D. Quantities in Chemical Reactions · Quantities in Chemical Reactions

Quantities in Chemical Reactions: Unit Review

A1.1 A1.4 A1.5 A1.6 A1.8 A1.10 A1.11 A1.12 A1.13 A2.1 A2.2 D1.1 D1.2 D2.1 D2.2 D2.3 D2.4 D2.5 D2.6 D2.7 D3.1 D3.2 D3.3 D3.4About 40 minutes

Learning goals

I can convert between mass, moles and particles.
I can find percentage composition, empirical formulas and molecular formulas.
I can solve stoichiometry problems, including limiting reagents and percentage yield.
I can explain why quantitative accuracy matters in real life.

Words to know

mole
6.02 × 10²³ entities (exactly 6.02214076 × 10²³).

molar mass
Mass of one mole of a substance, in g/mol.

empirical formula
Simplest whole-number ratio of atoms in a compound.

stoichiometry
Calculating quantities of reactants and products from a balanced equation.

limiting reagent
The reactant used up first; it decides the theoretical yield.

percentage yield
Actual yield ÷ theoretical yield × 100%.

Let's review!

In this unit you learned to count particles by weighing them. The mole (6.02 × 10²³ particles) and molar mass link grams, moles and particles. You used them to find percentage composition, and worked backwards from lab data to empirical and molecular formulas. Then you read balanced equations as mole ratios and used stoichiometry to predict masses of reactants and products, find limiting reagents, and compare theoretical, actual and percentage yields, both on paper and in investigations. Finally you saw why quantitative accuracy matters at home, at work and in industry.

This review has a few questions from every lesson. Try each one, then press Show answer to check.

How to use this review

Work through one or two sections at a time with a calculator and the molar-mass table from the Molar Mass lesson. For every calculation: write what you know and what you want, sketch the route on the mole map, show units, and round only the final answer (usually to three significant figures). For every question you miss, write down why (wrong ratio? forgot to balance? used the excess reagent? units?) and go back to that lesson's worked examples. Finish with the 10-question quiz.

How the big ideas of the unit connect. Every arrow is a calculation you can now do.

Just for fun

A chemist walks into a café and orders 6.02 × 10²³ coffees.

The barista says, "That's a mole of a lot of coffee." The chemist says, "Don't worry, I'll weigh them instead of counting."

The Mole and Avogadro's Constant

How many molecules are in 0.750 mol of nitrogen gas, N₂?

What amount of iron is 9.03 × 10²³ iron atoms?

How many oxygen atoms are in 0.400 mol of sulfur trioxide, SO₃?

Molar Mass: Linking Moles to Grams

Find the molar mass of calcium nitrate, Ca(NO₃)₂.

What is the mass of 0.250 mol of baking soda, NaHCO₃ (84.01 g/mol)?

What amount is 100. g of calcium chloride, CaCl₂ (110.98 g/mol)?

Mass, Moles and Particles: Using the Mole Map

How many molecules are in 4.00 g of methane, CH₄?

What is the mass of 3.01 × 10²² molecules of CO₂?

How many chlorine atoms are in 10.0 g of CCl₄ (153.81 g/mol)?

The Law of Definite Proportions and Percentage Composition

Find the percentage of nitrogen in potassium nitrate, KNO₃ (101.11 g/mol).

What mass of calcium is in 25.0 g of CaCO₃ (40.04% Ca)?

Two pure samples of sodium chloride, one of 1 g and one of 1 kg, are analysed. What does the law of definite proportions predict?

Investigation: Finding Percentage Composition in the Lab

A 2.000 g sample of CuSO₄·5H₂O weighs 1.278 g after heating to constant mass. Find the percentage of water and the percentage error (theoretical 36.08%).

In the magnesium oxide lab, some white smoke escaped. Is the % Mg too high or too low?

Why must a hydrate be heated to constant mass?

Empirical Formulas from Percentage Composition and Mass Data

A compound is 63.5% Ag, 8.25% N and 28.2% O. Find its empirical formula.

A mole ratio comes out as 1 : 1.67. What should you multiply by, and what is the whole-number ratio?

2.43 g of Mg combines with 7.09 g of Cl. Find the empirical formula.

Molecular Formulas

A compound has the empirical formula CH₂ and a molar mass of 70.15 g/mol. Find its molecular formula.

Styrene is 92.3% C and 7.7% H, with a molar mass of 104.16 g/mol. Find its molecular formula.

True or false: the empirical and molecular formulas of water are the same.

Mole Ratios in Balanced Chemical Equations

4NH₃ + 5O₂ → 4NO + 6H₂O. What amounts of O₂ react with, and of H₂O form from, 2.00 mol of NH₃?

Balance C₄H₁₀ + O₂ → CO₂ + H₂O, then find the amount of CO₂ from 0.500 mol of butane.

In 2CO + O₂ → 2CO₂, there are 3 mol of reactants and 2 mol of products. Is mass conserved? Explain.

Stoichiometry: Calculating Masses and Particles in Reactions

2KClO₃ → 2KCl + 3O₂. What mass of oxygen is produced from 12.25 g of KClO₃ (122.55 g/mol)?

CH₄ + 2O₂ → CO₂ + 2H₂O. What mass of CO₂ is produced from 8.00 g of methane?

C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. How many water molecules form from 2.00 g of propane?

Choosing the right tool. Match what you know and what you want to the calculation you need.

Limiting and Excess Reagents

2H₂ + O₂ → 2H₂O. 10.0 g of H₂ and 64.0 g of O₂ react. Find the limiting reagent, the mass of water and the excess left.

Zn + S → ZnS. 0.30 mol Zn and 0.25 mol S react. What amount of ZnS forms?

Why is the limiting reagent not always the reactant with the smaller mass?

Theoretical, Actual and Percentage Yield

Theoretical yield 15.0 g, actual yield 12.6 g. Find the percentage yield.

A process has a 75.0% yield. What theoretical yield is needed to collect 30.0 g of product?

Two steps have yields of 90% and 60%. What is the overall yield?

Investigation: Measuring Actual, Theoretical and Percentage Yield

Filter paper: 0.850 g. Paper + dry precipitate: 2.050 g. Theoretical yield: 1.250 g. Find the percentage yield.

A precipitate was weighed while still damp. How does this affect the percentage yield?

In the kitchen CO₂ investigation, 8.40 g of baking soda gives a mass loss of 3.9 g. Find the percentage yield (theoretical 4.40 g) and one likely reason it is under 100%.

Chemical Quantities at Home, at Work and in Industry

Imaginary Medicine W: 12 mg/kg. The liquid contains 100 mg per 5 mL. What volume is right for a 25 kg child?

Why must bleach never be mixed with ammonia-based cleaners?

A scrubber removes 128 kg of SO₂ (64.07 g/mol) with CaCO₃ + SO₂ → CaSO₃ + CO₂. What mass of CaCO₃ is needed?

Check your understanding

Quiz

How many molecules are in 2.00 mol of water?

What is the molar mass of magnesium hydroxide, Mg(OH)₂?

What is the mass of 0.500 mol of CO₂?

What is the percentage of hydrogen in methane, CH₄ (16.05 g/mol)?

A compound's mole ratio is C : H = 1 : 2.5. What is its empirical formula?

A compound has empirical formula NO₂ (46.01 g/mol) and molar mass 92.02 g/mol. What is its molecular formula?

N₂ + 3H₂ → 2NH₃. What amount of NH₃ forms from 6.0 mol of H₂ (with excess N₂)?

2Mg + O₂ → 2MgO. 3.0 mol Mg and 1.0 mol O₂ react. Which is limiting?

Theoretical yield 40.0 g, actual yield 34.0 g. What is the percentage yield?

10.

Why does a unit or ratio error matter in industry?

For parents and teachers

Use the review to find what to revisit:

Multiplying when they should divide by Avogadro's constant, or not saying which particle is counted → Lesson 1 (The Mole and Avogadro's Constant).
Errors with brackets, hydrates or units of g/mol → Lesson 2 (Molar Mass).
Trying to go straight from grams to particles, or forgetting subscripts when counting atoms → Lesson 3 (Mass, Moles and Particles).
Dividing by the wrong total, or confusing percent by atoms with percent by mass → Lesson 4 (Law of Definite Proportions and Percentage Composition).
Lab calculations, constant mass, percentage error and error direction → Lesson 5 (Investigation: Percentage Composition).
Rounding 1.5 or 1.33 instead of multiplying, or comparing grams instead of moles → Lesson 6 (Empirical Formulas).
Finding n with the wrong mass, or not multiplying every subscript → Lesson 7 (Molecular Formulas).
Using an unbalanced equation or flipping the mole ratio → Lesson 8 (Mole Ratios).
Skipping the mole step in mass-to-mass problems → Lesson 9 (Stoichiometry).
Using the reactant with the smaller mass as limiting, or calculating from the excess reagent → Lesson 10 (Limiting and Excess Reagents).
Dividing actual and theoretical the wrong way, or multiplying instead of dividing in backward problems → Lesson 11 (Theoretical, Actual and Percentage Yield).
Vague error analysis ("human error") or too many significant figures → Lesson 12 (Investigation: Percentage Yield).
Forgetting the "per" in units, or unsafe ideas about mixing cleaners → Lesson 13 (Chemical Quantities at Home, at Work and in Industry).

The mole concept is used again in the next unit, Solutions and Solubility (concentration in mol/L), in Gases and Atmospheric Chemistry, and throughout Grade 12 Chemistry (SCH4U).