Proton Transfer Reactions (DP IB Chemistry)

False.

A proton in aqueous solution can be represented as both H⁺ (aq) and H₃O⁺ (aq).

A conjugate acid-base pair is two species that are different from each other by a H⁺ ion.

In the following reaction, the acid and conjugate base pair is H₂CO₃ (aq) and HCO₃^- .

H₂CO₃ (aq) + OH^- (aq) ⇌ HCO₃^- (aq) + H₂O (l)

In the following reaction, the base and conjugate acid pair is CN^- (aq) and HCN (aq) .

H₂PO₄⁻ (aq) + CN^- (aq) ⇌  HCN (aq)  + HPO₄²⁻ (aq)

In the following reaction, the acid and conjugate base pair is HSO₄^- (aq) and SO₄^- (aq).

HSO₄^- (aq) + OH^- (aq) ⇌ SO₄^- (aq) + H₂O (l)

In the following reaction, the acid and conjugate base pair is HCO₃^- (aq) and CO₃^2- (aq).

HCO₃^- (aq) + H₂O (l)  ⇌ CO₃^2- (aq) + H₃O⁺ (aq)

False.

In the following reaction the acid on the left hand side has a conjugate base on the right hand side.

H₂SO₄ + HNO₃ ⇋  H₂NO₃⁺ + HSO₄^-

The conjugate acid in the reaction is CH₃COOH₂⁺.

CH₃COOH + HCl  ⇌ CH₃COOH₂⁺ + Cl^-

The conjugate base of NH₄⁺ is NH₃.

True.

The conjugate base of a strong acid is a weak base.

The relationship between the strength of an acid and its conjugate base is:

The stronger the acid, the weaker its conjugate base.

Amphiprotic is a term describing species that can act both as proton donors and acceptors.

True.

Water is an example of an amphiprotic species.

An amphoteric compound has both basic and acidic character.

True.

All amphiprotic substances are amphoteric.

False.

Not all amphoteric substances are amphiprotic.

For example, aluminium oxide is amphoteric but not amphiprotic.

False.

Amphiprotic substances can both accept or donate protons.

In the following equation is Al₂O₃ is acting as an acid.

Al₂O₃ (s) + 2NaOH (aq) + 3H₂O (l) → 2NaAl(OH)₄ (aq)     

True.

HCO₃⁻ is amphiprotic.

pH is the negative logarithm of the hydrogen ion concentration in mol dm^-3

pH = -log[H⁺]

The pH scale measures how acidic or basic a substance is.

False.

The pH scale is logarithmic.

The pH range for acidic solutions is 0-6 .

The pH range for alkaline solutions is 8-14.

The most accurate way of measuring pH is using a pH meter.

The pH of a neutral solution at 298K is 7.

As [H⁺] increases, pH decreases.

True.

A solution with pH 4 is 10 times more acidic than a solution with pH 5.

The pH of a solution that has an H⁺ concentration of 0.001 mol dm^-3 at 298 Kis pH 3.

In alkaline solutions OH^- ions are in a greater concentration than H⁺ ions.

False.

Using universal indicator is a less accurate method is to measure the pH of a substance than a using a pH meter.

The pH of HCl with a concentration of 0.43 mol dm^-3 is 0.37.

pH = -log[0.43]

pH = 0.37

An equation for the equilibrium established in water is:

H₂O (l) ⇌ H⁺ (aq) + OH^- (aq).

The equation for the ion product of water is:

K_w = [H⁺][OH^-]

False.

The value of K_w changes with temperature.

As temperature increases, K_w increases.

In pure water at 298K, [H⁺] = [OH^-].

True.

In acidic solutions, [H⁺] > [OH^-].

True.

In alkaline solutions, [H⁺] < [OH^-].

As temperature increases, the pH of water decreases.

The ionisation of water is endothermic.

The pH of a solution of sodium hydroxide, NaOH (aq) of concentration 1.0 × 10⁻⁴ mol dm⁻³ is 10.

• K_w = [H⁺] [OH^-]

• Rearranging gives [H⁺]  = K_w ÷ [OH^-]

• [H⁺] = (1.0 × 10⁻¹⁴) ÷ (1.0 × 10⁻⁴) = 1.0 × 10⁻¹⁰ mol dm⁻³

• pH = – log₁₀[1.0 × 10⁻¹⁰ (aq)]

• So the pH = 10

[H₂O (l)] not included in the expression for K_w a s the concentration of water molecules remains constant.

A strong acid is an acid that dissociates almost completely in aqueous solutions.

A weak acid is an acid that partially (or incompletely) dissociates in aqueous solutions.

True.

Weak acids are proton donors and their solutions are poor conductors.

A strong base is a base that dissociates almost completely in aqueous solutions.

True.

HCl, HBr, HI, HNO₃ and H₂SO₄ are all strong acids.

Examples of strong bases include Group 1 hydroxides (e.g. NaOH, KOH)

HI is a stronger acid than HBr.

This is because the HI bond is a longer bond than HBr and is weaker. Therefore H⁺ ions are more easily released.

Strong acids conduct electricity better than weak acids.

A stronger acid has a higher concentration of H⁺ so it conducts electricity better.

CH₃CH₂COOH has a higher pH than HNO₃.

Strong acids react more vigorously because the concentration of H⁺ is greater in strong acids compared to weak acids.

The equation for the ionisation of ethanoic acid is:

CH₃COOH (aq)CH₃COO^- (aq) + H⁺ (aq)

A strong acid has a higher concentration of H⁺ ions so it conducts electricity better than a weak acid.

Weak acids are less acidic than strong acids as they contain fewer H⁺ in solution. This is because they only partially ionise.

Weak bases include:

• NH₃

• Amines

• Some transition metal hydroxides

False.

In general, strong acids produce weak conjugate bases and weak acids produce strong conjugate bases.

The ionic equation to show how a base accepts hydrogen ions is:

OH^– (aq) +  H⁺ (aq)  ⇌  H₂O (l)

HNO₃ dissociates to form H⁺ and NO₃^- ions.

The NO₃^- ion is a weak conjugate base.

True.

A solution of 10 mol dm^-3 HCN is a concentrated weak acid.

A neutralisation reaction is one in which an acid and a base/alkali react together to form water and a salt

A sulfate salt is produced when sulfuric acid is reacted with an alkali.

A nitrate salt is produced when nitric acid is reacted with an alkali.

A chloride salt is produced when hydrochloric acid is reacted with an alkali.

Ethanoic acid is used to make ethanoate salts.

The general equation for the reaction between an acid and a metal hydroxide is:

acid + metal hydroxide → salt + water

The equation for the reaction between hydrochloric acid  + potassium hydrogencarbonate is:

HCl (aq) + KHCO₃ (s) → KCl (aq) + H₂O (l) + CO₂ (g)

The equation for the reaction between nitric acid + copper carbonate is:

2HNO₃ (aq) + CuCO₃ (s) → Cu(NO₃)₂ (aq) + H₂O (l) + CO₂ (g)

The equation for the reaction between sulfuric acid + barium hydroxide is:

H₂SO₄ (aq) + Ba(OH)₂ (aq) → BaSO₄ (s) + 2H₂O (l)

Ammonium sulfate is formed when sulfuric acid reacts with ammonia.

Bases that could be used to form copper(II) chloride are copper(II) oxide, copper(II) hydroxide, copper(II) carbonate.

During a titration, a pH meter can be used and a pH curve plotted.

X is the equivalence or stoichiometric point.

The pH at the equivalence point of a titration between HCl and NaOH is 7.

From a pH curve, you can determine:

• The initial pH of the acid

• The pH at the equivalence point

• The volume of base at the equivalence point

• The range of pH at the vertical section of the curve

The pH at the start of the titration is pH 3.

The pH at equivalence point is pH 8.

The volume of NaOH (aq) that was added to reach the equivalence point is 30 cm³.