4.1 Atomic structure and the periodic table — notes on one page

4.1.1.1 · Atoms, elements and compounds

Definition. An element is a substance made of only one type of atom, and an atom is the smallest part of an element that can exist. A compound contains two or more elements chemically combined in fixed proportions, and can be separated into those elements only by chemical reactions.

• About 100 elements are shown in the periodic table; every atom of an element shares one chemical symbol — O for oxygen, Na for sodium. You must know the symbols of the first 20 elements, Groups 1 and 7, and others named in the specification.
• Compounds form from elements by chemical reactions, which make new substances, often with a detectable energy change. The elements in a compound are chemically combined in a fixed ratio, so a formula records which elements are present and how many atoms of each.
• That fixed ratio separates a compound from a mixture, whose substances keep their own properties.
• Reactions are shown by word equations, or by balanced symbol equations with equal atoms of each element on both sides; only coefficients may change, never the subscripts in a formula.
• Higher tier: AQA restricts "write balanced half equations and ionic equations where appropriate" to Higher tier; Foundation answers use word or balanced symbol equations.

What the examiner wants. A classification answer that names the deciding feature, not the appearance: chemically combined or not, and fixed proportions or not. In equation questions the examiner accepts multiples and ignores state symbols, but never a changed formula.

Exam tip. For a ‘state the difference’ question, contrast both chemical combination and the method of separation.

4.1.1.2 · Mixtures

Definition. A mixture consists of two or more elements or compounds that are not chemically combined together. The chemical properties of each substance in a mixture are unchanged, so a mixture can be separated by physical processes alone.

• The specification names five physical separation processes: filtration, crystallisation, simple distillation, fractional distillation and chromatography. None involves a chemical reaction or makes a new substance, which is why they work on mixtures and not on compounds.
• Filtration separates an insoluble solid from a liquid, because the solid particles are too large to pass through the filter paper. Crystallisation separates a soluble solid from its solution by evaporating the solvent until crystals form; heat gently near the end to avoid decomposing the solid.
• Simple distillation separates a solvent from a solution by boiling it off and condensing the vapour. Fractional distillation separates miscible liquids using the difference in their boiling points, the column giving repeated evaporation and condensation so the lowest-boiling liquid reaches the top first.
• Chromatography separates substances by how they distribute between a stationary phase and a mobile phase.
• You must describe and explain each named process and choose a suitable technique for an unfamiliar mixture.

What the examiner wants. The property the separation exploits, named explicitly — particle size, solubility, or a difference in boiling point — and then the technique. A technique named with no property attached is usually worth nothing.

Exam tip. In a method question, name the physical property that makes the chosen separation work.

4.1.1.3 · The development of the model of the atom (common content with physics)

• Atomic models changed when new experimental evidence could not be explained by the existing model. Atoms were first pictured as indivisible spheres.
• Discovering the electron produced the plum pudding model: negative electrons embedded in a ball of positive charge. In alpha-particle scattering, most particles passed straight through, a few were deflected and very few rebounded.
• This showed that the atom is mostly empty space and that nearly all its mass and positive charge are concentrated in a tiny nucleus, so the nuclear model replaced plum pudding.
• Bohr then proposed electrons at specific distances; later work identified protons and Chadwick supplied evidence for neutrons.
• Examiners expect evidence linked explicitly to each conclusion.

Exam tip. For ‘explain why the model changed’, pair each observation with the conclusion it supports.

4.1.1.4 · Relative electrical charges of subatomic particles

• Protons have relative charge $+1$, neutrons have charge $0$ and electrons have charge $-1$.
• Protons and neutrons are in the nucleus; electrons occupy energy levels around it.
• An atom has no overall electrical charge because it has equal numbers of protons and electrons, so the positive and negative charges cancel.
• The atomic number is the number of protons, and it identifies the element: every atom of one element has the same proton number, while atoms of different elements have different proton numbers.
• In questions, use the atomic number to find protons and, for a neutral atom, electrons; do not infer neutron number unless a mass number is also supplied.

Exam tip. When the particle is an atom, explicitly use ‘protons = electrons’ before concluding that its charge is zero.

4.1.1.5 · Size and mass of atoms

Definition. The mass number of an atom is the sum of its protons and neutrons. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they have the same atomic number and different mass numbers.

• Atoms have a radius of about 0.1 nm, which is 1 x 10⁻¹⁰ m, and the nucleus has a radius less than one ten-thousandth of that, about 1 x 10⁻¹⁴ m. Almost all the mass is in the nucleus, because protons and neutrons each have a relative mass of 1 while an electron's is very small (about 1/1836).
• The atomic number is the number of protons and fixes which element the atom is; the mass number is the total of protons and neutrons, written at the top left of the symbol.
• Subtracting the atomic number from the mass number gives the neutrons, and a neutral atom has as many electrons as protons.
• For an ion, add one electron for each unit of negative charge and remove one for each unit of positive charge, but never change the protons — that would change the element.
• Isotopes differ only in neutrons, so they have identical electronic structures and identical chemical properties.

What the examiner wants. Both halves of the isotope definition in one answer — the same protons and the different neutrons. Answers that give only the difference in mass number, without naming the particle responsible, do not score the second mark.

Exam tip. Write ‘neutrons = mass number − atomic number’ before substituting values.

4.1.1.6 · Relative atomic mass

• Relative atomic mass, $A_r$, is the weighted mean mass of an atom of an element compared with one-twelfth of the mass of a carbon-12 atom.
• It accounts for both the masses of the isotopes and their relative abundances, so the periodic-table value is often not a whole number.
• Multiply each isotope mass by its abundance, add the products, then divide by the total abundance: $A_r=\dfrac{\sum(\text{isotope mass}\times\text{abundance})}{\sum\text{abundances}}$.
• Percentages total $100$, but ratios may have another total.
• In an exam, show the complete weighted calculation because an unsupported rounded answer can lose method marks.

Exam tip. Keep the abundance products visible so the examiner can award the weighted-mean method mark.

4.1.1.7 · Electronic structure

• Electrons occupy the lowest available energy levels, also called shells. For the first 20 elements, use the simple filling pattern $2,8,8,2$: fill an inner shell before starting the next.
• Electronic structure may be written as numbers separated by commas or drawn as electrons on concentric shells.
• For example, sodium has atomic number $11$, so a neutral atom has $11$ electrons and structure $2,8,1$.
• The number of occupied shells gives the period, while the outer-shell electron count explains the group and similar chemical properties of main-group elements.
• Examiners expect the correct total and distribution, not merely the right number of shells.

Exam tip. After drawing shells, add the electrons and check that the total equals the atomic number.

4.1.2.1 · The periodic table

• The modern periodic table arranges elements in increasing atomic, or proton, number. Similar properties recur at regular intervals, so elements with related chemical behaviour are placed in vertical columns called groups.
• Main-group elements in the same group have the same number of outer-shell electrons, which explains their similar reactions.
• The period indicates the number of occupied electron shells.
• Use an electronic structure to locate an element, then use its position to predict likely reactions and reactivity.
• The modern ordering is not by relative atomic mass: atomic number gives each element one unambiguous position and explains apparent mass-order anomalies caused by isotopes.

Exam tip. For a prediction, cite the shared group or outer-electron arrangement before stating the similar property.

4.1.2.2 · Development of the periodic table

• Before subatomic particles were known, scientists tried to classify elements mainly in order of atomic weight. Early tables were incomplete, and keeping a strict weight order sometimes placed elements with different properties together.
• Mendeleev prioritised repeating chemical properties: he left gaps for elements not yet discovered and changed the order in a few places.
• Newly discovered elements filled those gaps and had properties close to his predictions, providing strong evidence for his arrangement.
• Later knowledge of isotopes explained why atomic-weight order can appear inconsistent, while atomic number gives the correct modern sequence.
• Examiners want a chronological account showing how predictions were tested by later evidence.

Exam tip. In a history question, use the chain ‘prediction → discovery → matching properties → support’.

4.1.2.3 · Metals and non-metals

• Metals are elements that react to form positive ions; non-metals do not form positive ions. Most elements are metals, found mainly on the left and towards the bottom of the periodic table, while non-metals lie towards the right and top.
• Typical metals conduct heat and electricity, are strong and malleable, and often have high melting points.
• Solid non-metals are generally brittle and poor conductors, although the specification includes important structure-based exceptions such as graphite.
• Link chemical behaviour to electron arrangement: metal atoms tend to lose outer electrons, while non-metal atoms tend to gain or share electrons.
• In comparisons, state a property and then explain it using particles or electrons where requested.

Exam tip. For ‘explain the difference’, pair a characteristic property with the relevant electron behaviour.

4.1.2.4 · Group 0

• Group 0 elements are the noble gases. Their atoms have stable outer-electron arrangements: eight outer electrons, except helium, whose only shell is complete with two.
• This stability makes them very unreactive, so they do not easily form molecules and exist as single atoms.
• Down the group, relative atomic mass increases and boiling point increases.
• A data trend can be used to predict the boiling point or physical state of another noble gas, while recognising that a prediction is an estimate.
• For explanation marks, low reactivity must be connected directly to the complete outer shell rather than merely describing noble gases as ‘stable’.

Exam tip. When explaining unreactivity, write ‘complete outer shell’ before stating that atoms do not readily react.

4.1.2.5 · Group 1

• Group 1 elements are alkali metals with one electron in their outer shell. They react by losing this electron to form $+1$ ions, so their reactions are similar.
• Lithium, sodium and potassium react with oxygen to form oxides, with chlorine to form chlorides, and with water to form a metal hydroxide plus hydrogen.
• The water reactions become more vigorous down the group.
• Reactivity increases because the outer electron is farther from the nucleus and experiences more shielding, so the attraction to the nucleus is weaker and the electron is lost more easily.
• Examiners may ask for observations, products, balanced equations or a prediction for a lower element.

Exam tip. A reactivity-trend explanation needs distance, shielding, nuclear attraction and ease of electron loss.

4.1.2.6 · Group 7

• Group 7 elements are halogens: non-metals whose atoms have seven outer-shell electrons. They exist as diatomic molecules such as $\mathrm{Cl}_2$, and form ionic halides with metals or covalent compounds with non-metals.
• Down the group, relative molecular mass, melting point and boiling point increase, but reactivity decreases.
• A halogen reacts by gaining one electron; farther down the group, the outer shell is farther from the nucleus and more shielded, so attracting an incoming electron is harder.
• A more reactive halogen displaces a less reactive halogen from an aqueous halide solution.
• Examiners expect correct trends, observations or equations and an electron-based explanation.

Exam tip. For displacement, compare the two halogens’ positions before writing the products.

4.1.3.1 · Comparison with Group 1 elements (chemistry only)

• Transition elements are metals, including chromium, manganese, iron, cobalt, nickel and copper. Compared with Group 1 metals, they generally have higher melting points and densities and are stronger and harder.
• They are also much less reactive with water, oxygen and halogens.
• These differences make transition metals suitable where a material must keep its shape, withstand force or avoid rapid reaction, whereas Group 1 metals are soft and highly reactive.
• Questions often provide unfamiliar data and require an element to be classified or compared.
• Conclusions must use the data and stated general patterns, with a named transition element supporting broad statements when requested.

Exam tip. In a comparison, use paired language such as ‘higher density but lower reactivity than Group 1’.

4.1.3.2 · Typical properties (chemistry only)

• Transition elements show characteristic chemical properties as well as typical metallic ones. Many form ions with different charges, form coloured compounds and are useful catalysts.
• For example, iron can form $\mathrm{Fe}^{2+}$ and $\mathrm{Fe}^{3+}$ ions, and iron is the catalyst in the Haber process.
• Compounds of chromium, manganese, iron, cobalt, nickel and copper provide examples of coloured transition-metal compounds.
• A catalyst increases reaction rate without being used up overall; being coloured does not itself make a compound a catalyst.
• In exams, identify the property shown by information in the question and give a named example rather than assuming every transition element shows every property in the same way.

Exam tip. When asked for a typical property, state it precisely and attach it to a named transition element or compound.