Grade 10
  1. Matter and its properties  1.1. States of matter  1.2. Physical and chemical properties of matter  1.3. Classification of substances (elements, compounds, mixtures)  1.4. Changes in Matter (Physical and Chemical Changes)  1.5. Law of conservation of mass and law of definite proportions
  2. Atomic Structure  2.1. Historical development of the atomic model  2.2. Subatomic particles (protons, neutrons, electrons)  2.3. Atomic number, mass number and isotopes  2.4. Electronic Configuration and Energy Levels  2.5. Valency, ion formation and oxidation state
  3. Periodic table  3.1. History and Development of the Periodic Table  3.2. Modern Periodic Law and Periodic Trends  3.3. Groups and Periods in the Periodic Table  3.4. Trends in the Periodic Table  3.5. Metals, non-metals, metalloids and their properties  3.6. Noble gases and their chemical inertness
  4. Chemical bond  4.1. Formation of Chemical Bonds (Octet Rule)  4.2. Ionic Bonding and Properties of Ionic Compounds  4.3. Covalent Bond and Properties of Covalent Compounds  4.4. Metallic bond and its properties  4.5. Molecular geometry and VSEPR theory  4.6. Polarity and dipole moment of molecules  4.7. Intermolecular forces
  5. Chemical Reactions and Equations  5.1. Types of Chemical Reactions  5.2. Writing and balancing chemical equations  5.3. Law of conservation of mass in chemical reactions  5.4. Factors Affecting Chemical Reactions  5.5. Catalysts and their role in chemical reactions  5.6. Exothermic and Endothermic Reactions
  6. Stoichiometry and Chemical Calculations  6.1. Mole concept and Avogadro number  6.2. Molar Mass and Gram-Mole Conversions  6.3. Empirical and molecular formula  6.4. Stoichiometric calculations and chemical yield  6.5. Limiting and Excess Reactants
  7. Acids, Bases and Salts  7.1. Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)  7.2. pH Scale, pOH, and Indicators  7.3. Neutralization reactions and salt formation  7.4. Strength of Acids and Bases (Strong vs. Weak Acids and Bases)  7.5. Common Acids and Bases in Daily Life  7.6. Salts, their solubility and applications
  8. Metals and Nonmetals  8.1. Physical and chemical properties of metals  8.2. Physical and Chemical Properties of Non-Metals  8.3. Reactivity Series of Metals and Displacement Reactions  8.4. Extraction of metals from ores  8.5. Corrosion, its causes and prevention  8.6. Alloys, their composition and uses
  9. Carbon and its compounds  9.1. Allotropes of Carbon  9.2. Hydrocarbons and their classification (alkanes, alkenes, alkynes)  9.3. Functional groups in organic compounds  9.4. Nomenclature of organic compounds (IUPAC system)  9.5. Isomerism in organic chemistry (structural and geometrical)  9.6. Important organic reactions (combustion, addition, substitution, polymerization)  9.7. Applications of organic compounds in industry and daily life
  10. Environmental Chemistry  10.1. Air Pollution (Sources, Effects and Control)  10.2. Water Pollution (Causes, Effects and Remedies)  10.3. Greenhouse effect and global warming  10.4. Ozone layer depletion and its effects  10.5. Green Chemistry and Its Applications  10.6. sustainable chemistry practice
  11. Thermochemistry  11.1. Heat and Energy Changes in Chemical Reactions  11.2. Exothermic and Endothermic Reactions  11.3. Enthalpy, enthalpy change, and heat of reaction  11.4. Specific heat capacity and calorimetry  11.5. Energy Changes in Combustion Reactions
  12. Electrochemistry  12.1. Electrolytes and non-electrolytes  12.2. Electrolysis and its applications (electroplating, extraction of metals)  12.3. Redox reactions (oxidation and reduction)  12.4. Galvanic cell and electrochemical series  12.5. Corrosion and electrochemical prevention methods
  13. Chemical kinetics and equilibrium  13.1. Rate of chemical reactions and its measurement  13.2. Factors affecting the reaction rate (catalyst, temperature, concentration)  13.3. Reversible and irreversible reactions  13.4. Dynamic equilibrium and equilibrium constant  13.5. Le Chatelier's principle and its applications
  14. Gases and Gas Laws  14.1. Properties of gases and kinetic molecular theory  14.2. Boyle's Law (Pressure-Volume Relation)  14.3. Charles's Law  14.4. Gay-Lussac's Law  14.5. Combined Gas Law and Ideal Gas Law  14.6. Real Gases vs Ideal Gases
  15. Nuclear Chemistry  15.1. Introduction to Radioactivity and Nuclear Reactions  15.2. Types of radioactive decay (alpha, beta, gamma)  15.3. Half-life and applications of radioactive isotopes  15.4. Nuclear fission and fusion reactions  15.5. Nuclear power generation and its environmental impact

Grade 10 → Acids, Bases and Salts ↓


Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)


Understanding acids and bases is fundamental in chemistry. Acids and bases have unique properties that make them important in a variety of chemical reactions and applications. In this comprehensive guide, we will explore the three major theories that define acids and bases: Arrhenius, Bronsted-Lowry, and Lewis.

Arrhenius theory

In 1887, Svante Arrhenius proposed one of the first modern definitions of acids and bases. According to Arrhenius, an acid is a substance that increases the concentration of hydrogen ions, H +, in an aqueous solution, while a base increases the concentration of hydroxide ions, OH -.

Properties of Arrhenius acid

  • They taste sour. For example, lemons contain citric acid.
  • They turn blue litmus paper red.
  • They react with metals such as zinc to produce hydrogen gas.
  • They increase the concentration of H + ions in water.

Properties of Arrhenius bases

  • They taste bitter and are slippery, like soap.
  • They turn red litmus paper blue.
  • They increase the concentration of OH - ions in water.

Example

HCl (aq) → H + (aq) + Cl - (aq) NaOH (aq) → Na + (aq) + OH - (aq)
Acid: HCl Base: NaOH

Bronsted–Lowry theory

Johannes Bronsted and Thomas Lowry independently proposed a more general theory in 1923. In the Bronsted-Lowry theory, an acid is a proton donor, and a base is a proton acceptor. This theory expands on Arrhenius by including reactions outside of aqueous solution.

Acid-base reaction examples

NH 3 + H 2 O ⇌ NH 4 + + OH -

In this reaction, water acts as a Bronsted-Lowry acid by donating a proton to ammonia (NH 3), forming an ammonium ion (NH 4 +) and a hydroxide ion (OH -).

Visual representation

H2O NH 3

Lewis theory

Gilbert N. Lewis proposed an even more comprehensive theory in 1923. According to Lewis, acids are electron pair acceptors, while bases are electron pair donors. This definition incorporates all the previous definitions and emphasizes the transfer of electron pairs.

Example of a Lewis acid-base reaction

BF 3 + NH 3 → F 3 B←NH 3

In this reaction, boron trifluoride (BF 3) acts as a Lewis acid by accepting an electron pair from ammonia (NH 3), a Lewis base, forming a coordinate covalent bond.

Visual representation

bf 3 NH 3

Comparison of theories

Each principle of acids and bases enhances our understanding of chemical reactions:

  • The Arrhenius theory is limited to aqueous solutions and involves basic neutralization reactions.
  • The Brønsted–Lowry theory covers a broad range of acidity and basicity by defining acids and bases via proton transfer.
  • Lewis theory covers all acids and bases because it considers the transfer of electron pairs, thus explaining many other reactions, including those that do not occur in aqueous solution.

Here is a summary table that shows the various definitions and their examples:

written Acid Base Example response
Arrhenius generates H + produces OH- HCl + NaOH → NaCl + H 2 O
Bronsted-Lowry Proton donor Proton acceptor NH 3 + H 2 O ⇌ NH 4 + + OH -
Lewis electron pair acceptor electron pair donor BF 3 + NH 3 → F 3 B←NH 3

Applications of acids and bases

Acids and bases play important roles in a variety of industrial, biological and environmental processes:

  • In industrial applications, sulfuric acid (H 2 SO 4 is used in the manufacture of fertilizers and chemicals.
  • In biological systems, hydrochloric acid (HCl) in the stomach aids digestion.
  • In environmental science, understanding acid rain caused by sulfuric and nitric acids helps solve pollution problems.

Industrial example

2 NH 3 + H 2 SO 4 → (NH 4 ) 2 SO 4

This equation shows the production of ammonium sulfate, a common fertilizer.

Example from everyday life

NaHCO 3 + CH 3 COOH → CO 2 + H 2 O + CH 3 COONa

In this reaction, baking soda reacts with vinegar to produce carbon dioxide gas—a common school experiment to create a volcano with bubbling lava.

Conclusion

Understanding acids and bases through various principles enhances scientific progress and our ability to predict and control various chemical reactions essential to everyday life. From manufacturing to biology and environmental science, the principles of acids and bases are important for many applications in various fields.


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Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)

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