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3. Chemical bonding

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  1. 3.1 Electronegativity and bonding

    1. 3.1.1Electronegativity

      • Define electronegativity as the power of an atom to attract electrons to itself

    2. 3.1.2Factors influencing the electronegativities

      • Explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells

    3. 3.1.3The trends in electronegativity across

      • State/explain: the trends in electronegativity across a period and down a group of the Periodic Table

    4. 3.1.4The differences in Pauling electronegativity

      • Use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)

  2. 3.2 Ionic bonding

    1. 3.2.1Ionic bonding

      • Define ionic bonding as the attraction between oppositely charged ions (positively charged cations and negatively charged anions)

    2. 3.2.2Ionic bonding, e.g. sodium chloride, magnesium

      • Describe ionic bonding, e.g. sodium chloride, magnesium oxide and calcium fluoride

  3. 3.3 Metallic bonding

    1. 3.3.1

      • Define metallic bonding as the attraction between positive metal ions and delocalised electrons

  4. 3.4 Covalent bonding and coordinate (dative covalent) bonding

    1. 3.4.1Covalent bonding examples

      • Define covalent bonding as attraction between two nuclei and a shared electron pair. - Describe covalent bonding in H2, O2, N2, Cl2, HCl, CO2, NH3, CH4, C2H6 and C2H4 - Period 3 expanded octets: SO2, PCl5 and SF6 - Coordinate bonding: NH3 + HCl -> NH4+ and Al2Cl6

    2. 3.4.2Sigma and pi bonds

      • Describe covalent bonds by orbital overlap. - Sigma bonds: direct overlap between bonding atoms - Pi bonds: sideways overlap of adjacent p orbitals above and below the sigma bond - Apply to H2, C2H6, C2H4, HCN and N2; use sp, sp2 and sp3 hybridisation

    3. 3.4.3Terms

      • Gaseous state - (a) define the terms: - bond energy as the energy required to break one mole of a particular covalent bond in the - bond length as the internuclear distance of two covalently bonded atoms - (b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

  5. 3.5 Shapes of molecules

    1. 3.5.1The shapes of, and bond angles in, molecules

      • State/explain: the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples: - BF3 (trigonal planar, 120°) - CO2 (linear, 180°) - CH4 (tetrahedral, 109.5°) - NH3 (pyramidal, 107°) - H2O (non-linear, 104.5°) - SF6 (octahedral, 90°) - PF5 (trigonal bipyramidal, 120° and 90°)

    2. 3.5.2The shapes of, and bond angles in, molecules

      • Predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

  6. 3.6 Intermolecular forces and bond properties

    1. 3.6.1Hydrogen bonding

      • Ammonia and water as simple examples - (a) describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including - (b) use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water): - its relatively high melting and boiling points - its relatively high surface tension - the density of solid ice compared with liquid water

    2. 3.6.2Electronegativity and bond polarity

      • Use the concept of electronegativity to explain bond polarity and dipole moments of molecules

    3. 3.6.3Van der Waals’ forces as the intermolecular

      • Those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces - (a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than - (b) describe the types of van der Waals’ forces: - instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces - permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding - (c) describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom

    4. 3.6.4That, in general, ionic, covalent and metallic

      • State that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

  7. 3.7 Dot-and-cross diagrams

    1. 3.7.1

      • Use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)