3. Chemical bonding

Electronegativity is the power of an atom to attract the bonding pair of electrons towards itself in a covalent bond. It is a comparative property of bonded atoms, not the charge on an isolated ion and not the energy needed to remove an electron.

When two bonded atoms have different electronegativities, the bonding electrons are attracted more strongly towards the atom with the higher value. This unequal attraction can make the bond polar; similar values indicate more equal sharing.

Use electronegativity to describe the direction and relative strength of attraction within a bond. Periodic trends and the factors that cause them belong to the neighbouring objectives; numerical differences used to predict ionic or covalent bonding also belong to the later application card.

Do not treat electronegativity as a fixed whole-atom charge, confuse it with first ionisation energy, or infer the polarity of an entire molecule from one bond without considering the other bonds and molecular shape. This card establishes the definition only.

Greater nuclear charge tends to strengthen the attraction between the nucleus and a bonding electron pair. More protons increase the positive pull, but this effect must be considered with the electron's distance from the nucleus and the shielding provided by inner electrons.

A larger atomic radius places the bonding electron pair farther from the nucleus, so attraction is weaker and electronegativity is lower. A smaller radius brings the bonding electrons closer and strengthens the attraction.

Inner-shell electrons shield the outer region from some nuclear attraction. Greater shielding reduces the effective pull on bonding electrons; therefore proton number alone cannot determine electronegativity. Use nuclear charge, radius and shielding together.

Do not confuse shielding with the number of bonding electrons, or state that every increase in proton number automatically gives a stronger attraction. The combined factor explanation supports the periodic trends on the neighbouring card; numerical bond classification belongs to the later application card.

Across a period, electronegativity generally increases from left to right. Nuclear charge increases while the added electrons enter the same principal shell, so shielding changes relatively little; the stronger nuclear attraction and generally smaller atomic radius pull bonding electrons more strongly.

Down a group, electronegativity generally decreases. Although nuclear charge increases, new occupied shells increase atomic radius and inner-shell shielding, so the bonding electron pair is farther from the nucleus and feels a weaker effective attraction.

Explain a trend by linking direction to the three factors: nuclear charge, distance/atomic radius and shielding. State the overall trend first, then identify which factor strengthens or weakens attraction and why the dominant effect gives the observed direction.

Use ‘generally’ rather than claiming a perfectly smooth rule for every comparison. Do not reverse the down-group direction, explain the trend with proton number alone, or use bond-polarity classification before the neighbouring application objective.

Use the supplied Pauling electronegativity values in a fixed sequence: compare the two bonded atoms, determine the size of the difference, and then classify the bond using the stated syllabus rule. Similar values support equal or nearly equal sharing; a larger difference means the bonding pair is attracted more strongly towards one atom.

A small difference is consistent with a non-polar covalent bond, an intermediate difference with a polar covalent bond, and a sufficiently large difference with ionic bonding under the assessed classification. In a polar covalent bond, show the partial-charge direction towards the more electronegative atom rather than describing full electron transfer.

The electronegativity difference predicts how unevenly the bonding electrons are shared: the greater the difference, the greater the unequal attraction and bond polarity. Keep the bond-level conclusion tied to the supplied values and the stated classification boundary.

Do not invent a universal numerical cutoff when values are not supplied, confuse bond polarity with the overall polarity of a molecule, or treat a polar covalent bond as complete ionic electron transfer. Covalent character in ionic compounds is outside the assessed scope here.

Ionic bonding is the strong electrostatic attraction between oppositely charged ions: positive cations and negative anions. The attraction acts in all directions between neighbouring ions in the ionic lattice.

A metal transfers one or more outer-shell electrons to a non-metal. The metal becomes a positively charged cation and the non-metal becomes a negatively charged anion; each ion may thereby reach a stable outer-shell arrangement supported by the source.

The oppositely charged ions arrange in a regular repeating lattice. Each ion is attracted to oppositely charged ions around it, so the solid is overall electrically neutral and the strong attraction requires substantial energy to overcome.

Keep the mechanism distinct: electron transfer creates the ions, whereas the ionic bond is the electrostatic attraction between them. Do not describe an ionic compound as a single covalent molecule or add detailed lattice-energy calculations; specific charge-balance examples belong to the neighbouring card.

Use a fixed charge-balance method: identify the metal and non-metal, infer the ions formed from their outer electrons or group positions, write each ion charge, and choose the smallest whole-number combination whose total charge is zero.

For sodium chloride, Na⁺ and Cl⁻ combine 1:1 to give NaCl. For magnesium oxide, Mg²⁺ and O²⁻ combine 1:1 to give MgO. For calcium fluoride, Ca²⁺ requires two F⁻ ions, giving CaF₂. In every case, electron transfer creates the ions and electrostatic attraction holds the lattice together.

Check the formula by adding the ionic charges after writing the subscripts: the total must be zero. The subscript counts ions in the formula unit; it does not change the charge of an individual ion. The same method supports the SME examples Li₃N and Al₂O₃.

Do not write subscripts before balancing charge, confuse charge magnitude with the number of atoms, or treat the formula as a discrete molecule when it represents an ionic lattice. Detailed lattice-energy calculations and broader property explanations are outside this card’s objective.

Metal atoms are arranged in a regular, tightly packed lattice. When the atoms pack together, their outer-shell electrons become delocalised, so they are no longer associated with one particular atom.

The metallic structure is therefore a lattice of positive metal ions surrounded by a mobile sea of delocalised electrons. The positive ions remain in the lattice while the delocalised electrons move through the structure.

Metallic bonding is the strong electrostatic attraction between the positive metal ions and the surrounding delocalised electrons. This attraction holds the ions in place, counteracts repulsion between positive ions and maintains the stability of the metal structure.

Do not describe metallic bonding as attraction between neutral atoms or as one covalent electron pair shared by two atoms. The delocalised-electron model defines the bond here; detailed explanations of conductivity, malleability, alloys or property trends belong to later supported objectives.

A covalent bond is the electrostatic attraction between the positive nuclei of two atoms and a shared pair of outer-shell electrons. The atoms share rather than transfer electrons, allowing the specified non-metal atoms to reach a more stable outer-electron arrangement.

One shared pair gives a single covalent bond; two shared pairs give a double bond; three shared pairs give a triple bond. Use the outer-shell electrons to account for the shared pairs in the specified molecules, including H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆ and C₂H₄.

In a dot-and-cross representation, show only the relevant outer-shell electrons, use dots and crosses to distinguish the two atoms, and place shared electrons as pairs between the nuclei. Keep lone pairs and shared pairs visible when they are needed to justify the structure.

For the supported Period 3 cases SO₂, PCl₅ and SF₆, the central atom may have an expanded octet. A coordinate bond is still a shared pair, but both electrons originate from one donor atom: use NH₃ + H⁺ → NH₄⁺ and Al₂Cl₆ only at this source-supported level. Orbital overlap, hybridisation, bond energy and bond length belong to neighbouring objectives.

A covalent bond forms when atomic orbitals containing unpaired electrons overlap to form a molecular orbital containing the bonding pair. Greater effective overlap gives a stronger bond. The overlap model describes the shared electron pair without changing the covalent-bond definition.

A sigma (σ) bond forms by direct end-on overlap along the line joining the two nuclei, with electron density centred between them. A pi (π) bond forms by sideways overlap of adjacent p orbitals, placing electron density above and below the plane of the σ bond.

A double bond contains one σ and one π bond; a triple bond contains one σ and two mutually perpendicular π bonds. Use the supported cases: C₂H₄ is sp² at each carbon with one C–C σ and one π bond; C₂H₂/HCN/N₂ use sp hybridisation with one σ and two π components in the triple-bond region.

Hybridisation mixes orbitals from the same shell: sp³ gives four hybrid orbitals, sp² gives three, and sp gives two. Apply the source-supported mapping to tetrahedral, trigonal-planar and linear arrangements, while keeping detailed molecular-shape treatment for the neighbouring topic.

Bond energy is the energy required to break one mole of a particular covalent bond in the gaseous state, measured in kJ mol⁻¹. Bond length is the internuclear distance between two covalently bonded atoms.

Greater attraction between bonding electrons and nuclei pulls atoms closer, giving a shorter bond length and a stronger bond. Greater bond energy therefore indicates a stronger covalent bond.

Bond multiplicity changes electron density between nuclei: triple bonds are generally shortest and strongest, double bonds intermediate, and single bonds longest and weakest. Use supplied values and units when comparing bonds.

Use bond polarity, bond strength and bond type (σ/π) as the source-supported factors when comparing covalent-molecule reactivity. A longer, weaker bond may require less energy to break, but do not treat one factor as a universal complete reactivity prediction.

VSEPR theory predicts molecular shape from the repulsion between electron domains around a central atom. Bonding pairs and lone pairs arrange as far apart as possible in three-dimensional space to minimise repulsion.

Use the central atom to count electron domains: a single, double or triple bond counts as one domain, while a lone pair is also one domain. Identify the electron-domain arrangement, then name the molecular shape after accounting for which domains are lone pairs. Lone pairs repel more strongly than bonding pairs, so they compress the observed bond angles.

Apply the model to the assessed reference cases: BF₃ is trigonal planar (120°), CO₂ linear (180°), CH₄ tetrahedral (109.5°), NH₃ pyramidal (107°), H₂O non-linear (104.5°), SF₆ octahedral (90°), and PF₅ trigonal bipyramidal (90° and 120°).

Do not count the two or three electron pairs in a multiple bond as separate domains, and do not name a shape from the formula alone. Separate the electron-domain arrangement from the molecular shape: lone pairs occupy domains but are not shown as bonded atoms in the molecular shape.

Predict an unfamiliar molecule or ion in a fixed order: identify the central atom, count its bonding pairs and lone pairs, treat each single/multiple bond as one electron domain, choose the electron-domain arrangement, and then name the molecular shape after omitting lone pairs from the visible atom layout.

Use the reference VSEPR arrangements to transfer the method to analogous species. For example, a central atom with five bonding domains and no lone pair gives a trigonal-bipyramidal arrangement, while four bonding domains around carbon give a tetrahedral arrangement. A central atom with three bonds and one lone pair has a pyramidal molecular shape.

Start from the ideal angle for the electron-domain arrangement, then explain deviations: lone pair–lone pair repulsion is strongest, lone pair–bonding pair repulsion is next, and bonding-pair repulsion is weakest. Multiple bonds count as one domain for shape, but their greater electron density can affect repulsion and angle comparisons.

Do not select a shape from the molecular formula alone or treat a lone pair as a bonded atom. The conclusion must follow from the central atom's domain count and the supported VSEPR comparison; do not infer an exact angle when the relevant arrangement or source-supported reference is not established.

Hydrogen bonding is a strong intermolecular attraction and a special case of permanent dipole–permanent dipole attraction. For the assessed N–H and O–H examples, the donor molecule has a highly polar bond to hydrogen and the neighbouring molecule supplies a lone pair on an electronegative nitrogen or oxygen atom.

The electronegative atom pulls bonding electrons away from hydrogen, giving H a substantial δ+ charge. That hydrogen is attracted to a lone pair on N or O in another molecule; the attraction is strongest when the covalent and hydrogen bonds are close to linear. Ammonia and water therefore form intermolecular hydrogen-bond networks, with the number of possible links limited by available N–H/O–H hydrogens and lone pairs.

In water, hydrogen bonds hold molecules together, so more energy is needed to separate them: melting and boiling points are relatively high. At the surface, neighbouring molecules pull surface molecules inward, producing high surface tension. In ice, a rigid, more open hydrogen-bonded arrangement leaves molecules farther apart than in liquid water, so ice has a lower density.

Do not call the O–H or N–H covalent bond itself a hydrogen bond: the hydrogen bond is the intermolecular attraction to a neighbouring lone pair. Do not infer hydrogen bonding from the presence of hydrogen alone, and do not replace the network/packing explanation with the claim that all solids are denser than their liquids.

Electronegativity is an atom's tendency to attract the shared electron pair in a covalent bond. Equal or nearly equal attraction gives a non-polar bond; unequal attraction shifts electron density towards the more electronegative atom, leaving that end δ− and the other end δ+.

A bond dipole records the direction of this charge separation, with the dipole arrow pointing towards the partially negative end. A larger electronegativity difference gives a more polar bond, but bond polarity is a property of the individual bond before the whole molecular arrangement is considered.

To decide whether a molecule has an overall dipole, identify the polarity and direction of every polar bond, then use the molecular shape to see whether the dipoles cancel. In CH₃Cl they do not cancel, so the molecule is polar; in the symmetric CCl₄ arrangement they cancel, so the molecule is non-polar despite its polar C–Cl bonds.

Do not infer overall molecular polarity from one bond or from the presence of polar bonds alone. Keep δ+ and δ− on the correct atoms, and distinguish a polar bond from a polar molecule whose net dipole depends on three-dimensional arrangement.

Van der Waals’ forces are intermolecular attractions between molecular entities. The two assessed types are instantaneous dipole–induced dipole (London dispersion) forces and permanent dipole–permanent dipole forces; hydrogen bonding is a special permanent-dipole case and is treated separately for the detailed water consequences.

In every atom or molecule, electron motion can momentarily make the charge cloud uneven. The temporary dipole then induces an opposite dipole in a neighbour, so δ+ and δ− ends attract. These forces generally become stronger as electron number increases and as molecules have more close-contact positions.

A polar molecule has a lasting charge separation, so neighbouring molecules can align opposite partial charges and form permanent dipole–permanent dipole attractions. For small molecules with the same electron count, this interaction is stronger than London dispersion alone; for example, propanone has stronger intermolecular attraction and a higher boiling point than non-polar butane.

Keep the scale clear: covalent bonds are intramolecular, whereas van der Waals’ forces act between molecules. Do not call every intermolecular force a hydrogen bond, treat an instantaneous dipole as permanent, or compare force types without considering the stated electron count and molecular contact.

Intramolecular forces hold a particle together; intermolecular forces act between neighbouring molecules. Ionic bonding is attraction between oppositely charged ions in a lattice, covalent bonding holds shared electron pairs between atoms, and metallic bonding is attraction between positive metal ions and delocalised electrons.

Intermolecular forces include instantaneous dipole–induced dipole attraction, permanent dipole–permanent dipole attraction and hydrogen bonding as a special permanent-dipole case. These forces change how molecular entities attract one another but do not replace the covalent bonds within each molecule.

In general, the forces holding particles together within a substance are stronger than the attractions between molecules. The source-supported comparison places covalent bonding at the strong end and instantaneous dipole–induced dipole attraction at the weak end; use this as a general boundary, not as a universal numerical ranking for every substance.

When explaining a property or a change of state, first identify whether the relevant attraction is within a particle or between particles. Do not describe an intermolecular force as ionic, covalent or metallic bonding, and do not imply that breaking intermolecular attractions breaks the covalent molecule itself.

A dot-and-cross diagram shows only the outer-shell electrons of the atoms or ions. Use dots and crosses to keep the source atoms distinguishable; for an ion, enclose the electron arrangement in brackets and write the charge at the top right.

Choose the bonding model before drawing: show transfer of valence electrons for an ionic compound, shared pairs for a covalent compound, and a shared pair whose two electrons come from one donor for a coordinate bond. In a displayed formula, the coordinate-bond arrow points away from the donor lone pair.

After drawing, count the shown outer electrons, check each ion's charge and confirm that the shared or transferred electrons match the bonding model. The normal noble-gas/octet pattern is a useful check, but it is not an unconditional rule.

Keep the source-supported exception branches visible: some species have an expanded octet, some have an incomplete octet/electron-deficient centre, and some contain an odd number of valence electrons. Do not add inner-shell electrons or force these species into an eight-electron arrangement.