Grade 8
  1. Introduction to Chemistry  1.1. Nature and scope of chemistry  1.2. Importance of chemistry in daily life  1.3. Chemistry and its relation with other sciences  1.4. The role of chemistry in industry, medicine and the environment  1.5. Scientific investigation and experimental methods in chemistry
  2. Matter and its properties  2.1. Definition and classification of matter  2.2. Physical and chemical properties of matter  2.3. Law of conservation of matter  2.4. Extensive and Intensive Properties of Matter  2.5. Kinetic molecular theory of matter  2.6. States of matter: solids, liquids, gases, plasmas, and Bose-Einstein condensates  2.7. Changes in state and energy are involved  2.8. Diffusion and Brownian motion
  3. Elements, Compounds, and Mixtures  3.1. Definition and differences between elements, compounds, and mixtures  3.2. Atomic Structure and Atomic Number  3.3. Chemical symbols and formulas  3.4. Trends in the Periodic Table (Groups and Periods)  3.5. Classification of Elements: Metals, Nonmetals and Metalloids  3.6. Rare Earth Elements and Transition Metals  3.7. Types of mixtures: homogeneous and heterogeneous  3.8. Separation of Mixtures Using Physical and Chemical Methods
  4. Separation Techniques  4.1. Filtration, Sedimentation and Decantation  4.2. Evaporation and crystallization  4.3. Distillation and Fractional Distillation  4.4. Chromatography and its applications  4.5. Sublimation in the purification of compounds  4.6. The role of centrifugation in biochemical processes
  5. Atomic Structure  5.1. Dalton's atomic theory  5.2. Structure of the atom: protons, neutrons and electrons  5.3. Bohr's atomic model  5.4. Introduction to Quantum Mechanical Model  5.5. Electron Configuration and Energy Levels  5.6. Excitation and emission spectra  5.7. Isotopes and their applications
  6. Periodic Table and Chemical Trends  6.1. History and Development of the Periodic Table  6.2. Modern periodic law  6.3. Periodicity and trends in atomic and ionic sizes  6.4. Ionization Energy, Electron Affinity and Electronegativity  6.5. Introduction to Lanthanides and Actinides  6.6. Application of periodic trends in chemistry
  7. Chemical Bonding and Molecular Structure  7.1. Types of chemical bonds: ionic, covalent and metallic bonds  7.3. Polar and Nonpolar Molecules  7.4. Hydrogen bonding and its applications  7.5. Intermolecular forces: dipole-dipole, London dispersion force
  8. Chemical Reactions and Stoichiometry  8.1. Chemical Equations and Balance  8.2. Types of Chemical Reactions: Combination, Decomposition, Displacement and Redox  8.3. Introduction to stoichiometry and the mole concept  8.4. Limiting reactant in chemical equations  8.5. Reaction rate, catalyst, and factors affecting reaction rate
  9. Acids, Bases and Salts  9.1. Definition and Properties of Acids and Bases  9.2. Strong vs. Weak Acids and Bases  9.3. pH Scale and pH Calculation  9.4. Neutralization reactions  9.5. Buffer solutions and their importance  9.6. Applications of Acids, Bases and Salts in Daily Life
  10. Solutions and Solubility  10.1. Definition of Solution, Solute, and Solvent  10.2. Factors Affecting Solubility  10.3. Concentration units: molarity and molality  10.4. Saturated, unsaturated and supersaturated solutions  10.5. Concept of osmosis and reverse osmosis  10.6. Concentrative properties of solutions
  11. Gases and Gas Laws  11.1. Properties of gases  11.2. Introduction to Ideal and Real Gases  11.3. Boyle's Law, Charles' Law and Avogadro's Law  11.4. Ideal Gas Equation and Its Applications  11.5. Application of Gas Laws in Real Life
  12. Thermochemistry and energy transformation  12.1. Energy in Chemical Reactions  12.2. Exothermic vs. Endothermic Reactions  12.3. Heat and Enthalpy Changes  12.4. Calorimetry and heat transfer in reactions
  13. Metals and Nonmetals  13.1. Physical and Chemical Properties of Metals and Nonmetals  13.2. Reactivity Series of Metals  13.3. Corrosion and its prevention  13.4. Extraction of metals from ores  13.5. Uses of metals and non-metals in daily life
  14. Introduction to Organic Chemistry  14.1. Characteristics of carbon and organic compounds  14.2. Functional groups and their importance  14.3. Simple Hydrocarbons: Alkenes, Alkenes and Alkynes  14.4. Isomerism in organic compounds  14.5. Introduction to polymers and their uses
  15. Environmental Chemistry and Sustainability  15.1. Green Chemistry Principles  15.2. Effects of Chemical Pollution on Ecosystems  15.3. Acid rain and its effects  15.4. Ozone layer depletion and global warming  15.5. Waste Management and Recycling in Chemistry

Grade 8 ↓


Thermochemistry and energy transformation


Thermochemistry is the study of temperature changes that occur during chemical reactions. It is a branch of chemistry that combines the principles of thermodynamics and chemistry to understand how and why energy is transferred during chemical processes.

What is energy?

To understand thermochemistry, we first need to know what energy is. Energy is the capacity to do work or produce heat. It exists in various forms such as chemical energy, kinetic energy, potential energy, and thermal energy. In the context of thermochemistry, we are most concerned with chemical and thermal energy.

Types of energy

Let's look at some of the main types of energy:

  • Chemical energy: It is the energy stored in the bonds of chemical compounds such as molecules and atoms. During a chemical reaction, bonds are broken, and new bonds are formed, which can either absorb or release energy.
  • Thermal energy: This is energy that comes from heat. It is related to the temperature of a system, which is a measure of the kinetic energy or speed of its particles.
  • Kinetic energy: It is the energy of motion. When an object moves, it has kinetic energy. For example, a moving car or a person running.
  • Potential energy: This is the energy stored in an object because of its position. This can be gravitational potential energy, like a book on a shelf, or elastic potential energy, like a stretched spring.

Conservation of energy

The law of conservation of energy states that energy cannot be created or destroyed. It can only be transferred from one form to another. This means that the total energy of an isolated system remains constant.

For example, when you light a candle, the chemical energy in the wax is converted into heat energy (thermal energy) and light energy. Here's a simple visual example:

    Wax (chemical energy) --> Heat (thermal energy) + Light (light energy)
    Wax (chemical energy) --> Heat (thermal energy) + Light (light energy)

Exothermic and endothermic reactions

Chemical reactions can be classified into exothermic and endothermic reactions based on the energy changes that occur:

Exothermic reactions

Exothermic reactions are reactions that release energy in the form of heat or light. In these reactions, the energy of the products is less than the energy of the reactants. The excess energy is released into the surrounding environment, often causing the environment to feel warmer. Examples include:

  • Combustion of fuels, such as wood or gasoline.
  • Respiration in living organisms.
  • Reaction of acids with bases.

Visual example of an exothermic reaction:

    Reactants (high energy) --> Products (low energy) + Energy (heat/light)
    Reactants (high energy) --> Products (low energy) + Energy (heat/light)

Endothermic reactions

Endothermic reactions are reactions that absorb energy from the surrounding environment. This means that the energy of the products is greater than the energy of the reactants. The surrounding environment may seem colder after the reaction. Examples include:

  • Photosynthesis in plants.
  • Dissolving ammonium nitrate in water.
  • Melting ice cubes.

Visual example of an endothermic reaction:

    Reactants (low energy) + Energy (heat) --> Products (high energy)
    Reactants (low energy) + Energy (heat) --> Products (high energy)

Measuring energy changes

Energy changes in chemical reactions can be measured using a calorimeter. A calorimeter is an instrument that measures the amount of heat absorbed or released during a chemical reaction. The most commonly used unit of energy is the joule (J), but the calorie (cal) is also used.

Enthalpy

In thermochemistry, the term enthalpy (H) is used to describe the total energy of a system, including internal energy and the energy that can be exchanged with the surroundings (usually as a pressure–volume function).

The change in enthalpy (ΔH) is a measure of the heat change during a chemical reaction at constant pressure. It can be calculated using:

    ΔH = H_products - H_reactants
    ΔH = H_products - H_reactants

If ΔH is negative, the reaction is exothermic. If ΔH is positive, the reaction is endothermic.

Example calculation using enthalpy

Let's work out an example of calculating the enthalpy change for a reaction:

Suppose you have a reaction in which 2 moles of hydrogen gas react with 1 mole of oxygen gas to produce 2 moles of water vapor. The enthalpy change (ΔH) for this reaction is -483.6 kJ.

    2 H 2 (g) + O 2 (g) --> 2 H 2 O(g) ΔH = -483.6 kJ
    2 H 2 (g) + O 2 (g) --> 2 H 2 O(g) ΔH = -483.6 kJ

This means that when the reaction occurs, 483.6 kJ of energy is released into the surrounding environment.

Entropy

Entropy is a measure of the disorder or irregularity in a system. Generally, gases have higher entropy than liquids, and liquids have higher entropy than solids. Reactions that increase the disorder of a system tend to occur spontaneously.

Spontaneity of reactions

A spontaneous reaction is one that occurs naturally without the need for external energy. Spontaneity depends on both the change in enthalpy and the change in entropy. We use a quantity called Gibbs free energy (G) to measure spontaneity:

    ΔG = ΔH - TΔS
    ΔG = ΔH - TΔS

In this equation, ΔG is the change in Gibbs free energy, ΔH is the change in enthalpy, T is the temperature in Kelvin, and ΔS is the change in entropy. If ΔG is negative, the reaction is spontaneous.

Visual example using the reaction coordinate diagram

The reaction coordinate diagram can help depict energy changes during a reaction:

    Reaction Pathway: /-------- /  / (Products) / / / / ______/ (Reactants)/
    Reaction Pathway: /-------- /  / (Products) / / / / ______/ (Reactants)/

In this diagram:

  • The y-axis represents energy.
  • The x-axis represents the reaction progress.
  • The peak of the curve represents the energy of the intermediate state, called the transition state.

Exploring specific examples

Combustion of methane

Combustion is a common example of an exothermic reaction. When methane (CH_4) burns, it reacts with oxygen (O_2) to form carbon dioxide (CO_2) and water (H_2O), releasing energy:

    CH 4 + 2 O 2 --> CO 2 + 2 H 2 O ΔH = -890 kJ/mol
    CH 4 + 2 O 2 --> CO 2 + 2 H 2 O ΔH = -890 kJ/mol

This reaction releases 890 kJ of energy per mole of methane, indicating a significantly exothermic process.

Photosynthesis

Photosynthesis is an example of an endothermic reaction. Plants use sunlight to convert carbon dioxide and water into glucose and oxygen. The equation is:

    6 CO 2 + 6 H 2 O + light energy --> C 6 H 12 O 6 + 6 O 2
    6 CO 2 + 6 H 2 O + light energy --> C 6 H 12 O 6 + 6 O 2

In this process, energy from sunlight is absorbed to drive the reaction forward, making it endothermic.

Conclusion

Thermochemistry allows us to understand the energy changes in chemical reactions. By studying these changes, we can better understand how energy is transferred and transformed in various processes, from the burning of fuels to how plants produce their food. Concepts such as exothermic and endothermic reactions, enthalpy, and Gibbs free energy are essential to this understanding and are fundamental components of chemistry.


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Thermochemistry and energy transformation

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