Atomic structure: OCR A A-level Chemistry

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Scientists do not keep a model just because it is familiar. When new evidence does not fit, the model needs to change. For each model below, connect what scientists observed with what that tells us about the atom.

For example, most alpha particles passing through the foil suggests that most of the atom is empty space. The few large deflections suggest that the positive charge is concentrated in a very small region. Those are two different observations supporting two different conclusions.

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An atom has a tiny central nucleus containing protons and neutrons. Electrons occupy regions outside the nucleus. The table gives the relative masses and charges you need to know.

Protons and neutrons are much heavier than electrons, so almost all the atom's mass is in its nucleus. An electron's mass is very small, but it is not zero.

In a neutral atom, each positive proton is balanced by one negative electron. That is why a neutral atom has equal numbers of protons and electrons.

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The atomic number tells you which element you have: it is the number of protons. The mass number counts protons and neutrons, so subtract the atomic number to find the neutrons. Both numbers are whole numbers with no units.

1. Find the neutral atom first. Its proton number and electron number both equal the atomic number.
2. For a positive ion, remove electrons. A 2+ charge means two electrons have been lost.
3. For a negative ion, add electrons. A 1- charge means one electron has been gained.
4. Leave the proton and neutron numbers alone. Making an ion changes the electrons, not the nucleus.

Atomic number, Z

The number of protons in the nucleus. It identifies the element.

Mass number, A

The total number of protons and neutrons in one atom of an isotope.

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Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. The proton number keeps the element the same; the extra neutrons change its mass.

Neutral atoms of the isotopes have the same electron configuration. Since chemical reactions depend on electrons, the isotopes have the same chemical properties.

Their masses are different, so physical properties that depend on mass can differ. Keep mass number, which describes one isotope, separate from the element's weighted mean relative atomic mass.

Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons and different masses.

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A relative mass compares a particle's mass with a common reference: one-twelfth of the mass of one carbon-12 atom. It is a ratio of two masses, so it has no unit.

Read the particle named in the definition. Relative isotopic mass refers to one isotope. Relative atomic mass is a weighted mean that allows for an element's isotope abundances.

Relative molecular mass refers to molecules. For a substance without separate molecules, such as an ionic compound, use relative formula mass. Molar mass is a different quantity: the mass of one mole, usually in g mol⁻¹.

Relative isotopic mass

The mass of an atom of one isotope compared with one-twelfth of the mass of a carbon-12 atom.

Relative atomic mass, Aᵣ

The weighted mean mass of an atom of an element compared with one-twelfth of the mass of a carbon-12 atom.

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The formula tells you how many atoms to include. Add their Aᵣ contributions to find the relative molecular mass or relative formula mass.

Use relative molecular mass for substances made of separate molecules. Use relative formula mass when the formula describes a ratio in a giant structure, such as NaCl. Both relative masses have no units.

Use these terms and calculate the values; formal definitions of these two terms are not required here.

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Use the information on the spectrum. You do not need TOF instrument stages or kinetic-energy calculations here.

• Peak position: read m/z from the x-axis. For a singly charged monatomic ion, this is numerically equal to its relative isotopic mass.
• Peak height: read relative abundance or intensity from the y-axis. The tallest isotope peak represents the most abundant isotope.
• Check whether abundances are percentages, ratios or raw intensities before calculating the weighted mean.

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Relative atomic mass is an average, but the isotopes do not all contribute equally. A common isotope must count more than a rare one. That is what weighted mean means.

1. Multiply each isotope's relative mass by its abundance.
2. Add those contributions.
3. Divide by the total abundance. Use 100 for percentages, but add the values yourself for a ratio or raw intensities.

Your answer must lie between the lightest and heaviest isotope masses. It should be closer to the more abundant isotope. Relative atomic mass has no unit.

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This is an analytical-chemistry crossover, included to support the molecular-ion questions in your practice pack. It connects this topic to the later organic-analysis content.

A molecular ion is the intact molecule after losing an electron. For a singly charged molecular ion, its m/z gives the molecule's relative molecular mass.

The molecular-ion peak is usually the peak furthest to the right. A small M+1 isotope peak can appear just beyond it.

A molecule can also break into smaller ions called fragments. The tallest peak is the base peak; it need not be the molecular-ion peak. Identify the peak the question tells you to use.

Molecular ion, M⁺

The intact molecule after removal of one electron. For the identified singly charged molecular ion, m/z gives its relative mass.

Base peak

The most intense peak, conventionally scaled to 100. It may represent a fragment ion.

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A shell is a principal energy level. We number the shells n = 1, 2, 3 and 4, starting nearest the nucleus. Each shell contains subshells, which contain orbitals.

The expression 2n² gives the maximum capacity of a shell: 2, 8, 18 and 32 electrons for the first four shells. It does not tell you the order in which all the subshells fill.

That distinction matters when you reach 4s and 3d. Use shell capacities for maximum numbers, and the filling-order diagram for actual configurations.

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An orbital describes a region with a high probability of finding an electron. It is not a little circular track followed by the electron.

An s orbital is spherical. A p orbital has two lobes, often called a dumb-bell shape. The nucleus is at the centre.

A p subshell has three orbitals of the same energy. They point in three mutually perpendicular directions. Each orbital can hold up to two electrons, with opposite spins.

When showing where electrons go in a diagram, orbitals are usually drawn as boxes. One box represents one orbital; one arrow represents one electron.

Atomic orbital

A region around the nucleus that can hold up to two electrons with opposite spins.

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Count the orbitals first, then double the number to find the maximum electrons. Each orbital holds up to two electrons.

• An s subshell has one orbital, so it holds up to 2 electrons.
• A p subshell has three orbitals, so it holds up to 6 electrons.
• A d subshell has five orbitals, so it holds up to 10 electrons.

An orbital and a subshell are different levels of organisation. Only the s subshell happens to contain a single orbital.

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In an electrons-in-boxes diagram, each box is one orbital and each arrow is one electron. The arrow direction represents spin.

1. For equal-energy orbitals, place one electron in each box before adding a second electron to any box.
2. Draw the singly occupied orbitals with parallel spins in these ground-state arrangements.
3. When two electrons share a box, draw opposite arrows. One orbital holds no more than two electrons.

A p subshell therefore uses three boxes; a d subshell uses five. Check both the total electron count and the way they are distributed.

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Start by counting the electrons. A neutral atom has one electron for every proton, so its total equals its atomic number.

1. Follow the filling order shown in the diagram.
2. Use the capacities s², p⁶ and d¹⁰ as you fill each subshell.
3. Stop when all the atom's electrons have been placed.
4. Add the superscripts as a final check: their total must equal Z.

In a configuration, the number names the shell, the letter names the subshell and the superscript counts its electrons. [Ar] is shorthand for the complete 18-electron argon core.

The 4s and 3d energies are close and depend on occupancy. Use the stated filling order without assuming one fixed energy ordering applies to every atom and ion.

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Use Cr, Z = 24, and Cu, Z = 29, to check the total electrons. Learn their observed configurations without turning the half-filled/full-filled explanation into a rule for every element.

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Required scope: s- and p-block ions of elements up to Z = 36. Transition-metal ion configurations are outside this requirement.

1. Write the neutral atom's electron configuration.
2. Use the charge to work out how many electrons it has lost or gained.
3. For a positive ion, remove electrons from the highest occupied principal shell, taking p before s within that shell.
4. For a negative ion, add electrons to the available outer subshell.
5. Check the final electron total.

For Ga³⁺, remove the 4p electron and both 4s electrons, leaving [Ar] 3d¹⁰. Do not treat the ion as a new neutral atom and blindly refill its electrons.

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These questions connect several electron-structure skills. Keep returning to the electron count so you do not lose track of the atom or ion.

1. Use atomic number to count electrons in a neutral atom.
2. If an ion is given, use its charge to relate its electron count to the neutral atom.
3. Use the supplied atomic numbers to identify the element.
4. Write the configuration, then translate it into orbital boxes if asked. Paired electrons need opposite spins.

Show each step. A correct total alone does not prove the electrons occupy the correct subshells.

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Periodicity connection: ionisation energies are included here to show how electron structure explains experimental evidence.

First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions. Keep all those details in an exam definition.

Energy is needed to overcome the electron's attraction to the nucleus, so ionisation is endothermic. The unit is kJ mol⁻¹. A second ionisation starts from gaseous 1+ ions; later ionisations start from increasingly positive ions.

First ionisation energy

The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

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Periodicity connection: explain each trend through the attraction between the nucleus and the electron being removed.

Across a period: nuclear charge increases while electrons enter the same main shell. Inner-shell shielding changes little, the radius decreases and attraction strengthens. First ionisation energy generally increases.

Down a group: the outer electron is further away and more shielded. These effects outweigh the increased nuclear charge, so attraction weakens and first ionisation energy generally decreases.

For exceptions, check the subshell and electron pairing. Removing a higher-energy p electron, or an electron repelled by its partner in a p orbital, can require less energy than the simple trend suggests.

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Periodicity connection: a sequence of ionisation energies can reveal the atom's shell structure.

The energies rise as electrons are removed from an increasingly positive ion. A much larger jump marks the start of removal from an inner shell, where electrons are closer to the nucleus and less shielded.

1. Find the first large jump.
2. Count the electrons removed before that jump. This is the outer-shell electron count.
3. For a main-group element, link that number to the group.
4. Use any further large jumps to deduce the populations of inner shells.