Grade 9
  1. Matter and its nature  1.1. Introduction to matter  1.2. Physical and chemical properties of matter  1.3. States of matter  1.3.1. Solid state  1.3.2. Fluid state  1.3.3. Gaseous state  1.3.4. Plasma and Bose–Einstein condensate  1.4. Changes in the states of matter  1.4.1. Melting and freezing  1.4.2. Evaporation and Condensation  1.4.3. Sublimation and deposition  1.5. Classification of matter  1.6. Elements, Compounds, and Mixtures  1.7. Pure substances and impurities  1.8. Separation Techniques  1.8.1. Filtration  1.8.2. Distillation  1.8.3. Crystallization  1.8.4. Chromatography  1.8.5. Centrifugation  1.8.6. Sublimation  1.8.7. Decantation and sedimentation
  2. Atomic Structure  2.1. Discovery of atoms  2.2. Dalton's atomic theory  2.3. Structure of the atom  2.4. Subatomic particles  2.5. Atomic number and mass number  2.6. Isotopes and Isobars  2.7. Electronic configuration  2.8. Bohr's atomic model  2.9. Modern Atomic Model (Quantum Mechanical Model - Introduction)  2.10. Valency and chemical bonding
  3. C hemical reactions and equations  3.1. Introduction to Chemical Reactions  3.2. Chemical Equation Formulation  3.3. Balancing Chemical Equations  3.4. Types of Chemical Reactions  3.4.1. Combination Reactions  3.4.2. Decomposition reactions  3.4.3. Displacement reactions  3.4.4. Double displacement reactions  3.4.5. Redox reactions  3.4.6. Endothermic and Exothermic Reactions  3.5. Effects of chemical reactions  3.6. Energy Changes in Chemical Reactions  3.7. Catalysts and their role in reactions
  4. Periodic table and periodicity  4.1. History of Periodic Classification  4.2. Mendeleev's periodic table  4.3. Modern Periodic Table  4.4. Periods and groups in the periodic table and periodicity  4.5. Trends in the Periodic Table  4.5.1. Atomic size  4.5.2. Ionization Energy  4.5.3. Electron affinity  4.5.4. Electronegativity  4.5.5. Metallic and Non-Metallic Properties  4.6. Noble gases and their properties
  5. Chemical bond  5.1. Introduction to Chemical Bonding  5.2. Types of chemical bonds  5.2.1. Ionic bond  5.2.2. Covalent bond  5.2.3. Metallic bond  5.2.4. Hydrogen bonding  5.2.5. Van der Waals force  5.3. Properties of Ionic and Covalent Compounds  5.4. Molecular Structure and Geometry (VSEPR Theory - Basics)
  6. Acids, Bases and Salts  6.1. Introduction to Acids and Bases  6.2. Properties of Acids  6.3. Properties of bases  6.4. pH scale and its importance  6.5. Indicators and their uses  6.6. Neutralization reactions  6.7. Strength of Acids and Bases (Strong vs. Weak)  6.8. Preparation of salts  6.9. Uses of acids, bases and salts in daily life
  7. Metals and Nonmetals  7.1. Physical Properties of Metals  7.2. Physical Properties of Nonmetals  7.3. Chemical properties of metals  7.4. Chemical Properties of Nonmetals  7.5. Reactivity Series of Metals  7.6. Extraction of metals  7.7. Corrosion and its prevention  7.8. uses of metals and nonmetals
  8. Carbon and its compounds  8.1. Introduction to Carbon  8.2. Allotropes of carbon (diamond, graphite, fullerene)  8.3. Hydrocarbons  8.3.1. Hydrocarbons  8.3.2. Alkene  8.3.3. Alkynes  8.4. Functional groups in organic chemistry  8.5. Isomerism in organic compounds  8.6. Alcohols and Carboxylic Acids  8.7. Soaps and detergents  8.8. Polymers (Introduction to Natural and Synthetic Polymers)
  9. Solutions and Mixtures  9.1. Types of mixtures  9.2. Homogeneous and Heterogeneous Mixtures  9.3. Concentration of solutions  9.4. Solubility and factors affecting solubility  9.5. Suspensions and Colloids  9.6. Saturated, unsaturated and supersaturated solutions
  10. Air and atmosphere  10.1. Composition of air  10.2. Role of gases in the atmosphere  10.4. Acid rain and its effects  10.5. Greenhouse effect and global warming  10.6. Ozone layer and its depletion  10.7. Air pollution control measures
  11. Water and its importance  11.1. Composition and properties of water  11.2. Hardness of water (temporary and permanent hardness)  11.3. Causes of water pollution  11.4. Effects of Water Pollution  11.5. Water purification methods (filtration, boiling, chlorination)
  12. Industrial Chemistry  12.1. Introduction to Industrial Chemistry  12.2. Important industrial chemicals (sulfuric acid, ammonia, sodium hydroxide)  12.3. Chemical processes in industry  12.4. Uses of Industrial Chemistry in Daily Life  12.5. Environmental impact of industrial chemical processes

Grade 9 → C hemical reactions and equations ↓


Chemical Equation Formulation


Understanding chemical equations is a fundamental part of chemistry. They allow us to understand and predict the outcome of chemical reactions. In simple terms, a chemical equation is a symbolic representation of a chemical reaction where the reactants are on the left, the products are on the right, and an arrow separates the two sides. Let's dive into a detailed exploration of chemical equation representation.

Basic structure of chemical equations

A chemical equation shows the process of a chemical reaction using the chemical formulas of the substances involved. On the left side of the equation, we have the reactants, which are the substances that undergo a change. On the right side, we have the products, which are the new substances formed as a result of the reaction. For example:

 2H 2 + O 2 → 2H 2 O

In this equation, hydrogen gas (H2) and oxygen gas (O2) are the reactants, and water (H2O) is the product.

Components of a chemical equation

1. Reactants and products: These are the substances that are consumed and produced in a chemical reaction. 2. Coefficients: Numbers placed before chemical formulas to balance the equation and indicate the relative amounts of each substance. In the above equation, the coefficient "2" before H 2 and H 2 O means that the reaction involves two molecules of hydrogen gas and two molecules of water. 3. Phases of matter: Often, the physical states of reactants and products are indicated by symbols in parentheses: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution. Example: 

 NaCl(s) + AgNO 3 (aq) → NaNO 3 (aq) + AgCl(s)
In this reaction, solid sodium chloride reacts with aqueous silver nitrate, forming aqueous sodium nitrate and solid silver chloride.

Balancing chemical equations

Balancing chemical equations is essential because it respects the law of conservation of mass. This law states that matter cannot be created or destroyed in a normal chemical reaction. Therefore, a chemical equation must have the same number of each type of atom on both sides of the equation.

Steps to balance a chemical equation

Let's learn the steps to balance a chemical reaction using a simple example. Example reaction: Combustion of methane in oxygen: 

 CH 4 + O 2 → CO 2 + H 2 O

  1. Write the unbalanced equation: Write the skeletal equation of the chemical reaction with the correct formulas of all reactants and products.
  2. Count the number of atoms of each element: Count the number of atoms of each element on both sides of the equation.
    • Reactants: C=1, H=4, O=2
    • Product: C=1, H=2, O=3
  3. Add coefficients: Adjust the coefficients to get the same number of atoms of each element on both sides. Coefficients are placed in front of formulas. 
     CH 4 + 2O 2 → CO 2 + 2H 2 O
  4. Verify: Make sure the number of atoms for each element is balanced.
    • Reactants: C=1, H=4, O=4
    • Product: C=1, H=4, O=4

Balanced Reaction: 

 CH 4 + 2O 2 → CO 2 + 2H 2 O

Types of chemical reactions

There are many types of chemical reactions, each with its own characteristics and patterns. Here are some common types:

1. Combination (synthesis) reaction

In this type of reaction, two or more reactants combine to form a single product. It can be represented as: 

 A + B → AB
Example: 
 2H 2 (g) + O 2 (g) → 2H 2 O(l)

2. Decomposition reaction

A single compound breaks down into two or more simpler substances. General form: 

 AB → A + B
Example: 
 2HgO(s) → 2Hg(l) + O 2 (g)

3. Single displacement (substitution) reaction

In this reaction, one element displaces another element in a compound: 

 a + bc → ac + b
Example: 
 Zn(s) + 2HCl(aq) → ZnCl 2 (aq) + H 2 (g)

4. Double displacement (metathesis) reaction

Two compounds exchange ions to form two new compounds: 

 AB + CD → AD + CB
Example: 
 AgNO 3 (aq) + NaCl(aq) → AgCl(s) + NaNO 3 (aq)

5. Combustion reaction

It is a reaction in which a substance combines with oxygen, releasing a large amount of energy in the form of light and heat. Organic compounds are often burned to produce energy. Examples: 

 C 3 H 8 + 5O 2 → 3CO 2 + 4H 2 O
This is the combustion of propane.

Representation of chemical reactions: Visual example

For a clearer understanding of how atoms and molecules participate in a chemical reaction, it is often helpful to represent reactions with diagrams showing molecules as collections of atoms. Here's a simple example:

H2 , O2 H2O ,

Here, hydrogen gas and oxygen gas react to form water, which captures the essence of rearranging the bonds of molecules during a chemical reaction.

Conclusion

Understanding chemical equations is important for predicting the substances consumed and created in reactions. Mastering the skill of balancing chemical equations helps to effectively understand the nature of these reactions, which is essential for studying and exploring chemical processes.


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Chemical Equation Formulation

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