Chapter 6: Acid, Base and Salt

Overview

Acids, bases, and salts are fundamental concepts in chemistry that explain countless chemical reactions and processes in our daily lives. From the acids in our stomach that help digest food to the bases that clean our homes, and the salts that flavor our food, these substances play crucial roles in nature and industry. This chapter will guide you through the properties, reactions, and applications of acids, bases, and salts, including pH calculations, titration techniques, and qualitative analysis methods essential for SPM Chemistry success.

Learning Objectives

After studying this chapter, you should be able to:

Define acids and bases using different theories
Calculate pH and understand its significance
Distinguish between strong and weak acids/bases
Perform acid-base titration calculations
Prepare and identify various salts
Conduct qualitative analysis of salts
Apply acid-base concepts to real-world situations

6.1 The Role of Water in Showing Acidity and Alkalinity

Water as a Solvent

Water is called the "universal solvent" because it dissolves many ionic and polar compounds. In water, acids, bases, and salts dissociate into ions, making the solutions electrically conductive.

Ionization in Water

Acids in Water: Produce $H^+$ ions (or $H_3O^+$ hydronium ions) $\text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-$

Bases in Water: Produce $OH^-$ ions $\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$

Salts in Water: Produce both cations and anions $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$

Self-Ionization of Water

Water undergoes self-ionization: $\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-$

Ion Product Constant: $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$ at 25°C

Neutral Solution

In neutral water: $[\text{H}^+] = [\text{OH}^-] = 1.0 \times 10^{-7} \text{ M}$

Did You Know?

Pure water is actually slightly acidic due to dissolved carbon dioxide from the air, which forms carbonic acid. This is why even "pure" water has a pH of about 5.6 rather than exactly 7.0.

6.2 pH Value

Definition of pH

pH: Negative logarithm of the hydrogen ion concentration $\text{pH} =i -\log_{10} [\text{H}^+]$

pH Scale

pH Value	Solution Type	[H⁺] (M)	[OH⁻] (M)	Examples
0-3	Strongly acidic	$10^{-0}$ to $10^{-3}$	$10^{-14}$ to $10^{-11}$	HCl, $H_2SO_4$
4-6	Weakly acidic	$10^{-4}$ to $10^{-6}$	$10^{-10}$ to $10^{-8}$	Lemon juice, vinegar
7	Neutral	$1.0 \times 10^{-7}$	$1.0 \times 10^{-7}$	Pure water
8-10	Weakly basic	$10^{-8}$ to $10^{-10}$	$10^{-6}$ to $10^{-4}$	Sea water, soap
11-14	Strongly basic	$10^{-11}$ to $10^{-14}$	$10^{-3}$ to $10^{0}$	NaOH, bleach

pH Calculations

From [H⁺] to pH

Example: If $[H^+] = 1.0 \times 10^{-3}$ M $\text{pH} = -\log_{10} (1.0 \times 10^{-3}) = 3.0$

From pH to [H⁺]

Example: If pH = 2.5 $[\text{H}^+] = 10^{-\text{pH}} = 10^{-2.5} = 3.16 \times 10^{-3} \text{ M}$

Relationship between pH and pOH

$\text{pH} + \text{pOH} = 14 \quad (\text{at } 25°C)$

$\text{pOH} = -\log_{10} [\text{OH}^-]$

Common pH Values

Substance	pH	Approximate [H⁺] (M)
Battery acid	0	1
Gastric acid	1.5-3.5	$3.2 \times 10^{-2}$ to $3.2 \times 10^{-4}$
Lemon juice	2-3	$10^{-2}$ to $10^{-3}$
Vinegar	2.5-3.5	$3.2 \times 10^{-3}$ to $3.2 \times 10^{-4}$
Black coffee	5	$10^{-5}$
Milk	6.5-6.7	$2 \times 10^{-7}$ to $2 \times 10^{-7}$
Pure water	7	$1 \times 10^{-7}$
Sea water	8	$10^{-8}$
Soap	9-10	$10^{-9}$ to $10^{-10}$
Household ammonia	11-12	$10^{-11}$ to $10^{-12}$
Bleach	12.5	$3 \times 10^{-13}$

Measuring pH

pH Indicators

Indicator	pH Range	Color Change	Acid Color	Base Color
Methyl orange	3.1-4.4	Red → Yellow	Red	Yellow
Bromothymol blue	6.0-7.6	Yellow → Blue	Yellow	Blue
Phenolphthalein	8.2-10.0	Colorless → Pink	Colorless	Pink
Litmus	5.0-8.0	Red → Blue	Red	Blue

pH Measurement Devices

Device	Accuracy	Range	Principle
pH paper	±0.5	1-14	Color change
pH meter	±0.01	0-14	Electrochemical
Universal indicator	±1.0	1-14	Multiple dyes

SPM Exam Tips

For pH calculations:

Always show the formula before plugging in numbers

Use proper scientific notation for [H⁺]

Remember the relationship pH + pOH = 14
Be careful with significant figures in log calculations

6.3 Strength of Acids and Alkalis

Strong Acids

Definition: Completely dissociate in aqueous solution $\text{HA} \rightarrow \text{H}^+ + \text{A}^- \quad(\text{100\% dissociation})$

Common Strong Acids

Acid	Formula	Dissociation	pH of 0.1 M Solution
Hydrochloric acid	HCl	Strong	1.0
Sulfuric acid	$H_2SO_4$	Strong (first $H^+$ )	0.7
Nitric acid	$HNO_3$	Strong	1.0
Perchloric acid	$HClO_4$	Strong	1.0

Weak Acids

Definition: Partially dissociate in aqueous solution $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^- \quad{\text{<100\% dissociation}}$

Common Weak Acids

Acid	Formula	Dissociation	Ka at 25°C	pH of 0.1 M Solution
Acetic acid	$CH_3COOH$	Weak	$1.8 \times 10^{-5}$	2.87
Carbonic acid	$H_2CO_3$	Weak	$4.3 \times 10^{-7}$	3.68
Hydrofluoric acid	HF	Weak	$6.8 \times 10^{-4}$	2.08
Phosphoric acid	$H_3PO_4$	Weak	$7.5 \times 10^{-3}$	1.61

Strong Bases

Definition: Completely dissociate in aqueous solution $\text{MOH} \rightarrow \text{M}^+ + \text{OH}^- \quad (\text{100\% dissociation})$

Common Strong Bases

Base	Formula	Dissociation	pH of 0.1 M Solution
Sodium hydroxide	NaOH	Strong	13.0
Potassium hydroxide	KOH	Strong	13.0
Calcium hydroxide	$Ca(OH)_2$	Strong	12.4
Barium hydroxide	$Ba(OH)_2$	Strong	13.0

Weak Bases

Definition: Partially dissociate in aqueous solution $\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^- \quad (\text{<100\% dissociation})$

Common Weak Bases

Base	Formula	Dissociation	Kb at 25°C	pH of 0.1 M Solution
Ammonia	$NH_3$	Weak	$1.8 \times 10^{-5}$	11.13
Methylamine	$CH_3NH_2$	Weak	$4.4 \times 10^{-4}$	11.62
Pyridine	$C_5H_5N$	Weak	$1.7 \times 10^{-9}$	8.63

Concentration vs. Strength

Property	Strong Acid	Weak Acid
Dissociation	Complete	Partial
$[H^+]$ at 0.1 M	0.1 M	< 0.1 M
pH at 0.1 M	1.0	> 1.0
Conductivity	High	Low
Reaction rate	Fast	Slow
Examples	HCl, $H_2SO_4$	$CH_3COOH$ , $H_2CO_3$

Acid Dissociation Constants (Ka)

$K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$

Acid Strength Order

Acid	Ka	Relative Strength
HCl	Very large	Strongest
$H_2SO_4$	Very large	Very strong
$HNO_3$	Very large	Very strong
HF	$6.8 \times 10^{-4}$	Weak
$H_3PO_4$	$7.5 \times 10^{-3}$	Weak
$CH_3COOH$	$1.8 \times 10^{-5}$	Weaker
$H_2CO_3$	$4.3 \times 10^{-7}$	Very weak

Did You Know?

The strength of an acid is independent of its concentration. Concentrated acetic acid (vinegar) is still a weak acid because it only partially dissociates, while dilute hydrochloric acid is still a strong acid because it completely dissociates.

6.4 Chemical Properties of Acids and Bases

General Properties of Acids

Physical Properties

Property	Description	Example
Taste	Sour	Lemon juice, vinegar
Feel	Slippery (when dilute)	None (safely test with indicator)
pH	< 7	HCl: pH 1-2, $CH_3COOH$ : pH 2-3
Conductivity	Good (when in solution)	HCl solution conducts electricity
Reactivity	With metals, carbonates, bases	Reacts with Zn, $Na_2CO_3$ , NaOH

Chemical Properties

Reaction with Metals $\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen}$ $2\text{HCl} + \text{Zn} \rightarrow \text{ZnCl}_2 + \text{H}_2$
Reaction with Carbonates $\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Carbon dioxide} + \text{Water}$ $2\text{HCl} + \text{CaCO}_3 \rightarrow \text{CaCl}_2 + \text{CO}_2 + \text{H}_2\text{O}$
Reaction with Bases (Neutralization) $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$ $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$
Reaction with Metal Oxides $\text{Acid} + \text{Metal oxide} \rightarrow \text{Salt} + \text{Water}$ $2\text{HCl} + \text{CuO} \rightarrow \text{CuCl}_2 + \text{H}_2\text{O}$

General Properties of Bases

Physical Properties

Property	Description	Example
Taste	Bitter	Baking soda, soap
Feel	Slippery	Soap solution
pH	> 7	NaOH: pH 13-14, $NH_3$ : pH 11-12
Conductivity	Good (when in solution)	NaOH solution conducts electricity
Reactivity	With acids, ammonium salts, metals	Reacts with HCl, $NH_4Cl$ , Al

Chemical Properties

Reaction with Acids (Neutralization) $\text{Base} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$ $\text{NaOH} + \text{HCl} \rightarrow \text{NaCl} + \text{H}_2\text{O}$
Reaction with Ammonium Salts $\text{Base} + \text{Ammonium salt} \rightarrow \text{Salt} + \text{Ammonia} + \text{Water}$ $\text{NaOH} + \text{NH}_4\text{Cl} \rightarrow \text{NaCl} + \text{NH}_3 + \text{H}_2\text{O}$
Reaction with Metals $\text{Base} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen}$ $2\text{NaOH} + 2\text{Al} + 6\text{H}_2\text{O} \rightarrow 2\text{NaAl(OH)}_4 + 3\text{H}_2$
Thermal Decomposition $\text{Metal hydroxide} \rightarrow \text{Metal oxide} + \text{Water}$ $2\text{NaOH} \rightarrow \text{Na}_2\text{O} + \text{H}_2\text{O}$

Acid-Base Theories

Arrhenius Theory

Acid: Substance that produces $H^+$ ions in aqueous solution
Base: Substance that produces $OH^-$ ions in aqueous solution
Limitation: Only covers aqueous solutions

Brønsted-Lowry Theory

Acid: Proton ( $H^+$ ) donor
Base: Proton ( $H^+$ ) acceptor
Advantage: Covers non-aqueous solutions and can explain amphoteric behavior

Lewis Theory

Acid: Electron pair acceptor
Base: Electron pair donor
Advantage: Covers reactions without protons (e.g., $BF_3 + NH_3$ )

SPM Exam Tips

For acid-base reactions:

Memorize the general reaction patterns

Know the products of acid-metal, acid-carbonate, and base-ammonium reactions

Understand the different acid-base theories and their applications

Practice writing balanced equations for neutralization reactions

6.5 Concentration of Aqueous Solutions

Molarity

Definition: Number of moles of solute per liter of solution $\text{Molarity} = \frac{\text{moles of solute}}{\text{liters of solution}}$

Dilution Formula

$M_1V_1 = M_2V_2$

Example: Dilute 100 mL of 2.0 M HCl to 0.5 M $(2.0 \text{ M})(0.100 \text{ L}) = (0.5 \text{ M})(V_2)$ $V_2 = 0.4 \text{ L} = 400 \text{ mL}$

Concentration Calculations

Example 1: Preparing 0.1 M HCl

Given concentrated HCl is 12 M $\text{Moles needed} = 0.1 \text{ M} \times 0.1 \text{ L} = 0.01 \text{ mol}$ $\text{Volume of concentrated HCl} = \frac{0.01 \text{ mol}}{12 \text{ M}} = 0.00083 \text{ L} = 0.83 \text{ mL}$

Example 2: Neutralization Reaction

$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$ If 25.0 mL of 0.1 M HCl requires 20.0 mL of NaOH, find M of NaOH: $(0.1 \text{ M})(25.0 \text{ mL}) = M_{\text{NaOH}}(20.0 \text{ mL})$ $M_{\text{NaOH}} = \frac{2.5}{20.0} = 0.125 \text{ M}$

Did You Know?

The concentration of acids and bases can vary dramatically in nature. Stomach acid has a pH of about 1.5-3.5, while blood maintains a very precise pH of 7.35-7.45. Even small deviations from this range can be life-threatening.

6.6 Acid-Base Titration

Definition of Titration

Titration: Technique to determine the concentration of one solution by reacting it with a solution of known concentration.

Types of Titrations

Type	Indicator	Examples
Strong acid-Strong base	Phenolphthalein, methyl orange	HCl vs NaOH
Strong acid-Weak base	Methyl orange	HCl vs $NH_3$
Weak acid-Strong base	Phenolphthalein	$CH_3COOH$ vs NaOH
Weak acid-Weak base	None (pH meter needed)	$CH_3COOH$ vs $NH_3$

Titration Procedure

Equipment Needed

Burette, pipette, conical flask
Burette stand and clamp
Standard solutions
Appropriate indicator
White tile for color contrast

Step-by-Step Procedure

Fill Burette: Fill burette with standard solution (known concentration)
Pipette Sample: Pipette analyte solution into conical flask
Add Indicator: Add 2-3 drops of appropriate indicator
Titration: Slowly add standard solution from burette until color change
Record Volume: Note volume used at endpoint
Repeat: Perform multiple trials for accuracy

Indicator Selection

Titration Type	Suitable Indicator	Color Change
Strong acid-Strong base	Phenolphthalein	Colorless → Pink
Strong acid-Weak base	Methyl orange	Red → Yellow
Weak acid-Strong base	Phenolphthalein	Colorless → Pink

Titration Calculations

Example: Titration of HCl with NaOH

Given: 25.0 mL of unknown HCl requires 23.5 mL of 0.100 M NaOH $\text{Moles NaOH} = (0.100 \text{ M})(0.0235 \text{ L}) = 0.00235 \text{ mol}$ $\text{From balanced equation: HCl + NaOH → NaCl + H}_2\text{O}$ $\text{Moles HCl} = \text{Moles NaOH} = 0.00235 \text{ mol}$ $\text{Molarity HCl} = \frac{0.00235 \text{ mol}}{0.0250 \text{ L}} = 0.094 \text{ M}$

Endpoint vs Equivalence Point

Term	Definition	Importance
Endpoint	Point where indicator changes color	Visual detection
Equivalence Point	Point where moles acid = moles base	Theoretical completion
Titration Error	Difference between endpoint and equivalence point	Should be minimal

Common Titrations in Laboratory

Acid-Base Titrations

Reaction	Indicator	Color Change
HCl + NaOH	Phenolphthalein	Colorless → Pink
$CH_3COOH$ + NaOH	Phenolphthalein	Colorless → Pink
$H_2SO_4$ + NaOH	Bromothymol blue	Yellow → Blue

Oxidation-Reduction Titrations

Reaction	Indicator	Method
$KMnO_4$ + $Fe^{2+}$	Self-indicator (pink → colorless)	Permanganate titration
$K_2Cr_2O_7$ + $Fe^{2+}$	Diphenylamine sulfonate	Dichromate titration

SPM Exam Tips

For titration calculations:

Always write the balanced equation first

Use the mole ratio correctly

Remember the dilution formula $M_1V_1$ = $M_2V_2$

Practice with different types of titrations

Pay attention to significant figures

6.7 Salts, Crystals and Their Uses in Life

Salt Formation

Neutralization Reaction

$\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

Acid	Base	Salt	Name
HCl	NaOH	NaCl	Sodium chloride
$H_2SO_4$	NaOH	$Na_2SO_4$	Sodium sulfate
$HNO_3$	KOH	$KNO_3$	Potassium nitrate
$H_3PO_4$	$Ca(OH)_2$	$Ca_3(PO_4)_2$	Calcium phosphate

Double Displacement Reaction

$\text{Salt}_1 + \text{Salt}_2 \rightarrow \text{Salt}_3 + \text{Salt}_4$

Crystal Structures

Types of Crystals

Type	Structure	Properties	Examples
Ionic Crystals	Lattice structure	Hard, high melting point	NaCl, $CaCO_3$
Covalent Crystals	Network structure	Very hard, high melting point	Diamond, $SiO_2$
Metallic Crystals	Metallic bonding	Malleable, conductive	Fe, Cu, Al
Molecular Crystals	Molecules held by forces	Soft, low melting point	Ice, sugar

Common Crystals in Daily Life

Crystal	Formula	Uses	Properties
Sodium chloride	NaCl	Food seasoning, de-icing	Cubic crystals, soluble
Sugar (sucrose)	$C_{12}H_{22}O_{11}$	Sweetener, preservative	Hexagonal crystals, soluble
Copper sulfate	$CuSO_4$	Electroplating, fungicide	Blue crystals, soluble
Calcium carbonate	$CaCO_3$	Cement, antacid	Rhombohedral, insoluble
Sodium bicarbonate	$NaHCO_3$	Baking, fire extinguisher	Monoclinic, soluble

Uses of Salts and Crystals

Food Industry

Salt	Uses	Properties
NaCl	Seasoning, preservative	Enhances flavor, prevents spoilage
NaHCO_3	Baking, antacid	Leavening agent, neutralizes acid
NaNO_2	Curing meat	Preservative, color retention
CaCO_3	Calcium supplement	Source of calcium

Industrial Applications

Salt	Uses	Properties
Na_2CO_3	Glass manufacturing, detergent	Alkaline, cleaning
CaSO_4	Plaster of Paris, cement	Sets when hydrated
NH_4NO_3	Fertilizer, explosives	High nitrogen content
CuSO_4	Electroplating, fungicide	Blue color, toxic

Medical and Pharmaceutical

Salt	Uses	Properties
NaCl	IV fluids, irrigation	Isotonic, maintains osmotic balance
MgSO_4	Epsom salts, laxative	Muscle relaxant, osmotic laxative
KCl	Potassium supplement, electrolyte	Replaces potassium
CaCl_2	Calcium supplement, de-icer	Source of calcium

Did You Know?

Salt crystals are actually cube-shaped in their perfect form, but we often see them as irregular chunks. The cubic shape comes from the ionic bonds between sodium and chloride ions arranging themselves in a repeating three-dimensional pattern.

6.8 Qualitative Analysis

Purpose of Qualitative Analysis

Qualitative Analysis: Systematic identification of ions present in a substance.

Test for Cations

Group I Cations ( $Ag^+$ , $Pb^{2+}$ , $Hg_2^{2+}$ )

Ion	Test	Observation
$Ag^+$	Add HCl	White precipitate (AgCl)
$Pb^{2+}$	Add HCl	White precipitate ( $PbCl_2$ )
$Hg_2^{2+}$	Add HCl	White precipitate ( $Hg_2Cl_2$ )

Group II Cations ( $Cu^{2+}$ , $Fe^{2+}$ , $Fe^{3+}$ , $Al^{3+}$ , $Zn^{2+}$ )

Ion	Test	Observation
$Cu^{2+}$	Add NaOH	Blue precipitate
$Fe^{2+}$	Add NaOH	Green precipitate
$Fe^{3+}$	Add NaOH	Rust-brown precipitate
$Al^{3+}$	Add NaOH	White gelatinous precipitate
$Zn^{2+}$	Add NaOH	White precipitate (soluble in excess)

Group III Cations ( $Ca^{2+}$ , $Sr^{2+}$ , $Ba^{2+}$ )

Ion	Test	Observation
$Ca^{2+}$	Add $(NH_4)_2C_2O_4$	White precipitate
$Sr^{2+}$	Add $(NH_4)_2SO_4$	White precipitate
$Ba^{2+}$	Add $K_2CrO_4$	Yellow precipitate

Group IV Cations ( $Mg^{2+}$ , $NH_4^{+}$ )

Ion	Test	Observation
$Mg^{2+}$	Add NaOH + magneson	Blue precipitate
$NH_4^{+}$	Add NaOH + heat	Ammonia gas (pungent smell)

Test for Anions

Carbonate ( $CO_3^{2-}$ )

$\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{CO}_2 + \text{H}_2\text{O}$ Test: Add dilute acid → effervescence ( $CO_2$ turns limewater milky)

Sulfate ( $SO_4^{2-}$ )

$\text{SO}_4^{2-} + \text{Ba}^{2+} \rightarrow \text{BaSO}_4 \downarrow$ Test: Add $BaCl_2$ → white precipitate

Chloride ( $Cl^-$ )

$\text{Cl}^- + \text{Ag}^+ \rightarrow \text{AgCl} \downarrow$ Test: Add $AgNO_3$ → white precipitate (soluble in $NH_3$ )

Nitrate ( $NO_3^-$ )

Test: Brown ring test $\text{NO}_3^- + \text{Fe}^{2+} + 4\text{H}^+ \rightarrow \text{Fe}^{3+} + \text{NO} + 2\text{H}_2\text{O}$ $\text{NO} + \text{FeSO}_4 \rightarrow \text{[Fe(NO)]SO}_4 \text{ (brown ring)}$

Flame Tests

Ion	Flame Color
$Li^+$	Crimson red
$Na^+$	Golden yellow
$K^+$	Lilac
$Ca^{2+}$	Brick red
$Sr^{2+}$	Crimson red
$Ba^{2+}$	Apple green
$Cu^{2+}$	Blue-green

Systematic Analysis Procedure

Step 1: Preliminary Tests

Physical appearance
Solubility in water
Odor
pH

Step 2: Confirmatory Tests

Test for specific ions using reagents
Observe color changes and precipitates

Step 3: Identification

Compare observations with known reactions
Confirm presence of ions

SPM Exam Tips

For qualitative analysis:

Memorize the characteristic tests for each ion

Learn the color changes and observations

Understand the systematic approach to analysis

Practice flame test colors

Know the solubility rules

Laboratory Practical Exercise: Acid-Base Titration

Objective

To determine the concentration of an unknown acid solution using titration.

Materials Needed

Burette, pipette, conical flask
Standard NaOH solution (0.100 M)
Unknown HCl solution
Phenolphthalein indicator
Burette stand
White tile

Procedure

Rinse burette with standard NaOH solution
Fill burette with NaOH solution
Pipette 25.0 mL of unknown HCl into conical flask
Add 2-3 drops of phenolphthalein
Titrate until colorless → pink endpoint
Record volume used
Repeat for accuracy

Calculations

$M_{\text{HCl}} = \frac{M_{\text{NaOH}} \times V_{\text{NaOH}}}{V_{\text{HCl}}}$

Expected Outcomes

Skill in titration technique
Accuracy in concentration calculations
Understanding of acid-base stoichiometry

Summary

This chapter has covered the fundamental concepts of acids, bases, and salts:

pH and Acidity/Alkalinity: Understanding pH scale and water's role
Acid-Base Theories: Arrhenius, Brønsted-Lowry, and Lewis theories
Strength of Acids/Bases: Strong vs weak acids and bases
Chemical Properties: Reactions with metals, carbonates, and each other
Concentration: Molarity and dilution calculations
Titration: Technique for determining unknown concentrations
Salts and Crystals: Formation, properties, and uses
Qualitative Analysis: Systematic identification of ions

Mastering these concepts is essential for understanding chemical reactions and their applications in various fields.

Practice Tips for SPM Students

Practice pH calculations regularly

Work through titration problems step by step

Memorize qualitative tests for common ions
Create flashcards for acid-base reactions
Review laboratory applications and safety

Chapter 6: Acid, Base and Salt

Overview

Learning Objectives

6.1 The Role of Water in Showing Acidity and Alkalinity

Water as a Solvent

Ionization in Water

Self-Ionization of Water

Neutral Solution

6.2 pH Value

Definition of pH

pH Scale

pH Calculations

From [H⁺] to pH

From pH to [H⁺]

Relationship between pH and pOH

Common pH Values

Measuring pH

pH Indicators

pH Measurement Devices

6.3 Strength of Acids and Alkalis

Strong Acids

Common Strong Acids

Weak Acids

Common Weak Acids

Strong Bases

Common Strong Bases

Weak Bases

Common Weak Bases

Concentration vs. Strength

Acid Dissociation Constants (Ka)

Acid Strength Order

6.4 Chemical Properties of Acids and Bases

General Properties of Acids

Physical Properties

Chemical Properties

General Properties of Bases

Physical Properties

Chemical Properties

Acid-Base Theories

Arrhenius Theory

Brønsted-Lowry Theory

Lewis Theory

6.5 Concentration of Aqueous Solutions

Molarity

Dilution Formula

Concentration Calculations

Example 1: Preparing 0.1 M HCl

Example 2: Neutralization Reaction

6.6 Acid-Base Titration

Definition of Titration

Types of Titrations

Titration Procedure

Equipment Needed

Step-by-Step Procedure

Indicator Selection

Titration Calculations

Example: Titration of HCl with NaOH

Endpoint vs Equivalence Point

Common Titrations in Laboratory

Acid-Base Titrations

Oxidation-Reduction Titrations

6.7 Salts, Crystals and Their Uses in Life

Salt Formation

Neutralization Reaction

Double Displacement Reaction

Crystal Structures

Types of Crystals

Common Crystals in Daily Life

Uses of Salts and Crystals

Food Industry

Industrial Applications

Medical and Pharmaceutical

6.8 Qualitative Analysis

Purpose of Qualitative Analysis

Test for Cations

Group I Cations (Ag+Ag^+Ag+, Pb2+Pb^{2+}Pb2+, Hg22+Hg_2^{2+}Hg22+​)

Group II Cations (Cu2+Cu^{2+}Cu2+, Fe2+Fe^{2+}Fe2+, Fe3+Fe^{3+}Fe3+, Al3+Al^{3+}Al3+, Zn2+Zn^{2+}Zn2+)

Group III Cations (Ca2+Ca^{2+}Ca2+, Sr2+Sr^{2+}Sr2+, Ba2+Ba^{2+}Ba2+)

Group IV Cations (Mg2+Mg^{2+}Mg2+, NH4+NH_4^{+}NH4+​)

Test for Anions

Group I Cations ( $Ag^+$ , $Pb^{2+}$ , $Hg_2^{2+}$ )

Group II Cations ( $Cu^{2+}$ , $Fe^{2+}$ , $Fe^{3+}$ , $Al^{3+}$ , $Zn^{2+}$ )

Group III Cations ( $Ca^{2+}$ , $Sr^{2+}$ , $Ba^{2+}$ )

Group IV Cations ( $Mg^{2+}$ , $NH_4^{+}$ )

Carbonate ( $CO_3^{2-}$ )

Sulfate ( $SO_4^{2-}$ )

Chloride ( $Cl^-$ )

Nitrate ( $NO_3^-$ )