1. Atomic structure

An atom is mostly empty space. A very small, dense, positively charged nucleus contains protons and neutrons, with electrons arranged in shells around it.

The nucleus contains nearly all of an atom’s mass because protons and neutrons are much more massive than electrons. In a neutral atom, the positive charge of the protons is balanced by the negative charge of the electrons.

The number of protons identifies the element; changing the number of neutrons gives an isotope, while changing the number of electrons gives an ion. These changes do not move the nucleus from the centre of the atom.

Shells describe allowed electron energy levels or regions, not fixed miniature planetary orbits. Detailed electron configurations, orbital shapes and atomic-radius trends belong to the neighbouring objectives.

Atoms contain three main subatomic particles: protons and neutrons in the nucleus, and electrons in shells around the nucleus. The particles differ in relative charge, relative mass and location.

Protons have relative charge +1 and relative mass 1; neutrons have charge 0 and relative mass 1; electrons have charge −1 and a much smaller relative mass. Protons and neutrons are nuclear particles, whereas electrons occupy the space around the nucleus.

Because protons and neutrons are much more massive than electrons, nearly all atomic mass is concentrated in the nucleus. In a neutral atom, equal numbers of protons and electrons make the total charge zero.

Relative values are used to compare the particles; they are not exact masses. A change in electron number makes an ion, but changing the proton number changes the element. Detailed electron configurations belong to the neighbouring objectives.

The atomic number, Z, is the number of protons in the nucleus. It identifies the element because every atom of that element has the same proton number.

The mass number, A, is the total number of protons and neutrons in the nucleus. Therefore, neutron number = A − Z. Electrons are not included in the mass number because their mass is negligible relative to nuclear particles.

A nuclide can be written as ^A_ZX: A is the mass (nucleon) number at the upper left, Z is the atomic (proton) number at the lower left, and X is the element symbol. Read both numbers before identifying the particle counts.

Changing Z changes the element; changing A while keeping Z fixed changes the isotope. Do not confuse mass number for one nuclide with relative atomic mass, which is an average based on isotope abundances.

An atom is mostly empty space, but its mass and charge are not spread evenly. Nearly all the mass and all the positive charge are concentrated in a very small nucleus.

Protons and neutrons each have relative mass about 1, whereas an electron has a much smaller relative mass. This is why the nucleus contains nearly all the atom’s mass even though electrons occupy the surrounding region.

The nucleus is positive because it contains protons; the surrounding electrons contribute negative charge. In a neutral atom these charges balance overall, although the positive charge remains concentrated at the centre.

Do not describe an atom as a solid sphere or place electrons inside the nucleus. Mass concentration and charge concentration are related but distinct ideas; particle identities and electron arrangements are handled by neighbouring objectives.

In an electric field, a charged particle experiences a force toward the oppositely charged plate; a neutral particle is not deflected by this electric force. Comparing beams at the same speed reveals how charge and mass affect the deflection.

A proton has relative charge +1 and relative mass 1, so it bends toward the negative plate. An electron has charge −1 and a much smaller mass, so it bends toward the positive plate and is deflected more strongly. A neutron has zero charge and continues undeflected in the electric field.

The direction of deflection identifies the sign of charge, while the amount of bending depends on the force relative to the particle’s mass. The electron’s very small mass gives it a much larger acceleration than a proton in the same field.

The comparison assumes beams with the same speed and a stated electric-field arrangement. Do not infer particle identity from deflection alone without checking the field direction and the experimental conditions; magnetic-field details are not added beyond the supported objective evidence.

Read isotope notation ^A_ZX by taking A as the mass (nucleon) number and Z as the atomic (proton) number. X identifies the element symbol; the two numbers give the nuclear particle counts.

Proton number = Z. Neutron number = A − Z. The mass number counts only protons and neutrons, so it is the total number of nucleons in the nucleus.

For a neutral atom, electron number = Z. For an ion, adjust the electron count for the stated charge: a positive charge means electrons have been lost, while a negative charge means electrons have been gained.

Use a fixed sequence: read A and Z, calculate A − Z for neutrons, then use the charge only to adjust electrons. Keep the element identity tied to Z; changing electron number changes charge, not the element.

Atomic radius describes the size of an atom, while ionic radius describes the size after electrons have been gained or lost. Both are controlled by the attraction between the nucleus and the outer electron region.

Across a period, proton number increases while electrons are added to the same principal shell. Nuclear attraction therefore becomes stronger overall, so atomic radius generally decreases across the period.

Down a group, each step adds an occupied electron shell. The outer electrons are farther from the nucleus and more shielded by inner shells, so atomic radius generally increases down the group.

A positive ion is usually smaller than its atom because electron loss reduces the outer electron region; a negative ion is usually larger because added electrons increase electron–electron repulsion. Use the stated species and a consistent radius definition when comparing values.

Isotopes are atoms of the same element. They have the same number of protons, so they have the same atomic number, but they contain different numbers of neutrons.

Because isotope proton numbers are the same but neutron numbers differ, their mass (nucleon) numbers are different. The nuclear identity can therefore be compared using proton number and neutron number separately.

In isotope notation, the lower-left atomic number stays constant for isotopes of one element, while the upper-left mass number changes. This lets the isotope be distinguished without changing the element symbol.

Changing neutrons makes an isotope; changing electrons makes an ion. Do not use a different mass number as evidence that the proton number, and therefore the element, has changed.

Read isotope notation in the form ^A_ZX by locating the mass (nucleon) number A at the upper left and the atomic (proton) number Z at the lower left. X is the chemical symbol, so Z identifies the element.

Use a fixed sequence: (1) copy A and Z, (2) calculate neutron number as A − Z, and (3) check the symbol against Z. The mass number counts protons plus neutrons; it is not the number of electrons.

For isotopes of one element, Z and X stay the same while A changes because the neutron number changes. Thus a different upper-left number does not by itself mean a different element.

Only change the electron count when a charge is shown: a positive ion has lost electrons and a negative ion has gained electrons. Keep this charge adjustment separate from the nuclear counts A, Z and A − Z.

Isotopes of one element have the same number of electrons and the same electron arrangement. Because chemical behaviour depends mainly on electron configuration, especially the outer-shell electrons, the isotopes have similar chemical properties.

The shared electron arrangement means isotopes form the same types of bonds, react with the same elements and show essentially the same chemical reactivity. Their different neutron numbers do not change the bonding electron arrangement.

Isotopes can differ in physical properties because different neutron numbers change the mass of the nucleus without adding charge. This supports differences in relative atomic mass and slight differences in density.

Keep the two explanations separate: electron arrangement explains similar chemical behaviour, whereas neutron number explains mass-related physical differences. Do not infer identical physical properties or different chemical reactivity from isotope mass alone.

Isotopes have different numbers of neutrons. Neutrons add mass to the nucleus but carry no charge, so isotope mass can change physical properties without changing the element or its bonding electron arrangement.

The different isotope masses give different relative atomic masses and can cause slight differences in density. These are physical differences linked to nuclear mass, not to a change in chemical identity.

All isotopes of one element have the same number and arrangement of electrons, so they show similar chemical properties even though their physical properties are not identical.

Do not treat a different mass number as a different element or as evidence of different bonding. Keep the isotope-mass explanation for physical-property variation separate from the electron-arrangement explanation for chemical reactivity.

Use this electron-structure sequence for ground-state atoms and ions within the assessed range hydrogen to krypton. The scope tells you which cases to handle; it is not a new chemical rule.

Within that range, describe electrons using principal shells, s/p/d/f sub-shells and orbitals, and use full or noble-gas shorthand configurations as required. Fill lower-energy sub-shells first, while keeping the stated 4s/3d ordering and any supported configuration exception in view.

For an ion, adjust the electron count for the charge and then write the configuration. For transition-metal ions, remove 4s electrons before 3d electrons; do not apply the neutral-atom filling order mechanically when forming the ion.

Do not import excited-state arrangements or examples beyond krypton into a scope-limited answer. Conversely, an out-of-scope example is not evidence that a broader chemical statement is false; it only lies outside the cases this card assesses. Ionisation-energy trends belong to the next Level-3 topic.

Use the hierarchy in order: a principal shell is an energy level labelled by n; each shell contains sub-shells labelled s, p, d and, where supported, f; each sub-shell contains atomic orbitals. An orbital is an allowed region, not a fixed circular path.

Each orbital holds at most two electrons. Therefore s has 1 orbital and holds 2 electrons, p has 3 orbitals and holds 6, d has 5 and holds 10, and f has 7 and holds 14. For example, 2p means the p sub-shell in shell n = 2.

Higher principal shells are generally higher in energy, while sub-shell energy follows the supported order s < p < d < f with overlap at higher shells, including 4s/3d. s orbitals are spherical; p sub-shells contain three perpendicular dumbbell-shaped orbitals, pₓ, pᵧ and p_z.

Do not confuse shell, sub-shell and orbital, or treat an orbital as a track around the nucleus. The detailed filling and electron-count method belongs to the neighbouring electron-configuration objective; this card supplies the location, capacity and model boundaries.

A sub-shell is made of atomic orbitals: an s sub-shell contains 1 orbital, a p sub-shell contains 3 orbitals, and a d sub-shell contains 5 orbitals. The three p orbitals are conventionally labelled pₓ, pᵧ and p_z.

Each orbital can hold a maximum of 2 electrons. Therefore the corresponding sub-shell capacities are s = 1 × 2 = 2 electrons, p = 3 × 2 = 6 electrons and d = 5 × 2 = 10 electrons. Where included, f has 7 orbitals and a capacity of 14 electrons.

To determine a sub-shell capacity, identify its orbital count and multiply by two. In a ground-state p sub-shell, orbitals of the same sub-shell have equal energy; electrons occupy separate orbitals before pairing, subject to the electron-configuration rules.

Do not confuse the number of orbitals with the number of electrons, and do not infer a new shell from the s/p/d label alone. Detailed filling sequences, ions and exceptions belong to the neighbouring electron-configuration objective; detailed d-orbital shape is not required at AS Level.

Build a ground-state configuration by filling available sub-shells from lower to higher energy. Shell number alone does not give the complete order because sub-shell energies overlap at higher levels; in particular, 4s is filled before 3d for a neutral atom.

Use a fixed method: count the electrons, follow the supported increasing-energy sequence, place no more than the allowed number in each sub-shell, and check that the exponents total the required electron count. A shorthand configuration uses a preceding noble-gas core where appropriate.

For potassium, the argon core is followed by 4s¹, so its shorthand configuration is [Ar] 4s¹ rather than [Ar] 3d¹. The 4s/3d overlap is the reason the simple numerical shell order is not a safe filling rule.

Do not transfer neutral-atom filling order unchanged to a transition-metal ion: when such ions form, electrons are removed from 4s before 3d. Keep filling order, ion formation and special exceptions distinct; orbital capacities are covered by the neighbouring sub-shell objective.

Read an electron configuration from left to right: the leading number identifies the principal shell, the letter identifies the sub-shell, and the superscript gives the number of electrons in that sub-shell. For example, 2p⁴ means four electrons in the p sub-shell of shell n = 2.

Determine the species’ electron count before checking the notation: a neutral atom has electrons equal to its atomic number, a positive ion has lost electrons, and a negative ion has gained electrons. The superscripts in a complete configuration must add to that count.

A full configuration lists occupied sub-shells from the beginning; shorthand replaces the inner-electron part with the nearest noble-gas symbol in brackets. The outer occupied sub-shells identify the valence-electron region relevant to the species’ outer electronic structure.

Use the notation to report electron arrangement, not as a substitute for the separate filling-order rules. Check charge and total superscripts before accepting an ion configuration, and do not confuse a sub-shell superscript with an orbital count or with the atomic number.

In a ground-state sub-shell, orbitals of the same sub-shell have equal energy. Electrons occupy these orbitals singly before any pairing, and the single electrons have parallel spins. This is Hund’s rule.

Once every equivalent orbital contains one electron, additional electrons pair in the orbitals. The two electrons in a paired orbital must have opposite spins; this arrangement limits the effect of electron–electron repulsion within the sub-shell.

Apply the rule by drawing one box per orbital: p³ is ↑, ↑, ↑, whereas p⁴ is ↑↓, ↑, ↑. The arrows show spin and the number of unpaired electrons, not merely the total electron count.

Do not pair electrons in the first orbital while equivalent orbitals remain empty, and do not give paired electrons the same spin. Keep Hund’s rule for equal-energy orbitals distinct from the separate order in which sub-shells are filled.

Use this sequence for a ground-state configuration: (1) determine the species’ electron count, (2) fill sub-shells in increasing energy, (3) respect each sub-shell’s orbital capacity and Hund’s rule, and (4) add the superscripts to check the total. Use full or noble-gas shorthand notation as requested.

For an ion, start from the neutral atom’s proton number and adjust only the electron count for the charge. A positive ion has fewer electrons and a negative ion has more; the proton number and element identity do not change.

When forming a transition-metal ion, remove electrons from 4s before 3d, even though 4s was filled before 3d in the neutral atom. For example, Fe is [Ar] 4s² 3d⁶, whereas Fe²⁺ is [Ar] 3d⁶.

Reject a configuration if its superscripts do not match the species’ electron count or if it violates capacity, filling order or spin occupancy. Do not remove 3d before 4s or treat ion formation as a proton-number change; special Cr/Cu arrangements remain separate supported exceptions.

In electron-box notation, each box represents one atomic orbital and each arrow represents one electron. Opposite arrow directions in one box show a pair with opposite spins.

Draw the required boxes for the sub-shell, place electrons in the supported energy order, and use one arrow in each equivalent orbital before pairing. The diagram therefore displays both the number of electrons and their orbital occupancy.

For a 2p⁴ sub-shell, draw three boxes as ↑↓, ↑, ↑. There are four arrows in total, one paired box and two unpaired electrons; the arrow directions make the spin information visible.

A box diagram and superscript notation represent the same configuration in different formats. Do not pair electrons while an equivalent orbital is empty, and do not interpret a box as a shell or a fixed planetary path.

An s sub-shell contains one orbital with a spherical shape. The size of the s orbital increases with principal shell number, but its shape remains spherical in this model.

A p sub-shell contains three orbitals, labelled pₓ, pᵧ and p_z. Each has a dumbbell-like shape and is oriented along one of three mutually perpendicular axes; in a ground-state atom the three p orbitals in one sub-shell have equal energy.

Use the shapes to describe allowed regions and spatial orientation, not fixed electron tracks. The label 2p identifies shell n = 2 and the p sub-shell; it does not change the orbital shapes into separate shells.

At this level, recognise and compare s and p shapes without adding detailed d-orbital drawings. Shape, orbital count and electron occupancy are related but distinct properties.

A free radical is a species containing one or more unpaired electrons. The defining evidence is the electron arrangement, not whether the species is neutral or charged.

In an electron-box diagram, an unpaired electron is shown by a single arrow in an orbital; in a configuration, inspect the outer sub-shell occupancy. A dot can also represent the unpaired electron in a radical symbol.

A chlorine radical has a 3p⁵ outer arrangement: two 3p orbitals contain paired electrons and one 3p orbital contains one unpaired electron. The single electron is the evidence for radical status.

Do not equate radical with ion: a neutral species can be a radical, and charge alone does not establish an unpaired electron. The SME also links radical formation to homolytic fission, where a covalent-bond electron pair separates evenly; do not extend this card beyond that supported boundary.

Use ground-state electron configurations for the assessed ionisation-energy scope: atoms and ions are treated in their lowest-energy available arrangement, using the supported range from hydrogen to krypton.

Ionisation energy refers to removing one mole of electrons from one mole of gaseous species, one electron-removal step at a time, under the stated standard conditions. First ionisation uses gaseous atoms; successive ionisation uses the increasingly positive gaseous ions formed after earlier removals.

When successive-ionisation data are interpreted, use the ground-state configuration to locate the outer-shell electrons. A large jump means the next electron is being removed from an inner shell with stronger nuclear attraction; the number removed before that jump indicates the outer-shell count for the supported s- and p-block cases.

Do not use an excited-state promotion or an element outside the assessed range as the default configuration, and do not remove several electrons in one ionisation step. The scope limit defines the cases to apply; it does not claim that other states or elements cannot exist. Detailed trend explanations remain in the neighbouring ionisation-energy objectives.

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous 1+ ions. It is measured under the stated standard conditions, in kJ mol⁻¹, and is positive because electron removal is endothermic.

Write the process for one atom as X(g) → X⁺(g) + e⁻, while the definition refers to one mole of gaseous atoms and one mole of electrons removed. The electron removed is the outer electron in the ground-state atom, so its attraction to the nucleus and the shielding by inner electrons affect the energy required.

Keep every state and amount in the definition: the atom and ion are gaseous, the product is a gaseous 1+ ion, and only one electron is removed in the first step. Do not write a solid ion, remove several electrons at once, or confuse first ionisation energy with a successive ionisation energy of an already positive gaseous ion.

Use this card for the definition and equation of first ionisation energy. Periodic trends, successive-ionisation jumps and detailed electronic-configuration deductions belong to the neighbouring objectives, even though they use the same gaseous-ion convention.

Successive ionisation energies are the energies required to remove one mole of electrons, one electron-removal step at a time, from the same element as its gaseous ion becomes more positive. The second step starts from X⁺(g), the third from X²⁺(g), and so on.

Write each gaseous species explicitly: X⁺(g) → X²⁺(g) + e⁻ for the second ionisation, then X²⁺(g) → X³⁺(g) + e⁻ for the third. As electrons are removed, shielding decreases and the proton-to-electron ratio increases, so the remaining electrons are generally more strongly attracted to the nucleus and successive values increase.

Interpret the sequence by locating a large jump between successive values. The jump means the next electron is being removed from a new inner shell, closer to the nucleus and more strongly attracted; the number of electrons removed before the jump gives the outer-shell count in the supported s- and p-block inference.

Do not compare successive values as if each came from a fresh neutral atom, and do not remove several electrons in one step. Keep the gaseous state and changing positive charge in every equation; use the jump as evidence about shell structure, not as a standalone claim about an unsupported element or configuration.

Across a period, first ionisation energy generally increases. Proton number rises while added electrons enter the same principal shell, so shielding changes relatively little and the outer electron experiences a stronger overall attraction to the nucleus.

The rising nuclear charge pulls the outer electron closer, while atomic radius generally decreases across the period. The stronger attraction makes the outer electron harder to remove, so more energy is required; the trend is general rather than perfectly smooth.

Explain a small dip from the electron arrangement, not from a fall in nuclear charge. A p electron can be higher in energy than an s electron, as for Al compared with Mg, or paired p electrons can repel each other, as for S compared with P; either makes removal slightly easier.

When reading a period trend, compare the relevant outer-electron sub-shell, distance, shielding and spin-pairing evidence together. Do not claim a perfectly straight increase or explain every dip only by atomic number; detailed successive-ionisation jumps and group deductions belong to the neighbouring objectives.

Successive ionisation energies generally increase because each electron is removed from a more positively charged gaseous ion. Shielding decreases and the proton-to-electron ratio increases, so the remaining electrons are held more strongly by the nucleus.

A large jump means the next electron is being removed from a new inner shell. It is closer to the nucleus and experiences stronger attraction, so much more energy is required; smaller changes can occur within a shell or sub-shell.

To infer the outer-shell count, find the first large change in the successive-ionisation data and count the electrons removed before it. For example, if the third value is the first very large value, two outer electrons were removed before the shell change, supporting a Group 2 inference for the supported case.

Do not treat every increase as a shell jump, or read a jump without checking the gaseous-ion sequence and the stated period or element information. The method infers shell structure from a change of scale; it does not replace the separate first-ionisation trend explanation or justify unsupported configurations.

Ionisation energy reflects the attraction between the nucleus and the electron being removed. A larger ionisation energy means that, under the same definition and gaseous-state conditions, more energy is needed to overcome that attraction.

Compare four linked factors: greater nuclear charge strengthens attraction; greater distance from the nucleus weakens it; inner-shell electrons shield the outer electron; and spin-pair repulsion can make a paired electron easier to remove. The observed value is the combined result, not a proton count alone.

Use the electron’s actual shell and sub-shell before predicting a value. A 3p electron is generally easier to remove than a 2p electron because it is farther from the nucleus and more shielded; within a paired orbital, electron–electron repulsion also lowers the energy needed for removal.

Do not say that more protons always give a higher ionisation energy without checking distance, shielding, sub-shell energy and pairing. Keep these attraction factors distinct from the separate definition of first ionisation energy and from the successive-ionisation jump method.

For a first-ionisation-energy comparison, identify the electron removed in each species and check four linked factors: nuclear charge, distance from the nucleus, shielding by inner shells, and spin-pair repulsion.

State how each factor changes the attraction: greater nuclear charge tends to increase ionisation energy, whereas greater distance or shielding tends to reduce it. Pairing can lower the value because two electrons repel within one orbital. Then decide which effect dominates for the actual comparison.

Use the sub-shell and pairing evidence to explain exceptions to a general period trend. A 3p electron can be easier to remove than a 3s electron because it is higher in energy, while a paired 3p electron can be easier to remove than an unpaired counterpart; these explain the Mg/Al and P/S-type dips supported by the SME.

Do not list four factors without linking them to the electron being removed, and do not claim that proton number alone decides the result. Keep this comparison method separate from the definition of first ionisation energy, the across-period overview and the successive-ionisation jump method.

Use successive-ionisation data to infer an atom’s outer-shell configuration: locate the largest jump, then count the electrons removed before that jump. Those electrons belonged to the outer shell; the next value begins removal from a new inner shell.

Combine the outer-electron count with the stated period or other supplied information. The period identifies the occupied principal shell, so the count can then be used to infer the relevant outer sub-shells and write the supported ground-state configuration.

For example, a first large jump after seven removals supports seven outer electrons; with Period 3 information, the outer arrangement is 3s²3p⁵. Treat this as evidence-led deduction from the jump and the supplied period, not as a guess from the jump alone.

A jump gives the number of outer electrons, not automatically the element. Check the period, the gaseous-ion sequence and the supported configuration scope before naming an element; do not confuse a shell jump with the smaller within-shell or sub-shell changes.

Read successive-ionisation data in sequence: compare adjacent values, locate the largest change of scale, and count the electrons removed before that jump. The count gives the number of outer-shell electrons before the next electron is taken from an inner shell.

Use the outer-electron count as a group pattern for the supported s- and p-block cases, then combine it with the stated period or configuration to identify the outer sub-shells and possible element. A large jump between the seventh and eighth values therefore indicates seven outer electrons; Period 3 information supports the chlorine configuration.

Treat the jump as evidence of a shell-depth change: the next electron is closer to the nucleus and more strongly attracted. Preserve the order of the successive values and compare their scale, rather than calling any larger next value a shell jump.

The jump alone does not name an element. Check the period, supplied information and ground-state scope before writing a configuration, and keep this data-reading method distinct from explaining the four attraction factors or the general across-period trend.