MYP Integrated Sciences Balancing Equations Notes

What Does the Law of Conservation of Mass Say?

The Law of Conservation of Mass

The law of conservation of mass states that:

In a chemical reaction, mass is neither created nor destroyed.

Instead:

1. The atoms present at the start are rearranged to form new substances.
2. The total mass of the reactants equals the total mass of the products.

We can change how atoms are bonded, but we do not create or destroy atoms in ordinary chemical reactions.

Experimental Evidence

We can test this law in the lab by weighing substances before and after a reaction.

Why Does Mass Sometimes Appear Not to Be Conserved?

1. In everyday experiments, we often use open systems (e.g. beakers open to the air).
2. In such cases, gases can escape or enter, making it seem like mass has been lost or gained.

Escape of Gases – Apparent Loss of Mass

When a reaction produces a gas that escapes into the air, the measured mass of the container + contents decreases, even though the total mass of the system (including the surrounding air) is unchanged.

Absorption from the Surroundings – Apparent Gain of Mass

Sometimes a substance can gain mass because it reacts with something from the air (often oxygen or water vapour).

Closed vs Open Systems

1. In a closed system, no reactants or products can enter or leave.
   1. It is easier to show that mass is conserved.
2. In an open system, gases can escape or be absorbed.
   1. The mass of what you are weighing may change, even though the overall mass of the universe is unchanged.

How Does Conservation of Mass Help Us Understand Reactions?

The idea that mass is conserved helps us to:

1. Balance chemical equations.
2. Predict how much product will form from given reactants.
3. Check whether our reaction descriptions make sense.

Balancing Chemical Equations

1. Because atoms are neither created nor destroyed, a correct chemical equation must have:
2. The same number of each type of atom on both sides.
3. For example, the reaction of hydrogen and oxygen to form water:
   1. Unbalanced: $$\mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
   2. Reactants: 2 H atoms, 2 O atoms
   3. Products: 2 H atoms, only 1 O atom
4. To obey conservation of mass, we balance it: $$2 \mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
   1. Reactants: 4 H, 2 O
   2. Products: 4 H, 2 O → balanced

Predicting Reactant and Product Quantities (Stoichiometry)

Once an equation is balanced, we can use it to calculate how much product will form from given masses of reactants.

Why Must Chemical Equations Be Balanced?

1. When chemists describe a chemical reaction, they write a chemical equation.
2. This is like a recipe: it shows which substances react (reactants) and what new substances are formed (products).
3. For the equation to accurately represent what happens, it must be balanced.

Link to the Law of Conservation of Mass

From the previous article, we know about the law of conservation of mass:

Mass is neither created nor destroyed during a chemical reaction. The total mass of the reactants equals the total mass of the products.

1. Because atoms are the building blocks of mass:
   1. The same number of each type of atom must appear before and after the reaction.
   2. In other words, atoms are rearranged, not created or destroyed.
2. A balanced equation is one where the number of each kind of atom on the left (reactants) = the number of that same atom on the right (products).

What Does a Balanced Equation Reveal?

1. Equality of atoms:
   1. Each element has the same number of atoms on both sides.
2. Predicting quantities:
   1. The coefficients (big numbers in front of formulas) tell us the ratios in which substances react and are produced.

How Do We Systematically Balance Chemical Equations?

Balancing equations is a step-by-step process. Here’s a reliable method you can use.

Step 1: List the Atoms on Each Side

Write the unbalanced equation, then list how many atoms of each element appear on the reactant side and on the product side.

Step 2: Start with the Most Complex Substance

Begin with the formula that:

1. Has the most different elements, or
2. Appears least often in the equation.

This is usually a compound, not a simple element like O₂ or H₂.

Step 3: Use Coefficients, Not Subscripts

1. Coefficients are the big numbers in front of formulas (e.g. 2H₂O).
2. Subscripts are the small numbers inside formulas (e.g. H₂O).

Step 4: Balance Polyatomic Ions as Units (When Possible)

If a polyatomic ion (like SO₄²⁻, NO₃⁻, CO₃²⁻) appears unchanged on both sides, you can balance it as a whole unitinstead of balancing each atom separately.

Step 5: Double-Check Your Work

• Make sure every element is balanced.
• Confirm that coefficients have no common factor (simplify if needed).
• Check that you have not changed any subscripts.

What Do Coefficients Tell Us About Amounts of Reactants and Products?

The coefficients in a balanced equation do more than just balance atoms. They tell us the relative amounts of substances that react and form.

Particles and Moles

1. In the equation: $$2 \mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
2. The coefficients mean:
   1. Particle view: 2 molecules of H₂ react with 1 molecule of O₂ to form 2 molecules of H₂O.
   2. Mole view: 2 moles of H₂ react with 1 mole of O₂ to form 2 moles of H₂O.
3. Because 1 mole = $6.02×10^{23}$ particles (Avogadro’s number), the same ratio applies at the particle level and at the macroscopic (mole) level.

From Moles to Masses

Using the molar mass (mass of 1 mole), we can convert between moles and grams.

Steps:

1. Use the balanced equation to get the mole ratio.
2. Convert the given mass to moles.
3. Use the mole ratio to find the moles of another substance.
4. Convert back to mass if needed.

Scaling Reactions

1. Balanced equations can be scaled up or down:
2. If the equation says: $$\mathrm{CH}_4+2 \mathrm{O}_2 \rightarrow \mathrm{CO}_2+2 \mathrm{H}_2 \mathrm{O}$$ then doubling everything: $$2 \mathrm{CH}_4+4 \mathrm{O}_2 \rightarrow 2 \mathrm{CO}_2+4 \mathrm{H}_2 \mathrm{O}$$ describes burning twice as much methane while keeping the same ratios.
3. This is essential in industry, where reactions must be scaled to produce kilograms or tonnes of product.