NJC 5 Chemical Bonding (H1)

National Junior College SH1 H1 Chemistry 1 CHEMICAL BONDING 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 an d 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 A l2Cl6 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); NH 3 (trigonal pyramidal); H 2O ( bent); SF 6 (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 (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 induce d dipoles, as in CHC l3(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 intermolecular forces to the liquefaction of gases when subjected to high pressure and/or low temperature (j) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (k) explain the terms bond energy and bond length for covalent bonds (l) compare the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity (m) 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 (n) describe, interpret and/or predict the effect of different types of structure and bonding on the physical properties of substances (o) suggest the type of structure and bonding present in a substance from given information References • Cambridge International AS and A Level Chemistry by Peter Cann and Peter Hughes, published by Hodder Education • A Level Chemistry (4th Edition) by E. N. Ramsden, published by Oxford University Press Copyright © 2022 National Junior College All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording or any other information storage and retrieval system, without prior permission in writing from the copyright owner.

National Junior College SH1 H1 Chemistry 2 Note: Strong covalent bond between O and H atoms. Weaker intermolecular forces between H2O molecules. A. INTRODUCTION All matter on earth are made of atomic particles bonded together to form molecules or compounds. These binding forces (chemical bonds) result in the substance achieving a lower energy arrangement. The re -distribution of the outer electrons (valence electrons) of the atoms is required for the formation of a bond. Chemical Bonding is at the heart of Chemistry. Generally, there are two categories of chemical bonds --- inter-atomic bonds and intermolecular forces. Table 1: Type of chemical bonds and their approximate strengths Type of bond Strength of bond Bond Energy/ kJ mol–1 Ionic bond Strong 120 – 450 Covalent bond Strong 120 – 550 Metallic bond Strong 80 – 580 Type of Intermolecular Forces Strength of IMF Bond Energy/ kJ mol–1 Instantaneous dipole - induced dipole (id-id) interactions Very weak <4 Permanent dipole - permanent dipole (pd-pd) interactions Weak <13 Hydrogen bonds Weak 20 – 100 Example: Water Chemical Bonds • Instantaneous dipole - induced dipole interactions • Permanent dipole - permanent dipole interactions • Hydrogen bondings • Metallic • Ionic • Covalent Strong Inter-atomic Bonds between atoms/ions Weak Intermolecular Forces between molecules Intermolecular forces of attraction between the water molecules Covalent bond between the O and H atoms in a H2O molecule

National Junior College SH1 H1 Chemistry 3 Learning Objective: • understand that all chemical bonds and forces of attraction are electrostatic in nature. Name of Bond Diagrammatic Representation Electrostatic Forces of Attraction between Inter-atomic Bonding Metallic Bond (+) metal cations And (–) sea of delocalised electrons Ionic Bond (+) Cations And (–) Anions Covalent Bond (+) Nucleus (proton) And (–) Shared Electrons Pair Intermolecular Bonding Instantaneous dipole - induced dipole interactions (+) Instantaneous dipole (+) And (–) Induced dipole (–) Permanent dipole – Permanent dipole interactions (+) Permanent dipole (+) And (–) Permanent dipole (–) Hydrogen bonds (+) Permanent dipole (+) on H bonded to O, N or F And (–) Lone pair electrons on O, N or F Ion - Molecule Ion – dipole interaction (+) Cations Permanent dipole (+) And And (–) Permanent Dipole (–) Anions

National Junior College SH1 H1 Chemistry 4 B. INTER–ATOMIC BONDS Learning Objective: • describe metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons • describe the metallic lattice structure of a crystalline solid as in copper • describe, interpret and/or predict the effect of different type s of structure and bonding on the physical properties of substances 1 METALLIC BONDS The metallic bond is the electrostatic forces of attraction between metal cations and the sea of delocalised electrons. 1.1 Structure and Bonding of Metals 1. Metals consist of crystalline structures of closely packed atoms with 8-12 atoms surrounding any given atom and at equal distance from it. 2. The metallic lattice structure comprises metal cations immersed within a sea of delocalised electrons. This is due to each metal atom contributing its loosely bound valence electrons to form a sea of delocalised electrons. The metallic cations are attracted to the sea of delocalised electrons. The sea of delocalised electrons prevents the repulsion between the cations. 3. The electrons are no longer bound to a particular metal cation but are free to move throughout the metal (delocalised electrons). Delocalised electrons Metal cations The “electron sea” model of a metallic crystal 4. The metallic bonding is non-directional. i.e. no preferred attraction for delocalised electrons in any particular direction. 1.2 Factors affecting Strength of Metallic Bonds For Metals with different no. of delocalised electrons Metallic Bond strength ∝ No. of delocalised electrons For Metals with same no. of delocalised electrons Metallic Bond strength ∝ 1 ionic radius Inter–Atomic Bonds

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