Welcome to the starting line of Chemistry. This LO covers the scientific method, measurement, significant figures, dimensional analysis, and matter classification. These aren't "easy intro topics" – they're the tools you'll use in every single lab, calculation, and exam question for the rest of the year. Master them now, and Chemistry becomes logical. Skip them, and you'll struggle everywhere.
A systematic, evidence-based approach to understanding the natural world through observation, experimentation, and logical reasoning. From the Latin scientia = knowledge.
Science isn't just a collection of facts in a textbook – it's a process. A way of thinking. A method for turning "I wonder why..." into "Now I know because..."
Study the physical and natural world. Split into two main categories:
Physical Sciences – deal with non-living matter and energy:
Life Sciences – deal with living organisms:
Formal sciences provide the tools (equations, models, algorithms) used by all other sciences.
Chemistry is sometimes called the "Central Science" because it bridges physics (which explains why atoms behave as they do) and biology (which studies the living systems built from molecules). You can't fully understand biology without chemistry, and you can't fully understand chemistry without physics.
The scientific method is the backbone of all scientific discovery. It's not a rigid recipe – it's a cyclical, self-correcting process that scientists use to investigate questions about the natural world.
Arab scientist Hassan Ibn al-Haytham (known in the West as Alhazen) was the first to establish the modern scientific method. In his groundbreaking Book of Optics (Kitab al-Manazir), he insisted that theories must be verified through systematic experimentation – not just philosophical reasoning. He proved that light travels in straight lines and enters the eye (rather than the eye emitting light), disproving the Greek emission theory. He is considered the father of modern optics and the pioneer of the scientific method.
English philosopher Francis Bacon published Novum Organum, formalising the inductive method – drawing general conclusions from specific observations. He emphasised that science must be based on empirical evidence, not on ancient authority.
French philosopher René Descartes published Discourse on the Method, establishing a framework of systematic doubt and deductive reasoning. Together with Bacon's inductive approach, this formed the foundation of modern scientific inquiry.
Notice something interesting in the natural world. Use your senses or instruments to gather information. Good science starts with careful, detailed observation.
Example: You notice that a metal spoon left in hot soup becomes hot at the handle, but a wooden spoon does not.
Ask a specific, testable question about your observation.
Example: "Does the material a spoon is made from affect how quickly heat travels through it?"
Propose a testable, falsifiable explanation – an educated guess based on existing knowledge. A good hypothesis makes a prediction that can be tested.
Example: "If the material affects heat transfer, then metal spoons will transfer heat faster than wooden spoons."
Design a controlled experiment to test your hypothesis. Change only the independent variable, measure the dependent variable, and keep everything else constant (controlled variables).
Example: Place metal, wooden, and plastic spoons in identical cups of hot water. Measure handle temperature at regular intervals.
Collect, organise, and analyse your results. Use tables, graphs, and statistical tools. Look for patterns and determine whether the data supports or refutes your hypothesis.
State whether the results support or refute your hypothesis. A hypothesis is never "proven" – it is either supported by evidence or not. If refuted, modify the hypothesis and test again.
Share your findings through publication, peer review, and replication. Other scientists must be able to repeat your experiment and get similar results. This is how science builds trust and knowledge.
Common exam trap: A hypothesis is never proven – it is either supported or refuted by evidence. Even after thousands of supporting experiments, it remains a hypothesis until it becomes part of a broader theory. Science is about evidence, not proof.
These two terms are often confused – even by adults. They are NOT the same thing, and one is NOT "better" than the other.
A concise statement (often mathematical) that describes what happens in nature under specific conditions. A law tells you the pattern but NOT why it happens.
Example: Newton's Law of Gravity tells you that all objects with mass attract each other – but it doesn't explain why mass creates gravitational pull.
Key traits:
A well-tested, comprehensive explanation of why something happens in nature. A theory is supported by a large body of evidence from multiple experiments and observations.
Example: Einstein's Theory of General Relativity explains why gravity works – it's the curvature of spacetime caused by mass.
Key traits:
Exam-critical distinction: A law describes what happens. A theory explains why it happens. A theory does NOT become a law when it's "proven enough." They are fundamentally different types of scientific knowledge. A theory can contain laws, but a law never becomes a theory.
| Feature | Scientific Law | Scientific Theory |
|---|---|---|
| Answers | What happens? | Why does it happen? |
| Form | Usually a mathematical equation | A detailed explanation |
| Scope | Narrow – specific phenomenon | Broad – explains many related phenomena |
| Example | Boyle's Law: | Kinetic Molecular Theory |
| Can be disproven? | Yes | Yes |
Understanding variables is essential – every lab, every experiment, every graph question on your exam depends on this.
The variable YOU change in an experiment. It's the cause – the factor you deliberately manipulate.
Also called: manipulated variable (المتغير المستقل)
How to identify it: Ask "What am I changing on purpose?" That's your independent variable.
On a graph: Always plotted on the x-axis (horizontal).
Example: In an experiment testing how temperature affects reaction rate, temperature is the independent variable – you control it.
When you plot independent vs dependent variables on a graph, you'll see one of these patterns:
As the independent variable increases, the dependent variable increases proportionally.
Graph shape: Straight line going up from left to right.
Where is a positive constant (the slope).
Example: Hooke's Law – the more force you apply to a spring, the more it stretches:
Correlation does NOT mean causation. Just because two variables change together doesn't mean one causes the other.
Classic example: Ice cream sales and drowning rates both increase in summer. Does ice cream cause drowning? No – both are caused by a third variable: hot weather.
In science, you can only claim causation if you've conducted a controlled experiment with proper variable management. Observational data can only show correlation.
Science depends on precise, universal measurement. That's why we use the SI system (Système International d'Unités) – the international standard for measurement used by scientists worldwide.
These are the foundation – every other unit in science is derived from these seven:
| Quantity | SI Unit | Symbol |
|---|---|---|
| Length (الطول) | metre | m |
| Mass (الكتلة) | kilogram | kg |
| Time (الزمن) | second | s |
| Temperature (درجة الحرارة) | kelvin | K |
| Amount of Substance (كمية المادة) | mole | mol |
| Electric Current (شدة التيار) | ampere | A |
| Luminous Intensity (شدة الإضاءة) | candela | cd |
Exam tip: The kilogram is the only base SI unit that includes a prefix (kilo-). This is a historical quirk – the gram was too small for practical use, so the kilogram became the standard. When converting, always convert to grams first, then to your target unit.
Derived units are combinations of base units. Some common ones:
| Quantity | Derived Unit | Symbol | In Base Units |
|---|---|---|---|
| Area | square metre | m² | m × m |
| Volume | cubic metre | m³ | m × m × m |
| Density | kg per cubic metre | kg/m³ | kg × m⁻³ |
| Velocity | metre per second | m/s | m × s⁻¹ |
| Force | newton | N | kg × m × s⁻² |
| Pressure | pascal | Pa | kg × m⁻¹ × s⁻² |
| Energy | joule | J | kg × m² × s⁻² |
SI prefixes let you express very large or very small numbers efficiently:
| Prefix | Symbol | Factor | Example |
|---|---|---|---|
| tera- | T | 1 TB = bytes | |
| giga- | G | 1 GHz = Hz | |
| mega- | M | 1 MW = W | |
| kilo- | k | 1 km = m | |
| hecto- | h | 1 hL = L | |
| deca- | da | 1 dam = 10 m | |
| – | – | base unit | |
| deci- | d | 1 dm = 0.1 m | |
| centi- | c | 1 cm = 0.01 m | |
| milli- | m | 1 mm = 0.001 m | |
| micro- | μ | 1 μm = m | |
| nano- | n | 1 nm = m | |
| pico- | p | 1 pm = m |
Memory trick: For the most common prefixes, remember "King Henry Doesn't Usually Drink Chocolate Milk" – Kilo, Hecto, Deca, Unit, Deci, Centi, Milli. Each step is a factor of 10.
The amount of matter in an object. Mass is a scalar quantity – it has magnitude only.
The gravitational force acting on an object's mass. Weight is a vector quantity – it has both magnitude and direction (always toward the centre of gravity).
Where = weight (N), = mass (kg), and = gravitational acceleration ( on Earth).
No measurement is perfectly exact. Every measurement has some degree of uncertainty – and a good scientist always reports it.
When you read an instrument (ruler, thermometer, graduated cylinder), you should:
A ruler has markings every 0.1 cm. You measure an object and see it falls between 2.3 cm and 2.4 cm, closer to the 2.3 side. You estimate the last digit as 2.34 cm.
Errors that are consistent and repeatable – they push all measurements in the same direction (always too high or always too low).
Causes:
Key trait: Can be identified and corrected by calibrating equipment or improving technique.
Example: A scale that reads 0.5 g too high – every measurement will be 0.5 g above the true value.
Errors that are unpredictable – they cause measurements to scatter both above and below the true value.
Causes:
Key trait: Cannot be eliminated, but can be reduced by taking multiple measurements and averaging.
Example: Timing a reaction by hand – sometimes you start the stopwatch slightly early, sometimes slightly late.
Percentage error tells you how far off your measurement is from the accepted (true) value:
You measure the density of aluminium as 2.85 g/cm³. The accepted value is 2.70 g/cm³.
When you combine measurements in calculations, errors propagate:
When adding or subtracting measurements, add the absolute uncertainties:
Example:
These two words sound similar, but they mean completely different things in science.
How close your measurement is to the true (accepted) value. Accuracy answers the question: "Did I hit the target?"
How close your repeated measurements are to each other. Precision answers the question: "Are my shots clustered together?"
Imagine shooting a basketball at a hoop:
| Scenario | What it looks like | Description |
|---|---|---|
| Accurate AND Precise | All shots go through the hoop, clustered tightly | Close to true value, consistent with each other |
| Accurate but NOT Precise | Shots scattered around the hoop – some left, some right – but centre on the basket | Average is close to true value, but individual shots are spread out |
| Precise but NOT Accurate | All shots hit the same spot – but it's the wrong spot (e.g., all hit the backboard) | Far from true value, but consistent with each other |
| Neither Accurate nor Precise | Shots are scattered everywhere, none near the hoop | Far from true value, inconsistent |
Exam trap: Precision does NOT guarantee accuracy. You can have a very precise set of measurements that are all wrong (systematic error). In the basketball analogy, hitting the same wrong spot every time = precise but not accurate.
Three students measure the boiling point of water (accepted value: 100.0 °C):
| Student | Trial 1 | Trial 2 | Trial 3 | Average | Assessment |
|---|---|---|---|---|---|
| A | 99.8 | 100.1 | 100.0 | 100.0 | Accurate AND Precise |
| B | 95.0 | 100.0 | 105.0 | 100.0 | Accurate, NOT Precise |
| C | 97.2 | 97.3 | 97.1 | 97.2 | NOT Accurate, Precise |
| D | 94.0 | 101.0 | 108.0 | 101.0 | NOT Accurate, NOT Precise |
Significant figures (sig figs) tell you the precision of a measurement. They include all certain digits plus one estimated digit.
Every digit from 1–9 counts.
Examples:
Zeros sandwiched between non-zero digits always count.
Examples:
Zeros that come before the first non-zero digit are just placeholders – they don't count.
Examples:
If a number has a decimal point, trailing zeros after the last non-zero digit are significant – they indicate measured precision.
Examples:
Trailing zeros in a whole number with no decimal point are generally considered NOT significant.
Examples:
Solution: Use scientific notation to remove ambiguity:
Exact numbers (defined quantities, counting numbers) have infinite precision and never limit your sig fig count.
Examples:
Rule: Your answer can have only as many significant figures as the measurement with the fewest sig figs.
Example:
Why? 4.56 has 3 sig figs. 1.4 has 2 sig figs. The answer is limited to 2 sig figs.
More examples:
Round down – keep the preceding digit the same.
Example: 3.44 rounded to 2 sig figs = 3.4
Round up – increase the preceding digit by 1.
Example: 3.46 rounded to 2 sig figs = 3.5
Round to the nearest EVEN number (banker's rounding). This eliminates the systematic upward bias of always rounding 5 up.
Examples:
Exam tip: Some textbooks always round 5 up (the "round half up" rule). In STEM exams, check which rule your teacher follows. The banker's rounding rule (round to even) is scientifically more accurate, but both may be accepted.
Dimensional analysis (also called the factor-label method or unit-factor method) is a powerful technique for converting between units. The core idea: multiply by conversion factors that equal 1.
A conversion factor is a fraction where the numerator and denominator represent the same quantity in different units:
Since the fraction equals 1, multiplying by it changes the unit without changing the value.
Write down what you're given, including the unit.
Arrange the conversion factor so the unwanted unit cancels out (appears in both numerator and denominator).
Cross out matching units, then multiply/divide the numbers.
Verify the unit is correct and the magnitude is reasonable.
Common exam mistake: When converting squared or cubed units, you must square or cube the entire conversion factor:
Temperature measures the average kinetic energy of particles in a substance. There are three scales you need to know:
Based on water:
Most commonly used scale in daily life and science (outside the US). Named after Swedish astronomer Anders Celsius.
The SI unit of temperature. Based on absolute zero – the lowest possible temperature where all particle motion theoretically stops.
Note: Kelvin uses NO degree symbol. Write "300 K" not "300 °K".
Based on brine and body temperature:
Used primarily in the United States. Named after German physicist Daniel Gabriel Fahrenheit.
| Point | Celsius | Kelvin | Fahrenheit |
|---|---|---|---|
| Absolute Zero | −273.15 °C | 0 K | −459.67 °F |
| Water Freezes | 0 °C | 273.15 K | 32 °F |
| Human Body | 37 °C | 310.15 K | 98.6 °F |
| Water Boils | 100 °C | 373.15 K | 212 °F |
Exam shortcut: For quick conversions, you can use 273 instead of 273.15. Most exam questions accept this approximation:
Convert 25 °C to Kelvin and Fahrenheit:
To Kelvin:
To Fahrenheit:
The mass per unit volume of a substance. It tells you how tightly packed the matter is in a given space.
Where = density, = mass, and = volume.
| System | Common Units |
|---|---|
| SI | kg/m³ |
| Lab/CGS | g/cm³ or g/mL |
Key relationship: 1 g/cm³ = 1 g/mL = 1000 kg/m³
The density of pure water at 4 °C = exactly 1.000 g/cm³ (or 1000 kg/m³). This is a reference value you should memorise.
A metal block has a mass of 156.0 g and a volume of 57.8 cm³. What is its density?
This density matches aluminium (2.70 g/cm³) – so the block is likely aluminium.
The ratio of a substance's density to the density of a reference substance (usually water at 4 °C). Since it's a ratio of two densities, specific gravity is dimensionless – it has no units.
Since the density of water = 1.000 g/cm³, the specific gravity of a substance is numerically equal to its density in g/cm³ – but without units.
Example: Aluminium has a density of 2.70 g/cm³, so its specific gravity = 2.70 (no units).
Practical rule:
Everything in the universe is made of matter – anything that has mass and occupies space (volume).
Examples: ice, iron, diamond, salt crystals
Anything with mass and volume. All matter is classified as either a pure substance or a mixture.
Has a fixed, definite composition and consistent properties throughout. Two types:
A combination of two or more substances where each retains its original properties. The components can be separated by physical means. Two types:
Properties that can be observed or measured without changing the substance's chemical identity.
Examples:
Properties that describe a substance's ability to undergo chemical changes – to transform into a different substance.
Examples:
A change that alters a substance's form or appearance but NOT its chemical identity. No new substance is formed.
Key indicators:
Examples:
A change that transforms a substance into a completely new substance with different properties.
Key indicators:
Examples:
Exam trap: Dissolving salt in water is a physical change – the NaCl molecules separate but remain NaCl. You can get the salt back by evaporating the water. But dissolving a metal in acid is a chemical change – a new substance (a salt) forms.
Since mixtures are held together by physical means (not chemical bonds), they can be separated using physical techniques:
Questions sourced from past STEM exams, reference materials, and chemistry question banks. Use these to test your understanding of everything covered in this lesson.
Chemistry: A Molecular Approach (Tro):
Chemistry: The Central Science (Brown, LeMay, Bursten):
STEM Curriculum References:
Scientific Method – Historical Sources:
Video Resources:
Example: