Comprehensive Guide to Acids, Bases, and Salts

Definition and General Properties of Acids

Definition of an Acid: An acid is a hydrogen ion donor ( $H^+$ ion donor). It is a substance that releases hydrogen ions when dissolved in water.
- Example of Ionization: $HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$
Common Examples of Acids:
- Hydrochloric acid: $HCl$
- Sulphuric acid: $H_2SO_4$
- Nitric acid: $HNO_3$
- Carbonic acid: $H_2CO_3$
- Ethanoic acid: $CH_3COOH$
- Tartaric acid
- Citric acid
- Phosphoric acid: $H_3PO_4$
Physical Properties of Acids:
- They possess a $pH$ value less than $7$ .
- They turn blue litmus paper red.
- They have a sour taste.
- Certain acids are corrosive.

Chemical Properties and Reactions of Acids

Reaction with Metals: Acids react with metals to produce a salt and hydrogen gas.
- General Equation: $\text{ACID} + \text{METAL} \rightarrow \text{SALT} + \text{HYDROGEN}$
- Reaction Requirements: Metals that react with dilute acids must be more reactive than hydrogen to displace it. These metals include Magnesium, Aluminium, Zinc, Iron, and Tin (collectively referred to by the acronym MAZIT).
- Reaction Examples:
  - $Mg(s) + H_2SO_4(aq) \rightarrow MgSO_4(aq) + H_2(g)$
  - $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$
- Inert Metals: Copper and Silver do not react with dilute acids because they are lower than hydrogen in the reactivity series.
- Test for Hydrogen Gas: Insert a burning splint into the gas. If hydrogen is present, a "pop" sound is produced.
Neutralization of Bases: Acids react with bases (metal oxides and metal hydroxides) to form salt and water.
- General Equation: $\text{ACID} + \text{BASE} \rightarrow \text{SALT} + \text{WATER}$
- Reaction Examples:
  - $H_2SO_4(aq) + CuO(s) \rightarrow CuSO_4(aq) + H_2O(l)$
  - $HNO_3(aq) + NaOH(aq) \rightarrow NaNO_3(aq) + H_2O(l)$
Reaction with Carbonates: Acids react with carbonates to produce salt, water, and carbon dioxide.
- General Equation: $\text{ACID} + \text{CARBONATE} \rightarrow \text{SALT} + \text{WATER} + \text{CARBON DIOXIDE}$
- Reaction Examples:
  - $H_2SO_4(aq) + Na_2CO_3(aq) \rightarrow Na_2SO_4(aq) + H_2O(l) + CO_2(g)$
  - $2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$
- Diagnostic Use: This reaction is a chemical test to confirm the presence of carbonate ions in a salt solution; the presence is indicated by the bubbling of carbon dioxide gas.
- Test for Carbon Dioxide Gas: Bubble the gas through lime water. If carbon dioxide is present, the lime water turns milky.

Acid Strength and Concentration

Strong Acids: These are substances that ionize completely when dissolved in water, producing a high concentration of hydrogen ions.
- Example: $HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$
- List of Strong Acids: Sulphuric acid, Hydrochloric acid, Nitric acid, Phosphoric acid.
Weak Acids: These are substances that ionize only partially when dissolved in water, resulting in a lower concentration of hydrogen ions.
- Example: $CH_3COOH(aq) \rightleftharpoons CH_3COO^-(aq) + H^+(aq)$
- List of Weak Acids: Tartaric acid, Carbonic acid, Citric acid, Sulphurous acid, Ethanoic acid.
Concentration vs. Strength:
- Concentration: Refers to the number of moles of acid present in one litre ( $1\,dm^3$ ) of solution. For example, $0.5\,mol/dm^3$ HCl is more concentrated than $0.1\,mol/dm^3$ HCl.
- Strength: Refers specifically to the extent of ionization in water. Hydrochloric acid is always strong regardless of whether it is concentrated or dilute, and ethanoic acid is always weak regardless of concentration.

Bases and Alkalis

Definition of a Base: A base is a hydrogen ion acceptor. Bases include metal oxides and metal hydroxides.
Types of Bases:
1. Soluble Bases (Alkalis): These substances ionize to produce hydroxide ions ( $OH^-$ ) when dissolved in water.
2. Insoluble Bases: These do not contain hydroxide ions in water and do not cause color changes in pH indicators. These are typically metal oxides.
Strength of an Alkali: This refers to the degree to which an alkali ionizes in water.
- Strong Alkali: Completely ionizes to form a high concentration of hydroxide ions ( $OH^-$ ).
- Weak Alkali: Partially ionizes to form a low concentration of hydroxide ions.
Properties of Bases and Alkalis:
- Alkalis turn red litmus paper blue.
- They have a $pH$ greater than $7$ .
- Alkalis have a soapy feel because they react with skin oils to form soap.
- Bases neutralize acids to produce salt and water.
- Alkalis react with ammonium salts ( $NH_4^+$ ) under gentle warming to form salt, water, and ammonia gas.
  - Example: $NH_4Cl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l) + NH_3(aq)$
- Test for Ammonia Gas: Use damp red litmus paper. Because ammonia gas is alkaline, it will turn the paper blue. When dissolved in water, it produces aqueous ammonia ( $NH_4OH$ ).

The pH Scale and Indicators

The pH Scale:
- Ranges from $1$ to $14$ .
- $1$ to $6$ : Increasing acidity (lower numbers are stronger acids).
- $7$ : Neutral.
- $8$ : to $14$ : Increasing alkalinity (higher numbers are stronger alkalis).
Indicators and Color Changes:
- Litmus: Red in acid, Red/Blue in neutral, Blue in alkali.
- Methyl Orange: Red in acid, Orange in neutral, Yellow in alkali.
- Phenolphthalein: Colourless in acid, Colourless in neutral, Pink in alkali.
Universal Indicator: A mixture of dyes that provides different colors across the pH range ( $1$ - $14$ ):
- $pH\,1-3$ : Red (Strong acid).
- $pH\,4-6$ : Orange-Yellow (Weak acid).
- $pH\,7$ : Green (Neutral).
- $pH\,8-11$ : Blue (Weak alkali).
- $pH\,12-14$ : Violet-Purple (Strong alkali).

Daily Life Applications of Acid-Base Reactions

Treatment of Acidic Soils: Soil pH can be raised by adding calcium oxide (quicklime), calcium hydroxide (slaked lime), or calcium carbonate (limestone).
Treatment of Indigestion: Excess Hydrochloric acid in the stomach is neutralized by antacids such as sodium hydrogen carbonate (bicarbonate of soda), Milk of Magnesia, or ENO.
Brushing Teeth: Toothpaste contains a base to neutralize acids produced by microorganisms feeding on food in the mouth.
Insect Stings:
- Bee stings and ant bites inject methanoic acid. These are treated with calamine lotion (zinc carbonate) or bicarbonate of soda.
- Wasp stings inject an alkali and are treated with mild acids like ethanoic acid (vinegar).
Industrial Waste: Acidic liquid waste from factories is neutralized with quicklime, slaked lime, or limestone before discharge.

Classification of Oxides

Definition: An oxide is a compound consisting of oxygen and one other element.
Acidic Oxides: Non-metal oxides that neutralize bases. If soluble, they form acidic solutions (e.g., acid rain).
- Examples: $SO_2, SO_3, CO_2, NO_2, SiO_2$ .
- $H_2O(l) + CO_2(g) \rightarrow H_2CO_3(aq)$ (Carbonic acid).
- $H_2O(l) + SO_2(g) \rightarrow H_2SO_3(aq)$ (Sulphurous acid).
Basic Oxides: Metal oxides that neutralize acids. If soluble, they form alkaline solutions.
- Examples: $Na_2O, CuO, MgO, CaO$ .
- $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$ .
- $Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$ .
- Insoluble metal oxides like $CuO$ can dissolve in hot acids: $CuO(s) + 2HCl(aq) \rightarrow CuCl_2(aq) + H_2O(l)$ .
Amphoteric Oxides: Metal oxides that exhibit both acidic and basic properties and can react with both acids and bases.
- Examples: Aluminium oxide ( $Al_2O_3$ ), Zinc oxide ( $ZnO$ ), and Lead (II) oxide ( $PbO$ ).
Neutral Oxides: Non-metal oxides that show neither acidic nor basic properties.
- Examples: Water ( $H_2O$ ), carbon monoxide ( $CO$ ), dinitrogen oxide ( $N_2O$ ), and nitrogen monoxide ( $NO$ ).

Composition and Solubility of Salts

Salt Components:
- Cation (metal part or ammonium): Derived from the base.
- Anion (non-metal part): Derived from the acid.
- Example: $ZnO(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2O(l)$ . Zinc comes from the base; sulphate comes from the acid.
Solubility Rules:
- Soluble Salts: All nitrates, all Group 1 metal salts, and all ammonium salts. Most sulphates (except $CaSO_4$ , $PbSO_4$ , $BaSO_4$ ). Most chlorides (except $AgCl$ , $PbCl_2$ ).
- Insoluble Salts: Most carbonates (except Group 1 and ammonium carbonates). Specific sulphates and chlorides mentioned above.

Preparation of Soluble Salts

Methods for Soluble Salts: Generally involves crystallization.
Method 1: Acid + Metal (MAZIT metals):
- General: $\text{Acid}(aq) + \text{Metal}(s) \rightarrow \text{Salt}(aq) + \text{Hydrogen}(g)$ .
- Example: Preparing Zinc Sulphate:
  1. Gently warm $50\,cm^3$ of dilute sulphuric acid.
  2. Add zinc powder and stir until no more dissolves (ensures all acid is used).
  3. Filter to remove excess zinc; collect filtrate (zinc sulphate solution).
  4. Heat filtrate in an evaporating dish until saturated.
  5. Cool to allow crystal formation.
  6. Squeeze crystals between filter paper.
Method 2: Acid + Insoluble Base (Neutralization):
- Used for metals like Copper that do not react with dilute acids.
- Example: Preparing Copper (II) sulphate using copper (II) oxide and sulphuric acid. The procedure follows the same warming, stirring, filtering, and crystallization steps as the metal method.
Method 3: Acid + Insoluble Metal Carbonate:
- Procedure is identical to the insoluble base method but uses a carbonate, releasing carbon dioxide.
Method 4: Acid + Alkali (Titration):
- Used for reactive Group 1 metals. Uses precise volumes so neither reactant is in excess.
- Example: Preparing Sodium Sulphate:
  1. Pipette $25.0\,cm^3$ sodium hydroxide into a conical flask.
  2. Add methyl orange indicator.
  3. Titrate with sulphuric acid from a burette until the color shifts from yellow to orange.
  4. Calculate volume of acid used.
  5. Repeat the process without the indicator using the exact volumes.
  6. Evaporate and crystallize.

Preparation of Insoluble Salts

Precipitation Method:
- General Equation: $AB(aq) + CD(aq) \rightarrow AD(s) + CB(aq)$ .
- Procedure: Mix two aqueous salt solutions. Filter the resulting precipitate, wash with distilled water to remove impurities, and dry.
- Example: Preparing Copper (II) Carbonate:
  - $CuSO_4(aq) + Na_2CO_3(aq) \rightarrow CuCO_3(s) + Na_2SO_4(aq)$ .
- Other Examples: Preparation of Lead (II) sulphate, Silver chloride, Barium sulphate, and Lead (II) iodide.

Water of Crystallisation and Hydrates

Hydrates: Salts containing water of crystallisation.
- Copper (II) sulphate: $CuSO_4 \cdot 5H_2O$ (Blue).
- Iron (II) sulphate: $FeSO_4 \cdot 7H_2O$ (Green).
- Sodium carbonate: $Na_2CO_3 \cdot 10H_2O$ (White).
- Calcium sulphate: $CaSO_4 \cdot 2H_2O$ (White).
- Magnesium sulphate: $MgSO_4 \cdot 7H_2O$ (White).
- Cobalt (II) chloride: $CoCl_2 \cdot 6H_2O$ (Pink).
Anhydrous Salts: Formed when hydrates are heated gently and lose their water of crystallisation.

Identification of Anions (Negative Ions)

Negative Ion	Test	Result
Carbonate ( $CO_3^{2-}$ )	Add dilute $HCl$ to solid	Effervescence; $CO_2$ produced (turns limewater milky)
Chloride ( $Cl^-$ )	Acidify with Nitric acid, add aqueous Silver Nitrate	White precipitate ( $AgCl$ ); soluble in ammonia
Bromide ( $Br^-$ )	Acidify with Nitric acid, add aqueous Silver Nitrate	Cream precipitate ( $AgBr$ ); slightly soluble in ammonia
Iodide ( $I^-$ )	Acidify with Nitric acid, add Silver Nitrate OR Lead Nitrate	Yellow precipitate; insoluble in ammonia
Sulphate ( $SO_4^{2-}$ )	Acidify with $HCl$ , add Barium Chloride OR acidify with Nitric acid, add Barium Nitrate	White precipitate ( $BaSO_4$ )
Nitrate ( $NO_3^-$ )	Add $NaOH$ , add Aluminium foil (or Devarda's alloy) and warm	Ammonia gas produced (turns damp red litmus blue)

Identification of Cations (Positive Ions)

Positive Ion	Effect of Sodium Hydroxide ( $NaOH$ )	Effect of Ammonia Solution ( $NH_3$ )
Ammonium ( $NH_4^+$ )	Warm with $NaOH$ ; ammonia gas produced	-
Copper (II) ( $Cu^{2+}$ )	Light blue precipitate; insoluble in excess	Light blue precipitate; soluble in excess giving deep blue solution
Iron (II) ( $Fe^{2+}$ )	Green precipitate; insoluble in excess	Green precipitate; insoluble in excess
Iron (III) ( $Fe^{3+}$ )	Red-brown precipitate; insoluble in excess	Red-brown precipitate; insoluble in excess
Zinc ( $Zn^{2+}$ )	White precipitate; soluble in excess (colourless solution)	White precipitate; soluble in excess (colourless solution)
Calcium ( $Ca^{2+}$ )	White precipitate; insoluble in excess	No precipitate or slight precipitate
Aluminium ( $Al^{3+}$ )	White precipitate; soluble in excess (colourless solution)	White precipitate; insoluble in excess