Atomic structure and the periodic table

1. Your periodic table: a map of the elements

The periodic table puts the elements in a useful order. You do not need to memorise every box: start by learning how to read it.

Enlarge the table here — scroll sideways to explore

1. The order: count the protons

Elements are arranged by increasing atomic number. This is the number of protons in an atom’s nucleus. Hydrogen has 1 proton, helium has 2 and lithium has 3. Each step to the next element adds one proton.

In the sodium box, 11 is the atomic number, Na is the chemical symbol and sodium is the element’s name. A neutral sodium atom also has 11 electrons. The number at the top of this table’s boxes is not the relative atomic mass.

2. Rows are called periods

A period runs horizontally across the table. Its number tells you how many electron shells are occupied. Sodium is in Period 3 and has three occupied shells: 2,8,1.

3. Columns are called groups

A group runs vertically down the table. Elements in the main groups studied at GCSE have the same number of outer-shell electrons. This gives them similar chemical reactions.

• Group 1: lithium, sodium and potassium each have one outer electron.
• Group 7: fluorine, chlorine and bromine each have seven outer electrons.
• Group 0: the noble gases have full outer shells. Helium has two electrons; the others have eight in their outer shell.

4. The shape gives you clues

Most metals are on the left and in the centre. Non-metals are mainly towards the top right. Hydrogen is a non-metal, even though it sits on the left. The two rows underneath belong in Periods 6 and 7; moving them down stops the table becoming extremely wide.

Element names and modern numbering checked against IUPAC. This learning table shows atomic numbers, symbols and names; it does not include relative atomic masses.

2. Atoms, elements and compounds

An atom is the smallest part of an element that can exist. An element contains only one type of atom. Oxygen gas contains pairs of oxygen atoms, O₂, but it is still an element: both atoms are oxygen.

A compound contains atoms of different elements chemically joined in a fixed ratio. Water, H₂O, contains two hydrogen atoms for every oxygen atom. The small subscript belongs to the symbol immediately before it. If there is no subscript, there is one atom of that element. Carbon dioxide, CO₂, contains one carbon and two oxygen atoms per molecule.

Compounds have different properties from their elements. Sodium is a reactive metal; chlorine is a poisonous gas; sodium chloride is a white crystalline compound. Making a compound is a chemical reaction: new substances form, often with a detectable energy change. Splitting a compound into its elements also requires a chemical reaction.

Other useful symbols include iron Fe, copper Cu, zinc Zn, silver Ag, gold Au, lead Pb, bromine Br and iodine I. Use the periodic table to check unfamiliar names.

3. Describing reactions with equations

A word equation gives the names of reactants on the left and products on the right. The arrow means “react to form”. For example: magnesium + oxygen → magnesium oxide.

A symbol equation uses formulae. Atoms are rearranged in a reaction; they are not created or destroyed. Balance the equation so that each element has the same number of atoms on each side. Change numbers in front of formulae, never the subscripts in correct formulae.

A binary compound contains two elements. Many names end in -ide: MgO is magnesium oxide, NaCl is sodium chloride and MgBr₂ is magnesium bromide. Names ending in -ate usually include oxygen as well as another element: calcium carbonate is CaCO₃. Learn common formulae from the course; an ending alone does not tell you every atom count.

4. Mixtures: choose a physical separation

In a mixture, substances are together but are not chemically combined. Their proportions can vary and each substance keeps its chemical properties. Air is a mixture of gases; salt water is a mixture of salt and water. Separating a mixture by physical methods does not make a new substance.

Recovering a dissolved solid

Heat a solution gently in an evaporating basin to remove some solvent. Stop before all the liquid has gone; allow the solution to cool and crystals to form. Filter the crystals and dry them. A dissolved substance will pass through ordinary filter paper, so filtering salt solution alone cannot remove the salt.

Recovering a solvent

In simple distillation, heat the solution. The vapour passes through a condenser and cools to a liquid, the distillate. A thermometer bulb at the entrance to the condenser measures vapour temperature. Cooling water enters the lower condenser connection and leaves the upper connection, keeping its jacket full. The apparatus must not be a sealed pressure vessel.

Separating liquids

A fractionating column sits between the heated flask and condenser. Vapour repeatedly condenses and evaporates. The liquid with the lower boiling point is collected first. A simple distillation can give a poor separation when liquid boiling points are close.

Paper chromatography

1. Draw a pencil baseline near the bottom of chromatography paper. Pencil graphite does not dissolve in the solvent.
2. Put small sample spots on the line. Stand the paper in a shallow solvent with the baseline above the solvent level.
3. Allow the solvent to rise, then remove the paper before it reaches the top. Mark the solvent front immediately.
4. Compare spots using the same solvent and conditions. More than one separated spot shows that the sample contains more than one soluble component.

These separation methods are practical skills in this topic. AQA’s numbered chromatography required practical belongs to Chemical analysis; it is not an additional required practical for section 5.1. Use eye protection and follow teacher instructions for hot glassware and solvents. Keep flammable liquids away from naked flames.

5. Why the atomic model changed

A scientific model is an explanation or representation supported by evidence. It can change when new observations cannot be explained by the old model.

1. Solid-sphere model: early models treated atoms as tiny indivisible spheres.
2. Discovery of electrons: negatively charged particles showed that atoms had smaller parts. In the plum-pudding model, negative electrons were embedded in a spread-out ball of positive charge.
3. Alpha scattering: most positively charged alpha particles passed straight through thin gold foil. Some changed direction; a very small number were deflected through large angles or backwards.
4. Nuclear model: most of an atom is empty space. Almost all its mass and positive charge are concentrated in a tiny nucleus. A small, concentrated positive region could strongly repel an alpha particle; a diffuse positive ball could not explain the rare large deflections.
5. Bohr’s model: electrons occupy particular energy levels, or shells, at specific distances from the nucleus. This improved agreement with observations.
6. Protons and neutrons: later evidence identified the proton. Chadwick’s work provided evidence for uncharged neutrons in the nucleus.

6. Subatomic particles, ions and scale

The atomic number is the number of protons. It identifies the element. Every carbon atom has six protons. A neutral atom has equal numbers of protons and electrons, so the positive and negative charges cancel.

The mass number is protons + neutrons. Therefore neutrons = mass number − atomic number. In isotope notation, the mass number is above the atomic number to the left of the symbol.

An ion is a charged atom or group of atoms. Losing negative electrons gives a positive ion; gaining negative electrons gives a negative ion. A chloride ion, Cl⁻, has one more electron than a chlorine atom. Ion formation in ordinary chemical reactions does not change the nucleus.

How small is an atom?

A typical atomic radius is about 0.1 nm = 1 × 10⁻¹⁰ m. Nano means one billionth: 1 nm = 10⁻⁹ m. A nucleus has a radius of roughly 10⁻¹⁴ m, less than about one ten-thousandth of an atom’s radius. Almost all the mass is in this tiny nucleus.

To compare sizes, divide the larger by the smaller. Using the approximate powers of ten, 10⁻¹⁰ ÷ 10⁻¹⁴ = 10⁴, or 10,000. If an atom’s radius were enlarged to 100 m, a nucleus on that approximate scale would have a radius of about 1 cm. Diagrams cannot usually show both clearly at the same scale.

7. Isotopes and relative atomic mass

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Carbon-12 and carbon-14 both have six protons; they have six and eight neutrons respectively. Their mass numbers differ, but they are still carbon.

The relative atomic mass, Aᵣ, is a weighted average that accounts for the natural abundances of the isotopes. It is not necessarily a whole number. Do not confuse it with the whole-number mass number of one isotope.

If abundances are counts or a ratio rather than percentages, divide by the total count or sum of the ratio parts. For three atoms of mass 10 and one of mass 11, the weighted average is (3 × 10 + 1 × 11) ÷ 4 = 10.25.

8. Electron arrangements and the periodic table

Electrons fill the lowest available energy levels first. For the first 20 elements, use shell arrangements of up to 2 in the first shell, then 8 in the second, then 8 before starting the fourth. Use this pattern for the first 20 elements; larger atoms can have more complicated arrangements.

Sodium has atomic number 11, so a neutral sodium atom has 11 electrons: 2,8,1. Chlorine has 17: 2,8,7. Calcium has 20: 2,8,8,2. Count the dots as well as the shells when checking a drawing.

The modern periodic table is arranged by increasing atomic number. A horizontal row is a period; a vertical column is a group. For the main groups studied here, elements in a group have the same number of outer-shell electrons and similar reactions.

• Period number: number of occupied electron shells.
• Group 1: one outer electron; Group 7: seven outer electrons.
• Group 0: full outer shell, with eight electrons except helium, which has two.

Potassium, 2,8,8,1, is in Period 4 and Group 1. Magnesium, 2,8,2, is in Period 3 and Group 2. This allows you to predict chemical behaviour from an element’s position.

9. A short history: spotting patterns, testing predictions

The periodic table developed as scientists found better ways to organise the elements.

1. Early tables — order by mass. Before subatomic particles were known, chemists used atomic weights (relative atomic masses). Some elements were still undiscovered, and strict mass order did not always group similar properties together.
2. Newlands, 1860s — repeating patterns. He noticed that properties often repeated at every eighth element: his “law of octaves”. He left no gaps, so the pattern forced some unlike elements together.
3. Mendeleev, 1869 — leave room for discoveries. He used mass and chemical properties. He sometimes changed the mass order to keep similar elements together, and left gaps for missing elements. He predicted their properties. Later discoveries, including gallium, matched several predictions and supported his table.
4. Moseley, 1913 — atomic number. His X-ray evidence supported ordering elements by atomic number rather than mass. Today we understand atomic number as proton number.

Why can mass order differ? Isotopes have different neutron numbers; average mass depends on their abundances. Argon (18, Aᵣ ≈ 39.9) comes before potassium (19, Aᵣ ≈ 39.1), despite its larger average mass.

Explore the history with the Royal Society of Chemistry.

10. Metals and non-metals

Most elements are metals, found mainly on the left and towards the bottom of the periodic table. Non-metals are mainly towards the top right; hydrogen is a non-metal despite its position on the left.

Metals usually conduct heat and electricity well, are shiny when freshly cut and can be shaped without breaking. Most are solid at room temperature, although mercury is liquid. Non-metals often conduct poorly and, when solid, tend to be brittle; graphite is an important conducting exception.

Metals react by losing electrons to form positive ions. Non-metals do not normally form positive ions in these reactions: they may gain electrons to form negative ions or share electrons in covalent bonds. Outer-shell arrangement helps explain these differences. A Group 1 atom can lose its one outer electron; a Group 7 atom can gain one to complete its outer shell.

11. Group 0: the noble gases

The noble gases include helium, neon and argon. They exist as separate atoms and are very unreactive under ordinary conditions. Their outer shells are full: helium has two electrons; the others have eight outer electrons. These stable arrangements mean the atoms do not readily gain, lose or share electrons.

Going down Group 0, relative atomic mass and boiling point increase. You can use data to predict an unknown value or describe a trend. All the noble gases are gases at room temperature, but they can be liquefied at low temperatures.

Remember that −153 °C is higher than −186 °C. “Higher boiling point” does not necessarily mean a positive temperature.

12. Group 1: the alkali metals

Lithium (Li), sodium (Na) and potassium (K) are soft metals with one outer-shell electron. Rubidium (Rb), caesium (Cs) and francium (Fr) are further down Group 1. They form ions with charge +1 by losing that outer electron.

Reactions to recognise

• With water: metal + water → metal hydroxide + hydrogen. For sodium: 2Na + 2H₂O → 2NaOH + H₂. The hydroxide solution is alkaline.
• With chlorine: metal + chlorine → metal chloride. For lithium: 2Li + Cl₂ → 2LiCl. The product is a white ionic solid.
• With oxygen: the metals react rapidly with oxygen, giving oxide-containing products and tarnishing in air. The simple oxide equation for lithium is 4Li + O₂ → 2Li₂O. At this level, learn the metal + oxygen → metal oxide pattern; actual burning products of sodium and potassium can include other oxygen compounds.

On water, lithium floats and fizzes relatively gently. Sodium usually melts into a moving ball because of heat released. Potassium reacts more vigorously and may burn with a lilac flame. These are teacher demonstrations: do not attempt them yourself.

Reactivity increases down Group 1. The outer electron is farther from the nucleus and more shielded by inner electrons. The attraction to it is weaker, so it is lost more easily, even though the nucleus has more protons.

13. Group 7: the halogens

The halogens have seven outer electrons and occur as diatomic molecules: F₂, Cl₂, Br₂ and I₂. Fluorine is a pale yellow gas; chlorine a green gas; bromine a red-brown liquid; iodine a grey-black solid that produces purple vapour when warmed. Astatine (At) lies below iodine in the group.

Going down Group 7, relative molecular mass, melting point and boiling point increase. At room temperature, chlorine is a gas, bromine a liquid and iodine a solid.

Halogens react with metals to form ionic metal halides. For example, Mg + Cl₂ → MgCl₂. With non-metals, they form covalent compounds: hydrogen + chlorine → hydrogen chloride, H₂ + Cl₂ → 2HCl.

Reactivity decreases down Group 7. A halogen atom reacts by gaining an electron. Down the group, the outer region is farther from the nucleus and more shielded. Attraction for an incoming electron is weaker, so gaining it is harder.

Displacement: use the reactivity order

A more reactive halogen displaces a less reactive halogen from its salt solution. Chlorine displaces bromine from potassium bromide: Cl₂ + 2KBr → 2KCl + Br₂. Bromine displaces iodine from potassium iodide: Br₂ + 2KI → 2KBr + I₂. Iodine cannot displace chlorine from a chloride solution.

Use solutions only under teacher instructions with eye protection; halogens and their vapours can be harmful. A colour change can provide evidence of a displacement, but name the new halogen and explain the reactivity order rather than relying on colour alone.

Plan a separation

A flask contains water, dissolved sodium chloride and insoluble sand. You must collect clean sand, salt crystals and a sample of water. Use the separation illustration in Learn to plan an order of methods. Explain where each substance is collected and include one safety precaution.

Interpret evidence: chromatography

On one chromatogram, sample A produces a blue spot 2 cm above the baseline and a red spot 5 cm above it. Pure dye B produces a blue spot at 2 cm using the same solvent and paper.

Draw it, then check it

On paper, draw the electron shells of oxygen (atomic number 8) and magnesium (12). Draw every electron once. Then write each arrangement as numbers.

Your answer sheet

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One final thought

Use a scientific idea and a specific example in your reflection.