Comprehensive Study Notes: Acids, Bases, and Salts

Scope of Syllabus and General Definitions

• Simple Definitions: Concepts are defined in terms of molecules and their characteristic properties.
• Ions in Substances: Acids, alkalis, and salts contain specific ions in their mineral and solution forms.
• Acid Properties: Acids form hydronium ions ($H_3O^+$) as the only positive ions in water. These ions turn blue litmus red.
• Alkali Properties: Alkalis form hydroxyl ions ($OH^-$) as the only negative ions in water. These turn red litmus blue.
• Salt Definition: Salts are formed by the partial or complete replacement of the ionizable hydrogen ion of an acid by a metal or an ammonium ion.
• Introduction to pH Scale: The pH scale, pH paper, and Universal indicators are used to test for acidity ($pH < 7$), neutrality ($pH = 7$), and alkalinity ($pH > 7$).
• Types of Salts:
  • Normal Salts: Formed by complete replacement of hydrogen ions.
  • Acid Salts: Formed by partial replacement of hydrogen ions.
  • Basic Salts: Formed by partial replacement of hydroxyl ions of a base.

Physical Properties and States of Acids

• Sour Taste: Most acids, like lemonade, have a sour taste. However, inorganic/mineral acids like sulphuric acid ($H_2SO_4$) should never be tasted as they are harmful.
• Physical States of Acids:
  • Solid Acids: Boric acid ($H_3BO_3$), Oxalic acid ($(COOH)_2$), Tartaric acid ($C_4H_6O_6$), Citric acid ($C_6H_8O_7$).
  • Liquid Acids: Acetic acid ($CH_3COOH$), Formic acid ($HCOOH$), Carbonic acid ($H_2CO_3$), Hydrochloric acid ($HCl$), Nitric acid ($HNO_3$), Phosphoric acid ($H_3PO_4$), Sulphurous acid ($H_2SO_3$), Sulphuric acid ($H_2SO_4$).
• Effect on Skin: Strong mineral acids have a corrosive action and cause painful burns. Organic acids and Carbonic acid are notably non-corrosive.
• Electrolytic nature: Acids in aqueous solution conduct electricity, acting as electrolytes. When a power supply and bulb are connected to a circuit with an acid like $HCl$, the bulb will glow.

Chemical Definition and Ionization of Acids

• Definition: An acid is a compound which, when dissolved in water, yields hydronium ions ($H_3O^+$) as the only positively charged ions.
• Ionization Process for Hydrochloric Acid:
  • $HCl(aq) \rightarrow H^+ + Cl^-$
  • $H^+ + H_2O \rightarrow H_3O^+$
  • Overall: $HCl + H_2O \rightarrow H_3O^+ + Cl^-$
• Ionization of Sulphuric Acid:
  • $H_2SO_4(aq) \rightarrow 2H^+ + SO_4^{2-}$
  • $2H^+ + 2H_2O \rightarrow 2H_3O^+$
  • Overall: $H_2SO_4 + 2H_2O \rightarrow 2H_3O^+ + SO_4^{2-}$
• Ionization of Nitric Acid:
  • $HNO_3(aq) \rightarrow H^+ + NO_3^-$
  • $H^+ + H_2O \rightarrow H_3O^+$
  • Overall: $HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$
• Formation of Hydronium Ion: A water molecule ($H_2O$) has lone pairs of electrons. When an acid provides a hydrogen ion ($H^+$), the lone pair on oxygen forms a coordinate bond with the $H^+$ to create the hydronium ion ($H_3O^+$).

Bases and Alkalis

• Definition of a Base: A compound which reacts with the hydronium ions of an acid to produce salt and water as the only products. Bases are typically oxides or hydroxides of metals (including ammonium hydroxide).
• Examples of Bases: Copper oxide ($CuO$), Copper hydroxide ($Cu(OH)_2$), Magnesium oxide ($MgO$), Calcium hydroxide ($Ca(OH)_2$), Sodium hydroxide ($NaOH$), Potassium hydroxide ($KOH$), Ferrous hydroxide ($Fe(OH)_2$).
• Reaction Examples:
  • $CuO + 2HCl \rightarrow CuCl_2 + H_2O$
  • $Zn(OH)_2 + H_2SO_4 \rightarrow ZnSO_4 + 2H_2O$
  • $NH_4OH + HNO_3 \rightarrow NH_4NO_3 + H_2O$
• Exceptions: Not all metal oxides are bases. For instance, Lead (IV) oxide ($PbO_2$) and Manganese (IV) oxide ($MnO_2$) react with $HCl$ to produce salt, water, AND chlorine gas:
  • $PbO_2 + 4HCl \rightarrow PbCl_2 + 2H_2O + Cl_2$
  • $MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$
• Definition of Alkali: A water-soluble base is called an alkali. When dissolved in water, it yields hydroxyl ions ($OH^-$) as the only negatively charged ions.
  • Soluble Bases (Alkalis): $NaOH$, $Na_2O$, $KOH$, $K_2O$, $NH_4OH$. $Ca(OH)_2$ is sparingly soluble.
  • Insoluble Bases: $ZnO$, $Zn(OH)_2$, $MgO$, $Fe(OH)_2$, $Cu(OH)_2$, $Fe(OH)_3$.
• Concept Rule: All alkalis are bases, but all bases are not alkalis.

Classification of Acids and Bases

• Based on Source (Acids):
  • Organic Acids: Derived from plants (e.g., Citric acid, Acetic acid).
  • Inorganic/Mineral Acids: Derived from minerals (e.g., $HCl$, $H_2SO_4$, $HNO_3$).
• Based on Molecular Composition (Acids):
  • Hydracids: Contain hydrogen and a non-metal other than oxygen (e.g., $HCl$, $HF$, $HBr$).
  • Oxy-acids: Contain hydrogen, oxygen, and another element (e.g., $HNO_3$, $H_2SO_4$, $H_3PO_4$, $H_2CO_3$).
• Based on Strength:
  • Strong Acid: Dissociates almost completely in aqueous solution, yielding high concentration of $H_3O^+$ (e.g., $HCl$, $H_2SO_4$, $HNO_3$).
  • Weak Acid: Dissociates only partially, yielding low concentration of $H_3O^+$ (e.g., $CH_3COOH$, $H_2CO_3$).
  • Strong Alkali: Dissociates almost completely in aqueous solution, yielding high concentration of $OH^-$ (e.g., $NaOH$, $KOH$, $LiOH$).
  • Weak Alkali: Dissociates only partially, yielding low concentration of $OH^-$ (e.g., $NH_4OH$, $Ca(OH)_2$).
• Based on Concentration:
  • Concentrated: High percentage of acid/alkali in aqueous solution.
  • Dilute: Low percentage of acid/alkali in aqueous solution.
• Basicity of Acids: Defined as the number of hydrogen ions produced per molecule in aqueous solution.
  • Monobasic: Produces one $H^+$ (e.g., $HCl$, $HNO_3$, $CH_3COOH$). Forms only normal salts.
  • Dibasic: Produces two $H^+$ (e.g., $H_2SO_4$, $H_2CO_3$, $H_2SO_3$, $(COOH)_2$, $H_3PO_3$). Forms acid salts and normal salts.
  • Tribasic: Produces three $H^+$ (e.g., $H_3PO_4$, Citric acid). Forms two acid salts and one normal salt.
• Acidity of Bases: Defined as the number of hydroxyl ions produced per molecule.
  • Monoacidic: Yields one $OH^-$ (e.g., $NaOH$, $KOH$).
  • Diacidic: Yields two $OH^-$ (e.g., $Ca(OH)_2$, $Mg(OH)_2$).
  • Triacidic: Yields three $OH^-$ (e.g., $Al(OH)_3$, $Fe(OH)_3$).

Safety during Dilution of Acids

• Caution: Dissolution of acid in water is highly exothermic.
• The Procedure: Concentrated acid must be added drop-wise to cold water with constant stirring.
• Why: If water is added to acid, the large amount of heat generated turns water into steam instantly, causing the acid to spurt and result in accidents. Cold water absorbs the excess heat if acid is added to it.

Indicators and pH Applications

• Indicators: Weak organic compounds that change color based on the pH or hydrogen ion concentration of the medium.
• Types:
  • Natural: Litmus (red/blue), Red cabbage extract.
  • Synthetic: Phenolphthalein (colorless in acid, pink in alkali), Methyl orange (red in acid, yellow in alkali).
• Color Chart for Common Indicators:
  • Red Litmus: Remains red in acid, turns blue in alkali.
  • Blue Litmus: Turns red in acid, remains blue in alkali.
  • Methyl Orange: Turns red in acidic, turns yellow in alkaline.
  • Phenolphthalein: Colorless in acidic, turns pink in alkaline.
• Universal Indicator: A mixture of organic dyes used to determine the exact pH or strength. It matches specific colors to pH values (1–14).
• pH Scale Definition: Calculated as the negative log of hydrogen ion concentration: $pH = -\text{log}[H^+]$.
• Applications of pH:
  • Agriculture: Soil testing. Citrus fruits need alkaline soil; Rice needs acidic medium; Sugarcane needs neutral medium.
  • Acid Rain: Defined by pH less than $5.6$.
  • Human Health: Blood pH range is $7.34$ to $7.45$. Milk turning sour shows a change from pH $6.6$.
  • Medicine: Urine/blood pH used for diagnosis.

Preparation of Acids

1. Synthesis/Direct Combination: Non-metal + Hydrogen. Example: $H_2 + Cl_2 \rightarrow 2HCl$ (in diffused sunlight).
2. Acidic Oxides + Water:
   • $CO_2 + H_2O \rightarrow H_2CO_3$
   • $SO_2 + H_2O \rightarrow H_2SO_3$
3. Non-volatile acid + Salt: Utilizing the lower volatility of conc. $H_2SO_4$ to liberate volatile acids like $HCl$ or $HNO_3$:
   • $NaCl + \text{conc. } H_2SO_4 \rightarrow NaHSO_4 + HCl$ (at $< 200^{\circ}C$).
   • $KNO_3 + \text{conc. } H_2SO_4 \rightarrow KHSO_4 + HNO_3$ (at $< 200^{\circ}C$).
4. Oxidation of Non-metals: Non-metals react with concentrated nitric acid to form oxyacids.
   • $C + 4HNO_3 \rightarrow CO_2 + 2H_2O + 4NO_2$
   • $S + 6HNO_3 \rightarrow H_2SO_4 + 2H_2O + 6NO_2$

Preparation of Bases

1. Metal + Oxygen: $4Na + O_2 \rightarrow 2Na_2O$; $2Ca + O_2 \rightarrow 2CaO$.
2. Soluble metallic oxides + Water: $K_2O + H_2O \rightarrow 2KOH$. Also Ammonia + Water: $NH_3 + H_2O \rightarrow NH_4OH$.
3. Water + Active Metal: $2K + 2H_2O \rightarrow 2KOH + H_2$. $Zn + H_2O \rightarrow ZnO + H_2$.
4. Precipitation: $FeCl_2 + 2NH_4OH \rightarrow 2NH_4Cl + Fe(OH)_2$.
5. Thermal Decomposition:
   • Carbonates: $CuCO_3 \rightarrow CuO + CO_2$.
   • Nitrates: $2Pb(NO_3)_2 \rightarrow 2PbO + 4NO_2 + O_2$.

Typical Chemical Properties of Acids

• Reaction with Metals: Active metals (above Hydrogen in activity series) + Dilute Acid ($HCl$ or $H_2SO_4$) $\rightarrow$ Soluble Salt + $H_2$ gas.
  • $Fe + H_2SO_4 \rightarrow FeSO_4 + H_2$
  • $Zn + 2HCl \rightarrow ZnCl_2 + H_2$
  • Observation: Colorless, odorless gas that burns with a "pop" sound ($H_2$).
• Nitric Acid ($HNO_3$) Exception: It is a strong oxidizing agent. Usually, it produces water and a different nitrogen oxide instead of hydrogen, as the nascent oxygen released oxidizes $H_2$ to $H_2O$.
  • $3Cu + 8HNO_3 \rightarrow 3Cu(NO_3)_2 + 4H_2O + 2NO$
  • Exemption: Only extremely dilute $HNO_3$ reacts with Magnesium ($Mg$) and Manganese ($Mn$) to liberate $H_2$.
• Neutralization: Acid + Base $\rightarrow$ Salt + Water only.
  • $2HCl + CuO \rightarrow CuCl_2 + H_2O$
  • $H_2SO_4 + Mg(OH)_2 \rightarrow MgSO_4 + 2H_2O$
• Reaction with Carbonates/Bicarbonates: Salt + $H_2O$ + $CO_2$.
  • $Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2$
  • Observation: Gas turns lime water milky but has no effect on acidified potassium dichromate.
• Reaction with Sulphites/Bisulphites: Salt + $H_2O$ + $SO_2$.
  • $Na_2SO_3 + H_2SO_4 \rightarrow Na_2SO_4 + H_2O + SO_2$
  • Observation: Gas turns lime water milky AND turns acidified potassium dichromate from orange to green.
• Reaction with Sulphides: Metal sulphide + dilute acid $\rightarrow$ Salt + $H_2S$ gas.
  • $FeS + 2HCl \rightarrow FeCl_2 + H_2S$
  • Observation: Moist lead acetate paper turns silvery black ($PbS$ formation).
• Precipitation Reactions:
  • Dilute $HCl + AgNO_3 \rightarrow AgCl \text{ (white ppt)} + HNO_3$. $AgCl$ is soluble in $NH_4OH$.
  • Dilute $HCl + Pb(NO_3)_2 \rightarrow PbCl_2 \text{ (white ppt)} + 2HNO_3$. $PbCl_2$ is soluble in hot water.

Detailed Salt Classification and Laboratory Preparation

• Type comparison:
  • Normal Salts: No replaceable hydrogen (e.g., $NaCl$, $Na_2SO_4$). Do not show acid properties.
  • Acid Salts: Contain replaceable hydrogen (e.g., $NaHSO_4$, $NaHCO_3$). Show acid properties in solution.
  • Basic Salts: Contain hydroxyl groups (e.g., $Zn(OH)Cl$).
• Salt Solubility Rules:
  • Soluble: All Nitrates; all Ammonium, Sodium, and Potassium salts ($NaHCO_3$/$KHCO_3$ are exceptions); most Sulphates (except $Pb$, $Ba$, $Ca$); most Chlorides (except $AgCl$, $PbCl_2$ is soluble only in hot water).
  • Insoluble: Most Carbonates (except $Na$, $K$, $NH_4$); most Sulphides (except $Na$, $K$, $NH_4$); most Oxides and Hydroxides (except $Na$, $K$, $NH_4$).
• Preparation Methods:
  1. Direct Combination: For anhydrous salts like $FeCl_3$. $2Fe + 3Cl_2 \rightarrow 2FeCl_3$. Stored with anhydrous $CaCl_2$ as it is deliquescent.
  2. Simple Displacement: Metal + acid. Used for soluble salts of active metals.
  3. Neutralization (Titration): Neutralization of an alkali with an acid using a burette/indicator. Example: $NaOH + HCl \rightarrow NaCl + H_2O$.
  4. Precipitation (Double Decomposition): For insoluble salts. Soluble salt + Soluble salt $\rightarrow$ Insoluble salt + Soluble salt.
     • Example: $BaCl_2 + H_2SO_4 \rightarrow BaSO_4 \text{ (ppt)} + 2HCl$.
  5. Two-Step Preparation: Conversion of an insoluble reactant to a nitrate first (using $HNO_3$), then precipitating the target salt.
     • Preparation of $PbCl_2$ from $PbO$:
       • Step 1: $PbO + 2HNO_3 \rightarrow Pb(NO_3)_2 + H_2O$
       • Step 2: $Pb(NO_3)_2 + 2NaCl \rightarrow PbCl_2 \text{ (ppt)} + 2NaNO_3$

Analytical Chemistry: Action of Alkalis on Salt Solutions

• Reactions with $NaOH$ and $NH_4OH$: Used to identify cations ($Ca^{2+}$, $Cu^{2+}$, $Fe^{2+}$, $Fe^{3+}$, $Pb^{2+}$, $Zn^{2+}$).
• Cation Observations:
  • $Cu^{2+}$: Pale blue ppt [$Cu(OH)_2$]. Insoluble in excess $NaOH$. Soluble in excess $NH_4OH$ to form an inky blue [$Cu(NH_3)_4(OH)_2$] solution.
  • $Fe^{2+}$: Dirty green ppt [$Fe(OH)_2$]. Insoluble in excess of either alkali.
  • $Fe^{3+}$: Reddish brown ppt [$Fe(OH)_3$]. Insoluble in excess of either alkali.
  • $Zn^{2+}$: Gelatinous white ppt [$Zn(OH)_2$]. Soluble in excess of BOTH alkalis forming colorless solutions.
  • $Pb^{2+}$: Chalky white ppt [$Pb(OH)_2$]. Soluble in excess $NaOH$, but INSOLUBLE in excess $NH_4OH$.
  • $Ca^{2+}$: Curdy white ppt [$Ca(OH)_2$]. Sparingly soluble in excess $NaOH$. No precipitate formed with $NH_4OH$.
• Amphoteric Nature: The oxides and hydroxides of $Zn$, $Pb$, and $Al$ react with BOTH acids and concentrated alkalis ($NaOH, KOH$).
  • $Zn + 2NaOH \rightarrow Na_2ZnO_2 \text{ (Sodium Zincate)} + H_2$
  • $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 \text{ (Sodium Aluminate)} + H_2O$

Hydrolysis of Salts

• Salt of Strong Acid + Strong Base: Do not hydrolyze. Solution is neutral (pH 7). Examples: $NaCl$, $Na_2SO_4$, $NaNO_3$.
• Salt of Strong Acid + Weak Base: Hydrolyzes to form an acidic solution. Examples: $NH_4Cl$, $(NH_4)_2SO_4$, $NH_4NO_3$. Turns blue litmus red.
• Salt of Weak Acid + Strong Base: Hydrolyzes to form an alkaline solution. Examples: $Na_2CO_3$, $CH_3COONa$. Turns red litmus blue.

Questions & Discussion

• Question: The basicity of acetic acid is?
  • Answer: 1 (Only the hydrogen in the hydroxyl group is replaceable).
• Question: Solution A has pH = 1, B has pH = 4.
  • Answer: A is a much stronger acid than B.
• Question: Distinguish between $FeSO_4$ and $FeCl_3$ solutions.
  • Answer: Use $NaOH$. $FeSO_4$ gives a dirty green ppt; $FeCl_3$ gives a reddish brown ppt.
• Question: What gas is evolved if Magnesium is put in an acidic solution?
  • Answer: Hydrogen gas ($H_2$).
• Question: Identify the cation: Salt Z added to calcium hydroxide and heated gives a pungent smelling gas turning moist red litmus blue.
  • Answer: Ammonium ($NH_4^+$); the gas is Ammonia ($NH_3$).