Atomic structure and the periodic table

Explanation

• All substances are made of atoms, and an atom is the smallest part of an element that can exist. Required chemical symbols include the first 20 elements, Groups 1 and 7, and other elements named in the specification.
• Compounds form when elements react and become chemically combined in fixed proportions, so a formula records the elements and their ratio.
• A compound has different properties from its constituent elements and can be separated into elements only by chemical reactions.
• Exam questions require names, symbols and formulae to be distinguished, then specified reactions to be represented using word equations or balanced symbol equations.
• Higher tier: write balanced half equations and ionic equations where appropriate.

Explanation

• A mixture contains two or more elements or compounds that are not chemically combined, so each substance keeps its chemical properties and no new substance is formed.
• Separation therefore uses physical properties.
• Filter an insoluble solid from a liquid; crystallise a dissolved solid from solution; use simple distillation to recover a solvent; use fractional distillation for miscible liquids with different boiling points; and use chromatography when substances move differently because of their solubilities and attractions.
• When asked to suggest a method, identify the relevant property and describe what is collected.
• Examiners reward a linked explanation, not just the name of a technique.

Explanation

• 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.

Explanation

• 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.

Explanation

• An atom has a radius of about $0.1\ \mathrm{nm}$, or $1\times10^{-10}\ \mathrm{m}$. Its nucleus has a radius below one ten-thousandth of the atom’s radius, about $1\times10^{-14}\ \mathrm{m}$, yet contains almost all the mass.
• Protons and neutrons each have relative mass $1$; an electron’s relative mass is very small.
• The mass number is the total number of protons and neutrons, so neutron number equals mass number minus atomic number.
• Isotopes are atoms of the same element with the same proton number but different neutron numbers.
• Exam questions may require scale comparisons, particle-count calculations or an explanation of why isotopes remain the same element.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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’.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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.

Explanation

• 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.