The Atom
An atom is the smallest particle of an element that can still take part in a chemical reaction. Every atom is built from three subatomic particles: the proton and neutron, which sit together in the tiny, dense nucleus at the centre, and the electron, found in the mostly-empty space around it.
An atom is about 10⁻¹⁰ m across, but its nucleus is only 10⁻¹⁴ to 10⁻¹⁵ m across: more than 10,000 times smaller. Almost all of an atom's mass is packed into that tiny nucleus.
Behaviour in an Electric Field
Behaviour in an electric field reveals the three particles' charges directly. Fire beams of protons, neutrons and electrons, all at the same speed, between two charged plates:
- The neutron beam is undeflected: it carries no charge.
- The proton beam curves toward the negative plate.
- The electron beam curves toward the positive plate, through a much larger angle.
The direction of deflection depends only on the sign of the charge. The size of the deflection depends on the charge-to-mass ratio:
A proton and an electron carry the same magnitude of charge, but the electron is about 1840× lighter, so it swings through a far bigger angle for the same field.
Given a deflection angle for one particle, you can work out the expected angle for any other charged particle by comparing their charge-to-mass ratios directly.
Example
If ¹H⁺ deflects at an angle of +40°, what angle will ²H⁺ deflect to?
²H⁺ has twice the mass of ¹H⁺, but the same charge, so its charge-to-mass ratio is half as big, and it deflects at roughly half the angle: +20°.
Nuclides
For a neutral atom, number of electrons = number of protons = Z.
For an ion, adjust for the charge: an anion has Z + |charge| electrons, a cation has Z − |charge| electrons.
Isotopes
Isotopes
Atoms of the same element (same Z) with different numbers of neutrons, so a different A.
Because isotopes have identical electron arrangements, they have the same chemical properties. But because their masses differ, their physical properties differ too.
Example
Cl3517 and Cl3717 both form NaCl the same way: same chemical properties.
Cl3717₂ is very slightly denser than Cl3517₂: different physical properties.
Electrons
Electrons in an atom don't occupy fixed positions around the nucleus or move in fixed orbits (circular or elliptical paths) around it.
- Electrons are found in regions of space around the nucleus, called atomic orbitals.
- Each electron has energy that is quantised (energy restricted to certain values) and occupies a discrete energy level.
Electronic structure
The arrangement of electrons in an atom or ion.
Electronic structure tells you three things about an atom or ion:
- The number of electrons it has.
- How those electrons are distributed.
- The energies of those electrons.
It matters because it determines how the atom reacts with other atoms to form ions or molecules.
Electrons around the nucleus are organised in a hierarchy:
Outside the nucleus → shells → subshells → orbitals.
Electronic Shells
Electrons in atoms are arranged in a series of shells (or energy levels) surrounding the nucleus. Each electron shell is described by a number known as the principal quantum number, n.
The principal quantum number n is a positive integer: n = 1, 2, 3, …, numbered outward from the nucleus.
The greater the value of n:
- The further the electron shell (or electron) is from the nucleus.
- The higher the energy level of the electron shell (or electron).
- The weaker the electrostatic attraction between the nucleus and the electron.
- The larger the size of the orbital (or the more diffuse the orbital becomes).
Subshells
Each shell contains one or more subshells: s, p, d, f.
| Principal quantum number, n | Types of subshell |
|---|---|
| 1 | 1s |
| 2 | 2s, 2p |
| 3 | 3s, 3p, 3d |
| 4 | 4s, 4p, 4d, 4f |
The electron shell with principal quantum number n contains n subshells.
Within the same shell, subshell energy increases in the order s, p, d, f. So within shell n = 4, energy rises 4s → 4p → 4d → 4f.
Each subshell is itself made up of one or more orbitals. An orbital holds a maximum of 2 electrons.
Add up every subshell's orbitals and a shell's total capacity works out to 2n² electrons: 2 for n = 1, 8 for n = 2 (2 + 6, as above), 18 for n = 3.
Orbitals
An orbital represents a region of space with a high probability (>95%) of containing an electron.
- Each orbital can hold a maximum of 2 electrons.
- Each orbital has its own distinctive geometrical shape.
- The energy of an orbital is the energy of the electron occupying it.
Every subshell is made up of one or more orbitals: s has 1, p has 3, d has 5, f has 7.
Writing Electronic Configurations
An atom's electronic configuration describes how its electrons are distributed among orbitals. There are three equivalent ways to write it: shell notation, full spdf notation, or noble-gas-core notation.
Example
Silicon (Z = 14)
Shell notation: 2.8.4 (rarely used)
Full spdf notation: 1s² 2s² 2p⁶ 3s² 3p²
Noble-gas-core notation: [Ne] 3s² 3p²
Aufbau ("building-up") principle: electrons fill the lowest-energy orbital available first. Filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → … Note that 4s fills before 3d, even though 3d is written after 4s in the final configuration.
Hund's rule: within a subshell (degenerate orbitals), electrons occupy orbitals singly, with parallel spins, before any pairing occurs.
Pauli exclusion principle: an orbital holds at most 2 electrons, and if it holds 2, they must have opposite spins.
Example
Vanadium (Z = 23): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s²
Anomalous Electronic Configurations
Anomalies: Cr and Cu. Both break the expected pattern because a half-filled (d⁵) or fully-filled (d¹⁰) d subshell is more stable: the electron cloud is spread out as symmetrically as possible, minimising repulsion.
Example
Expected Cr: [Ar] 3d⁴ 4s²
Actual Cr: [Ar] 3d⁵ 4s¹
Expected Cu: [Ar] 3d⁹ 4s²
Actual Cu: [Ar] 3d¹⁰ 4s¹
Ground State vs Excited State
The ground state is an atom's lowest-energy, most stable configuration: the one most atoms are in at room temperature, and the one every electronic configuration covered so far describes.
If one or more electrons absorb energy and are promoted to a higher energy level, the atom enters an excited state instead. Excited atoms are unstable, and emit energy to fall back down to the ground state.
Ions
To find an ion's electronic configuration, start from the parent atom's configuration, then adjust for the charge it carries.
Anions
Formed when an atom gains electrons. The extra electrons go into the next available orbital.
Cations
Formed when an atom loses electrons. Electrons are removed from the highest-energy occupied orbital first, typically the outermost shell. For d-block cations specifically, 4s electrons are removed before 3d electrons, since once electrons are present, 4s sits at a higher energy than 3d.
Example
S: 1s²2s²2p⁶3s²3p⁴ → S²⁻: 1s²2s²2p⁶3s²3p⁶
Al: [Ne] 3s²3p¹ → Al³⁺: 1s²2s²2p⁶
Fe: [Ar] 3d⁶4s² → Fe²⁺: [Ar] 3d⁶, not [Ar] 3d⁴4s²
Why does 4s fill before 3d, but empty first?
While an atom is being "built up" from scratch, 4s is briefly the lower-energy option: it's more diffuse and spreads out further, reducing electron-electron repulsion, so it fills first.
But once electrons are actually present in both 4s and 3d, the compact, closer-to-the-nucleus 3d electrons end up more strongly held than the diffuse 4s ones. So 4s becomes the higher-energy orbital, and the first one removed, when forming a cation.
Isoelectronic Species
Isoelectronic species have the same total number of electrons, though not necessarily the same configuration.
Example
N³⁻, O²⁻, F⁻, Ne and Na⁺ are all isoelectronic with 10 electrons each.
All share the configuration 1s²2s²2p⁶.
Electronic Configuration and the Periodic Table
The last subshell an element's electrons fill determines which block it sits in: s, p, d or f.
Horizontal rows in the Periodic Table are called periods. Elements in the same period have the same number of electron shells. The period number of an element is the number of electron shells occupied with electrons in an atom of that element.
Vertical columns in the Periodic Table are called groups. Elements in the same group have the same number of valence electrons, and hence similar outer electron configurations.
Example
Bromine is Group 17, Period 4.
Group 17 → valence shell: 4s²4p⁵
Period 4 → 4 electron shells occupied
Group number (s/p block) fixes valence configuration; period number fixes the number of occupied electron shells; block = subshell of the last electron added.
Electrostatic Effects
Electrostatic effects
The attraction of opposite charges and the repulsion of like charges. These effects play a major role in determining the energy states of atoms, and so influence atomic properties such as atomic size and ionisation energy.
Three main factors affect the strength of the electrostatic attraction between the nucleus and the electrons:
- Number of electron shells: more shells means a higher principal quantum number (n) for the outermost shell, and more inner-shell electrons. Electrostatic attraction between the nucleus and the outermost electron decreases.
- Nuclear charge, Z: more protons means a greater nuclear charge. Electrostatic attraction between the nucleus and the outermost electron increases.
- Shielding (screening) effect: shielding by electrons in the same shell (including the outermost shell) is small; shielding from inner-shell electrons is far more significant.
Across a period, it's nuclear charge and shielding (factors 2 and 3) that are considered. Down a group, all three factors apply, but shielding is the most significant, because the number of electron shells increases: more inner-shell electrons means more shielding for the outermost electron, so electrostatic attraction to the nucleus decreases.
These three factors explain the trends covered next: atomic radius and ionic radius, first ionisation energy across a period and down a group, successive ionisation energies, and electronegativity.
Nuclear Charge and Shielding
Effective nuclear charge
The net attraction actually felt by an outer electron, roughly nuclear charge minus shielding. When both nuclear charge and shielding increase at once, you have to judge which effect wins.
Shielding ability of electrons in the same shell decreases in the order s > p > d > f. d and f electrons are especially poor at shielding.
- More shells → weaker attraction. More protons → stronger attraction. More (inner) shielding → weaker attraction.
- Same-shell electrons barely shield each other; inner-shell electrons shield strongly.
- "Effective nuclear charge" = the net pull an outer electron actually feels, after accounting for shielding.
Atomic Radius
Atomic radius
Half the shortest inter-nuclear distance found in the structure of the element.
This isn't a single fixed number: it depends on how the atom is bonded to its neighbours, so chemists define it three different ways.
- Metallic radius: half the distance between neighbouring atoms in a metallic lattice. Only applies to metals.
- Covalent radius: half the distance between two covalently-bonded atoms.
- Van der Waals radius: half the distance between two non-bonded, touching atoms. The only option for the noble gases, which don't form bonds.
Variation in Atomic Radius
Across a period, atomic radius decreases.
- The number of electron shells stays the same.
- The number of protons increases, so nuclear charge increases.
- The added electrons go into the same outermost shell, so shielding stays roughly constant, since same-shell electrons barely shield each other.
- Electrostatic attraction between the nucleus and the valence electrons increases, pulling the electron cloud in tighter.
Down a group, atomic radius increases.
- The number of electron shells increases.
- The number of protons also increases, so nuclear charge increases.
- But the extra electron shells mean significantly more shielding from inner-shell electrons.
- This increase in shielding outweighs the increase in nuclear charge, so electrostatic attraction to the outermost electron decreases, and the size of the electron cloud increases.
Ionic Radius
Ionic radius
The radius of an ion within an ionic compound.
Cations are always smaller than their parent atom.
- Both the cation and its parent atom have the same number of protons, and hence have the same nuclear charge.
- However, the cation has one less electron shell than its parent atom, so the outermost electrons are shielded to a smaller extent, as there are fewer inner-shell electrons.
- Electrostatic attraction between the nucleus and the outermost electrons is stronger in the cation, resulting in the cation having a smaller electron cloud and a smaller radius than its parent atom.
Anions are always larger than their parent atom.
- Both the anion and its parent atom have the same number of protons, and hence have the same nuclear charge.
- But the anion has more electrons packed into the same outer shell as its parent atom, increasing electron-electron repulsion between them.
- That extra repulsion (a same-shell shielding effect) is enough to expand the electron cloud, even without adding a new shell, resulting in the anion having a larger radius than its parent atom.
Ionic Radius of Isoelectronic Species
Ionic radii decrease from Na⁺ to Al³⁺.
- Na⁺, Mg²⁺ and Al³⁺ are isoelectronic species, so their outermost electrons experience the same shielding effect.
- Nuclear charge increases from Na⁺ to Al³⁺.
- Electrostatic attraction between the nucleus and the outermost electrons increases from Na⁺ to Al³⁺, decreasing the size of the electron cloud.
Example
Na⁺, Mg²⁺ and Al³⁺ all have 10 electrons.
Na⁺ (0.095 nm) > Mg²⁺ (0.065 nm) > Al³⁺ (0.050 nm)
Example
P³⁻, S²⁻ and Cl⁻ all have 18 electrons.
P³⁻ (0.212 nm) > S²⁻ (0.184 nm) > Cl⁻ (0.181 nm)
The ionic radius of P³⁻ is bigger than that of Al³⁺.
Al³⁺ (0.050 nm) → P³⁻ (0.212 nm)
- From Al³⁺ to P³⁻, the number of electron shells increases.
- Shielding experienced by the outermost electrons increases significantly, as there are now more inner-shell electrons.
- Despite the increasing nuclear charge, electrostatic attraction between the nucleus and the outermost electrons decreases, increasing the size of the electron cloud.
Ionisation Energy
First ionisation energy (1st IE)
The energy required to remove 1 mole of electrons from 1 mole of gaseous M atoms to form 1 mole of gaseous M⁺ ions.
M(g) → M⁺(g) + e⁻
Second ionisation energy (2nd IE)
The energy required to remove 1 mole of electrons from 1 mole of gaseous M⁺ ions to form 1 mole of gaseous M²⁺ ions.
M⁺(g) → M²⁺(g) + e⁻
Ionisation is an endothermic process, and ionisation energies are positive values, since energy is absorbed during ionisation to overcome the attraction between the electron and the nucleus.
Ionisation energies affect the type of bond an atom forms with other atoms. Elements with low ionisation energies find it easy to lose an electron to form a cation, resulting in ionic bonds being formed.
Trends in Ionisation Energy
Across a period, 1st ionisation energy generally increases.
- The number of electron shells remains the same.
- The number of protons increases, so nuclear charge increases.
- The number of electrons also increases, but these electrons are added to the same outermost electron shell, so shielding remains approximately constant.
- Electrostatic attraction between the nucleus and the valence electrons increases, resulting in an increase in the energy required to remove the valence electron from an atom.
First ionisation energy of Al is lower than that of Mg.
Mg: 1s² 2s² 2p⁶ 3s²
Al: 1s² 2s² 2p⁶ 3s² 3p¹
- The 3p electron removed from Al is at a higher energy level than the 3s electron removed from Mg.
- Less energy is required to remove the 3p electron in Al, as it is less attracted to the nucleus than the 3s electron in Mg.
First ionisation energy of S is lower than that of P.
P: [Ne] 3s² 3p³
S: [Ne] 3s² 3p⁴
- The 3p electron removed from S is a paired electron, while that removed from P is an unpaired electron.
- Due to greater electron-electron repulsion (interelectronic repulsion) between paired electrons in the same orbital, less energy is required to remove the paired 3p electron from S.
Down a group, 1st ionisation energy decreases.
- The number of electron shells increases.
- Shielding experienced by valence electrons increases significantly, as there are more inner-shell electrons.
- Despite the increasing nuclear charge, electrostatic attraction between the nucleus and the valence electrons decreases, resulting in a decrease in the energy required to remove a valence electron.
First ionisation energy of Na is much lower than that of Ne, and of K much lower than that of Ar.
- Na and K are each the first element of a new period, so each has one more electron shell than the noble gas immediately before it (Ne and Ar).
- That extra shell means significantly more shielding for the valence electron.
- Despite the increasing nuclear charge, electrostatic attraction between the nucleus and the valence electron decreases, so much less energy is needed to remove it.
On a graph of nth IE against atomic number, the element at the lowest point belongs to Group n (for n = 1 or 2) or Group n + 10 (for n > 2).
- For any nth ionisation energy, the element with n valence electrons still has an outer-shell electron left to lose.
- The preceding element, with n − 1 valence electrons, has already lost all its valence electrons, so its nth electron comes from an inner shell instead, needing much more energy.
- The following element has a higher nth IE too, from the usual increase in nuclear charge across a period.
Trend in Successive Ionisation Energies of an Element
Second IE is always greater than first IE, because more energy is required to remove each subsequent electron.
Successive ionisation energies of one element always increases.
K(g) → K⁺(g) → K²⁺(g) → K³⁺(g) → K⁴⁺(g) → K⁵⁺(g) → K⁶⁺(g)
- Number of protons remains the same, so nuclear charge remains the same.
- Number of electrons decreases, so shielding experienced by the remaining outermost electrons decreases.
- Electrostatic attraction between the nucleus and the remaining electrons increases, resulting in an increase in the energy required to remove each subsequent electron.
A genuinely large jump between two successive IEs signals that the next electron is being pulled from a shell closer to the nucleus, much less shielded and much more strongly attracted. A small jump just signals a change of subshell within the same shell.
Example
Al (Z = 13): 1s²2s²2p⁶ 3s²3p¹
A big jump after the 3rd electron removed, and again after the 11th.
13 electrons split into groups of 3, 8, 2, matching n = 3, n = 2, n = 1.
3 electrons before the first jump → 3 valence electrons → Group 13.
The number of electrons before the first big jump gives the element's group: that number directly for Group 1 or 2, or that number plus 10 for Groups 13 to 18.
- A large jump between two successive ionisation energies means the next electron removed comes from a completely different, inner electron shell, much less shielded and much more strongly attracted to the nucleus.
- The number of electrons removed before that jump equals the number of valence electrons the element has.
The same idea extends to the whole dataset: counting every jump in a full set of successive IEs reveals the element's complete electron configuration, since each jump marks a shell boundary and the number of electrons removed within each segment gives that shell's electron count.
Electronegativity
Electronegativity
The electronegativity of an atom in a molecule is a relative measure of its ability to attract bonding electrons.
Across a period, electronegativity increases.
- The number of electron shells remains the same.
- The number of protons increases, so nuclear charge increases.
- The number of electrons also increases, but these electrons are added to the same outermost electron shell, so shielding remains approximately constant.
- Electrostatic attraction between the nucleus and the bonding electrons increases.
Down a group, electronegativity decreases.
- The number of electron shells increases.
- Shielding experienced by bonding electrons increases significantly, as there are more inner-shell electrons.
- Despite the increasing nuclear charge, electrostatic attraction between the nucleus and the bonding electrons decreases.
When two atoms of similar electronegativity are covalently bonded, a non-polar bond is formed. If the difference between the electronegativities of the two bonding atoms (ΔEN) increases, the covalent bond becomes more polar, and its ionic character also increases. With a sufficiently large difference, the two atoms form an ionic bond instead.
| ΔEN | Bond type |
|---|---|
| < 0.5 | Nonpolar covalent |
| 0.5 – 1.6 | Polar covalent |
| > 2.0 | Ionic |
Metals have low electronegativities and tend to lose electrons, while non-metals have high electronegativity and tend to gain electrons. Electronegativity is closely tied to whether an element behaves as a metal, metalloid or non-metal.