Acids, Bases and Salt Preparations (2d)

Solubility: Sodium, Potassium and Ammonium Salts

All soluble

Solubility: Nitrates

All soluble

Solubility: Chlorides

most, except Silver and Lead(II)

Solubility: Sulfates

most, except Barium, Calcium and Lead(II)

Solubility: Carbonates

most insoluble, except Sodium, Potassium and Ammonium

Solubility: Hydroxides

most insoluble, except Sodium, Potassium and Ammonium

(Calcium Hydroxide is slightly soluble)

Proton Transfer: Acids

proton donors

• ionize in solutions → produce protons, H⁺ ions

• H⁺ ions make aqueous solution acidic

• ex: HCl

HCl (aq)   →    H⁺ (aq)    +    Cl^– (aq)

Proton Transfer: Bases

proton acceptors

• ionize in solutions → produce OH^- ions that accept protons

• OH^- ions make aqueous solution alkaline

• ex: NaOH

NaOH (s)   →    Na⁺ (aq)    +    OH^– (aq)

Metal + Acid Reactions

• only metals above hydrogen react with dilute acids

• more reactive metal = more vigorous reaction

• K and Na: dangerous, explosive reaction

• reaction forms salt and hydrogen gas

Metal + Acid general equation

metal + acid → salt + hydrogen gas

Acid + Base Reactions

• produces salt and water

• identity of salt depends on acid used and base’s positive ions

2HCl + CuO ⟶ CuCl₂ + H₂O

H₂SO₄ + 2NaOH ⟶ Na₂SO₄ + 2H₂O

HNO₃ + KOH ⟶ KNO₃ + H₂O

Acids + Metal Carbonates

• forms metal salt, carbon dioxide and water

• distinguishable because of effervescence caused by carbon dioxide gas

2HCl + Na₂CO₃ ⟶ 2NaCl + H₂O + CO₂

H₂SO₄ + CaCO₃⟶ CaSO₄ + H₂O + CO₂

Bases vs Alkali

• water soluble base = alkali

• alkaline pH: >7

• turn red litmus paper blue

• bases: usually oxides, hydroxides or carbonates of metals

• OH^- ions makes aqueous solutions alkaline

• ammonia is an exceptional base because it produces OH^- ions

Soluble Salt Preparation

• reaction of acid and insoluble base

• insoluble reactant added in excess to ensure all the acid has reacted

• if not, unreacted acid becomes dangerously concentrated during evaporation and crystallisation

• excess reactant is filtered to ensure only salt and water remain

• since all acid has reacted, excess solid base removed → solution left can only be salt and water

• water evaporated until small crystals appear

  • usually when half the water is left

  • evaporating slowly over many days = larger crystals

• if carbonate used, CO₂ released into atmosphere

• ex: preparing copper(II) sulfate from copper(II) oxide and dilute sulfuric acid

CuO (s) + H₂SO₄ (aq) ⟶ CuSO₄ (aq) + H₂O (l)

Soluble Salt Preparation with a Metal

acid can also be reacted with a metal to produce a salt, where:

• metal is above hydrogen in the reactivity series

• no too reactive that a dangerous reaction takes place

Soluble Salt Preparation using Titration

• use pipette to measure alkali into conical flask

• add drops of indicator (phenolphthalein or methyl orange)

• add acid into burette, note starting volume

• add acid very slowly from burette until indicator changes to correct colour

• note and record final volume, calculate volume of acid added (starting - final acid volume)

• add this volume of acid into same volume of alkali without indicator

• heat to partially evaporate, leaving saturated solution

• leave to crystallise decant excess solution, allow crystals to dry

Insoluble Salt Preparation

• precipitation reaction

• solid salt is precipitate → solid salt formed must be insoluble

• pattern:

  • soluble salt 1 + soluble salt 2 ⟶  insoluble salt + soluble salt 3

  • AB + CD ⟶ AD + CB

• measure out fixed volume of one solution and add second salt solution until in slight excess → ensures max. amount of precipitate obtained

• precipitate recovered by filtration, washed with distilled water to remove residue (recovering solid) contaminating reactants → left to dry

• good way to prepare silver and lead(II) salts which are often insoluble → starting material = usually nitrates of silver or lead(II) since all nitrates are soluble

Practical (Preparing Copper(II) Sulfate): Aim

to prepare a pure, dry sample of hydrated copper(II) sulfate crystals

Practical (Preparing Copper(II) Sulfate): Materials

• 1.0 mol / dm³ dilute sulfuric acid

• copper(II) oxide

• spatula & glass rod

• measuring cylinder & 100 cm³ beaker

• Bunsen burner

• tripod, gauze & heatproof mat

• filter funnel & paper, conical flask

• evaporating basin and dish.

Practical (Preparing Copper(II) Sulfate): Diagram

Practical (Preparing Copper(II) Sulfate): Method

• add 50 cc dilute acid into beaker

• warm gently with a Bunsen burner

• add copper(II) oxide slowly into hot dilute acid

• stir until base is in excess (stops dissolving, suspension forms in acid)

• filter mixture into evaporating basin → remove excess bases

• gently het solution in water bath or with electric heater → evaporate water and saturate solution

• check solution saturation by dipping cold glass rod and seeing if crystals form on the end

• leave filtrate in a warm place to dry and crystallise

• decant excess solution and allow to dry

Practical (Preparing Copper(II) Sulfate): Excess Base Added

• added in excess to use up all the acid

  • acid would become dangerously concentrated during evaporation and crystallisation

Practical (Preparing Copper(II) Sulfate): Results

hydrated copper(II) sulfate crystals should be bright blue, regularly shaped

Practical (Preparing Lead(II) Sulfate): Aim

• prepare a dry sample of lead(II) sulfate

• solid salt obtained = precipitate → solid salt formed must be insoluble, reactants must be soluble

Practical (Preparing Lead(II) Sulfate): Method

• add 25 cc of 0.5 mol dm³ lead(II) nitrate solution to a small beaker

• add 25 cc of 0.5 mol dm³ potassium nitrate to the beaker and mix with stirring rod

• filter to remove precipitate from mixture

• wash filtrate with distilled water to remove traces of other solutions

• leave in oven to dry

Soluble salt 1 = lead(II) nitrate       

Soluble salt 2 = potassium sulfate

lead(II) nitrate + potassium sulfate → lead(II) sulfate + potassium nitrate

Pb(NO₃)₂  (aq) + K₂SO₄ (aq) → PbSO₄ (s) + 2KNO₃ (aq)