Lesson Notes By Weeks and Term v5 - Grade 11

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Performance objectives

• Define acids and bases using three different theories.
• Write equations for neutralization and calculate pH.
• Define oxidation, reduction, and their agents.
• Balance redox reactions using the half-reaction method.
• Explain corrosion (rusting) as a redox process.

Lesson summary

Chemical reactions are fundamental to life and industry. Understanding different types of reactions allows us to predict their outcomes, control them, and harness their power for various applications. In this week, we'll focus on two major types of chemical reactions: acid-base reactions and redox (reduction-oxidation) reactions. These reactions are crucial in numerous everyday processes, from cleaning with household chemicals to the extraction of metals from ores and even the functioning of our own bodies. In South Africa, understanding these reactions is particularly relevant to industries like mining, agriculture (soil pH management), and water treatment.

Lesson notes

2.1 Acid-Base Reactions 2.1.1 Defining Acids and Bases: Arrhenius Definition: An Arrhenius acid is a substance that produces hydrogen ions (H+) in aqueous solution, while an Arrhenius base produces hydroxide ions (OH-) in aqueous solution. For example, HCl(aq) → H+(aq) + Cl-(aq) (acid) and NaOH(aq) → Na+(aq) + OH-(aq) (base). This definition is limited to aqueous solutions.

Bronsted-Lowry Definition: A Bronsted-Lowry acid is a proton (H+) donor, and a Bronsted-Lowry base is a proton acceptor. This definition is broader than the Arrhenius definition. For example, in the reaction NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq), NH3 acts as a Bronsted-Lowry base (accepts H+) and H2O acts as a Bronsted-Lowry acid (donates H+).

Lewis Definition: A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. This is the broadest definition and includes reactions where no protons are transferred. An example is the reaction between BF3 (Lewis acid) and NH3 (Lewis base). 2.1.2 Conjugate Acid-Base Pairs: In a Bronsted-Lowry acid-base reaction, the acid donates a proton to form its conjugate base, and the base accepts a proton to form its conjugate acid. For example, in the reaction HCl(aq) + H2O(l) ⇌ H3O+(aq) + Cl-(aq), HCl is the acid, Cl- is its conjugate base, H2O is the base, and H3O+ is its conjugate acid.

Acid/Conjugate Base: HCl/Cl- Base/Conjugate Acid: H2O/H3O+ 2.1.3 Neutralization Reactions: Neutralization is the reaction between an acid and a base, resulting in the formation of salt and water (in most cases).

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) 2.1.4 pH Scale and Calculation: pH is a measure of the acidity or basicity of a solution. It is defined as pH = -log[H+], where [H+] is the concentration of hydrogen ions in mol/

L. Example 1: Calculate the pH of a 0.01 M solution of HCl (a strong acid, assuming complete dissociation). Since HCl is a strong acid, [H+] = 0.01 M. pH = -log(0.01) = -log(1 x 10^-2) = 2 Example 2: Calculate the pH of a 0.005 M solution of NaOH (a strong base, assuming complete dissociation). Since NaOH is a strong base, [OH-] = 0.005

M. First, calculate pOH: pOH = -log[OH-] = -log(0.005) = 2.30 Then, use the relationship: pH + pOH = 14 pH = 14 - 2.30 = 11.70 2.2 Redox Reactions 2.2.1 Oxidation and Reduction: Oxidation: Loss of electrons OR increase in oxidation number.

Reduction: Gain of electrons OR decrease in oxidation number.

OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). 2.2.2 Oxidation Numbers: Oxidation numbers are assigned to atoms in a molecule or ion to keep track of electron distribution. Some rules for assigning oxidation numbers: The oxidation number of an element in its elemental form is 0. (e.g., Na(s), O2(g), Fe(s) all have an oxidation number of 0) The oxidation number of a monatomic ion is equal to its charge. (e.g., Na+ has an oxidation number of +1, Cl- has an oxidation number of -1) Oxygen usually has an oxidation number of -2, except in peroxides (like H2O2), where it is -1, or when bonded to fluorine (OF2) where it is +

2. Hydrogen usually has an oxidation number of +1, except when bonded to metals, where it is -1 (e.g., NaH). The sum of the oxidation numbers in a neutral compound is

0. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

Example: Determine the oxidation number of Mn in KMnO

4. K has an oxidation number of +

1. O has an oxidation number of -

2. Let the oxidation number of Mn be x. +1 + x + 4(-2) = 0 x - 7 = 0 x = +7 Therefore, the oxidation number of Mn in KMnO4 is +7. 2.2.3 Oxidizing and Reducing Agents: Oxidizing Agent: A substance that causes oxidation by accepting electrons (it gets reduced).

Reducing Agent: A substance that causes reduction by donating electrons (it gets oxidized). 2.2.4 Balancing Redox Reactions (Half-Reaction Method): Write the unbalanced equation. Separate the equation into two half-reactions: an oxidation half-reaction and a reduction half-reaction.

Balance each half-reaction: Balance all elements except H and O. Balance O by adding H2O to the side that needs oxygen. Balance H by adding H+ to the side that needs hydrogen. Balance charge by adding electrons (e-) to the side with the more positive charge. Multiply each half-reaction by an integer so that the number of electrons lost in the oxidation half-reaction equals the number of electrons gained in the reduction half-reaction. Add the two half-reactions together. The electrons should cancel out. Simplify the equation. Make sure all coefficients are in the lowest possible ratio.

For basic solutions: Add OH- ions to both sides of the equation to neutralize the H+ ions, forming water (H2O). Cancel out any water molecules that appear on both sides.