Difficulty: Beginner | Prerequisites: None (this is where the course starts)
Source: Chem 1 Final Study Guide, General Chemistry (Purdue University, Chhabra)
Tags: chemistry, intro to chemistry, data analysis, significant figures, scientific notation, dimensional analysis, properties of matter, states of matter, atomic structure, nuclear chemistry, radioactivity, half-life, Dalton, Rutherford, Thomson
These chapters lay the groundwork for everything else in the course. You start with what chemistry even is, learn how to handle measurements and units properly, then move into classifying matter and understanding its properties. From there, you get into the structure of the atom itself, from the earliest philosophical models through to nuclear chemistry. If you are behind, this is the place to catch up, because every later topic assumes you are comfortable with sig figs, dimensional analysis, and the basic anatomy of an atom.
Chemistry is the study of matter and its changes. Measuring and reporting data correctly (SI units, sig figs, scientific notation) is foundational. Matter is classified as pure substances or mixtures, and its properties are either physical or chemical. Atoms have a nucleus of protons and neutrons surrounded by electrons, and unstable nuclei undergo radioactive decay measured in half-lives.
Chemistry
The study of matter and the changes it undergoes. In simple terms, it is the science of what stuff is made of and how it transforms.
SI units
The International System of Units, the standard measurement system used in science. The seven base units are second (s), metre (m), kilogram (kg), kelvin (K), mole (mol), ampere (A), and candela (cd).
Derived unit
A unit formed by combining two or more base SI units. Think of it as a "compound unit." Volume (cm³) and density (g/cm³) are common examples.
Density
The ratio of an object's mass to its volume: D = m/V. In simple terms, how much stuff is packed into a given space.
Scientific notation
A way of writing very large or very small numbers as a coefficient (between 1 and 10) multiplied by a power of 10. For example, 602,200,000,000,000,000,000,000 becomes 6.022 x 10²³.
Dimensional analysis
A problem-solving method that uses conversion factors to cancel unwanted units and arrive at the desired unit. Think of it as a unit-tracking strategy.
Accuracy
How close a measured value is to the accepted (true) value.
Precision
How close a series of measurements are to one another, regardless of whether they are close to the true value.
Significant figures (sig figs)
The digits in a measurement that carry meaning about the precision of that measurement. In simple terms, the digits you can trust.
Pure substance
Matter with a uniform and unchanging chemical composition, such as pure water or pure table salt.
Physical property
A characteristic that can be observed or measured without changing the substance's composition (colour, density, melting point).
Chemical property
The ability of a substance to combine with or change into one or more other substances (flammability, reactivity with acid).
Physical change
A change that alters a substance's form without changing its chemical composition (cutting paper, melting ice).
Chemical change
A process that transforms one or more substances into entirely new substances (rusting, combustion).
Law of Conservation of Mass
Mass is neither created nor destroyed during a chemical reaction. Discovered by Antoine Lavoisier. Mass of reactants = mass of products.
Mixture
A combination of two or more pure substances where each retains its own chemical identity.
Heterogeneous mixture
A mixture that does not blend smoothly; individual substances remain visibly distinct (sand and water).
Homogeneous mixture (solution)
A mixture with a constant composition throughout; appears uniform (saltwater).
Element
A pure substance that cannot be broken down into simpler substances by physical or chemical means.
Compound
A combination of two or more different elements combined chemically. Can be broken down by chemical means.
Law of Definite Proportions
A compound is always composed of the same elements in the same proportion by mass, regardless of sample size.
Law of Multiple Proportions
When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers. Think H₂O vs H₂O₂.
Atom
The smallest particle of an element that retains the chemical properties of that element.
Cathode ray experiment
J.J. Thomson's experiment that discovered the electron and led to the "plum pudding" model of the atom.
Gold foil experiment
Ernest Rutherford's experiment that proved atoms have a tiny, dense, positively charged nucleus with mostly empty space around it.
Nucleus
The small, dense, positively charged centre of an atom, containing protons and neutrons.
Proton
A positively charged subatomic particle found in the nucleus.
Neutron
A neutral (no charge) subatomic particle found in the nucleus. Discovered by James Chadwick.
Electron
A negatively charged subatomic particle found outside the nucleus. Its charge was determined by Robert Millikan.
Atomic mass
The weighted average mass of an element's naturally occurring isotopes, calculated as: (% abundance x mass of isotope 1) + (% abundance x mass of isotope 2) + ...
Radioactivity
The spontaneous emission of radiation from an unstable nucleus.
Alpha radiation
Radiation composed of alpha particles (helium nuclei, ⁴₂He). Charge: 2+. Blocked by paper.
Beta radiation
Radiation composed of beta particles (electrons, ⁰₋₁β). Charge: 1-. Blocked by metal foil.
Gamma radiation
High-energy electromagnetic radiation (photons). No mass, no charge. Not completely blocked by lead or concrete.
Half-life
The time required for half of a radioactive sample's nuclei to decay. Amount remaining = (initial amount)(1/2)^n, where n = number of half-lives passed.
Nuclear fission
The splitting of a heavy nucleus into smaller fragments, releasing neutrons and enormous amounts of energy. This is what powers nuclear reactors.
Chain reaction
A self-sustaining series of fission events, where neutrons released from one fission trigger further fissions in nearby nuclei.
Critical mass
The minimum mass of fissile material needed to sustain a chain reaction.
Transmutation
The conversion of one element into another, either spontaneously (radioactive decay) or by bombardment (induced transmutation).
Strong nuclear force
The force that holds protons and neutrons together in the nucleus, overcoming the electrostatic repulsion between protons.
Band of stability
The region on a graph of neutrons vs protons where stable nuclei are found. Above the band: too many neutrons (beta decay). Below: too many protons (positron emission or electron capture).
Chemistry is the study of matter and the changes it undergoes
Lab safety is a significant portion of this chapter (19 rules), but the key ones tested include: always wear goggles, never eat or drink in the lab, add acid to water (not the other way round), use the fume hood for toxic gases, and neutralise spills appropriately
SI base units
Time: second (s)
Length: metre (m)
Mass: kilogram (kg)
Temperature: kelvin (K)
Amount of substance: mole (mol)
Electric current: ampere (A)
Luminous intensity: candela (cd)
Derived units and density
Volume: the space occupied by an object, measured in cm³ or mL (1000 cm³ = 1000 mL = 1 L)
Density = mass / volume, commonly in g/cm³
Rearranged: mass = density x volume, or volume = mass / density
Temperature conversions
K = °C + 273
°C = K – 273
Scientific notation operations
Adding/subtracting: adjust exponents to match, then add/subtract coefficients
Multiplying: multiply coefficients, add exponents
Dividing: divide coefficients, subtract exponents
Always convert the result back to proper scientific notation (coefficient between 1 and 10)
Dimensional analysis
Multiply the given quantity by conversion factors so that unwanted units cancel
Example: 550 m/s to km/min
(550 m / 1 s) x (1 km / 1000 m) = 0.55 km/s
(0.55 km / 1 s) x (60 s / 1 min) = 33 km/min
Percent error
Percent error = |observed value – expected value| / expected value x 100
Significant figures rules
Non-zero digits are always significant
Zeros between non-zero digits are significant
Trailing zeros after the decimal point are significant
Leading zeros (placeholders) are not significant
Counting numbers and defined constants have infinite sig figs
Addition/subtraction: round to the fewest decimal places
Multiplication/division: round to the fewest total sig figs
Classifying matter
Pure substances have uniform, unchanging composition
Elements: cannot be broken down further
Compounds: two or more elements chemically combined
Mixtures retain each component's identity
Heterogeneous: visibly distinct components
Homogeneous (solutions): uniform throughout
Mixtures, colloids, and solutions compared
Mixtures (suspensions): heterogeneous, settle out, show Tyndall effect
Colloids: both heterogeneous and homogeneous traits, do not settle out, show slight Tyndall effect
Solutions: totally homogeneous (clear), do not settle out, no Tyndall effect
Separation techniques
Filtration: separates solids from liquids using a porous barrier
Distillation: separates by boiling point differences
Crystallisation: forms pure solid from a dissolved substance
Chromatography: separates by how quickly components travel across a surface
Decantation: pours off liquid from a settled precipitate
States of matter
Solid: definite shape and volume (not defined by hardness)
Liquid: definite volume, takes shape of container
Gas: fills entire container, no definite shape or volume
Vapour specifically refers to the gaseous state of a substance that is normally a solid or liquid at room temperature
Conservation of mass
Mass(reactants) = Mass(products)
Discovered by Antoine Lavoisier
Early atomic models
Democritus: atoms are solid, indivisible, move through empty space
Aristotle: rejected the idea of empty space
John Dalton: all matter is composed of atoms; atoms of a given element are identical; atoms combine in whole-number ratios; atoms are rearranged (not created or destroyed) in reactions
Key experiments
J.J. Thomson (cathode ray): discovered the electron; proposed the "plum pudding" model
Robert Millikan: determined the charge of the electron
Ernest Rutherford (gold foil): discovered the nucleus; most of the atom is empty space
James Chadwick: discovered the neutron
Calculating atomic mass
Weighted average: Σ (percent abundance x isotope mass) for all isotopes
Types of radioactive decay
Beta decay: emission of an electron; occurs when a nucleus has too many neutrons relative to protons (above the band of stability)
Alpha decay: emission of a helium nucleus (⁴₂He); usually from nuclei with more than 83 protons
Positron emission: a proton converts to a neutron, emitting a positron (opposite charge of an electron)
Electron capture: the nucleus absorbs an inner-shell electron, combining it with a proton to form a neutron
Half-life calculations
Amount remaining = (initial amount)(1/2)^n, where n = number of half-lives
Or: amount remaining = (initial amount)(1/2)^(t/T), where t = elapsed time, T = half-life duration
Nuclear fission
Splitting a heavy nucleus (e.g. U-235) into smaller fragments plus neutrons and energy
The released neutrons can trigger further fissions, creating a chain reaction
A sample needs critical mass to sustain the chain reaction
Ionising radiation
Radiation energetic enough to knock electrons off atoms it collides with
Density: D = m / V
Temperature: K = °C + 273
Percent error: |observed – expected| / expected x 100
Scientific notation multiplication: (a x 10^m)(b x 10^n) = (a x b) x 10^(m+n)
Half-life: Amount remaining = (initial)(1/2)^(t/T)
Atomic mass: Σ (fractional abundance x isotope mass)
Dimensional analysis is used constantly in engineering and pharmacy to convert between unit systems and avoid dosing errors.
Half-life calculations underpin carbon dating (archaeology), medical imaging (technetium-99m has a 6-hour half-life, making it useful for scans without prolonged radiation exposure), and nuclear waste management.
The Law of Conservation of Mass is why chemical equations must be balanced, and it is the entire basis of stoichiometry later in the course.
Students often think that a solid must be hard or rigid. It does not have to be. Wax sculptures are solids despite being soft.
"Vapour" and "gas" are not interchangeable. Vapour specifically means the gaseous form of something that is normally a solid or liquid at room temperature.
Students frequently confuse accuracy and precision. You can be very precise (measurements close together) but completely inaccurate (far from the true value).
A common error with sig figs: students count leading zeros as significant. They are not. The number 0.0045 has only two significant figures.
In half-life problems, students sometimes divide the initial amount by the number of half-lives instead of using the (1/2)^n formula. Each half-life halves the remaining amount, it does not subtract a fixed quantity.
⚠️ Sig fig rules are tested constantly, often embedded inside other problems (stoichiometry, molarity). Get these wrong and you lose marks even when your method is correct.
⚠️ Dimensional analysis appears in nearly every quantitative chapter. Practise setting up conversion factors until it is automatic.
⚠️ Know the three key experiments (Thomson, Rutherford, Millikan) and what each one proved. This is a favourite multiple-choice topic.
⚠️ Half-life calculations show up on nearly every general chemistry final. Be comfortable with both forms of the equation.
⚠️ Know the difference between physical and chemical changes, and be able to classify examples of each.
True or False: The number 0.00320 has four significant figures.
Fill in the blank: The SI base unit for temperature is __________.
True or False: In Rutherford's gold foil experiment, most alpha particles were deflected.
Fill in the blank: Beta decay occurs when a nucleus has too many __________ relative to protons.
True or False: A compound can be separated into simpler substances by physical means.
Q: A sample of carbon-14 has a half-life of 5730 years. If you start with 80 g, how much remains after 17,190 years?
A: 17,190 / 5730 = 3 half-lives. 80 x (1/2)³ = 80 x 1/8 = 10 g.
Q: Convert 25°C to kelvin.
A: K = 25 + 273 = 298 K.
Q: How many significant figures are in 4.050 x 10³?
A: Four. The trailing zero after the 5 is significant because it is to the right of the decimal point.
Q: What is the density of an object with a mass of 45.0 g and a volume of 15.0 cm³?
A: D = 45.0 / 15.0 = 3.00 g/cm³.
Q: Which type of radiation is stopped by a sheet of paper?
A: Alpha radiation.
Q: What is the key difference between an element and a compound?
A: An element cannot be broken down into simpler substances by any chemical or physical means. A compound is made of two or more elements chemically combined and can be broken down by chemical means.
Sig figs and dimensional analysis reappear in every quantitative chapter, especially stoichiometry (Ch. 12), gas laws (Ch. 14), and solution chemistry (Ch. 15).
Atomic structure directly feeds into electron configurations (Ch. 5), periodic trends (Ch. 6-7), and bonding (Ch. 8-9).
Nuclear chemistry connects to thermochemistry (Ch. 16) through the concept of energy released in nuclear vs chemical reactions.
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