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O Levels Chemistry Notes - Mole Concept

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Summary

Key Points

Definition

Example

Relative Atomic Mass (Ar)

Ar is the average mass of one atom of an element compared to 1/12 of the mass of a carbon-12 atom.

Ar = Sum of (% abundance × isotopic mass)

Relative Molecular Mass (Mr)

Mr is the average mass of one molecule compared to 1/12 of the mass of a carbon-12 atom.

Mr of compound = Total Ar of all atoms

Empirical and Molecular Formulae

Empirical formula is the simplest ratio of atoms in a compound; 

Molecular formula shows the actual number of atoms in a molecule.

The Mole

A mole represents a quantity of 6.02 x 1023 particles (Avogadro’s number).

1 mole = 6.02 x 1023 particles



Mole Formulas

Mole =no. of particles1023

Mole = volume of gas (dm3)24cm3/mol

Mole = mass of substance (g)Mr

conc. (mol/dm3)= no. moles of solute (mol)volume of solvent (dm3)

conc. (g/dm3)= mass of solute (g)volume of solvent (dm3)

1dm3=1000cm3

Limiting reagent

Limiting reagent is the reagent that limits the amount of product formed 

Percentage Yield and Purity

Percentage yield measures how much product is actually obtained, while percentage purity measures the purity of a substance.

% Yield = (Actual yield / Theoretical yield) × 100%

Let us look at what you are required to know:

(a) define relative atomic mass, Ar

(b) define relative molecular mass, Mr, and calculate relative molecular mass (and relative formula mass) as the sum of relative atomic masses

(c) define the term mole in terms of the Avogadro constant

(d) calculate the percentage mass of an element in a compound when given appropriate information

(e) calculate empirical and molecular formulae from relevant data

(f) calculate stoichiometric reacting masses and volumes of gases (one mole of gas occupies 24 dm3 at room temperature and pressure); calculations involving the idea of limiting reactants may be set (knowledge of the gas laws and the calculations of gaseous volumes at different temperatures and pressures are not required)

(g) apply the concept of solution concentration (in mol/dm3 or g/dm3) to process the results of volumetric experiments (e.g. titration) and to solve simple problems (appropriate guidance will be provided where unfamiliar reactions are involved)

(h) calculate % yield and % purity.

What are some common pitfalls?

1. Molar Mass Calculation:

Pitfall: Incorrectly calculating molar mass by summing atomic masses without considering subscripts.

Elaboration: Ensure that you understand how to calculate molar mass by multiplying the atomic mass of each element by its subscript in the chemical formula.

Example: Calculate the molar mass of calcium carbonate (CaCO3).

Incorrect Calculation: Molar Mass = Ca (40.08 g/mol) + C (12.01 g/mol) + O (16.00 g/mol) = 68.09 g/mol

Correct Calculation: Molar Mass = Ca (1 × 40.08 g/mol) + C (1 × 12.01 g/mol) + O (3 × 16.00 g/mol) = 100.09 g/mol

 

2. Balanced Equations:

Pitfall: Using unbalanced chemical equations for stoichiometric calculations.

Elaboration: Remember to balance chemical equations before performing any stoichiometric calculations. Make it a regular practice

Example: Consider the reaction: N2 + H2 → NH3

Unbalanced Equation: N2 + H2 → NH3

Correct Balanced Equation: N2 + 3H2 → 2NH3

 

3. Stoichiometric Calculations:

Pitfall: Misinterpreting coefficients as mole ratios.

Elaboration: Remember that the coefficients in a balanced equation represent the mole ratios of reactants and products. Use these ratios correctly in stoichiometric calculations.

Example: In the balanced equation 2H2 + O2 → 2H2O, find the number of moles of O2 needed to react with 4 moles of H2.

Incorrect Calculation: 4 moles of H2 = 4 moles of O2 (misinterpreting coefficients)

Correct Calculation: 4 moles of H2 × (1 mole of O2 / 2 moles of H2) = 2 moles of O2

 

4. Limiting Reactants:

Pitfall: Failing to identify the limiting reactant, leading to incorrect yield calculations.

Elaboration: Determine the limiting reactant by comparing the stoichiometric coefficients and the initial amounts of reactants. Limiting reactants determine the maximum amount of product.

Example: Consider the reaction of 3 moles of hydrogen (H2) and 2 moles of oxygen (O2). Determine the limiting reactant and the moles of water formed.

Incorrect Calculation: 3 moles of H2 react with 2 moles of O2, so O2 is limiting, and 2 moles of water are formed.

Correct Calculation: To find the limiting reactant, calculate the moles of water formed by each reactant. Moles of H2 produce 3 moles of water, and moles of O2 produce 1 mole of water. Since only 1 mole of water can be formed from O2, O2 is the limiting reactant.

 

5. Concentration Units:

Pitfall: Mixing up units when working with different concentration units (mol/dm³, g/dm³).

Elaboration: Clearly define the units and provide examples of converting between them. Use consistent units in calculations.

Example: Calculate the concentration of glucose (C6H12O6) in a solution containing 20 g of glucose in 500 cm³ of water.

Incorrect Calculation: Concentration = 20 g / 500 cm³ = 0.04 g/dm³

Correct Calculation: First, convert the volume to dm³ (500 cm³ = 0.5 dm³), then calculate the concentration in g/dm³: Concentration = 20 g / 0.5 dm³ = 40 g/dm³

Study Guide

Studying the topic of atomic structure effectively requires a combination of strategies that cater to the conceptual understanding and application of key concepts. Here are some specific studying tips for the topic of atomic structure:

1. Master the Basics First:

Begin by understanding the basic components of an atom, including protons, neutrons, electrons, and their respective properties (charge, mass, location).

2. Study the Periodic Table:

Familiarize yourself with the periodic table, paying special attention to group numbers, periods, and the arrangement of elements based on atomic number.

3. Electron Configurations:

Focus on writing electron configurations for different elements. Practice writing configurations for elements across the periodic table to become proficient.

4. Electronic Structure:

Learn how to draw electronic structures to represent the distribution of electrons in an atom. Practice drawing these diagrams for various elements.

Practice drawing ionic structures too! These are commonly drawn wrongly by many students.

5. Isotopes:

Learn how to identify isotopes and calculate the average atomic mass of an element considering isotopic abundances.

6. Practice Problem Solving:

Solve problems related to atomic structure, including determining the number of protons, neutrons, and electrons in various elements or ions.

7. Use Visual Aids:

Utilize diagrams, charts, and models to visualize atomic structure concepts. These visual aids can help you grasp abstract ideas more easily.

8. Real-World Applications:

Connect atomic structure to real-world applications, such as the use of radioactive isotopes in medicine, electron configuration in chemical reactions, and atomic models in technology.

 

By following this study plan, you can confidently approach questions related to the atomic structures in your ‘O’ levels chemistry exams. Understanding the core concepts will set a strong foundation for your success.

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