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 ↓


Stoichiometry and Chemical Calculations


Stoichiometry is one of the fundamental concepts of chemistry. It involves the quantitative relationship of reactants and products in chemical reactions. This subject helps us understand how much of a chemical substance is involved or produced in a given reaction. Stoichiometry is like a chemistry recipe, ensuring that the correct proportions of ingredients are used to obtain the desired product.

Understanding chemical equations

To fully understand stoichiometry, it is important to first understand chemical equations. A chemical equation represents a chemical reaction using symbols and formulas. The substances that react are called reactants, and the substances that are produced are called products.

2 H 2 + O 2 → 2 H 2 O (Water formation)

The equation above shows that two molecules of hydrogen gas react with one molecule of oxygen gas to form two molecules of water. It is important to understand that chemical equations need to be balanced, which means that there must be the same number of atoms of each element on both sides of the equation.

Law of conservation of mass

Balancing chemical equations is guided by the law of conservation of mass, which states that mass is neither created nor destroyed in a chemical reaction. Therefore, the total mass of the reactants must equal the total mass of the products.

Balancing chemical equations

Balancing a chemical equation means adjusting the coefficients (the numbers in front) of chemical formulas so that there are equal numbers of each type of atom on both sides of the equation. For example, consider the combustion of methane:

CH 4 + O 2 → CO 2 + H 2 O

Balancing the equation:

  1. Balance the carbon atoms: CH4 has 1 carbon. CO2 must have 1 as well, so there is already a balance for the carbon.
  2. Balance the hydrogen atoms: CH4 has 4 hydrogens. H2O has 2 hydrogens, so we need 2 H2O (2 x 2 = 4 hydrogens).
  3. Balance the oxygen atoms: On the right side, there are 2 oxygens in CO 2 and 2 oxygens in H 2 O, so a total of 4 oxygens are needed on the left side. Thus, use 2 O 2 molecules.
CH 4 + 2 O 2 → CO 2 + 2 H 2 O

This equation is now balanced.

Mole concept in stoichiometry

The mole is a central unit in chemistry that serves as a bridge between the atomic and macroscopic worlds. A mole is equal to Avogadro's number, which is approximately 6.022 x 1023, indicating the number of atoms or molecules.

Example: moles of water

How many moles are there in 36 grams of water (H2O)?

  • The molar mass of H2O is 18 g/mol (2 g/mol for H and 16 g/mol for O).
  • Use the formula: moles of H 2 O = mass (g) ÷ molar mass (g/mol).
Moles of H 2 O = 36 g ÷ 18 g/mol = 2 moles

Stoichiometric calculations

Stoichiometric calculations use balanced chemical equations to find the amounts of reactants needed or products formed.

Example

Using the balanced equation for the combustion of methane:

CH 4 + 2 O 2 → CO 2 + 2 H 2 O

If you have 5 moles of CH4, how many moles of CO2 and water will be produced?

  • The equation shows a 1:1 ratio of CH4 to CO2.
  • Thus, 5 moles of CH4 will produce 5 moles of CO2.
  • The ratio of water (H 2 O) is 1:2, which gives 10 moles of water.

Practical example: The baking analogy

Understand stoichiometry in terms of baking cookies. If your recipe calls for 2 cups of flour for every dozen cookies, and you want to make 3 dozen cookies, you know you'll need 6 cups of flour. Chemical reactions work similarly, using balanced equations to ensure the correct proportions of reactants.

Limiting and excess reactants

In chemical reactions, the limiting reactant is the substance that is completely consumed when the chemical reaction is complete. Without it, the reaction cannot proceed. Other reactants that are not used up are called excess reactants.

Example: making a sandwich

If needed to make a sandwich:

2 slices of bread + 1 slice of cheese → 1 sandwich

And you have 10 slices of bread and 4 slices of cheese:

  • You can make 4 sandwiches, because cheese sets your limit (limiting reactant).
  • This will leave you with two slices of bread, which will become additional reactants.

Yield in chemical reactions

Yield measures the amount of product formed in a reaction. The theoretical yield is the maximum amount that is expected, while the actual yield is what you actually get. Percent yield shows how efficient a reaction is:

Percent Yield = (Actual Yield / Theoretical Yield) x 100%

Example

If the theoretical yield of a reaction is 20 g, but the actual yield is 15 g:

Percent Yield = (15 g / 20 g) x 100% = 75%

Conclusion

Stoichiometry is a vital tool for scientists and industry, ensuring that reactions are efficient and that resources are used effectively. By understanding moles, balancing equations, and identifying limiting reactants, we can accurately predict and measure the outcomes of chemical reactions.


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Stoichiometry and Chemical Calculations

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