Lesson Notes By Weeks and Term v5 - Grade 11

Lesson Video

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Performance objectives

• Define acids and bases using Arrhenius and Bronsted-Lowry theories.
• Differentiate between strong and weak acids/bases.
• Identify oxidation, reduction, and assign oxidation numbers.
• Relate these reactions to real-world applications.

Lesson summary

Chemical reactions are fundamental to life and industry. Understanding the different types of chemical reactions allows us to predict, control, and utilize them. This week, we will delve into two important classes of chemical reactions: acid-base reactions and redox (reduction-oxidation) reactions. These reactions are crucial for understanding everything from the digestion of food in our bodies to the production of electricity in batteries, and even the corrosion of metal structures, which has significant economic implications for South Africa's infrastructure.

Lesson notes

2.1 Acid-Base Reactions 2.1.1 Definitions of Acids and Bases Arrhenius Theory: Arrhenius defined acids as substances that produce hydrogen ions (H+) when dissolved in water, and bases as substances that produce hydroxide ions (OH-) when dissolved in water.

For example: Acid: HCl(aq) → H+(aq) + Cl-(aq)

Base: NaOH(aq) → Na+(aq) + OH-(aq)

Limitation: This theory only applies to aqueous solutions and doesn't explain the basic properties of substances like ammonia (NH3).

Bronsted-Lowry Theory: This theory provides a broader definition. A Bronsted-Lowry acid is a proton (H+) donor, and a Bronsted-Lowry base is a proton acceptor.

For example: NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) In this reaction, water acts as an acid (proton donor) and ammonia acts as a base (proton acceptor).

Conjugate Acid-Base Pairs: In a Bronsted-Lowry acid-base reaction, the acid and base are related by the transfer of a proton. After an acid donates a proton, it becomes its conjugate base. After a base accepts a proton, it becomes its conjugate acid. In the example above, NH3 and NH4+ are a conjugate acid-base pair, and H2O and OH- are another conjugate acid-base pair. 2.1.2 Strength of Acids and Bases Strong Acids/Bases: Strong acids and bases completely dissociate (ionize) in water. This means they break apart into their ions almost entirely.

Examples include: Strong acids: HCl, H2SO4, HNO3 Strong bases: NaOH, KOH Weak Acids/Bases: Weak acids and bases only partially dissociate in water. An equilibrium is established between the undissociated acid/base and its ions.

Examples include: Weak acids: CH3COOH (acetic acid - vinegar), HF Weak bases: NH3 The extent of dissociation is quantified by the acid dissociation constant (Ka) for acids and the base dissociation constant (Kb) for bases. A smaller Ka or Kb value indicates a weaker acid or base. For example, consider the dissociation of acetic acid: CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq) Ka = [H+][CH3COO-] / [CH3COOH] A small Ka value (around 1.8 x 10-5) indicates that only a small fraction of acetic acid dissociates in water. 2.1.3 Neutralization Reactions Acid-base reactions typically result in the formation of a salt and water. These are called neutralization reactions. Acid + Base → Salt + Water For example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) The salt formed depends on the acid and base used. 2.2 Redox Reactions 2.2.1 Oxidation and Reduction Oxidation: The loss of electrons.

Reduction: The gain of electrons.

Remember the mnemonic OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). Oxidation and reduction always occur together in a redox reaction. 2.2.2 Oxidation Numbers Oxidation numbers are a way to keep track of electron transfer in redox reactions. Here are some rules for assigning oxidation numbers: The oxidation number of an element in its elemental form is 0. (e.g., Na(s), O2(g), Cu(s) all have oxidation number 0). The oxidation number of a monatomic ion is equal to its charge. (e.g., Na+ has oxidation number +1, Cl- has oxidation number -1). The sum of oxidation numbers in a neutral compound is

0. The sum of oxidation numbers in a polyatomic ion equals the charge of the ion. In compounds, alkali metals (Group 1) have oxidation number +1, and alkaline earth metals (Group 2) have oxidation number +

2. Fluorine always has oxidation number -

1. Oxygen usually has oxidation number -2 (except in peroxides like H2O2, where it's -1, and with fluorine). Hydrogen usually has oxidation number +1 (except in metal hydrides like NaH, where it's -1).

Example: Assign oxidation numbers to each element in KMnO

4. K is in Group 1, so its oxidation number is +

1. O usually has oxidation number -

2. Since there are 4 oxygen atoms, the total oxidation number for oxygen is -

8. The sum of oxidation numbers must be

0. Therefore, +1 + Mn + (-8) = 0, so Mn = +7. 2.2.3 Oxidizing and Reducing Agents Oxidizing Agent: The substance that causes oxidation and is itself reduced. It accepts electrons.

Reducing Agent: The substance that causes reduction and is itself oxidized. It donates electrons. 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. Balance H by adding H+. (If the reaction takes place in basic conditions, add OH- to both sides equal to the number of H+, then combine H+ and OH- to form H2O. Cancel out any water molecules that appear on both sides). Balance charge by adding electrons (e-) to the side with the more positive charge. Multiply each half-reaction by a factor 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 balanced half-reactions together, cancelling out electrons and any common species (H+, H2O, OH-) on both sides.