Comprehensive Guide to Acids, Bases, and Salts

Introduction to Acids, Bases, and Salts

• Classification of Compounds: Elements combine in fixed ratios to form compounds, which are fundamentally classified into three distinct categories: acids, bases, and salts.
• Basic Sensory Identification:     * Acids: Identified by their sour taste.     * Bases: Identified by a bitter taste and a soapy/slippery touch.     * Salts: Formed through the chemical interaction (combination) of acids and bases.
• Safety Warning: Not all acids and bases should be tasted or touched, as many are corrosive and can cause severe harm to the body.
• Scientific Identification: Acids and bases are differentiated using chemical indicators—substances that exhibit a sharp change in color depending on the nature (acidity or alkalinity) of a solution—and through their characteristic chemical reactions.
• Historical Context:     * Lavoisier and Davy: These scientists studied these terms on scientific lines.     * Scientific Findings on Acids: All acids must contain hydrogen as a constituent element in the form of a proton ($H^+$). Examples include Hydrochloric acid, sulphuric acid, and nitric acid.     * Scientific Findings on Bases: Bases are defined as metallic oxides ($O^{2-}$), metallic hydroxides ($OH^-$), or ammonia solutions. Examples include Copper oxide, sodium hydroxide, and ammonium hydroxide.     * Scientific Findings on Salts: These are chemical compounds formed by replacing the hydrogen ions in an acid with metallic ions or non-metallic ammonium ions ($NH_4^+$).

The Concept and Definition of Acids

• Etymology: The word 'acid' is derived from the Latin word acidus, which means "sour."
• Chemical Definition: Acids are compounds containing one or more hydrogen atoms that, when dissolved in water, produce hydronium ions ($H_3O^+$) as the only positively charged ions.
• Proton Donation: An acid molecule furnishes a hydrogen ion ($H^+$), also known as a proton, in aqueous solution.     * Equation: $HCl \xrightarrow{H_2O} H^+ + Cl^-$
• Hydronium Ion Formation: The $H^+$ ion cannot exist independently in water. It combines with a water molecule to form the hydronium ion ($H_3O^+$).     * Equation: $H^+ + H_2O \rightarrow H_3O^+$
• Acidic Properties: The characteristic properties of an acid (sourness, reaction with metals, etc.) are actually the properties of the hydronium ions present in the solution.
• Simplified Notation: In chemistry, $H^+$ is often used as a simplified notation for the hydronium ion ($H_3O^+$). Occasionally, the ion is written as $OH_3^+$ because the positive charge originates from the extra $H^+$ attached to the neutral water molecule.
• Examples of Dissociation:     * Sulphuric acid: $H_2SO_4 + 2H_2O \rightarrow 2H_3O^+ + SO_4^{2-}$     * Nitric acid: $HNO_3 + H_2O \rightarrow H_3O^+ + NO_3^-$     * Acetic acid: $CH_3COOH + H_2O \rightarrow H_3O^+ + CH_3COO^-$

Classification of Acids

• Based on Sources:     1. Organic Acids: Obtained from plants. They contain carbon atoms along with hydrogen. They are "weak" acids, meaning they do not ionise completely in solution (leaving both ions and molecules present). Notable natural examples include:         * Acetic acid: Found in Vinegar.         * Citric acid: Found in Citrus fruits (oranges, lemons).         * Butyric acid: Found in Rancid butter.         * Formic acid: Found in the sting of bees and ants ($HCOOH$).         * Lactic acid: Found in sour milk and curd.         * Malic acid: Found in Apples.         * Oleic acid: Found in Olive oil.         * Oxalic acid: Found in Tomatoes ($(COOH)_2$).         * Stearic acid: Found in Fats.         * Tartaric acid: Found in Grapes, tamarind, and apples.         * Uric acid: Found in Urine.     2. Inorganic (Mineral) Acids: Obtained from minerals. Usually, they do not contain carbon. They are "strong" acids that ionise completely in solution, producing a high concentration of hydronium ions. Examples include:         * Hydrochloric acid ($HCl$).         * Sulphuric acid ($H_2SO_4$).         * Nitric acid ($HNO_3$).         * Phosphoric acid ($H_3PO_4$).         * Carbonic acid ($H_2CO_3$): A notable exception as it is an inorganic acid containing carbon; however, it is a weak mineral acid used in soft drinks because it is non-corrosive.
• Based on Molecular Composition:     1. Oxy-acids: Contain oxygen along with hydrogen and another element ($HNO_3, H_2SO_4$).     2. Hydracids: Contain hydrogen and a non-metallic element but no oxygen ($HCl, HBr$).
• Based on Basicity:     * Basicity is the number of hydronium ions ($H_3O^+$) produced by the ionisation of one molecule of the acid in aqueous solution. It depends on the number of ionisable hydrogen atoms, not the total count of hydrogens.     1. Monobasic Acids: Produce one $H_3O^+$ ion per molecule. They ionise in one step and form only one normal salt.         * Examples: $HCl, HBr, HNO_3, HI$.         * Acetic Acid Special Case: $CH_3COOH$ contains 4 hydrogens but is monobasic because only the hydrogen in the carboxyl group ionises: $CH_3COOH + H_2O \rightarrow H_3O^+ + CH_3COO^-$.     2. Dibasic Acids: Produce two $H_3O^+$ ions per molecule. They ionise in two steps and can form one acid salt and one normal salt.         * Examples: $H_2SO_4, (COOH)_2, H_2SO_3, H_2CO_3$.         * Stepwise Ionisation (e.g., Sulphuric Acid):             * Step 1: $H_2SO_4 + H_2O \rightarrow H_3O^+ + HSO_4^-$             * Step 2: $HSO_4^- + H_2O \rightarrow H_3O^+ + SO_4^{2-}$         * Phosphorous Acid Case: $H_3PO_3$ is dibasic. Only hydrogens attached to oxygen are replaceable; those bonded directly to phosphorus are not.     3. Tribasic Acids: Produce three $H_3O^+$ ions per molecule. They ionise in three steps and can form two different acid salts and one normal salt. Examples: Phosphoric acid ($H_3PO_4$), Citric acid ($C_3H_5O(COOH)_3$).     4. Tetrabasic Acids: Silicic acid ($H_4SiO_4$) has a basicity of four.

Strength and Concentration of Acids

• Concentration: Refers to the amount of acid relative to the amount of water in an aqueous solution.     * Concentrated Acid: Contains a very small amount of water or no water.     * Dilute Acid: Contains far more water than the mass of the acid itself.
• Strength: The measure of the concentration of hydronium ions produced. More $H^+$ ions released equates to a stronger acid.     * Comparative Example: Dilute $HCl$ is a stronger acid than highly concentrated acetic acid because $HCl$ releases more $H^+$ ions.
• Degree of Ionisation ($\alpha$): The strength is determined by the percentage of molecules that ionise in solution.     * Formula: $\alpha = \frac{\text{No. of acid molecules ionised}}{\text{Total no. of acid molecules present in solution}} \times 100$     * Strong Acid/Alkali: $\alpha > 30\%$.     * Weak Acid/Alkali: $\alpha < 30\%$.
• Dilution Procedure: To dilute an acid, one must always pour acid into water in small amounts and stir constantly.     * Exothermic Warning: Water must never be added to acid. This process is highly exothermic; the produced heat can cause the acidic solution to splash, or the container to break, which can be fatal.

Preparation of Acids

• Synthesis: Binary acids (containing two elements) are prepared by direct combination of elements under appropriate conditions.     * $H_2 + Cl_2 \rightarrow 2HCl$     * $H_2 + Br_2 \rightarrow 2HBr$     * $H_2 + S \rightarrow H_2S$
• Action of Water on Acidic Oxides (Acidic Anhydrides): Most oxy-acids are prepared by dissolving acidic oxides in water.     * Sulphuric acid: $SO_3 + H_2O \rightarrow H_2SO_4$     * Sulphurous acid: $SO_2 + H_2O \rightarrow H_2SO_3$     * Carbonic acid: $CO_2 + H_2O \rightarrow H_2CO_3$     * Phosphoric acid: $P_2O_5 + 3H_2O \rightarrow 2H_3PO_4$     * Nitric acid: $N_2O_5 + H_2O \rightarrow 2HNO_3$     * Note: $NO_2$ is a "mixed" or "double" acid anhydride because it forms two acids (nitrous and nitric) when reacting with water: $2NO_2 + H_2O \rightarrow HNO_2 + HNO_3$.
• Oxidation of Non-metals: Concentrated nitric acid is used to oxidise non-metals like sulphur or phosphorus to produce sulphuric or phosphoric acid.     * $S + 6HNO_3 \rightarrow H_2SO_4 + 2H_2O + 6NO_2$     * $P + 5HNO_3 \rightarrow H_3PO_4 + H_2O + 5NO_2$
• Displacement: A less volatile or non-volatile acid (like sulphuric acid) can displace a more volatile acid from its normal salts.     * $NaCl + H_2SO_4 \rightarrow NaHSO_4 + HCl$     * $NaNO_3 + H_2SO_4 \rightarrow NaHSO_4 + HNO_3$

Properties and Indicators of Acids

• Physical Properties:     * Taste: Sour.     * Corrosive Action: Mineral acids like $H_2SO_4$ and $HNO_3$ are highly corrosive. They char skin black ($H_2SO_4$), turn it yellow ($HNO_3$), or amber ($HCl$). Organic acids and carbonic acid are non-corrosive.     * Physical State: Some are solids (Boric, Oxalic, Tartaric, Citric, Phosphoric) while others are liquids (Acetic, Formic, Carbonic, $HCl, HNO_3, H_2SO_4$).     * Volatility: Most liquid acids are volatile (vaporize easily) except for Sulphuric acid ($H_2SO_4$), which is non-volatile.     * Conductivity: Acids in the aqueous state conduct electricity and are therefore electrolytes.
• Chemical Indicators:     * Litmus: Changes from Red (Acidic) to Purple (Neutral) to Blue (Basic).     * Methyl Orange: Changes from Red/Pink (Acidic) to Orange (Neutral) to Yellow (Basic).     * Phenolphthalein: Remains Colorless in Acidic and Neutral media, becomes Pink in Basic media.     * Universal Indicator: A mixture of indicators showing a spectrum: Red (Strong Acid) $\rightarrow$ Yellow (Weak Acid) $\rightarrow$ Green (Neutral) $\rightarrow$ Blue/Purple (Alkaline).     * Olfactory Indicators: Substances like onion, vanilla, and clove change their smell in different media. Basic solutions (e.g., $NaOH$) destroy the smell of onion extract, whereas acidic solutions do not.     * Natural Indicators: Red cabbage extract remains red in acid and turns green in base.
• Chemical Reactions:     1. With Active Metals: Forms metallic salt and hydrogen gas ($H_2$).         * $Mg + 2HCl \rightarrow MgCl_2 + H_2$         * Exception: $HNO_3$ is a strong oxidising agent and usually produces water instead of hydrogen. Only very dilute ($1\%$) $HNO_3$ produces hydrogen with $Mg$ or $Mn$.     2. Neutralisation: Acids react with basic oxides or hydroxides to form salt and water ONLY.         * $CuO + H_2SO_4 \rightarrow CuSO_4 + H_2O$     3. Decomposition of Carbonates/Bicarbonates: Liberates carbon dioxide ($CO_2$) gas.         * $CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2$     4. Decomposition of Sulphites/Bisulphites: Liberates sulphur dioxide ($SO_2$) gas.         * $CaSO_3 + 2HCl \rightarrow CaCl_2 + H_2O + SO_2$     5. Decomposition of Sulphides: Liberates hydrogen sulphide ($H_2S$) gas (smell of rotten eggs).         * $FeS + H_2SO_4 \rightarrow FeSO_4 + H_2S$     6. Reaction with Chlorides and Nitrates: Most require concentrated Sulphuric acid and warming.         * $NaCl + H_2SO_4 \xrightarrow{< 200^{\circ}C} NaHSO_4 + HCl$

The Concept and Definition of Bases

• Chemical Definition: A base is a metallic oxide, metallic hydroxide, or ammonium hydroxide that reacts with the hydronium ions of an acid to form salt and water only.
• Basic Oxide: A metallic oxide containing the $O^{2-}$ ion.     * Counterexample: Lead (IV) oxide ($PbO_2$) reacts with $HCl$ to produce salt, water, and chlorine gas, so it is NOT considered a base because it produces more than just salt and water.
• Basic Hydroxide: A metallic hydroxide containing the $OH^-$ group.
• Alkalis: These are bases that are soluble in water. They produce hydroxyl ions ($OH^-$) as the only negative ions in aqueous solution.     * Key Principle: All alkalis are bases, but all bases are not alkalis ($Fe(OH)_3$ and $Cu(OH)_2$ are bases but insoluble in water).     * Common Alkalis: Sodium hydroxide (Caustic soda), Potassium hydroxide (Caustic potash), Calcium hydroxide (Slaked lime), Ammonium hydroxide.

Classification of Bases

• Based on Strength:     1. Strong Alkalis: Undergo almost complete dissociation to produce a high concentration of $OH^-$ ions. Examples: $NaOH, KOH$.     2. Weak Alkalis: Undergo partial dissociation to produce a low concentration of $OH^-$ ions. Examples: $NH_4OH, Ca(OH)_2$.
• Based on Acidity:     * Acidity is the number of hydroxyl ions ($OH^-$) produced per molecule of the base in aqueous solution. It also represents the number of $H^+$ ions with which one molecule of the base will react.     1. Monoacidic Base: Produces one $OH^-$ ion. Examples: $NaOH, KOH, NH_4OH$.     2. Diacidic Base: Produces two $OH^-$ ions. Examples: $Ca(OH)_2, Cu(OH)_2, Mg(OH)_2$.     3. Triacidic Base: Produces three $OH^-$ ions. Examples: $Al(OH)_3, Fe(OH)_3$.

Preparation of Bases

1. From Metals: Metals react with oxygen to form basic oxides.     * $4Na + O_2 \rightarrow 2Na_2O$
2. Metal + Water: Reactive metals (Na, K, Ca) react with water to form alkalis and hydrogen.     * $2K + 2H_2O \rightarrow 2KOH + H_2$
3. Dissolving Basic Oxides: Soluble metallic oxides dissolved in water form alkalis.     * $CaO + H_2O \rightarrow Ca(OH)_2$
4. Double Decomposition: Aqueous solutions of salts react with alkalis to precipitate insoluble metallic hydroxides.     * $FeCl_3 + 3NaOH \rightarrow Fe(OH)_3 \downarrow + 3NaCl$
5. Oxidation of Metal Sulphides: Heating metal sulphides in oxygen.     * $2ZnS + 3O_2 \xrightarrow{\Delta} 2ZnO + 2SO_2$
6. Decomposition of Salts: Heating metal carbonates or nitrates (excluding Na and K).     * $CaCO_3 \xrightarrow{\Delta} CaO + CO_2$     * $2Zn(NO_3)_2 \xrightarrow{\Delta} 2ZnO + 4NO_2 + O_2$
7. Ammonia + Water: Dissolving ammonia gas in water produces ammonium hydroxide ($NH_4OH$).

Properties and Reactions of Bases and Alkalis

• Physical Properties: Sharp and bitter taste, soapy/slippery touch, strong electrolytes, and mild corrosive action on skin.
• Indicators: Turn Red litmus blue, Methyl orange yellow, and Phenolphthalein pink (Insoluble bases do not affect indicators).
• Chemical Reactions:     1. Absorption of $CO_2$: Strong alkalis absorb carbon dioxide from the air to form carbonates.         * $2NaOH + CO_2 \rightarrow Na_2CO_3 + H_2O$     2. Neutralisation: React with acids to form salt and water.     3. Amphoteric Nature: Hydroxides of Zinc, Aluminium, and Lead are amphoteric; they dissolve in excess $NaOH$ or $KOH$.         * $Zn(OH)_2 + 2NaOH \rightarrow Na_2ZnO_2 + 2H_2O$ (Sodium zincate).     4. Warming with Ammonium Salts: Liberates ammonia gas ($NH_3$).         * $NH_4Cl + NaOH \xrightarrow{\Delta} NaCl + H_2O + NH_3$

The pH Scale and Its Importance

• Definition: Devised by Sorensen in 1909, the pH scale represents the hydronium ion concentration in an aqueous solution. The 'p' stands for 'Potenz' (Power).
• Mathematical Expression: $pH = -\log_{10}[H^+]$ or $pH = \log_{10} \frac{1}{[H^+]}$.
• Scale Range: 0 to 14.     * Neutral: $pH = 7$ (e.g., pure water).     * Acidic: $pH < 7$ (acidity increases as pH decreases).     * Alkaline/Basic: $pH > 7$ (alkalinity increases as pH increases).
• Ionic Product of Water ($K_w$): At $25^{\circ}C$, $K_w = [H_3O^+][OH^-] = 1 \times 10^{-14}\,mol^2\,litre^{-2}$.
• Everyday Importance of pH:     * Human Body: Works within a range of $7.0$ to $7.8$.     * Agriculture: Specific crops require specific pH. Rice (slightly acidic), Sugarcane (neutral), Citrus fruits (alkaline).     * Acid Rain: If rain water pH falls below $5.6$, it is acid rain (caused by $SO_2$ and $NO_X$), damaging monuments and reducing soil fertility.     * Digestion: Stomach produces $HCl$; excess causing irritation is treated with antacids like Milk of Magnesia ($Mg(OH)_2$).     * Tooth Decay: Decay starts if mouth pH falls below $5.5$. Tooth enamel (calcium phosphate) is corroded.     * Bees/Ants: Stings leave acid. Baking soda (a base) provides relief.

Definition and Classification of Salts

• Chemical Definition: A salt is a compound formed by the partial or total replacement of ionisable hydrogen atoms of an acid by a metallic or ammonium ion; an ionic compound yielding ions other than $H^+$ and $OH^-$.
• Types of Salts:     1. Normal Salts: Formed by complete replacement of hydrogen atoms. No ionisable hydrogen remains ($NaCl, Na_2SO_4$).     2. Acid Salts: Formed by partial replacement of hydrogen atoms from a polybasic acid. They contain ionisable hydrogen and show acidic properties in solution ($NaHSO_4, NaH_2PO_4$).     3. Basic Salts: Formed by partial replacement of hydroxyl groups of a diacidic or triacidic base. Contain metallic cations, hydroxyl ions, and acid anions ($Pb(OH)Cl, Cu(OH)NO_3$).     4. Double Salts: Formed by the union of two simple salts that crystallize together ($K_2SO_4 \cdot Al_2(SO_4)_3 \cdot 24H_2O$ - Potash alum).     5. Mixed Salts: Contain more than one basic or acid radical ($CaOCl_2$ - Bleaching powder).     6. Complex Salts: Dissociate to give one simple ion and one complex ion, the latter written in square brackets ($Na[Ag(CN)_2]$ - Sodium argentocyanide).

Preparation of Salts

• Soluble Salts (Prepared in solution and recovered by evaporation/crystallisation):     * Direct Combination: Heating elements ($2Fe + 3Cl_2 \rightarrow 2FeCl_3$).     * Simple Displacement: Active metal + dilute acid ($Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2$).     * Neutralisation (Soluble Base): Known as Titration. Used for Na, K, and $NH_4$ salts.     * Neutralisation (Insoluble Base): Acid + Insoluble oxide/carbonate ($CuO + H_2SO_4 \rightarrow CuSO_4 + H_2O$).
• Insoluble Salts (Prepared by precipitation/double decomposition):     * Precipitation: Mixing two appropriate soluble salt solutions ($AgNO_3 + HCl \rightarrow AgCl \downarrow + HNO_3$).     * Conversion Rule: To prepare an insoluble salt from another insoluble one, convert the first into a soluble form (usually nitrate) before precipitation.

General Properties of Salts

• Nature: Electrovalent compounds, non-volatile solids, crystallisable.
• Salt Hydrolysis: The phenomenon where a salt of a weak acid/strong base (or vice versa) reacts with water to produce an alkaline or acidic solution.     * Strong base + Weak acid $\rightarrow$ pH $> 7$.     * Weak base + Strong acid $\rightarrow$ pH $< 7$.
• Water of Crystallisation: A definite quantity of water in chemical combination with crystalline salts. It can be removed by heating above $100^{\circ}C$.     * Notation: $CuSO_4 \cdot 5H_2O$ (Blue) $\xrightarrow{\Delta}$ $CuSO_4$ (White powder) + $5H_2O$.
• Salt Behaviors:     1. Efflorescence: Salts losing water of crystallisation to dry air, becoming powdery (e.g., Washing soda, Glauber's salt).     2. Hygroscopy: Absorbing atmospheric moisture without dissolving (e.g., Conc. $H_2SO_4$, Quicklime, Silica gel).     3. Deliquescence: Absorbing moisture until dissolving into a saturated solution (e.g., Caustic soda, $FeCl_3, CaCl_2$).         * Note: Table salt ($NaCl$) is not deliquescent when pure, but commercial salt contains $MgCl_2$ or $CaCl_2$ impurities which are deliquescent.
• Agents:     * Drying Agent: Removes moisture physically (e.g., used in desiccators).     * Dehydrating Agent: Removes chemically combined water/elements of water physically or via chemical change ($H_2O$ in ratio 2:1) (e.g., Conc. $H_2SO_4$).

Questions & Discussion

• What is the positive ion formed when an acid dissolves in water?: The hydronium ion ($H_3O^+$).
• Why is water never added to acid to dilute it?: The process is highly exothermic; adding water to concentrated acid can cause the solution to boil instantly and splash, potentially causing fatal burns.
• What is meant by the basicity of an acid?: It is the number of ionisable hydrogen atoms per molecule of acid that can produce hydronium ions in aqueous solution.
• Explain why dil. HCl is stronger than concentrated acetic acid: Strength depends on the degree of ionisation ($\alpha$). $HCl$ ionises almost completely, providing a much higher concentration of $H^+$ ions compared to the partial ionisation of acetic acid.
• Why does Lead (IV) Oxide ($PbO_2$) not qualify as a base?: By definition, a base reacts with an acid to form salt and water only. $PbO_2$ reacts with $HCl$ to produce salt ($PbCl_2$), water, and chlorine gas, thus excluding it from the class of bases.