Higher Chemistry Notes: Periodicity and Related Trends

The Periodic Table

  • Elements are arranged in order of increasing atomic number (number of protons). The periodic table allows chemists to predict physical properties and chemical behaviour from an element’s position.

  • Features of the table:

    • Groups: vertical columns; elements in the same group have similar chemical properties due to a common number of valence electrons in their outer shell.

    • Periods: horizontal rows; increasing atomic number; show an increasing number of outer electrons and a move from metallic to non-metallic characteristics.

  • The first 20 elements are categorised by bonding/structure:

    • Metallic: Li, Be, Na, Mg, Al, K, Ca

    • Covalent molecular: H2, N2, O2, F2, Cl2, P4, S8, and fullerenes (e.g. C60)

    • Covalent network: B, C (diamond, graphite), Si

    • Monatomic (noble gases)

The first 20 elements: bonding and structure (summary)

  • Covalent radius is a measure of atom size; trends explained by occupied electron shells and nuclear charge.

  • Ionisation energy trends depend on atomic size and shielding; increasing nuclear charge generally raises IE across a period and can lower IE down a group due to shielding and larger atomic radius.

  • Electronegativity trends: electronegativity increases across a period and decreases down a group; explained by covalent radius, nuclear charge, and inner-shell shielding.

  • Noble gases (group 0/8) have no covalent radii values because they do not form bonds with themselves.

Atomic structure: nuclear charge and shielding

  • Nuclear charge: total positive charge of the nucleus (the sum of protons). It attracts electrons in orbit around the nucleus.

  • Electron shells: electrons occupy different energy levels; the number of electrons per shell depends on position in the table and ionisation state.

  • Shielding (screening): inner electrons shield outer electrons from the full pull of the nucleus.

    • As more inner shells are present, outer electrons feel less attraction to the nucleus.

    • Shielding is cumulative with each additional shell.

  • Consequences:

    • Covalent radius tends to increase down a group due to greater shielding and more electron shells.

    • Across a period, covalent radius tends to decrease because nuclear charge increases (more protons) with the same number of electron shells, pulling outer electrons closer.

Covalent radius

  • Definition: a measure of the size of an atom; equal to ½ the distance between centres of two covalently bonded atoms.

  • Units: picometres (pm) or angstroms (Å).

  • Trends:

    • Down a group: covalent radius increases (more shells and shielding).

    • Across a period: covalent radius decreases (increasing nuclear charge with constant shielding).

  • Noble gases: no covalent radii values because they do not form covalent bonds with themselves.

  • Covalent vs ionic radii:

    • Ionic radius refers to the radius of an ion (size changes with electron loss/gain).

    • Comparing covalent vs ionic radii shows that ions often have smaller radii than neutral atoms due to electron loss reducing the number of occupied shells; e.g., Na vs Na⁺.

  • Notable exceptions:

    • Aluminium vs gallium: despite being lower in the group, Ga can have a larger actual size due to shielding effects from d-block electrons preceding Ga; the extra protons in Ga do not fully offset the shielding, leading to non-simple trend deviations.

Electron configuration and trends within groups

  • Look at electronic structure in groups to explain reactivity trends:

    • Group 1 (alkali metals): reactivity increases down the group due to easier loss of the outer electron as shielding increases.

    • Group 7 (halogens): reactivity decreases down the group due to increasing shielding and larger atomic size, making attraction for electrons weaker.

    • Group 8 (noble gases): nonreactive (inert) due to full outer shells.

Ionisation energy (IE)

  • Definition: the energy required to remove one mole of electrons from one mole of gaseous atoms. First IE refers to the removal of the first electron; second IE refers to removing a second electron, etc.

  • Ionisation equation (1st IE):
    (X(g)<br>ightarrowX+(g)+e−)(X(g) <br>ightarrow X^{+}(g) + e^{-})

  • 2nd IE (conceptual):
    (X+(g)<br>ightarrowX2+(g)+e−)(X^{+}(g) <br>ightarrow X^{2+}(g) + e^{-})

  • Trends:

    • Across a period: 1st IE generally increases from left to right.

    • Down a group: 1st IE generally decreases.

  • Why these trends occur:

    • Across a period: increasing nuclear charge (more protons) pulls outer electrons more strongly; same shielding; higher IE.

    • Down a group: adding electron shells increases shielding; outer electrons are farther from the nucleus and less strongly attracted; IE decreases.

  • 2nd and subsequent IEs: removal from a more positively charged ion; typically much larger than the previous IE because the remaining electrons are held more tightly (often from a shell closer to the nucleus).

  • Important notes:

    • These trends are explained using atomic size, nuclear charge, and inner-shell shielding.

    • Some questions ask to write ionisation equations for specific elements; refer to the data booklet for standard forms.

Electronegativity

  • Definition: a measure of the attraction an atom involved in a bond has for the electrons of that bond (shared electrons).

  • Pauling scale: fluorine is the most electronegative element (assigned a high reference value); electronegativities of other elements are given relative to fluorine.

  • Trends:

    • Across a period: electronegativity increases (more protons and smaller atomic radius strengthen attraction for bonding electrons).

    • Down a group: electronegativity decreases (increasing shielding and larger atomic radius reduce the nucleus’ pull on bonding electrons).

  • Implications:

    • Higher electronegativity elements tend to be oxidising agents; lower electronegativity elements tend to be reducing agents.

    • Example: Li (EN ≈ 1.0) tends to be reducing; F (EN ≈ 4.0) tends to be oxidising.

  • Note: noble gases do not have meaningful electronegativity values because they do not form bonds in their elemental state.

Other trends in the periodic table

  • Density:

    • Generally increases from group 1 toward the centre of a period, then falls toward the right edge; density trends depend on metallic bonding and the arrangement of atoms.

    • Across a period, increased nuclear charge and decreasing covalent radius can lead to higher density for metals because more atoms fit into a given volume with greater mass.

  • Electron affinity:

    • The energy change when an electron attaches to a neutral atom in the gaseous state to form an anion.

    • Across a period: electron affinity generally becomes more exothermic (more negative) as atoms attract added electrons more strongly, though the exact pattern is nuanced by orbital structure.

    • Down a group: electron affinity generally decreases due to increased shielding and distance from the nucleus; noble gases have no electron affinity values because they are reluctant to gain electrons.

Carbon and its allotropes (context from Unit 1 content)

  • Carbon exhibits multiple forms:

    • Covalent networks: diamond, graphite (each carbon atom forms covalent bonds in a vast lattice/network).

    • Covalent molecules: fullerenes (e.g., C60) and other multi-atom molecules like P4 and S8.

    • Graphite features a sea of delocalised electrons that enable conductivity in one plane, while diamond is an extremely hard 3D network.

  • Allotropes explain how the same element can have very different properties depending on bonding and structure.

Exam-style and study tips (summarised from the content)

  • Use the data booklet values for trends, but understand the underlying concepts: nuclear charge, shielding, and covalent radius.

  • Be able to explain why trends occur (not just state them): link to electron configuration, energy levels, and inter-electronic effects.

  • Know the key equations and when to apply them, e.g.:

    • Ionisation: (X(g)<br>ightarrowX+(g)+e−)(X(g) <br>ightarrow X^{+}(g) + e^{-}) and higher orders for subsequent ionisations.

    • Electron affinity: F(g)+e−<br>ightarrowF−(g)F(g) + e^{-} <br>ightarrow F^{-}(g) (illustrative example).

  • For covalent radius, remember the half-distance concept and the ½ rule for bonded atoms: $rc = frac{1}{2} d{ ext{bond}}$.

  • Remember exceptions and special cases (e.g., aluminium vs gallium size anomaly; noble gases lack EN values; density trends in metals).

  • Practice exam-style questions that require explaining trends with reference to shielding, nuclear charge, and electron shells.

Quick reference: key definitions and terms

  • Atomic number: number of protons (and electrons in a neutral atom).

  • Covalent radius $r_c$: half the distance between two covalently bonded atoms; measured in pm or Å.

  • Ionic radius: radius of an ion; differs from the covalent radius.

  • Nuclear charge: total positive charge in the nucleus from protons.

  • Shielding/Screening: inner electrons reduce the effective nuclear charge felt by outer electrons.

  • First ionisation energy: energy required to remove one electron from a gaseous atom; equation as above.

  • Second/third ionisation energies: energy required to remove subsequent electrons from ions.

  • Electronegativity: attraction of an atom for bonding electrons; Pauling scale; increases across a period, decreases down a group.

  • Electron affinity: energy change when an electron is added to a neutral atom in the gaseous state; trends are nuanced but generally increase across a period and decrease down a group.

Connections to other topics and relevance

  • Periodic trends underpin predictions of element properties and reactivity in inorganic chemistry.

  • The concepts of shielding and nuclear charge connect to deeper quantum/atomic theory and orbital structure.

  • Carbon allotropy shows how bonding and structure govern material properties (hardness, conductivity).

  • Understanding ionisation energy and electronegativity informs why certain elements form ions and the type of bonding (ionic, covalent polar, or covalent non-polar) in compounds.

Notes on sources and practice resources

  • The material references SQA specifications and data book values (e.g., data book p.12 for ionisation/electronegativity data) and common exam-style question prompts.

  • Practice areas include: identifying bonding/structure for first 20 elements, explaining trends (IE, EN, covalent radius), and applying concepts to ionic/covalent/radical species and their properties.