Lesson Notes By Weeks and Term v5 - Grade 10

Lesson Video

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Performance objectives

• Differentiate between physical and chemical changes.
• Identify reactants and products in a reaction.
• Represent reactions using balanced chemical equations.
• Explain and apply the law of conservation of mass.
• Describe the evidence of a chemical change.

Lesson summary

Welcome to Week 5 of Physical Sciences! This week, we delve into the fascinating world of chemical change. Understanding chemical changes is crucial, not just in the lab, but in everyday life. From the food we digest to the fuel that powers our cars, chemical reactions are constantly at work. In South Africa, understanding these changes is particularly relevant to industries like mining (e.g., refining metals) and agriculture (e.g., fertilizer production). Ignoring the principles of chemical change can lead to environmental disasters, health risks, and economic losses. This topic lays the foundation for more advanced chemistry topics in the coming years.

Lesson notes

2.1 Physical vs. Chemical Changes The fundamental difference between physical and chemical changes lies in whether the composition of the substance changes.

Physical Change: A physical change alters the form or appearance of a substance, but not its chemical composition. The substance is still fundamentally the same. These changes are often reversible.

Examples include: Melting ice (H₂O(s) → H₂O(l)) - Water is still water, just in a different state. Boiling water (H₂O(l) → H₂O(g)) - Again, water remains water. Dissolving sugar in water (C₁₂H₂₂O₁₁(s) → C₁₂H₂₂O₁₁(aq)) - The sugar molecules are dispersed in the water, but are still sugar. Crushing a rock - The rock is in smaller pieces, but its chemical composition is unchanged.

Chemical Change: A chemical change involves the formation of new substances with different chemical compositions and properties. This occurs through a chemical reaction where existing bonds are broken and new bonds are formed. Chemical changes are typically irreversible (though the reverse reaction may be possible under different conditions).

Examples include: Burning wood (complex organic molecules + O₂ → CO₂ + H₂O + ash) - Wood is converted into entirely different substances: carbon dioxide, water, and ash. Rusting of iron (Fe(s) + O₂(g) → Fe₂O₃(s)) - Iron reacts with oxygen to form iron oxide (rust), a completely new substance. Cooking an egg (proteins denature and coagulate) - The proteins in the egg undergo a change in structure, forming a solid. Photosynthesis (CO₂ + H₂O → C₆H₁₂O₆ + O₂) - Plants convert carbon dioxide and water into glucose and oxygen.

Evidence of a Chemical Change: Several observations can indicate that a chemical change has taken place: Change in colour: A new colour appears (e.g., rust forming on iron).

Formation of a precipitate: A solid forms when two solutions are mixed (e.g., mixing silver nitrate and sodium chloride solutions, forming silver chloride, a white precipitate).

Gas production: Bubbles are produced (e.g., reacting an acid with a metal).

Change in temperature: Heat is either released (exothermic reaction, e.g., burning) or absorbed (endothermic reaction, e.g., dissolving ammonium nitrate in water).

Change in odour: A new smell is detected.

Light production: Light is emitted (e.g., burning). 2.2 Representing Chemical Change: Chemical Equations A chemical equation is a symbolic representation of a chemical reaction. It shows the reactants (the substances that react) on the left-hand side and the products (the substances formed) on the right-hand side, separated by an arrow (→). Reactants → Products

Example: The reaction of methane (CH₄) with oxygen (O₂) to produce carbon dioxide (CO₂) and water (H₂O) can be represented as: CH₄ + O₂ → CO₂ + H₂O (This equation is unbalanced.)

State Symbols: State symbols indicate the physical state of each substance: (s) - solid (l) - liquid (g) - gas (aq) - aqueous (dissolved in water) Adding state symbols to the previous reaction: CH₄(g) + O₂(g) → CO₂(g) + H₂O(g) (Still unbalanced.) 2.3 Balancing Chemical Equations The Law of Conservation of Mass states that matter cannot be created or destroyed in a chemical reaction. This means the number of atoms of each element must be the same on both sides of the chemical equation. Balancing a chemical equation ensures that this law is obeyed.

Steps to Balancing Chemical Equations: Write the unbalanced equation (as shown above). Count the number of atoms of each element on both sides of the equation. Start balancing by adding coefficients (numbers in front of the chemical formulas) to the side with fewer atoms of a particular element. It is usually best to start with elements that appear only once on each side of the equation. Adjust the coefficients until the number of atoms of each element is the same on both sides. Check your work to make sure the equation is balanced. Use the smallest whole number coefficients possible.

Example: Balancing the combustion of methane: Unbalanced equation: CH₄(g) + O₂(g) → CO₂(g) + H₂O(g)

Counting atoms: Left side: 1 C, 4 H, 2 O Right side: 1 C, 2 H, 3 O Balancing hydrogen: We need 4 H on the right side, so add a coefficient of 2 in front of H₂O: CH₄(g) + O₂(g) → CO₂(g) + 2H₂O(g)

Counting atoms again: Left side: 1 C, 4 H, 2 O Right side: 1 C, 4 H, 4 O Balancing oxygen: We need 4 O on the left side, so add a coefficient of 2 in front of O₂: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Counting atoms (final check): Left side: 1 C, 4 H, 4 O Right side: 1 C, 4 H, 4 O Balanced equation: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Example 2: Decomposition of Potassium Chlorate Potassium chlorate (KClO₃) decomposes upon heating to produce potassium chloride (KCl) and oxygen gas (O₂).

Unbalanced equation: KClO₃(s) → KCl(s) + O₂(g)

Counting atoms: Left side: 1 K, 1 Cl, 3 O Right side: 1 K, 1 Cl, 2 O Balancing oxygen: The lowest common multiple of 3 and 2 is 6.