABLE INDICATOR CHANGES COLOUR SOMEWHERE ON THE VERTICAL SECTION OF A pH TITRATION CURVE
The Half-neutralisation point
The point at which enough base has been added to neutralise exactly half of the acid
Half neutralisation= Ka = [H+]
working out pH when no numbers = -log(Ka)
Half neutralisation = half equivalence
