Acids, Bases and Salts (Salt Preparation)

Solubility Rules

• So, Po, Ammo, Ni are soluble

• For Chloride → lead (II) and silver

• For Sulphate → lead (II), calcium, barium

• Carbonates, Oxides, Hydroxides → INSOLUBLE

Reactivity Series

• K

• Na

• Ca

• Mg

• Al

• Zn

• Fe

• Pb

• H

• Cu

• Ag

• Au

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Formation of Salt

• Acid + Metal → Salt + Hydrogen Gas

• Acid + Carbonate → Salt + Water + Carbon Dioxide Gas

• Acid + Base → Salt + Water

• Ammonium Salt + Base → Salt + Water + Ammonia Gas

• 2 soluble (aqueous) reagents → Insoluble salt

Experimental Procedure for Reaction of acid + excess metal/insoluble base and carbonate (8)

↣ This procedure is used when preparing a soluble salt with insoluble reactants.

↣ It is not recommended to use excess metal due to the extreme reactivity or lack thereof of some metals.

↣ Saturation can be tested by dipping a cold, dry glass rod into the filtrate. If the filtrate is saturated, crystals will form on the glass rod.

1. Using the measuring cylinder, measure out 10 cm³ of (acid)  into a 100 ml beaker. 

2. Add excess (metal) to the acid in the beaker. Stir constantly using a glass rod until no more (metal) dissolves in the acid.

3. Filter the mixture into a 100ml beaker using a filter funnel and filter paper to separate the excess (metal) from the aqueous solution of the (salt). The salt is collected as the filtrate and the (metal) is collected as the residue.

4. With a Bunsen burner and evaporating dish, heat the filtrate to evaporate the water and concentrate the salt solution to obtain a saturated solution of (salt).

5. On cooling, (salt) crystals will form.

6. Filter the mixture with a filter funnel and filter paper to collect the (salt) crystals as residue.

7. Wash the crystals with 1–2 drops of cold distilled water to remove impurities.

8. Dry the crystals by pressing them between layers of filter paper.

Titration Experimental procedure (13)

↣ This procedure is used when preparing a soluble salt with soluble reactants.

1. Fill a burette with (base). Record the initial burette reading as V₁.

2. Using a pipette, transfer 25.0cm³ of (acid) into a 250cm³ conical flask.

3. Add 1 to 2 drops of methyl orange indicator to the (acid). The solution appears red.

4. Add (base) slowly from the burette, dropwise near the end point, until the end point colour change from red to orange is observed.

5. Record the final burette reading as V₂.

6. Record the volume of (base) required for complete neutralisation as V₂ - V₁.

7. Using a pipette, transfer 25.0cm³ of (acid) into a 250cm³ beaker.

8. Add  V₂ - V₁ cm³ of (base) from the burette to the (acid).

9. Heat the mixture with a Bunsen burner to evaporate off water to prepare a saturated solution.

10. Allow the saturated solution to cool down slowly. Crystals will start forming. Use an ice bath if necessary.

11. Filter off the crystals using a filter funnel and filter paper.

12. Wash the crystals with 1–2 drops of cold distilled water to remove impurities.

13. Dry the crystals by pressing them between layers of filter paper.

Why PbO(s) cannot be added directly to HI(aq) to prepare PbI2(s)?

Equation: PbO + 2HI → PbI2 + H2O

Reaction of PbO with HI forms insoluble PBI2 which coats the surface of PbO solid. This protective layer prevents PbO from further reaction with HI, leading to low yield of PbI2.

Experimental Procedure for Ionic Precipitation (with dissolving) (9)

↣ This procedure is used when preparing an insoluble salt.
↣ The reactants must be two suitable ionic compounds. It is compulsory to convert any insoluble starting material to a soluble salt by dissolving.

• The starting material, (insoluble salt) is insoluble. 

• Add 2g of (insoluble salt) to 10cm³ of (suitable acid) in a 250ml beaker. 

• Stir with a glass rod until no more (insoluble salt) dissolves.

• Filter the mixture with a filter funnel and filter paper to obtain (soluble salt) as the filtrate.

1. Using a 50 ml measuring cylinder, transfer 2cm³ of (soluble salt) into a clean 100 ml beaker.

2. Add (salt 2) solution in excess using a measuring cylinder and stir until no more precipitate forms.

3. Filter the mixture with a filter funnel and filter paper to obtain (salt) as the residue.

4. Wash the precipitate with 1–2 drops of cold distilled water to remove impurities.

5. Dry the precipitate by pressing them between layers of filter paper.

Salt preparation Mind Map

Which starting material should you NOT choose if you have a choice for reaction with acid and why?

If you have a choice of starting material, don’t choose the metal reaction (either lack of reactivity or violently react)

Solubility Rules for Salts

TYPE	SOLUBLE	INSOLUBLE
Salts with Group 1 Cations Li+, Na+, K+, Rb+, Cs+, Fr+ So, Po	All	-
Ammonium Salts NH4+ Ammo	All	-
Nitrate Salts NO3- Ni	All	-
Chloride Salts	Almost All	Lead (II)) Silver PbCl2 AgCl
Sulphate Salts	Almost All	Calcium Lead  Barium > everything under calcium
Carbonate Salts	Group 1 > Ammonium carbonate (NH4)2CO3	Almost All
Oxides/Hydroxides	Group 1	Almost All

Why do water of crystallisation happen?

• Regular ionic lattice → form regularly shaped crystals

• Many combine with water molecules to form crystals

• Water molecules are chemically bonded to the ions within crystals

• Water trapped inside due to extra space → space between anion and cation, water occupies it

Water of Crystallisation

water bonded chemically within the crystal

Anhydrous Salts

Salts that do not contain water of crystallisation

Hydrated Salts

Salts that contain water of crystallisation

Hydrated sodium carbonate = Na₂CO₃ . 10H₂O

Salts

• ionic compounds – NaCl, CaSO₄

• Contain a metal cation or ammonium ion and a nonmetal anion

• Usu. formed by replacement of hydrogen ions of an acid by a metallic ion or an ammonium ion

• Come from Acid-Base reaction

What are some neutral oxides?

• Carbon monoxide (CO)

• Water

• Nitric oxide (NO)

• Nitrous oxide (N₂O)

Neutral Oxides

Do not react with acids or bases

Amphoteric oxide reacting with acid: Al₂O₃ + 6HCl →

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

What does Al₂O₃ not react with?

water

Which oxides are amphoteric?

Aluminium, Lead (II), Zinc

• > Al₂O_3’

• > PbO (only 2+, 4+ cannot)

• > ZnO

IMPORTANT!! Aluminium Oxide Reaction with base to form sodium aluminate (amphoteric)

Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ [sodium aluminate]

Amphoteric Oxide

Reacts with both acids and bases

Silicon Oxide reaction (acidic) with NaOH

• SiO₂ is not soluble

• SiO₂ reacts with concentrated hot NaOH or molten NaOH at high temperature to form sodium silicate

• SiO₂ (s) + 2NaOH (l) → Na₂SiO₃ (l) + H₂O (g)
  molten sodium silicate

acidic reacting with base

• acidic oxide + base → salt + water

• P₄O₁₀ (s) + 12NaOH → 4Na₃PO₄(aq) + 6H₂O (l)
  sodium phosphate

• SO₂ (g) + 2NaOH (aq) → Na₂SO₃ (aq) + H₂O (l)
  sodium sulfite

• SO₃ (g) + 2NaOH (aq) → Na₂SO₄ (aq) + H₂O (l)
  sodium sulfate

axidic oxide reacting with water

• When acidic nature is mixed with water, it turns into an acid without any by-products.

• These reactions are non-redox. Oxidation states do not change.

• Forms pH < 7 

• SO₃ + H₂O → H₂SO₄
  sulfuric acid

• SO₂ + H₂O → H₂SO₃
  sulphurous acid

Which oxides are acidic in nature?

Group 14, Group 15, Group 16

basic oxide reacting with acid

acid + metal base → salt + water
Na₂O + 2HCl → 2NaCl + H₂O (sodium chloride)
K₂O → Potassium chloride
MgO → Magnesium chloride
CaO → Calcium chloride

basic oxide reacting with water

Na₂O + H₂O → 2NaOH (sodium hydroxide)
K₂O + H₂O → 2KOH (potassium hydroxide)
Na₂O and K₂O are soluble in water → give alkaline solution with pH > 7

Basic oxides are usually ___ in water

insoluble

Which oxides are basic in nature?

Group 1, Group 2

Metal oxides are ____

either basic or amphoteric

Non-metal oxides are ____

acidic

How does nature of oxide change across Period?

Nature of oxides change from basic (metal) to amphoteric to acidic (non-metal) across Period

basic (metal) → amphoteric → acidic (non-metal)

What is the haber process?

• Industrial method to manufacturer ammonia

• Reversible reaction

• N₂(g) + 3H₂(g) → ← 2NH₃(g)

• High pressure of 250 atm

• Moderate temperature of 450 degrees Celsius

• Iron catalyst

• Molar ratio of N₂ : H₂ = 1 : 3

What is synthesised in the haber process?

Ammonia

Neutralisation Reaction

• Base + acid → salt + water

KOH(aq) + HCl(aq) → KCl(aq) + H₂O(l)

Why can [acid] conduct electricity in water but not in organic solvent?: Answering Format (6)

• In water, [acid] ionises to form free mobile ions, H⁺ and [B]; [formula].

• Since H⁺ > [B] and pH < 7, the blue litmus paper turns red. The free mobile ions, H⁺ and [B] are able to act as mobile charge carriers to conduct electricity.

• In organic solvent, [acid] exists as a simple molecular compound.

• It does not ionise to form H⁺ and [B] but remains as a molecule, [acid].

• Hence, there is no change in the colour of litmus.

• There is no electrical conductivity as there are no free mobile ions to act as mobile charge carriers.

Properties of Alkalis (5)

• Bitter taste and feel soapy

• Turn damp red litmus paper blue

• Turn Universal Indicator blue or violet

• pH greater than 7

• Conduct electricity; can act as charge carriers as they exist as free mobile ions in aqueous solutions

  NaOH → Na⁺ + OH^-

Properties of Acids (6)

• Sour taste

• pH less than 7

• Turns damp blue litmus paper red

• Turn Universal Indicator orange or red

• Must be dissolved in water before they can act as acids

• Conduct electricity; can act as charge carriers as they exist as free mobile ions in aqueous solutions
  H₂SO₄ → 2H⁺ + SO₄^2-

Give me some indicators and their color changes.

— is a measure of the concentration of H+ in a solution. A strong acid of the same concentration will — to give a higher H+. A weak acid of the same concentration will — only to give a lower H+. Hence, the pH of the strong acid will be — than the weak acid of the same concentration.

pH is a measure of the concentration of H+ in a solution. A strong acid of the same concentration will dissociate fully to give a higher H+. A weak acid of the same concentration will dissociate partially only to give a lower H+. Hence, the pH of the strong acid will be lower than the weak acid of the same concentration.

Indicators

Dyes/Mixture of dyes which change colour when added to acids or alkalis

Universal indicator

mixture of dyes that can be used to estimate the pH of a solution by tallying the colour of the solution with the Universal colour chart

What must be kept constant when using pH values to compare strength of 2 acids?

1. concentration

2. temperature

Alkaline on pH scale

Alkaline
pH > 7
H⁺ < OH^-

Acidic on pH scale

Acidic
pH < 7
H⁺ > OH^-

Neutral on pH scale:

Neutral
pH = 7
H⁺ = OH^-

pH scale

pH scale – used to indicate whether a solution is acidic, neutral or alkaline; 25 degrees Celsius

Examples of finding pH value of acids

• The pH value of 0.001 mol dm^-3 hydrochloric acid is ________
  > Is it strong?
  > If yes, all of it becomes H⁺
  > Input into equation
  > Thus, the answer will be 3

• The pH value of 0.3 mol dm^-3 H₂SO₄ is ________
  > Is it strong?
  > If yes, all of it becomes H⁺
  > Since there are 2 H, multiply with 2 to become 0.6
  > Input into equation
  > Thus, the answer will be 0.222

How to find pH value of an acid?

1. Determine whether if it is strong or not

2. If it is strong, all parts of the acid will get turned into H⁺

3. If there are more than 1 H⁺ in equation, multiply with suitable number

4. Input into the equation: pH = –log₁₀[H⁺]

5. Find answer

Larger pH → ____ [H⁺]

larger pH → smaller [H⁺]

What does H⁺ mean?

[H⁺] means concentration in mol dm^-3

[H⁺] = 10^–pH

pH = –log₁₀[H⁺]

pH of solution

• pH of solution – negative logarithm to base 10 of the concentration of hydrogen ions in solution in mol dm^-3

• pH = –log₁₀[H⁺]

• [H⁺] = 10^–pH

• Strong acid + Low concentration → everything dissociated

• Strong acid + High concentration → everything dissociated

• Weak acid + Low concentration → left

• Weak acid + High concentration → left

Strength is ________ on concentration!

independent

Strength of an acid measures:

Strength of an acid – measure of the extent of dissociation of an acid in solution

A weak base:

	Weak Base
Definition	Dissociates partially in solution to give hydroxide ions, OH-
Degree of dissociation	< 100%
Type of arrow in dissociation equation	Double arrow
Example	NH3 + H2O → ← NH4+ + OH-

A weak acid:

	Weak Acid
Definition	Dissociates partially in solution to give protons, H+
Degree of dissociation	< 100%
Type of arrow in dissociation equation	Double arrow
E.g. of organic acid	HCOOH→ ← H+ + HCOO-
E.g of monobasic acid	HCN → ← H+ + CN-

A strong base:

	Strong Base
Definition	Dissociates fully in solution to give hydroxide ions, OH-
Degree of dissociation	Approximately 100%
Type of arrow in dissociation equation	→ single arrow
E.g. of monoacidic base	KOH → K+ + OH-
E.g of diacidic base	Ca(OH)2 → Ca2+ +2OH-

A strong acid:

	Strong Acid
Definition	Dissociates fully in solution to give protons, H+
Degree of dissociation	Approximately 100%
Type of arrow in dissociation equation	→ single arrow
E.g. of monobasic acid	HNO3 → H+ + NO3-
E.g of dibasic acid	H2SO4 → 2H+ + SO42-

How are the strength of Acids and Bases classified?

• Classified based on their extent of dissociation to give ions in solution

• Must be 100% to be considered strong

Common acids (6)

Common alkalis (4)

Alkali

•  bases that are soluble in water and produces OH^-

• [NaOH, KOH, Ca(OH)₂]

Base

• a species that can accept a proton, H⁺ and donates OH^-

• metal oxides, metal hydroxides [Na₂O, ZnO, Al(OH)₃]

Acid

a species that can donate a proton, H⁺

Bronsted-Lowry reaction

• involves the transfer of a proton from the acid (proton donor) to the base (proton acceptor) in water