Atomic Structure

An atom is the smallest particle of an element that still has the element's chemical properties. It is built from three sub-atomic particles. Protons and neutrons, together called nucleons, sit in the tiny central nucleus. Electrons move around it in shells (Figure 2.1).

• An atom is electrically neutral: number of protons = number of electrons.
• The nucleus holds almost all the mass and all the positive charge. Electrons are held by the electrostatic attraction of the protons.
• The atom's radius is at least 100 000 times the radius of its nucleus. If the nucleus were a 1 cm marble, the atom would be about 1 km across, so an atom is mostly empty space.

So the number of neutrons is A − Z. Nuclide notation writes A at the top left and Z at the bottom left of the symbol (Figure 2.2). An ion forms when an atom gains or loses electrons. Its protons and neutrons do not change. Each electron lost adds +1 to the charge and each electron gained adds −1. A positive ion is a cation and a negative ion is an anion. Particles with the same number of electrons, such as Na⁺, Ne and F⁻, are isoelectronic.

Isotopes have the same number of electrons, so they have the same electronic configuration and the same chemical reactions, because chemical reactions involve electrons and not the nucleus. Their masses differ, so their physical properties, such as density, melting point, boiling point and rate of diffusion, differ. Name an isotope by its element and nucleon number, as in chlorine-35, ³⁵Cl or ²₁H (Figure 2.3).

A beam of protons or electrons bends in an electric field, because they are charged. Neutrons carry no charge and go straight on. A positive particle bends towards the negative plate. A negative particle bends towards the positive plate (Figure 2.4).

An electron has the same size of charge as a proton but about 1/1836 of the mass, so its charge ÷ mass is far larger and it bends much more. Inside the field the path is curved. After the plates the beam travels in a straight line, and the angle is measured between this line and the original direction.

Electrons occupy shells at different distances from the nucleus. The inner shells have the lowest energy and fill first. For the first 20 elements, the first shell holds up to 2 electrons and the second and third shells up to 8 each. This gives the electronic configuration: sodium is 2.8.1 and potassium is 2.8.8.1.

The Periodic Table lists the elements in order of proton number (Figure 2.5). Its rows are periods and its columns are groups, numbered 1 to 18. Groups 3 to 12 begin only in Period 4.

• The outer shell is the valence shell, and its electrons are valence electrons. A full valence shell is stable, which is why the noble gases are unreactive.
• Period number = number of occupied shells. For Groups 1 and 2, group number = valence electrons. For Groups 13 to 18, valence electrons = group number − 10. Helium has 2.
• Elements in the same group have the same number of valence electrons and so similar chemical properties.

This 2.8.8 pattern breaks down after calcium. The finer description below explains why, and is what A Level and the olympiad use.

Each shell is a principal quantum shell, numbered n = 1, 2, 3, 4. Shell n is made of subshells, and each subshell is made of orbitals.

An electron has no fixed path. Marking its position again and again builds a cloud that is dense near the nucleus and thins out further away. The orbital is the region that contains most of that cloud.

An s subshell has 1 orbital, a p subshell 3 and a d subshell 5. Shell n holds up to 2n² electrons. The 4f subshell is not needed in this course.

Energies. Energy rises with n: 1s < 2s < 3s < 4s. Within a shell the order is s < p < d. Orbitals in the same subshell have the same energy and are called degenerate. The levels crowd together as n increases, so the 4s orbital is lower in energy than 3d and fills first (Figure 2.6).

Shapes. An s orbital is a sphere. A p orbital is a dumb-bell, and the three p orbitals point along the x, y and z axes. Of the five d orbitals, d_x²−y² and d_z² point along the axes, and d_xy, d_xz and d_yz point between the axes (Figure 2.7). Four of the d orbitals have four lobes. The d_z² orbital has two lobes and a ring around its middle. Orbitals of the same type get larger as n increases.

Fill the orbitals of an atom in its ground state, the state with every electron in the lowest available energy level, using three rules.

1. Aufbau principle: an electron goes into the empty orbital of lowest energy first. After 3p, the 4s orbital fills before 3d.
2. Hund's rule: in a subshell, orbitals fill singly first, with parallel spins, before any pairs form. This keeps repulsion between electrons low.
3. Pauli exclusion principle: an orbital holds at most 2 electrons, and they must have opposite spins.

Write each subshell with its number of electrons as a superscript, in order of increasing n (so 3d comes before 4s). Oxygen is 1s² 2s² 2p⁴: the 2 in 2p is the shell, the p is the subshell and the 4 is the number of electrons. The same configuration is drawn as boxes, with a box for each orbital and a half arrow for each electron (Figure 2.8). Label the boxes with the subshell only.

A shorter form shows the inner electrons as the symbol of the previous noble gas, as in bromine, [Ar] 3d¹⁰ 4s² 4p⁵. The position in the Periodic Table also shows the last subshell filled. Sodium is in Period 3 and the s block, so it ends 3s¹. Phosphorus is the third element of the p block in Period 3, so it ends 3p³.

2.2.4Chromium and copper

Chromium and copper break the pattern. Cr is [Ar] 3d⁵ 4s¹ and Cu is [Ar] 3d¹⁰ 4s¹. A half-filled or full 3d subshell is more stable than 3d⁴ 4s² or 3d⁹ 4s², and 3d and 4s are so close in energy that moving one electron costs little (Figure 2.9).

2.2.5Ions

Write the atom's configuration first. For a cation, remove electrons from the outermost shell, taking them from 4s before 3d. For an anion, add electrons to the empty or part-filled valence orbitals. This is the trap: 4s fills first but also empties first, because once 3d holds electrons, 4s (n = 4) is the outermost shell (Figure 2.10).

The second ionisation energy removes an electron from X⁺(g) to give X²⁺(g), and so on. Every ionisation energy is positive (endothermic: energy is taken in) because energy is needed to pull the electron away from the nucleus. Gaseous atoms are far apart, so the energy measured is only that of removing the electron. The (g) after each formula is a state symbol meaning gas, and it is part of the equation.

2.3.1Trends in first ionisation energy

Three factors decide how strongly the outer electron is held: the nuclear charge, the shielding by inner shells, and the distance from the nucleus. Inner electrons repel the outer electron and cancel part of the pull of the nucleus. What is left is the effective nuclear charge, Z_eff = Z − shielding. Electrons in the same shell shield each other very little (Figure 2.11).

• Across a period, the first ionisation energy generally increases. Nuclear charge rises while the electrons enter the same shell, so shielding stays about the same, the effective nuclear charge rises and the atom shrinks. The outer electron is held more strongly, so more energy is needed to remove it.
• Down a group, it decreases. The atom has more shells, so the outer electron is further out and more shielded. This largely cancels the rise in nuclear charge, so the outer electron is held more weakly and less energy is needed.

In an answer, state the effect on the attraction and then on the energy required. Figure 2.12 shows the pattern for the first 20 elements. Two small dips interrupt the rise across Periods 2 and 3.

2.3.2Successive ionisation energies

Successive ionisation energies always increase. The number of protons stays the same while the number of electrons falls, so each remaining electron is held more strongly. The size of each step shows the shell structure (Figure 2.14).

• Gradual increase: the electrons removed come from the same subshell.
• Small jump: the next electron comes from a subshell of lower energy in the same shell, for example from s after p.
• Big jump: the next electron comes from an inner shell, which is much closer to the nucleus, less shielded and held much more strongly.

The big jump gives the group. The number of electrons removed before it equals the valence electrons.