ASRJC H2 Chem 3. Chemical Bonding

2024/ASRJC/Chemistry 1 Anderson Serangoon Junior College 2024 JC1 H2 Chemistry Chemical Bonding Content • Ionic bonding, metallic bonding, covalent bonding and co–ordinate (dative covalent) bonding • Shapes of simple molecules and bond angles • Bond polarities and polarity of molecules • Intermolecular forces, including hydrogen bonding • Bond energies and bond lengths • Lattice structure of solids • Bonding and physical properties Learning Outcomes Candidates should be able to: (a) show understanding that all chemical bonds are electrostatic in nature and describe: (i) ionic bond as the electrostatic attraction between oppositely charged ions (ii) covalent bond as the electrostatic attraction between a shared pair of electrons and positively charged nuclei (iii) metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons (b) describe, including the use of ‘dot–and–cross’ diagrams, (i) ionic bonding as in sodium chloride and magnesium oxide (ii) covalent bonding as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene (iii) co–ordinate (dative covalent) bonding, as in formation of the ammonium ion and in the Al2Cl6 molecule (c) describe covalent bonding in terms of orbital overlap (limited to s and p orbitals only), giving σ and π bonds (d) explain the shapes of, and bond angles in, molecules such as BF3 (trigonal planar); CO2 (linear); CH4 (tetrahedral); NH3 (trigonal pyramidal); H2O (bent); SF6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory (e) predict the shapes of, and bond angles in, molecules analogous to those specified in (d) (f) explain and deduce bond polarity using the concept of electronegativity [quantitative treatment of electronegativity is not required] (g) deduce the polarity of a molecule using bond polarity and its molecular shape (analogous to those specified in (d)); (h) describe the following forces of attraction (electrostatic in nature): (i) intermolecular forces, based on permanent and induced dipoles, as in CHCl3(l); Br2(l) and the liquid noble gases (ii) hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups (i) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (j) explain the terms bond energy and bond length for covalent bonds (k) compare the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity (l) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride and magnesium oxide (ii) simple molecular, as in iodine (iii) giant molecular, as in graphite and diamond (iv) hydrogen–bonded, as in ice (v) metallic, as in copper [the concept of the ‘unit cell’ is not required] (m) describe, interpret and/or predict the effect of different types of structure and bonding on the physical properties of substances (n) suggest the type of structure and bonding present in a substance from given information References 1 Peter Cann and Peter Hughes. Chemistry for Advanced level. 2 Brown, LeMay, Bursten. Chemistry, The Central Science. 3 E. N. Ramsden. A–Level Chemistry 4 Jeanne Tan and Kim Seng Chan, Understanding Advanced Physical Inorganic Chemistry

2024/ASRJC/Chemistry 2 1. INTRODUCTION All chemical bonds are electrostatic in nature which involve attractions between a positive charge and a negative charge. Strong forces of attraction Type of chemical bond Electrostatic attraction between… (i) metallic bond metal ions and ‘sea’ of delocalised electrons (ii) ionic bond cations and anions (iii) covalent bond shared pair of electrons and positively charged nuclei The type of chemical bond in a substance is primarily determined by the outer –shell electronic structure of its atoms where atoms lose, gain or share valence electrons to achieve the stable noble gas configuration. Weak forces of attraction Some atoms are bonded together as a group by covalent bonds to form discrete molecules. Weak forces of attraction exist between these molecules. These are called intermolecular forces of attraction which include: (i) Instantaneous dipole–induced dipole attraction (ii) Permanent dipole–permanent dipole attraction (iii) Hydrogen bonding All these strong and weak electrostatic forces of attraction between particles bind them together to form matter. Lattice structures There are two major types of solid: crystalline and non–crystalline solids. In non–crystalline (or amorphous) solids, the particles are not arranged in any order (i.e. fixed in random positions). In crystalline solids, the particles are arranged in a regular three –dimensional pattern. This regular arrangement of points in space, occupied by the particles such as atoms, ions or molecules, is called a lattice. The different types of structures include: (i) giant metallic (ii) giant ionic (iii) simple molecular (iv) giant molecular The different types of structure and bonding present in substances largely determine the properties of the substances.

2024/ASRJC/Chemistry 3 2. GIANT METALLIC LATTICE 2.1 Metallic Bond Metallic bonds are strong electrostatic attractions that exist between the metal ions and the ‘sea’ of delocalised electrons in a giant metallic structure. ▪ Metallic bonds are formed when each metal atom (e.g. Mg) contributes its valence electrons to form a ‘sea’ of electrons that are delocalised throughout the lattice of metal ions (Mg2+). 2.2 Physical properties of metallic elements ▪ Generally high melting points Strong electrostatic forces of attraction exist between the metal ions and ‘sea’ of delocalised electrons. Large amount of energy is required to break the strong metallic bonds. ▪ High electrical conductivity Presence of delocalised valence electrons that are small and move fast gives rise to its high electrical conductivity. ▪ Malleable and ductile These properties are related to the ability of the cations to move over one another without breaking of metallic bonds. When stress is applied to the metal, the metal structure can change its shape without fracturing. The layers of metal ions are still held together by the negative sea of electrons without breaking the metallic lattice. The metal ions between the layers are not pushed apart by the repulsion as they are shielded by the delocalised electron cloud between them. Forces applied delocalised electrons cations Diagram showing how the structure in metal changes when pressure is applied source: John Wiley & Sons metallic bond in sodium metallic bond in magnesium

2024/ASRJC/Chemistry 4 2.3 Factors affecting the strength of metallic bond a) No. of valence electrons available for (metallic) bonding ❖ Greater the number of such electrons, the stronger the metallic bonds. b) Ionic radius of the metal cation ❖ Smaller the ionic radius of the cation, the stronger the metallic bonds. Note: When you are asked to compare 2 metals with contradicting effects from the two factors above (e.g. one metal has more valence electrons but also larger ionic radius than the other metal), see melting point data given by the question to decide which factor is more predominant. Checkpoint 1 metal Na Mg Al melting point / °C 98 650 660 no. of delocalised valence electrons 1 (Group 1) 2 (Group 2) 3 (Group 13) cationic charge, q+ +1 +2 +3 cationic radius, r+ / nm (data can be found in Data Booklet) 0.095 0.065 0.050 metallic bond strength Na < Mg < Al (a) Why is the melting point of K (64 C) lower than that of Na (98 C)? ▪ K+ has a larger ionic radius than Na+ due to an additional quantum shell. ▪ Weaker electrostatic forces of attraction exist betwe