Difficulty: Foundational to Intermediate | Prerequisites: Measurement and Notation study notes
Matter is classified as elements, compounds, or mixtures, and it can undergo physical changes (form changes, identity stays) or chemical changes (new substances form). Atoms are built from protons, neutrons, and electrons, and the way electrons arrange themselves around the nucleus, governed by the Aufbau principle, Pauli exclusion, and Hund’s rule, determines an element’s chemistry. The periodic table organises elements by atomic number, and its trends (atomic radius, ionization energy, electronegativity, electron affinity) follow logically from nuclear charge and electron shielding.
Extensive property
A property that depends on the amount of matter in a sample, such as mass or volume.
In simple terms, if you double the sample, the value doubles.
Intensive property
A property that depends only on the type of matter, not the amount. Density and boiling point are intensive.
In simple terms, a cup of water and a bathtub of water have the same density.
Physical change
A change in form or appearance that does not alter the substance’s chemical identity. Examples: boiling, melting, freezing, shredding.
Think of it as: the molecules are the same before and after; only their arrangement or state changed.
Chemical change
A transformation that produces one or more new substances with different chemical identities. Examples: rusting iron, leaves changing colour.
Think of it as: bonds break and new bonds form, so you end up with different molecules.
Element
A pure substance that cannot be broken down into simpler substances by ordinary chemical means.
Compound
A substance formed from two or more elements chemically combined in a fixed ratio. It can be decomposed back into its elements by breaking bonds.
Mixture
Two or more substances physically combined (not chemically bonded). Can be homogeneous (uniform throughout) or heterogeneous (visibly distinct regions).
Solution
A homogeneous mixture of two or more substances.
Atomic number (Z)
The number of protons in the nucleus of an atom. It defines the element.
Mass number (A)
The sum of protons and neutrons in the nucleus.
Isotopes
Atoms of the same element (same number of protons) with different numbers of neutrons.
In simple terms, isotopes are siblings: same element, different weight.
Radioisotope
An isotope with an unstable nucleus that undergoes radioactive decay.
Electron configuration
The arrangement of electrons in the orbitals around the nucleus, written using subshell notation (1s² 2s² 2p⁶, etc.).
Aufbau principle
Electrons fill orbitals starting from the lowest energy level and working upward.
Think of it as: electrons are lazy; they take the lowest-energy seat first.
Pauli exclusion principle
No two electrons in an atom can have the same set of four quantum numbers. Each orbital holds a maximum of two electrons, and they must have opposite spins.
Hund’s rule
Within a set of equal-energy orbitals, electrons occupy them singly (with parallel spins) before any orbital gets a second electron.
Think of it as: electrons are like bus passengers; they each take an empty seat before anyone doubles up.
Ionization energy
The energy required to remove a valence electron from a gaseous atom.
Electronegativity
The tendency of an atom to attract electrons toward itself in a chemical bond.
Electron affinity
The energy change when an atom gains an electron.
Extensive properties change with the amount of matter: mass, volume, length.
Intensive properties stay the same regardless of amount: density, temperature, colour, boiling point.
Density is the classic example of an intensive property. Double the sample and the mass doubles but so does the volume, so density stays constant.
Physical change: the substance’s identity is preserved. Boiling water turns liquid into vapour, but it is still H₂O.
Chemical change: one or more new substances form. Iron rusting produces iron oxide, a different compound entirely.
Signs of chemical change: colour change, gas production, precipitate formation, energy release or absorption.
Elements: pure substances that cannot be decomposed further by chemical means (e.g. gold, oxygen).
Compounds: two or more elements chemically bonded in a fixed ratio (e.g. H₂O, NaCl). Can be broken down by chemical means.
Mixtures: two or more substances physically combined, not chemically bonded.
Homogeneous (uniform composition throughout, e.g. saltwater).
Heterogeneous (visibly different regions, e.g. sand and iron filings).
Solutions: a specific type of homogeneous mixture.
Crystallisation: converting a dissolved substance into a solid crystalline form.
Decanting: carefully pouring off a liquid without disturbing the sediment.
Distillation: separating by boiling points; the liquid is vaporised and then condensed.
Chromatography: separating components based on differences in polarity between two phases.
Filtration: removing solid particles from a liquid or gas using a filter medium.
Evaporation: heating a homogeneous mixture to drive off the solvent, leaving the dissolved solid behind.
Scientist | Contribution |
|---|---|
Democritus | First to propose the idea of the atom (tiny, indivisible particles) |
Dalton | Atomic theory, law of multiple proportions |
Thomson (J.J.) | Discovered the electron via cathode ray tube; proposed the plum pudding model |
Chadwick | Discovered the neutron |
Rutherford | Discovered the nucleus via the gold foil experiment (alpha particles deflected) |
Bohr | Planetary model; electrons in discrete energy levels |
De Broglie | Proposed that electrons have wave-like properties |
Newlands | Law of octaves; early periodic organisation |
Mendeleev | Organised the periodic table by atomic mass |
Moseley | Reorganised the periodic table by atomic number |
Millikan | Measured the charge of an electron (oil-drop experiment) |
Lavoisier | Father of modern chemistry; law of conservation of matter |
Particle | Mass (amu) | Charge | Location |
|---|---|---|---|
Proton | 1 | +1 | Nucleus |
Neutron | 1 | 0 | Nucleus |
Electron | 0.0005 | −1 | Electron cloud (surrounding the nucleus) |
Atomic number (Z): equals the number of protons. It defines which element an atom is.
Atomic mass: the weighted average of the masses of all naturally occurring isotopes, each multiplied by its natural abundance.
Mass number (A): protons + neutrons in the nucleus.
Isotopes: same element (same proton count), different neutron count. Carbon-12 and carbon-14 are isotopes of carbon.
Radioisotopes: isotopes with unstable nuclei that undergo radioactive decay.
Derives from the Schrödinger wave equation. Unlike Bohr’s model, it does not assign fixed orbits to electrons. Instead it gives the probability of finding an electron at a given point around the nucleus.
The photoelectric effect: electrons are ejected from a metal surface when light of sufficient frequency strikes it. This demonstrated that light has particle-like behaviour (photons).
Quantised energy: in certain systems, energy can only take on specific discrete values, not any value along a continuum.
Electrons are distributed among orbitals labelled s, p, d, and f.
s: up to 2 electrons
p: up to 6 electrons
d: up to 10 electrons
f: up to 14 electrons
Aufbau principle: fill lowest-energy orbitals first. The filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.
Pauli exclusion principle: each orbital holds at most 2 electrons with opposite spins (½ and –½).
Hund’s rule: within a subshell of equal-energy orbitals, place one electron in each orbital (all with the same spin) before pairing any.
Shorthand (noble gas) electron configuration: replace the inner-shell configuration with the symbol of the preceding noble gas in brackets, e.g. [Ne] 3s² 3p⁵ for chlorine.
Arrangement:
Organised by increasing atomic number (Moseley’s contribution).
Elements in the same group (column) have the same number of valence electrons and therefore similar chemical properties.
Elements in the same period (row) have the same number of principal energy levels (shells).
Periodic Trends Summary:
Trend | Across a period (left to right) | Down a group (top to bottom) | Why |
|---|---|---|---|
Atomic radius | Decreases | Increases | More protons pull electrons closer across a period; more shells added going down |
Ionic radius | Decreases (for same-charge ions) | Increases | Same logic as atomic radius; cations are smaller than their parent atom, anions are larger |
Ionization energy | Increases | Decreases | Smaller atoms hold electrons more tightly; larger atoms have more shielding |
Electronegativity | Increases | Decreases | Smaller, more highly charged nuclei attract bonding electrons more strongly |
Electron affinity | Increases (generally) | Decreases (generally) | Atoms closer to a full shell gain electrons more readily |
Key reasoning for the trends:
Across a period, the principal energy level stays the same but nuclear charge (proton count) increases, pulling electrons inward. This shrinks the radius and raises ionization energy, electronegativity, and electron affinity.
Down a group, each new period adds a principal energy level (a new shell), increasing the distance between the nucleus and the outermost electrons. Electron shielding also increases, weakening the nuclear pull on valence electrons.
Ionic radius detail:
Cations (atoms that lose electrons) are smaller than their neutral atoms because removing electrons reduces electron-electron repulsion and the remaining electrons are pulled closer.
Anions (atoms that gain electrons) are larger than their neutral atoms because the extra electrons increase repulsion and the electron cloud expands.
Mass number: A = protons + neutrons
Atomic mass: weighted average = Σ (isotope mass × fractional abundance)
Max electrons in a shell: 2n² (where n is the principal quantum number)
Max electrons in a subshell: 4l + 2 (where l is the angular momentum quantum number)
Isotopes are used in medicine (radioactive tracers for imaging), archaeology (carbon-14 dating), and energy (uranium-235 in nuclear power). Periodic trends explain why fluorine is the most reactive non-metal (high electronegativity) and why caesium is the most reactive metal (low ionization energy). Separation techniques like distillation and chromatography are the backbone of industrial chemical purification and forensic analysis.
Students often think atomic mass and mass number are the same thing. They are not. Mass number is a whole-number count of protons plus neutrons in one specific isotope. Atomic mass is a weighted average across all naturally occurring isotopes.
Students often confuse electron shells with orbitals. A shell (n = 2, for example) contains multiple subshells (2s, 2p), and each subshell contains one or more orbitals.
Students often assume that periodic trends have no exceptions. There are notable exceptions (e.g. ionization energy dips between Be and B, and between N and O, because of subshell stability effects). Know the general trend and be aware exceptions exist.
Students often think a cation is larger than its parent atom. It is smaller: removing electrons reduces repulsion and pulls the remaining cloud inward.
⚠️ Know the contributions of each scientist listed in the early atomic theory table. These appear as matching or multiple-choice questions.
⚠️ Electron configuration is heavily tested. Be able to write both full and shorthand configurations, and identify violations of Aufbau, Pauli, or Hund’s rule.
⚠️ Periodic trends: expect questions asking you to rank a set of elements by atomic radius, ionization energy, or electronegativity, and to explain why.
⚠️ Intensive vs. extensive properties often appear as a quick identification question.
True or False: Density is an extensive property.
Fill in the blank: The atomic number equals the number of ___.
True or False: Isotopes of the same element have different numbers of protons.
Fill in the blank: Electron configuration fills orbitals in order of increasing ___ according to the Aufbau principle.
True or False: Electronegativity increases going down a group.
Answers: 1. False (density is intensive). 2. Protons. 3. False (they have different numbers of neutrons, same protons). 4. Energy. 5. False (it decreases down a group).
Q: An atom has 11 protons, 12 neutrons, and 11 electrons. What is its atomic number, mass number, and what element is it?
A: Atomic number = 11 (number of protons). Mass number = 11 + 12 = 23. The element is sodium (Na).
Q: Write the full electron configuration for chlorine (Z = 17).
A: 1s² 2s² 2p⁶ 3s² 3p⁵. Shorthand: [Ne] 3s² 3p⁵.
Q: Rank the following in order of increasing atomic radius: Na, Mg, K.
A: Mg < Na < K. Mg is to the right of Na in the same period (smaller radius). K is below Na in the same group (larger radius).
Q: Which scientist discovered the nucleus, and what experiment did they perform?
A: Rutherford, via the gold foil experiment. Alpha particles were directed at a thin gold foil; most passed through, but some were deflected, indicating a small, dense, positively charged nucleus.
Q: Explain why ionization energy generally increases across a period.
A: Across a period, nuclear charge increases while the principal energy level stays the same. The stronger nuclear pull makes it harder to remove an electron, so more energy is required.
Electron configuration underpins chemical bonding: how atoms share or transfer electrons (covered in the bonding and molecular geometry notes). Periodic trends explain reactivity patterns that drive the types of chemical reactions (synthesis, decomposition, single and double replacement). The quantum mechanical model connects to the light and energy equations used to calculate photon energies and wavelengths.
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