Learning goals Ionic and Covalent Bonding Models Dative Bonding and Common Examples Sigma and Pi Bonds (Orbital Overlap) Bond Energy, Bond Length and Bond Polarity Molecular Shapes and Bond Angles (VSEPR) Intermolecular Forces and Properties Solids: Structure and Physical Properties Chemical Bonding learning outcomes Chemical bonding connects electron behaviour to molecular shape, intermolecular forces and bulk properties. The central exam chain is: identify the particles and structure → name the electrostatic attraction → explain the observed property.
Before you start Beginner path: follow the Lessons (Recommended Order) from top to bottom.
Revision: Jump to: Quick Reference · What You Must Memorise · Common Exam Traps · Hub Quiz .
Prerequisites: skim the Prerequisites list first if you’re rusty.
Useful links: A Level portal · Exam Skills
What You’ll Learn
Describe ionic, covalent and metallic bonds as electrostatic attractions, and draw the required dot-and-cross examples.
Explain dative bonding in NH₄ + Al₂Cl₆ σ π
Predict the specified molecular shapes and bond angles using VSEPR, then deduce molecular polarity from bond dipoles and shape.
Explain permanent-dipole, induced-dipole and hydrogen-bond attractions, including their effects in water and ice.
Compare covalent-bond reactivity using bond energy, bond length and bond polarity.
Describe ionic, metallic, simple molecular, giant molecular and hydrogen-bonded lattices, and deduce structure from physical-property data.
Lessons (Recommended Order)
Ionic and Covalent Bonding Models Electrostatic definitions, metallic bonding, and the full syllabus set of ionic/covalent dot-and-cross examples.
Dative Bonding and Common Examples Electron-pair donation in ammonium ions and the aluminium chloride dimer.
Sigma and Pi Bonds (Orbital Overlap) Head-on and sideways overlap of s and p orbitals, giving sigma and pi bonds.
Bond Energy, Bond Length and Bond Polarity Bond energy, length, electronegativity, bond polarity, molecular polarity, and covalent-bond reactivity.
Molecular Shapes and Bond Angles (VSEPR) VSEPR reasoning for the six specified molecular examples and analogous species.
Intermolecular Forces and Properties Permanent and induced dipoles, hydrogen bonding, and the special physical properties of water and ice.
Solids: Structure and Physical Properties The five specified lattice types, structure–property links, and deduction from physical data.
Quick Reference
What You Must Memorise
Ionic bonding : electrostatic attraction between oppositely charged ions.Covalent bonding : electrostatic attraction between a shared pair and nuclei.Metallic bonding : attraction between positive ions and delocalised electrons.Dative covalent bond : both electrons in shared pair from one atom.Hydrogen bonding : attraction between δ⁺ N-H O-H F-H σ π σ π Common Exam Traps
Confusing intermolecular forces with covalent bonds (boiling breaks IMFs, not covalent bonds).
Conductivity explanations that don’t name charge carriers (mobile ions vs delocalised electrons).
Treating “polar bond” as the same as “polar molecule” (shape matters).
Using “stronger covalent bonds” to explain boiling points of simple molecules (it’s about IMFs).
For dative bonds, drawing the arrow from acceptor → donor (wrong direction).
Definitions that miss “electrostatic attraction ” (key phrase for bonds).
London forces explanations that don’t mention instantaneous dipole–induced dipole interactions.
Treating a dative bond as a different type after formation: its electron pair has a different origin, but the resulting bond is covalent.
Naming a structure from melting point alone instead of combining melting point, conductivity and solubility evidence.
Hub Quiz and Check Your Understanding
Use practice in two passes: first without notes, then return to the lesson covering the idea you found difficult.
Chemical Bonding Hub Quiz Check bonding definitions, molecular shape, intermolecular forces and structure–property reasoning.
Chemical Bonding Knowledge Check Find the idea that needs more work, use the feedback, then try a fresh question independently.
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Calcium (Group 2) reacts with fluorine (Group 17) to form calcium fluoride. Which statement about its formation and bonding is correct?
Each Ca atom transfers one electron to each of two F atoms, forming Ca²⁺ (2,8,8) and two F⁻ (2,8); the ions are held by electrostatic attraction in a giant lattice. One F atom gains both of calcium's outer electrons, forming Ca²⁺ and F²⁻. Ca and each F share a pair of electrons, forming two Ca–F covalent bonds. Ca²⁺ and F⁻ form, and the ionic bond is the process of electron transfer, which ends once the ions have formed. A dot-and-cross diagram of carbon dioxide, O=C=O, shows outer electrons only. How many shared pairs and lone pairs does it show in total?
4 shared pairs and 4 lone pairs 2 shared pairs and 6 lone pairs 4 shared pairs and 2 lone pairs 4 shared pairs and 0 lone pairs Ammonia reacts with a hydrogen ion: NH₃ + H⁺ → NH₄⁺. Which statement about the bond formed is correct?
N donates its lone pair into the empty orbital of H⁺, forming a co-ordinate bond; once formed, it is identical to the other three N–H bonds. H⁺ donates an electron pair to N. N transfers one electron to H⁺, forming an ionic bond. A co-ordinate bond forms, and it remains longer and weaker than the other three N–H bonds. Which statement describes the N≡N bond in N₂ in terms of orbital overlap?
One σ bond from head-on overlap of p orbitals along the internuclear axis and two π bonds from sideways overlap of the remaining p orbitals, in perpendicular planes. Three σ bonds from head-on overlap of three pairs of p orbitals. One π bond from sideways overlap and two σ bonds from head-on overlap. One σ bond and two π bonds, with both π bonds lying in the same plane. Bond data: H–Cl 0.127 nm, 431 kJ mol⁻¹; H–Br 0.141 nm, 366 kJ mol⁻¹; H–I 0.161 nm, 299 kJ mol⁻¹. A hot wire is placed in each gas. Which prediction is correct?
HCl has the most polar bond, but HI decomposes into its elements most readily because H–I is the longest and weakest bond. HI has the most polar bond and decomposes most readily. HCl has the most polar bond and therefore decomposes most readily. HCl has the most polar bond and decomposes most readily because the shortest bond needs the least energy to break. Oxygen (Group 16) forms OF₂, in which O has two bonding pairs. Using VSEPR theory, which shape and F–O–F bond angle are expected?
Bent, with an angle less than 109.5° Linear, with an angle of 180° Bent, with an angle of 120° Tetrahedral, with an angle of 109.5° CCl₄ and CHCl₃ are both tetrahedral about carbon. C–Cl bonds are polar and C–H bonds are almost non-polar. Which statement is correct?
CCl₄ is non-polar because its four identical bond dipoles cancel; CHCl₃ is polar because replacing one Cl with H leaves the dipoles unbalanced. Both are polar because both contain polar C–Cl bonds. Both are non-polar because both are tetrahedral. CHCl₃ is non-polar because it has fewer polar bonds than CCl₄. The boiling points of neon, argon and krypton are −246 °C, −186 °C and −153 °C. Which explanation of the trend is correct?
Down the group the atoms have more electrons, so their electron clouds are more easily distorted and instantaneous dipole–induced dipole attractions become stronger. Heavier atoms form covalent bonds between atoms, which must be broken on boiling. Larger atoms have larger permanent dipoles. The outer electrons are held more tightly down the group, so atoms attract each other more strongly. Ammonia, NH₃ (10 electrons), boils at −33 °C, whereas phosphine, PH₃ (18 electrons), boils at −88 °C. Which explanation is correct?
N is highly electronegative, so a δ+ H on one NH₃ is attracted to the lone pair on N of another NH₃, forming hydrogen bonds; PH₃ cannot form hydrogen bonds. PH₃ has more electrons, so ammonia's higher boiling point must come from stronger N–H covalent bonds breaking on boiling. NH₃ has no lone pair, so its molecules pack more closely. Hydrogen bonds form between the N atoms of neighbouring NH₃ molecules. Which statement correctly compares solid magnesium oxide with solid magnesium metal?
MgO is a giant lattice of Mg²⁺ and O²⁻ ions held by attraction between oppositely charged ions; magnesium is a lattice of Mg²⁺ ions held by attraction to delocalised electrons. MgO is a giant ionic lattice; in magnesium each delocalised electron is attracted only to the atom it came from. MgO contains delocalised electrons as well as ions; magnesium is a lattice of positive ions in delocalised electrons. MgO consists of discrete MgO molecules held by weak attractions; magnesium is a lattice of positive ions in delocalised electrons. Which description of the crystalline solids graphite, iodine and ice is correct?
Graphite: layers of hexagonally arranged C atoms, each covalently bonded to three others, with weak attractions between layers. Iodine: I₂ molecules held by weak attractions. Ice: H₂O molecules hydrogen-bonded in an open lattice. Graphite: each C covalently bonded to four others in a three-dimensional network. Iodine: I₂ molecules held by weak attractions. Ice: open hydrogen-bonded lattice. Graphite: layers of C atoms with weak attractions between layers. Iodine: I⁻ and I⁺ ions in a giant ionic lattice. Ice: open hydrogen-bonded lattice. Graphite: layers of C atoms with weak attractions between layers. Iodine: I₂ molecules held by weak attractions. Ice: H₂O molecules held only by instantaneous dipole–induced dipole attractions. Substance Y melts at 801 °C, does not conduct when solid but conducts when molten or dissolved in water, and is brittle. Which structure and explanation fit these data?
A giant ionic lattice: strong attraction between ions gives a high melting point, and ions are fixed in the solid but mobile when molten or dissolved. A metallic lattice: delocalised electrons carry the current. A giant molecular lattice: many strong covalent bonds give the high melting point. A giant ionic lattice: electrons become delocalised when the lattice melts, so the melt conducts.
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Calcium (Group 2) reacts with fluorine (Group 17) to form calcium fluoride. Which statement about its formation and bonding is correct?
Each Ca atom transfers one electron to each of two F atoms, forming Ca²⁺ (2,8,8) and two F⁻ (2,8); the ions are held by electrostatic attraction in a giant lattice. One F atom gains both of calcium's outer electrons, forming Ca²⁺ and F²⁻. Ca and each F share a pair of electrons, forming two Ca–F covalent bonds. Ca²⁺ and F⁻ form, and the ionic bond is the process of electron transfer, which ends once the ions have formed. A dot-and-cross diagram of carbon dioxide, O=C=O, shows outer electrons only. How many shared pairs and lone pairs does it show in total?
4 shared pairs and 4 lone pairs 2 shared pairs and 6 lone pairs 4 shared pairs and 2 lone pairs 4 shared pairs and 0 lone pairs Ammonia reacts with a hydrogen ion: NH₃ + H⁺ → NH₄⁺. Which statement about the bond formed is correct?
N donates its lone pair into the empty orbital of H⁺, forming a co-ordinate bond; once formed, it is identical to the other three N–H bonds. H⁺ donates an electron pair to N. N transfers one electron to H⁺, forming an ionic bond. A co-ordinate bond forms, and it remains longer and weaker than the other three N–H bonds. Which statement describes the N≡N bond in N₂ in terms of orbital overlap?
One σ bond from head-on overlap of p orbitals along the internuclear axis and two π bonds from sideways overlap of the remaining p orbitals, in perpendicular planes. Three σ bonds from head-on overlap of three pairs of p orbitals. One π bond from sideways overlap and two σ bonds from head-on overlap. One σ bond and two π bonds, with both π bonds lying in the same plane. Bond data: H–Cl 0.127 nm, 431 kJ mol⁻¹; H–Br 0.141 nm, 366 kJ mol⁻¹; H–I 0.161 nm, 299 kJ mol⁻¹. A hot wire is placed in each gas. Which prediction is correct?
HCl has the most polar bond, but HI decomposes into its elements most readily because H–I is the longest and weakest bond. HI has the most polar bond and decomposes most readily. HCl has the most polar bond and therefore decomposes most readily. HCl has the most polar bond and decomposes most readily because the shortest bond needs the least energy to break. Oxygen (Group 16) forms OF₂, in which O has two bonding pairs. Using VSEPR theory, which shape and F–O–F bond angle are expected?
Bent, with an angle less than 109.5° Linear, with an angle of 180° Bent, with an angle of 120° Tetrahedral, with an angle of 109.5° CCl₄ and CHCl₃ are both tetrahedral about carbon. C–Cl bonds are polar and C–H bonds are almost non-polar. Which statement is correct?
CCl₄ is non-polar because its four identical bond dipoles cancel; CHCl₃ is polar because replacing one Cl with H leaves the dipoles unbalanced. Both are polar because both contain polar C–Cl bonds. Both are non-polar because both are tetrahedral. CHCl₃ is non-polar because it has fewer polar bonds than CCl₄. The boiling points of neon, argon and krypton are −246 °C, −186 °C and −153 °C. Which explanation of the trend is correct?
Down the group the atoms have more electrons, so their electron clouds are more easily distorted and instantaneous dipole–induced dipole attractions become stronger. Heavier atoms form covalent bonds between atoms, which must be broken on boiling. Larger atoms have larger permanent dipoles. The outer electrons are held more tightly down the group, so atoms attract each other more strongly. Ammonia, NH₃ (10 electrons), boils at −33 °C, whereas phosphine, PH₃ (18 electrons), boils at −88 °C. Which explanation is correct?
N is highly electronegative, so a δ+ H on one NH₃ is attracted to the lone pair on N of another NH₃, forming hydrogen bonds; PH₃ cannot form hydrogen bonds. PH₃ has more electrons, so ammonia's higher boiling point must come from stronger N–H covalent bonds breaking on boiling. NH₃ has no lone pair, so its molecules pack more closely. Hydrogen bonds form between the N atoms of neighbouring NH₃ molecules. Which statement correctly compares solid magnesium oxide with solid magnesium metal?
MgO is a giant lattice of Mg²⁺ and O²⁻ ions held by attraction between oppositely charged ions; magnesium is a lattice of Mg²⁺ ions held by attraction to delocalised electrons. MgO is a giant ionic lattice; in magnesium each delocalised electron is attracted only to the atom it came from. MgO contains delocalised electrons as well as ions; magnesium is a lattice of positive ions in delocalised electrons. MgO consists of discrete MgO molecules held by weak attractions; magnesium is a lattice of positive ions in delocalised electrons. Which description of the crystalline solids graphite, iodine and ice is correct?
Graphite: layers of hexagonally arranged C atoms, each covalently bonded to three others, with weak attractions between layers. Iodine: I₂ molecules held by weak attractions. Ice: H₂O molecules hydrogen-bonded in an open lattice. Graphite: each C covalently bonded to four others in a three-dimensional network. Iodine: I₂ molecules held by weak attractions. Ice: open hydrogen-bonded lattice. Graphite: layers of C atoms with weak attractions between layers. Iodine: I⁻ and I⁺ ions in a giant ionic lattice. Ice: open hydrogen-bonded lattice. Graphite: layers of C atoms with weak attractions between layers. Iodine: I₂ molecules held by weak attractions. Ice: H₂O molecules held only by instantaneous dipole–induced dipole attractions. Substance Y melts at 801 °C, does not conduct when solid but conducts when molten or dissolved in water, and is brittle. Which structure and explanation fit these data?
A giant ionic lattice: strong attraction between ions gives a high melting point, and ions are fixed in the solid but mobile when molten or dissolved. A metallic lattice: delocalised electrons carry the current. A giant molecular lattice: many strong covalent bonds give the high melting point. A giant ionic lattice: electrons become delocalised when the lattice melts, so the melt conducts.
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Magnesium (Group 2) reacts with oxygen to form magnesium oxide, and nitrogen exists as N₂ molecules. Which statement describes the bonding in both substances correctly?
MgO: each Mg transfers two electrons to an O, giving Mg²⁺ and O²⁻ held by electrostatic attraction between oppositely charged ions. N₂: the atoms share three electron pairs, each N keeps one lone pair, and each shared pair is attracted to both nuclei. MgO: Mg and O share two electron pairs in a Mg=O double bond. N₂: three shared pairs attracted to both nuclei, with one lone pair on each N. MgO: one electron is transferred, giving Mg⁺ and O⁻ held by attraction between the ions. N₂: three shared pairs attracted to both nuclei, with one lone pair on each N. MgO: Mg²⁺ and O²⁻ held by attraction between oppositely charged ions. N₂: one shared pair, with two lone pairs on each N. MgO: Mg²⁺ and O²⁻ held by attraction between oppositely charged ions. N₂: three shared pairs and one lone pair on each N, with the atoms held by attraction between the two nuclei. Two AlCl₃ units join to form the Al₂Cl₆ molecule. Which statement correctly describes the co-ordinate bonds in Al₂Cl₆?
A Cl atom bonded to one Al donates a lone pair into an empty orbital of the other Al; there are two such bonds, giving each Al eight outer electrons. An Al atom donates a lone pair to a Cl atom of the other AlCl₃ unit. The two AlCl₃ units are joined by a single Al–Al co-ordinate bond. There is one co-ordinate bond, so one Al has eight outer electrons and the other has six. Which statement about the covalent bonds in propene, CH₂=CH–CH₃, is correct?
There are 8 σ bonds and 1 π bond; the π bond forms by sideways overlap of p orbitals on the two doubly bonded carbon atoms. There are 7 σ bonds and 2 π bonds, because the C=C double bond consists of two π bonds. There are 9 σ bonds and no π bonds, because both bonds of C=C form by head-on overlap. There are 8 σ bonds and 1 π bond; the π bond forms by head-on overlap of p orbitals along the C=C axis. Bond data: C–Cl 0.177 nm, 340 kJ mol⁻¹; C–I 0.214 nm, 240 kJ mol⁻¹. Iodoalkanes are hydrolysed faster than chloroalkanes, a reaction in which the carbon–halogen bond breaks. Which explanation is correct?
Cl is more electronegative than I, so C–Cl is more polar; but C–I is longer and has the lower bond energy, so less energy is needed to break it, and bond strength decides the rate. I is more electronegative than Cl, so C–I is the more polar bond and attracts the attacking species more strongly. C–I is shorter than C–Cl, so its bonding electrons are held less tightly and the bond breaks more easily. C–Cl releases more energy (340 kJ mol⁻¹) when it breaks, so breaking it is less favourable than breaking C–I. Phosphorus (Group 15) forms PCl₃, in which P has three bonding pairs and each P–Cl bond is polar. Using VSEPR theory, which description of its shape, Cl–P–Cl bond angle and polarity is correct?
Trigonal pyramidal; less than 109.5°; a polar molecule Trigonal planar; 120°; a non-polar molecule Tetrahedral; 109.5°; a non-polar molecule Trigonal pyramidal; less than 109.5°; a non-polar molecule because the three P–Cl bonds are identical Trigonal pyramidal; greater than 109.5°; a polar molecule Br₂ (boiling point 59 °C) and ICl (boiling point 97 °C) both have 70 electrons per molecule. ICl is a polar molecule. Why does ICl have the higher boiling point?
With equal numbers of electrons their instantaneous dipole–induced dipole attractions are similar, but ICl molecules also attract through permanent dipole–permanent dipole attractions. The I–Cl covalent bond is stronger than the Br–Br bond, and boiling breaks these bonds. ICl molecules form hydrogen bonds with each other. Br₂ is non-polar, so there are no attractions between Br₂ molecules. Water, H₂O, boils at 100 °C but hydrogen sulfide, H₂S, which has more electrons, boils at −60 °C. Which explanation is correct?
O is highly electronegative and small, so a δ+ H on one H₂O molecule is strongly attracted to a lone pair on O of another, forming hydrogen bonds; S–H bonds cannot do this, and more energy is needed to overcome hydrogen bonds. The O–H covalent bonds are stronger than S–H bonds, and they must be broken when water boils. H₂S has more electrons, so its molecules should attract more strongly; water's high boiling point shows it is an ionic compound. Hydrogen bonds form between the δ+ hydrogen atoms of neighbouring water molecules. Which description of the crystalline solids sodium chloride, copper, iodine, diamond and ice is correct?
NaCl: alternating Na⁺ and Cl⁻ ions, each with six oppositely charged neighbours. Copper: positive ions attracted to a sea of delocalised electrons. Iodine: I₂ molecules held by weak intermolecular attractions. Diamond: each C covalently bonded to four others tetrahedrally. Ice: H₂O molecules hydrogen-bonded in an open lattice. NaCl: separate Na–Cl molecules packed together. Copper: positive ions in delocalised electrons. Iodine: I₂ molecules held by weak attractions. Diamond: tetrahedral covalent network. Ice: open hydrogen-bonded lattice. NaCl: alternating ions with six neighbours. Copper: Cu atoms joined by localised shared electron pairs. Iodine: I₂ molecules held by weak attractions. Diamond: tetrahedral covalent network. Ice: open hydrogen-bonded lattice. NaCl: alternating ions with six neighbours. Copper: positive ions in delocalised electrons. Iodine: a giant lattice of I atoms each bonded to four others. Diamond: tetrahedral covalent network. Ice: open hydrogen-bonded lattice. NaCl: alternating ions with six neighbours. Copper: positive ions in delocalised electrons. Iodine: I₂ molecules held by weak attractions. Diamond: hexagonal layers held together by weak attractions. Ice: open hydrogen-bonded lattice. Substance X melts at 1710 °C, does not conduct electricity when solid or molten, and is insoluble in water. Which structure and explanation fit these data?
A giant molecular lattice: melting breaks many strong covalent bonds, and there are no mobile ions or delocalised electrons. A giant ionic lattice: strong electrostatic attraction between ions gives the high melting point. A simple molecular lattice: strong covalent bonds inside the molecules give the high melting point. A metallic lattice: strong metallic bonding gives the high melting point.
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