Stoichiometry

Stoichiometry: The Recipe Book of Chemistry 🧪

Imagine you’re baking the world’s most delicious cake. You need exact amounts of flour, sugar, and eggs. Use too little flour? Flat cake. Too many eggs? Rubbery mess. Chemistry works the same way!

The Big Idea

Stoichiometry is chemistry’s recipe system. It tells us exactly how much of each ingredient (chemical) we need and exactly how much product we’ll make.

Think of it like this:

• Recipe: 2 slices of bread + 3 slices of cheese = 1 grilled cheese sandwich
• Chemistry: 2 H₂ + O₂ → 2 H₂O (2 hydrogen molecules + 1 oxygen molecule = 2 water molecules)

1. Moles in Equations: Counting the Invisible

What’s a Mole?

A mole is just a counting word—like “dozen” means 12.

Why such a weird number? Because atoms are TINY. We need huge numbers just to have enough to see!

1 mole of anything = 6.02 × 10²³ of that thing

Reading Equations with Moles

Look at this equation:

2 H₂ + O₂ → 2 H₂O

The big numbers in front (called coefficients) tell us moles:

• 2 moles of hydrogen
• 1 mole of oxygen
• Makes 2 moles of water

Simple Example

Question: How many moles of water form from 4 moles of H₂?

Recipe says: 2 moles H₂ → 2 moles H₂O

So: 4 moles H₂ → 4 moles H₂O ✓

Double the ingredients, double the product!

2. Mole Ratios: The Secret Conversion Tool

What’s a Mole Ratio?

A mole ratio compares how much of one substance relates to another in a reaction.

2 H₂ + O₂ → 2 H₂O

From this equation, we get these ratios:

graph TD
    A["H₂ : O₂"] --> B["2 : 1"]
    C["H₂ : H₂O"] --> D["2 : 2 = 1 : 1"]
    E["O₂ : H₂O"] --> F["1 : 2"]

Using Mole Ratios

Question: If you have 3 moles of O₂, how many moles of H₂O can you make?

Step 1: Find the ratio → O₂ : H₂O = 1 : 2

Step 2: Multiply!

• 3 moles O₂ × (2 H₂O / 1 O₂) = 6 moles H₂O ✓

Real-Life Analogy

Making sandwiches:

• Recipe: 2 bread + 1 cheese = 1 sandwich
• Ratio of bread to sandwich = 2:1
• If you have 10 bread slices, you can make 5 sandwiches!

3. Limiting and Excess Reactants: Who Runs Out First?

The Party Problem

Imagine you’re setting tables for a party:

• Each table needs 4 chairs
• You have 5 tables and 12 chairs

Question: How many complete table setups can you make?

• 5 tables could seat 20 people (5 × 4)
• But you only have 12 chairs!
• 12 ÷ 4 = 3 complete setups

Chairs are the LIMITING factor! Tables are in EXCESS!

In Chemistry

graph TD
    A["Limiting Reactant"] --> B["Runs out first"]
    B --> C["Controls how much product forms"]
    D["Excess Reactant"] --> E["Leftover after reaction"]

Example Problem

Reaction: 2 H₂ + O₂ → 2 H₂O

Given: 5 moles H₂ and 3 moles O₂

Which runs out first?

For H₂: Need ratio 2:1 with O₂

• 5 moles H₂ needs 2.5 moles O₂
• We have 3 moles O₂ ✓ (enough!)

For O₂: Need ratio 1:2 with H₂

• 3 moles O₂ needs 6 moles H₂
• We only have 5 moles H₂ ✗ (not enough!)

H₂ is LIMITING → Controls the reaction O₂ is in EXCESS → Some left over

Product formed: 5 moles H₂O (same as H₂ because 1:1 ratio)

4. Empirical Formula: The Simplest Recipe

What Is It?

The empirical formula shows the simplest whole-number ratio of atoms in a compound.

Think of it as the “reduced fraction” of chemistry!

How to Find It

Example: A compound has 40% carbon and 60% oxygen. Find its empirical formula.

Step 1: Assume 100g sample

• Carbon: 40g
• Oxygen: 60g

Step 2: Convert to moles

• C: 40g ÷ 12g/mol = 3.33 mol
• O: 60g ÷ 16g/mol = 3.75 mol

Step 3: Divide by smallest

• C: 3.33 ÷ 3.33 = 1
• O: 3.75 ÷ 3.33 = 1.125 ≈ 1

Wait! 1.125 isn’t a whole number. Multiply both by 8:

• C: 1 × 8 = 8
• O: 1.125 × 8 = 9

Empirical Formula: C₈O₉

(Actually, let’s recalculate for cleaner numbers)

Better Example: 40% C, 6.7% H, 53.3% O

• C: 40 ÷ 12 = 3.33 mol
• H: 6.7 ÷ 1 = 6.7 mol
• O: 53.3 ÷ 16 = 3.33 mol

Divide by 3.33:

• C: 1, H: 2, O: 1

Empirical Formula: CH₂O ✓

5. Molecular Formula: The Complete Picture

Empirical vs. Molecular

The molecular formula is a multiple of the empirical formula!

graph TD
    A["Molecular Formula"] --> B["= Empirical Formula × n"]
    C["Find n"] --> D["n = Molar Mass ÷ Empirical Mass"]

Finding the Molecular Formula

Given:

• Empirical formula: CH₂O
• Molar mass: 180 g/mol

Step 1: Find empirical formula mass

• C: 12 × 1 = 12
• H: 1 × 2 = 2
• O: 16 × 1 = 16
• Total: 30 g/mol

Step 2: Find the multiplier (n)

• n = 180 ÷ 30 = 6

Step 3: Multiply empirical formula by n

• CH₂O × 6 = C₆H₁₂O₆ ✓

This is glucose—the sugar in your blood!

The Complete Picture

graph TD
    A["Chemical Equation"] --> B["Coefficients = Moles"]
    B --> C["Mole Ratios"]
    C --> D["Calculate Products"]
    D --> E{Do you have enough?}
    E -->|Yes| F["Reaction completes"]
    E -->|No| G["Find Limiting Reactant"]
    G --> H["Calculate actual product"]

    I["% Composition"] --> J["Empirical Formula"]
    J --> K["+ Molar Mass"]
    K --> L["Molecular Formula"]

Quick Reference

You’ve Got This! 💪

Remember: Stoichiometry is just cooking with atoms.

• Equations are recipes
• Coefficients tell you amounts
• Mole ratios help you convert
• Limiting reactants control your output
• Empirical formulas are simplified
• Molecular formulas are real

Now go balance some equations like the chemistry chef you are! 👨‍🍳🔬