Why chemistry matters
Chemistry is the study of matter: what it is made of, how it behaves and how it changes. It is in everything around you.
- Health: medicines like paracetamol and cancer drugs are made by chemists.
- Food: fertilisers and safe food preservatives come from chemistry.
- Materials: plastics, cement, steel, phone batteries and solar cells.
- Environment: chemists look for fuels and fridge gases that do not harm the ozone layer or warm the planet.
This chapter gives you the basic tools: how to describe matter, how to measure it and how to count atoms by weighing.
Nature and properties of matter
Matter is anything that has mass and occupies space. Air, water, a book and you are all matter.
Three states of matter
- Solid: particles are packed tightly in fixed places. Fixed shape and fixed volume.
- Liquid: particles are close but can slide. Fixed volume, but it takes the shape of the container.
- Gas: particles are far apart and move fast. No fixed shape or volume; it fills any container.
Heating or cooling changes one state into another: ice ⇌ water ⇌ steam.
Mixtures and pure substances
- Mixture: two or more substances mixed in any ratio. Homogeneous mixtures look the same everywhere (sugar in water, air). Heterogeneous mixtures do not (sand in water). Mixtures can be separated by physical methods such as filtering or distillation.
- Pure substance: fixed make-up. It is either an element (only one kind of atom, like copper or oxygen) or a compound (two or more elements joined in a fixed ratio by mass, like water H₂O or carbon dioxide CO₂). A compound can be broken into elements only by chemical methods, and its properties are different from those of its elements.
Physical and chemical properties
A physical property can be measured without changing what the substance is: colour, smell, melting point, boiling point, density. A chemical property shows up only when the substance changes into something new: burning, rusting, reacting with acid.
Measurement and SI units
Every measurement has a number and a unit. "5" means nothing; "5 g" means something. Scientists all over the world use the SI system with seven base units:
| Quantity | Unit | Symbol |
|---|---|---|
| Length | metre | m |
| Mass | kilogram | kg |
| Time | second | s |
| Electric current | ampere | A |
| Temperature | kelvin | K |
| Amount of substance | mole | mol |
| Luminous intensity | candela | cd |
Prefixes make units bigger or smaller: kilo (10³), centi (10⁻²), milli (10⁻³), micro (10⁻⁶), nano (10⁻⁹).
Mass and weight
Mass is the amount of matter; it is the same everywhere and is measured with a balance. Weight is the pull of gravity on it, so it changes from Earth to Moon.
Volume, density and temperature
Volume unit: m³. In the lab we use litre: 1 L = 1000 mL = 1000 cm³ = 1 dm³. Density = mass ÷ volume (kg m⁻³ or g cm⁻³). Temperature scales: K = °C + 273.15, and °F = (9/5)°C + 32. The kelvin scale never goes negative.
Uncertainty in measurement
Scientific notation
Very big or very small numbers are written as N × 10ⁿ, where N is between 1 and 10. Example: 0.00016 = 1.6 × 10⁻⁴; 232 500 = 2.325 × 10⁵. To multiply, multiply the N parts and add the powers. To add, first make the powers the same.
Precision and accuracy
Precision = how close repeated readings are to each other. Accuracy = how close a reading is to the true value. You can be precise but not accurate (a balance that is set wrong gives the same wrong answer every time).
Significant figures
Significant figures are the digits you are sure of plus one last guessed digit. Rules:
- All non-zero digits count: 285 → 3.
- Zeros between non-zero digits count: 2.005 → 4.
- Zeros before the first non-zero digit do not count: 0.0032 → 2.
- Zeros at the end count if there is a decimal point: 0.200 → 3; 100. → 3; but 100 → 1.
- Exact numbers (counted objects, defined values like 1 m = 100 cm) have unlimited significant figures.
Calculations with significant figures
Add / subtract: keep as many decimal places as the number with the fewest decimal places. 12.11 + 18.0 + 1.012 = 31.122 → 31.1.
Multiply / divide: keep as many significant figures as the number with the fewest. 2.5 × 1.25 = 3.125 → 3.1.
Rounding off
If the digit to drop is more than 5, raise the one before it by 1 (1.386 → 1.39). Less than 5: leave it (4.334 → 4.33). Exactly 5: make the previous digit even (6.35 → 6.4, 6.25 → 6.2).
Dimensional analysis (factor-label method)
To change units, multiply by a unit factor equal to 1, such as (1000 m / 1 km). Units cancel like numbers. Example: 3 h → s: 3 h × (60 min / 1 h) × (60 s / 1 min) = 10 800 s.
Try it at home
Measure the length of your notebook five times with a ruler and write every reading to 0.1 cm. Are your readings precise? Now compare with a friend's ruler. Are they accurate? Then type your reading into the 3D above and count its significant figures.
Key formulas and definitions
- Density = mass ÷ volume
- K = °C + 273.15; °F = (9/5)°C + 32
- 1 L = 1000 mL = 1000 cm³ = 1 dm³
- Scientific notation: N × 10ⁿ with 1 ≤ N < 10
- Add/subtract → fewest decimal places; multiply/divide → fewest significant figures
- Unit factor: (new unit / old unit) = 1
Worked examples
1. How many significant figures are in (a) 0.0025 (b) 208 (c) 5005 (d) 126 000 (e) 500.0 (f) 2.0034?
Line 1: (a) 0.0025 – front zeros never count → 2. Line 2: (b) 208 – middle zero counts → 3. Line 3: (c) 5005 → 4. Line 4: (d) 126 000 – no decimal point, end zeros do not count → 3. Line 5: (e) 500.0 – decimal point, so end zeros count → 4. Line 6: (f) 2.0034 → 5.
2. Write in scientific notation: (a) 0.0048 (b) 234 000 (c) 8008 (d) 500.0 (e) 6.0012.
(a) move the point 3 places right → 4.8 × 10⁻³. (b) move 5 places left → 2.34 × 10⁵. (c) 8.008 × 10³. (d) 5.000 × 10² (keep all 4 significant figures). (e) 6.0012 × 10⁰ = 6.0012.
3. Round each to 3 significant figures: 34.216, 10.4107, 0.04597, 2808.
34.216 → the next digit is 1 (less than 5) → 34.2. 10.4107 → next digit 1 → 10.4. 0.04597 → next digit 7 → 0.0460 (the zero at the end must stay). 2808 → next digit 8 → 2810, better written 2.81 × 10³.
4. Add 3.6 g + 12.35 g + 0.408 g and give the answer to the correct figures.
Line 1: raw sum = 16.358 g. Line 2: fewest decimal places is 3.6 (one place). Line 3: round to one decimal place → 16.4 g.
5. A block of metal has mass 25.0 g and volume 3.2 cm³. Find its density with correct significant figures.
Line 1: density = mass ÷ volume = 25.0 ÷ 3.2. Line 2: calculator gives 7.8125 g cm⁻³. Line 3: 3.2 has only 2 significant figures, so keep 2. Line 4: density = 7.8 g cm⁻³.
6. Convert 2 days into seconds using dimensional analysis.
Line 1: 2 d × (24 h / 1 d) = 48 h (days cancel). Line 2: 48 h × (60 min / 1 h) = 2880 min (hours cancel). Line 3: 2880 min × (60 s / 1 min) = 172 800 s. Answer: 1.728 × 10⁵ s (2 is a counted exact number).
7. Convert a body temperature of 37 °C to kelvin and to °F.
Line 1: K = 37 + 273.15 = 310.15 K ≈ 310 K. Line 2: °F = (9/5) × 37 + 32 = 66.6 + 32 = 98.6 °F.
8. Three students weigh a 2.00 g sample. A gets 1.95, 1.93 g; B gets 1.94, 2.05 g; C gets 2.01, 1.99 g. Who is accurate and precise?
Line 1: A's readings are close to each other (precise) but far from 2.00 (not accurate). Line 2: B's readings are far apart and their mean 1.995 ≈ 2.00: accurate on average but not precise. Line 3: C's readings are close to each other and to 2.00: both accurate and precise.
Common mistakes
- Counting the front zeros in 0.0045 as significant. They only show where the decimal point is; the answer is 2.
- Writing 100 when you mean 3 significant figures. Write 1.00 × 10² instead.
- Rounding in the middle of a calculation. Keep extra digits and round only the final answer.
- Mixing up mass and weight, or precision and accuracy. Mass stays the same on the Moon; precise readings can still all be wrong.