Book cover for Chemistry: The Molecular Nature of Matter

Chemistry: The Molecular Nature of Matter

Neil D. Jespersen, James E. Brady, Alison Hyslop

ISBN #9781118413920

7th Edition

3,064 Questions

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53,557 Students Helped

Homework Questions

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Summary
Learning Objectives
Key Concepts
Example Problems
Explanations
Common Mistakes

Summary

This chapter provides a comprehensive introduction to the mole concept, serving as a fundamental link between atomic-level measurements and practical laboratory calculations. Key topics include converting between mass and moles, determining empirical and molecular formulas from composition data, balancing chemical equations, conducting stoichiometric calculations, and evaluating limiting reactants and yields. Mastery of these topics is essential for both academic pursuits in chemistry and practical applications in industrial and research settings.

Learning Objectives

1

Explain the mole concept and its significance in connecting atomic-level measurements to laboratory-scale quantities.

2

Perform conversions between mass and moles using Avogadro's number and molar mass.

3

Determine empirical and molecular formulas from percentage composition data.

4

Balance chemical equations and conduct stoichiometric calculations, including identifying limiting reactants and calculating theoretical and percentage yields.

5

Apply problem-solving techniques using the mole concept to plan and interpret laboratory experiments and industrial processes.

Key Concepts

CONCEPT

DEFINITION

Mole

A unit that measures the amount of substance; one mole contains 6.022 x 10^23 particles (Avogadro’s number) of that substance.

Avogadro's Number

The number of particles in one mole of a substance, approximately 6.022 x 10^23.

Molar Mass

The mass of one mole of a substance, typically expressed in grams per mole (g/mol).

Empirical Formula

The simplest, whole-number ratio of elements in a compound.

Molecular Formula

The actual number of atoms of each element in a molecule, which may be a multiple of the empirical formula.

Stoichiometry

The area of chemistry that deals with the quantitative relationships that exist between reactants and products in chemical reactions.

Limiting Reactant

The reactant that is completely consumed first in a chemical reaction, thus limiting the amount of product formed.

Theoretical Yield

The maximum amount of product that could be produced from a given amount of reactant, based on stoichiometric calculations.

Percentage Yield

The efficiency of a reaction, calculated by comparing the actual yield to the theoretical yield, expressed as a percentage.

Example Problems

Example 1

What is the definition of the mole?

Example 2

Why are moles used, when all stoichiometry problems could be done using only atomic mass units?

Example 3

Write the following as balanced chemical equations. (a) Calcium hydroxide reacts with hydrogen chloride to form calcium chloride and water. (b) Silver nitrate and calcium chloride react to form calcium nitrate and silver chloride. (c) Lead(II) nitrate reacts with sodium sulfate to form lead(II) sulfate and sodium nitrate. (d) Iron(III) oxide and carbon react to form iron and carbon dioxide.

Example 4

Which contains more molecules: $2.5 \mathrm{~mol}$ of $\mathrm{H}_{2} \mathrm{O}$ or $2.5 \mathrm{~mol}$ of $\mathrm{H}_{2}$ ? Which contains more atoms? Which weighs more?

Example 5

How would Avogadro's number change if the atomic mass unit were to be redefined as $2 \times 10^{-27} \mathrm{~kg}$, exactly?
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Step-by-Step Explanations

QUESTION

How do you convert 18 grams of water (H2O) into moles?

STEP-BY-STEP ANSWER:

Step 1: Determine the molar mass of H2O. The atomic mass of Hydrogen is approximately 1 g/mol and Oxygen is approximately 16 g/mol. Therefore, molar mass = 2(1) + 16 = 18 g/mol.
Step 2: Use the conversion formula: moles = mass / molar mass.
Step 3: Substitute the given values: moles = 18 g / 18 g/mol.
Step 4: Calculate the result: 18 / 18 = 1 mole.
Final Answer: 18 grams of water is equivalent to 1 mole of H2O.

Mass to Moles Conversion

QUESTION

How can you determine the empirical formula for a compound given its percentage composition?

STEP-BY-STEP ANSWER:

Step 1: Assume a 100 g sample so that percentages can be directly converted to grams.
Step 2: Convert the mass of each element to moles by dividing by the respective atomic masses.
Step 3: Divide all the mole values by the smallest value to obtain a mole ratio.
Step 4: If necessary, multiply the ratios by a whole number to obtain whole numbers.
Final Answer: The resulting whole-number ratios give the empirical formula of the compound.

Determining Empirical Formula

QUESTION

In a reaction where 10 g of reactant A reacts with 15 g of reactant B, how would you determine the limiting reactant?

STEP-BY-STEP ANSWER:

Step 1: Write and balance the chemical equation for the reaction.
Step 2: Convert the masses of reactant A and B to moles using their molar masses.
Step 3: Compare the mole ratio from the balanced equation with the available moles of each reactant.
Step 4: The reactant with the smaller mole ratio (compared to its coefficient) is the limiting reactant.
Final Answer: The reactant that runs out first, based on stoichiometric calculations, is identified as the limiting reactant.

Identifying Limiting Reactant

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Common Mistakes

  • Confusing mass and moles due to overlooking the molar mass conversion factor.
  • Incorrectly balancing chemical equations, leading to errors in stoichiometric calculations.
  • Misidentifying the limiting reactant by not correctly comparing mole ratios from the balanced equation.
  • Forgetting to convert percentage composition into moles before determining empirical formulas.
  • Overlooking the difference between theoretical yield and actual yield when calculating percentage yield.