Acids, bases & pH ✅

200

3 types of bases

Metal oxides - NaO
Metal hydroxides - NaOH
Metal carbonates - Na₂CO₃

Neutralisation reaction

acid + base → salt + water

Can carry out titrations to find concentrations of acids and bases

Metal hydroxide + Acid
Metal oxide + Acid
Metal carbonate + Acid
Metal + Acid

Metal hydroxide + Acid → Salt + Water
Metal oxide + Acid → Salt + Water
Metal carbonate + Acid → Salt + Carbon dioxide + Water
Metal + Acid → Salt + Hydrogen gas

pH scale

1-6 Acidic
7 Neutral
8-12 Basic

Base vs Alkali

Alkali = A base that is soluble in water. All alkalis are bases
Bases = Insoluble in water

Are most hydroxides soluble or insoluble in water ?

Most hydroxides are SOLUBLE in water so are alkaline

E.g. the alkaline earth metals NaOH (aq)

Common acids, bases and alkalis

Acids: HCl, H₂SO₄, HNO₃
(Insoluble) Bases: Zn(OH)₂
(Soluble) Alkalis: NaOH, KOH, NH₃

What is a Bronsted-Lowry acid?

Proton donor
Releases H⁺ ions when mixed with H₂O

Which arrow sign should be put for weak acids and strong acids?

Weak acids partially dissociate so ⇌
Strong acids fully dissociate so →

How are protons in an aqueous solution represented ?

H₃O⁺
Hydronium ion

Concentration vs strength of acids

Concentration = The relative number of moles of acid in a given volume of water

Strength = How much the H⁺ dissociate

What is a Bronsted-Lowry base ?

Proton acceptor

Can H₂O act as a base?

Yes

What happens in a neutralisation reaction ?

● Acid donates a proton. HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
● Base accepts a proton, which then reacts with OH to form H₂O. KOH(aq) + H⁺(aq)→ K⁺(aq) + H₂O(l)

HNO₃ + KOH →

How are strong and weak acids/ bases different?
And so where does the equilibrium lie?
(check dis)

Strong acids/bases COMPLETELY dissociate/ ionise in aqueous solutions to release H⁺ ions

Weak acids/bases only SLIGHTLY dissociate/ ionise to release H⁺ ions

Name weak acids

Ethanoic acid CH₃COOH
Phosphoric acid H₃PO₄
and other carboxylic acids

Where does the equilibrium lie with weak acids?

Towards the LEFT so the backwards reaction is favoured so not many H⁺ is produced

Name strong acids

HCl hydrochloric acid
H₂SO₄ sulphuric acid
HNO₃ nitric acid

Where does the equilibrium lie with strong acids?

Towards the RIGHT so the forward reaction is favoured so lots of H⁺ is produced

Name Strong Bases

Group 1 metal hydroxides

NaOH, LiOH, KOH, Ca(OH)₂, Ba(OH)₂

Where does the equilibrium lie with strong bases ?

Towards the RIGHT so the forward reaction is favoured so lots of OH⁻ is produced

Name some weak bases

Ammonia (NH₃), Amines

Where does the equilibrium lie with weak bases?

Towards the LEFT so the backwards reaction is favoured so not many OH⁻ ions are produced

When acids and bases react with water, they form a ______________ reaction

Reversible

What ion is created when H+ ions react with water ?

H₃O⁺
Hydronium ion

Give the 2 equations for when water dissociates.

H₂O (l) ⇌ H⁺ (aq) + OH⁻ (aq)

Acid base equilibrium
H₂O (l) + H₂O (l) ⇌ H₃O⁺ (aq) + OH⁻ (aq)
The 1st water acts as a base as it accepts a proton. The 2nd water acts as an acid as it donates a proton.

Conjugate acid

Each acid is linked to a conjugate base on the other side of the equation

Base + H⁺ ⇌ Conjugate acid

Conjugate base

Each base is linked to a conjugate acid on the other side of the equation

Acid ⇌ Conjugate base + H⁺

What is the conjugate acid of OH⁻ ?

OH⁻ + H⁺ ⇌ H₂O

OH⁻ and H₂O are called a conjugate acid-base pair

What is the conjugate base of HF ?

F⁻

HF ⇌ F⁻ + H⁺

HF and F⁻ are called a conjugate acid-base pair

Whenever a substance acts as a Brønsted-Lowry base, we call the product it forms its...

Conjugate acid

Whenever a substance acts as a Brønsted-Lowry acid, we call the product it forms its...

Conjugate base

Conjugate pairs according to Brønsted-Lowry model

Every acid-base reaction has 2 conjugate pairs

NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq)
Conjugate pairs: (NH₃ and NH₄⁺) & (H₂O and OH⁻)

What happens when 2 acids are mixed together?

The stronger acid will act as a Brønsted-Lowry acid, donating a proton
The weaker acid will act as a Brønsted-Lowry base, accepting a proton

Complete the equation for HCl (aq) + CH₃COOH (aq)

HCl (aq) + CH₃COOH (aq) ⇌ Cl⁻ (aq) + CH₃C(OH)₂⁺ (aq)

The HCl is a stronger acid so acts as a proton donor
The ethanoic acid is a weaker acid so acts as a proton acceptor

H₂SO₄ (aq) + CH₃COOH (aq) → HSO₄⁻ (aq) + CH₃C(OH)₂⁺ (aq)
Identify the acid and base according to the Brønsted-Lowry model.

Brønsted-Lowry acid: H₂SO₄
Brønsted-Lowry base: CH₃COOH

Identify the acid and base according to the Brønsted-Lowry model.
NH₃ + NH₃ → NH₄⁺ + NH₂⁻
Give your answers as compound names, rather than chemical formulae.

Brønsted-Lowry acid: ammonia 1
Brønsted-Lowry base: ammonia 2

Ammonium ions can be used as a weak acid in organic reactions. Write an equation for the dissociation of this acid.

NH₄⁺ ⇌ H⁺ + NH₃

Consider the following acid-base reaction:
C₅H₅N + HCl → C₅H₅NH⁺ + Cl⁻
What are the conjugate pairs ?

Conjugate acid = HCl, Conjugate base = Cl⁻

Conjugate acid = C₅H₅NH⁺, Conjugate base = C₅H₅N

What is a monoprotic acid ? Give 3 examples

What is a monoprotic acid ? Give 3 examples

● An acid that donates ONE H⁺ ion for every acid molecule.
● So the concentration of the H⁺ ions is equal to the concentration of the acid HCl, HNO₃, HBr

What do square brackets represent? [ ]

Concentration

What is a diprotic acid ?

● An acid that donates 2 H⁺ ions for every acid molecule.
● So The concentration of H⁺ ions is 2x the concentration of the acid [H⁺]= 2[Acid]
E.g. H₂SO₄

How do diprotic acids dissociate? Use H₂SO₄ as an example.

First: H₂SO₄ → H⁺ + HSO₄⁻
Second: HSO₄⁻ ⇌ H⁺ + SO₄²⁻

Dissociates 1 proton at a time, and the 1st proton dissociates more fully than the 2nd

For 0.1 mol dm⁻³ HCl, what is the pH ?

-log₁₀[0.1] = 1.00

Acidity

The higher the concentration of H⁺ the higher the acidity

Define pH

-log₁₀[H⁺]

Where [H⁺] is the concentration of hydrogen ions in the solution.

How much should pH be rounded ?

ALWAYS to 2 d.p

2.00
11.67

How to find H⁺ ion concentration

[H⁺] = 10⁻ᵖʰ

What is the concentration of H+ ions in a solution with a pH of 3.5?

[H⁺] = 10⁻³.⁵
[H⁺] = 3.16 x 10⁻⁴

What do we -log the concentration with strong monoprotic acids to find the pH ?

Conc of the acid = the conc of H⁺ ions
-why? Because we assume it dissociates FULLY

How do you calculate the concentration of H⁺ ions when you know the pH?

Rearrange the pH equation
[H⁺]= 10⁻ᵖʰ

DONT FORGET THIS IS Inverse log (antilog), PRESS SHIFT THEN THE LOG BUTTON

Rearrange the pH equation
[H⁺]= 10⁻ᵖʰ

DONT FORGET THIS IS Inverse log (antilog), PRESS SHIFT THEN THE LOG BUTTON

When the pH increases by 1...

a tenfold difference in [H⁺]
E.g. pH 2 has 10x the H⁺ concentration of pH 3.

How much more acidic is pH 2 than pH 6 ?

10,000

Why is a logarithmic scale used for pH ?

The concentration of hydrogen ions in aqueous solution covers a very wide range.

What is the log of...
a) 100
b) 1,000,000
c) 10,000
d) 0.1
c) 0.001

a) 2
b) 6
c) 4
d) -1
e) -3

If y = 10ˣ
log(y) = x

Log(x) = 3
What is x equal to?

10³
= 1000

Log(a) = 1.2
What is a equal to?

10^1.2
= 15.8

Log(b) = -3.6
What is b equal to?

10^-36
= 2.51 x 10^-4

If the concentration of HNO₃ (aq) is 0.03 mol dm⁻³, then the concentration of H+ ions in the solution is...

0.03 mol dm⁻³

How to find pH of a solution that has been diluted with water?
Given Volume and Conc of acid.

● Find moles of acid
● Find moles of (strong) acid which is = moles of [H⁺]
● Find new volume (volume of acid + volume of water added)
● The number of [H⁺] stays the same, so find conc by doing moles of [H⁺] / new volume
● Find pH using -log₁₀[H⁺]

What happens to the pH and the volume of the solution if we add a solid base to a solution of acid?

pH increases
Volume stays the same (really small amount of water produced is negligible)

A student has 60mL of a perchloric acid (HClO₄) solution with concentration of 0.36mol dm⁻³. They add 50mL of a sodium hydroxide (NaOH) solution with a concentration of 0.25mol dm⁻³ to the acidic solution. The following reaction occurs:

HClO₄ (aq) + NaOH (aq) → NaClO₄ (aq) + H₂O (l)

What is the new pH of the solution after the reaction?

[H⁺] = [HBr]

Moles of H⁺ ions at start: 0.36 x 0.06 = 0.0216

Moles NaOH = 0.05 x 0.25 = 0.0125

Moles used = 0.0216 - 0.0125 = 0.0091

C = n / v

C = 0.0091 / 0.11 (the new volume)

C = 0.0827

pH = -log₁₀(0.0827)

pH = 1.08

A student has 50mL of a solution of hydrochloric acid (HCl), with a concentration of 5×10⁻³ mol dm⁻³.
How much water do they need to add to the HCl solution to raise its pH to 3?

Moles of H⁺ in initial solution:
0.05 x 0.005 = 0.00025

Calculate the concentration of H⁺ ions needed to get a pH of 3
[H⁺] = 10⁻ᵖʰ
[H⁺] = 10⁻³
[H⁺] = 0.001 mol dm⁻³

V = 0.00025 / 0.001
V = 0.25 dm³ = 250 mL

250 - 50 = 200 mL

Why is the concentration of water said to be a CONSTANT value ?

It only dissociates SLIGHTLY (equilibrium lies to the LEFT) so there are very little OH⁻ and H⁺ ions compared to the number of water molecules

What can be said about strong acids and the concentration of [H⁺] ions?

Strong acids fully dissociates so the concentration of the acid is equal to the concentration of [H⁺] ions

What is Kw? Expression and units

The ionic product of water, eqm lies to left as water hardly dissociates; so much water that it has a 'constant value'
Kw = [H⁺] [OH⁻]
Units mol²dm⁻⁶

How can you manipulate the Kw equation if using pure water?
Why does this work?

Kw = [H⁺]²
Because OH⁻ and H⁺ concentrations are equal
[H⁺] = [OH⁻]

How to find pH of water

[H+] = √Kw
pH = -log [H⁺]

Find the pH of pure water

Kw = [H⁺] [OH⁻]
Kw = [H⁺]²
[H⁺] = √Kw
[H⁺] 1 x 10⁻¹⁴ = 1 x 10⁻⁷ so pH = 7

What is the value of Kw at room temp ?

Always 1.00 X 10⁻¹⁴ mol²dm⁻⁶
at 25 degrees/ 298 K

At a given temp, Kw always has __________________ value in a solution

The SAME

Can the value of Kw change as the TEMP changes?

YES!

Why doesn't Kw include concentration of water?

[H₂O] is constant
[H₂O] is very high compared to [OH⁻] and [H⁺]

How can we predict the change in pH of pure water at different temperatures ?

● Use Le Chatelier's principle to predict the change in pH of pure water at different temperatures
● The dissociation of water is endothermic so increasing the temperature would push the equilibrium to the right
● Giving a bigger concentration of H⁺ ions and a lower pH

At 25°C, the Kw of pure water is 1.008×10⁻¹⁴ mol² dm⁻⁶.
Use the equation Kw = [H⁺]² to calculate the concentration of hydrogen ions in pure water at this temperature and the pH.

Kw = [H⁺]²
[H⁺] = √Kw
[H⁺] = √1.008×10⁻¹⁴
[H⁺] =1.004 x 10⁻⁷
pH = -log₁₀[1.004 x 10⁻⁷] = 7.00

Explain why the value of Kw increases as temperature increases

● Forward reaction is endothermic,
● Position of eqm shifts right to decrease temp so conc of H⁺ and OH⁻ increases
● So pH decreases, Kw increases ( Kw = [H⁺]² )

Give the effect of Kw if there is a temp increase/decrease for an endo/exo reaction

Endothermic reaction
● Temp increase = Kw increase
● Temp decrease = Kw decrease

Exothermic reaction
● Temp increase = Kw decrease
● Temp decrease = Kw increase

Why is the water still neutral ?

Even if pH decreases (due to temp ↑, forwards reaction endo, shifts right)
Equal conc of H⁺ and OH⁻ so is neutral

[H⁺] = [OH⁻]

What effect does changing concentration on a reactant or product have on Kw ?

No effect

● E.g. adding H⁺ ions to pure water will not affect Kw as position of eqm shifts to left decrease H⁺ ion conc, and Kw is the product pf H⁺ and OH⁻.
● H⁺ increases and OH⁻ decreases

If [H⁺] = 0.4 mol dm⁻³ and
Kw = 1.01×10⁻¹⁴ mol² dm⁻⁶,
what is [OH⁻] ?

b) Is this solution acidic, alkaline or neutral?

Kw = [H⁺] [OH⁻]
1.01×10⁻¹⁴ = [0.4] [OH⁻]
Rearrange to get OH⁻ = 2.53 x 10⁻¹⁴ mol dm⁻³

b) Solution is acidic as greater conc of H⁺ than OH⁻

What is the pH change if 20 cm³ of 0.1 mol dm⁻³ HCl has 30 cm³ water added ?

Strong acid fully ionises so [H⁺] = 0.1 mol dm⁻³
Starting pH: pH = -log₁₀[0.1] = 1.00
Adding water does not change moles HCl
n = c x v
n = 0.1 x 0.02 = 0.002

c = n / v
c = 0.002 / 0.05 = 0.04

Diluted pH = -log₁₀[0.03] = 1.40
pH change = 1.40 - 1.00 = 0.40

What do you have to do when calculating pH of a diprotic acid?

Multiply the concentration of the acid by 2

What happens when a small amount of sodium hydroxide is added to pure water?

● The concentration of hydroxide ions increases
● The concentration of hydrogen ions decreases
● The equilibrium constant of water dissociation stays constant

How can we work out the pH of a strong base ?

You use Kw!

Because Kw = 1.00 X 10⁻¹⁴ and Kw = [H⁺][OH⁻]
So, 1.00 X 10⁻¹⁴ = [H⁺][OH⁻]
You can use the pH given to work out the concentration of H⁺ then plug it into the equation,
[H⁺] = 10⁻ᵖʰ
Rearrange it to work out the concentration of OH⁻ ions

CONCENTRATION OF OH⁻ IONS = CONCENTRATION OF BASE (if monobasic - if dibasic multiply by 2)

Give the factors affecting the pH of a solution

● Amount of dissociation
● Solubility
● Conc of H⁺ ions
● Temperature

Group 2 hydroxides

Ba(OH)₂

Solubility of group 2 hydroxides

Solubility increases down the group, Ba(OH)₂ fully dissociates and is a strong base. Mg(OH)₂ is hardly soluble but the bit that does FULLY dissociates so it's still a strong base.

Be
Mg
Ca
Sr
Ba

Why does the conc of OH⁻ ions = the conc of the strong base ?

Because strong bases dissociate to produce 1 MOLE of OH⁻ ions for EVERY MOLE OF BASE

Which equilibrium constants always have the same units ?

Kw
Ka

What is Ka? What is it used for ? What does [HA] mean ?

● The acid dissociation constant for WEAK acids (same as Kc but for weak acids)
● is used to work out the pH of weak acids and bases
● Temperature dependent
● [HA]= conc of the weak acid

Kₐ = [H⁺] [A⁻] / [HA]

Write the expression for the acid dissociation constant (Ka) of methanoic acid.

Ka = [H⁺] [HCOO⁻] / [HCOOH]

What does a greater Ka indicate ?

● Higher Ka = Stronger the acid
● Because the acid dissociates into ions to a greater extent
● Eqm further to the right

How to calculate the pH of a weak acid

[H⁺] = √(Ka x [HA])
Find pH using -log₁₀[H⁺]

We assume [HA] is equal to initial [HA]
[HA] is the acid molecules