Introduction: The Chemistry "Bookkeeper"
Welcome to the world of Chemical Quantities! Have you ever wondered how scientists know exactly how much medicine to put in a tablet, or how much fuel a rocket needs to reach space? It isn't guesswork—it's chemistry "bookkeeping."
In this chapter, we will learn how to count atoms, weigh molecules, and predict exactly what happens during a chemical reaction. Don't worry if the idea of "maths in science" sounds scary; we will break it down into simple steps that anyone can follow!
4.5.2.1 Chemical Equations: The Recipe of Science
In chemistry, we use chemical equations to show what happens during a reaction. Think of an equation as a recipe: it tells you what you start with (the reactants) and what you end up with (the products).
The Grammar of Chemistry
We use chemical symbols from the Periodic Table to represent elements. For example, O is an atom of oxygen, and Na is an atom of sodium.
When atoms join together, they form formulae. Example: \(H_2O\) tells us there are two hydrogen atoms and one oxygen atom joined together.
State Symbols
Equations also tell us what "state" the chemicals are in. Look out for these letters in brackets:
- (s) = Solid
- (l) = Liquid
- (g) = Gas
- (aq) = Aqueous (this just means the substance is dissolved in water)
Key Takeaway
Reactants are on the left of the arrow, and products are on the right. The arrow means "reacts to produce."
4.5.2.2 The Law of Conservation of Mass
This is one of the most important rules in science: No atoms are lost or made during a chemical reaction.
Imagine you are building a Lego castle. If you take the castle apart and build a boat, you still have the exact same number of bricks. Chemistry is the same! The mass of the products must equal the mass of the reactants.
Why does the mass sometimes change?
If you perform an experiment in an open system (like a beaker without a lid), the mass might seem to change:
- Mass seems to increase: This usually happens because a gas from the air has reacted with the chemicals in the beaker.
- Mass seems to decrease: This usually happens because one of the products was a gas that escaped into the room.
Key Takeaway
Mass is always conserved. If the weight changes, it’s just because a gas moved in or out of your container.
4.5.2.3 Relative Formula Mass (\(M_r\))
Every element has a Relative Atomic Mass (\(A_r\))—this is the "weight" found on the Periodic Table. To find the Relative Formula Mass (\(M_r\)) of a compound, you just add up the masses of all the atoms in its formula.
How to calculate \(M_r\): A Step-by-Step Guide
Let's find the \(M_r\) of water (\(H_2O\)).
- Find the \(A_r\) of each element on the Periodic Table: Hydrogen (H) = 1, Oxygen (O) = 16.
- Count the atoms: There are 2 Hydrogens and 1 Oxygen.
- Add them up: \( (1 \times 2) + 16 = 18 \).
- So, the \(M_r\) of \(H_2O\) is 18.
Quick Review
To find the total mass of a formula: Add up the atomic masses for every atom shown.
4.5.2.4 Amounts in Moles (Higher Tier Only)
Because atoms are so tiny, scientists use a special unit to count them called a mole (abbreviation: mol).
Did you know? One mole of any substance contains exactly \(6.02 \times 10^{23}\) particles! This massive number is called the Avogadro constant.
The Simple Mole Rule
The mass of one mole of a substance in grams is exactly equal to its Relative Formula Mass (\(M_r\)).
Example: Since the \(M_r\) of water is 18, then 1 mole of water weighs exactly 18g.
Calculations Triangle
To switch between mass and moles, use this trick:
\( \text{Mass} = \text{moles} \times M_r \)
\( \text{Moles} = \frac{\text{Mass}}{M_r} \)
4.5.2.5 Calculations based on Equations (Higher Tier Only)
We can use balanced equations to calculate how much of a product we will get. Chemical equations can be read in moles.
Example: \(Mg + 2HCl \rightarrow MgCl_2 + H_2\) means 1 mole of Magnesium reacts with 2 moles of Hydrochloric Acid.
Limiting Reactants
Usually, one chemical gets used up before the others. This is the limiting reactant. Once it’s gone, the reaction stops.
Analogy: If you have 10 slices of bread and 1 slice of cheese, you can only make 1 cheese sandwich. The cheese is the limiting reactant because it limits how many sandwiches you can make.
Key Takeaway
The amount of product you get is always determined by the limiting reactant.
4.5.2.6 Concentrations of Solutions (Higher Tier Only)
Concentration tells us how much of a solid (the solute) is dissolved in a certain volume of liquid (the solvent). We measure this in grams per cubic decimetre (\(g/dm^3\)).
Note: \(1 dm^3\) is the same as 1 litre (1000 \(cm^3\)).
The Formula
\( \text{concentration} (g/dm^3) = \frac{\text{mass of solute } (g)}{\text{volume of solvent } (dm^3)} \)
Common Mistake: If the exam gives you the volume in \(cm^3\), you must divide it by 1000 to turn it into \(dm^3\) before doing the calculation!
Key Takeaway
A "strong" or concentrated solution has a lot of solid in a little bit of water. A "weak" or dilute solution has a little bit of solid in a lot of water.
Final Chapter Summary
- Conservation of Mass: Atoms cannot be created or destroyed. Mass stays the same.
- \(M_r\): Sum of atomic masses in a formula.
- Moles (HT): A way to count atoms. Mass = Moles \(\times M_r\).
- Limiting Reactants (HT): The chemical that runs out first and stops the reaction.
- Concentration (HT): Mass divided by volume. Always check your units (\(dm^3\))!