2021 H1 Chemical Bonding Lecture Tutor

St Andrew’s Junior College JC1 H1 Chemistry 2021 1 St Andrew’s Junior College H1 Chemistry 2021 Lecture Notes 3: Chemical Bonding (TUTOR) Assessment Objectives: (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 bonds (see also topic on Introduction to Organic Chemistry) (d) explain the shapes of, and bond angles in, molecules such as BF 3 (trigonal planar); CO 2 (linear); CH 4 (tetrahedral); NH3 (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 [quantitative treatment of electro-negativity 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 liquid and gaseous CHCl3, Br2 and the noble gases (ii) hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups (i) outline the importance of intermol ecular 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 a nd 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

St Andrew’s Junior College JC1 H1 Chemistry 2021 2 (m) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride, magnesium oxide (ii) simple molecular, as in iodine (iii) giant molecular, as in graphite; diamond (iv) hydrogen-bonded, as in ice (v) metallic, as in copper [the concept of the ‘unit cell’ is not required] (n) describe, interpret and/or predi ct 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 Introduction to Organic Chemistry: (f) (i) describe the shapes of the ethane, ethene and benzene molecules (ii) explain the shapes of, and bond angles, in the ethane, ethene and benzene molecules in relation to σ and carbon-carbon bonds (iii) predict the shapes of, and bond angles in, molecules analogous to those specified in (f)(ii) Content 1 Metallic Bonding 5 Covalent Bond Strength 2 Ionic bonding 6 Effect of electronegativity difference on bond type 3 Covalent bonding 7 Physical properties of Simple Covalent Molecules 4 Shapes of Molecules and Molecular Ions 8 Giant Molecular Structure Recommended Materials: 1 Cann, Peter and Hughes, Peter, Chapter 3, Page 52, Chemical Bonding in simple m olecules, Chemistry for advanced level, 2002 2 Covalent Bonding: http://chemguide.co.uk/atoms/bonding/covalent.html

St Andrew’s Junior College JC1 H1 Chemistry 2021 3 Introduction The reaction of two or more elements results in the formation of a chemical bond between atoms. The chemical bond formed between atoms results in the formation of different types of structures. 1 Metallic Bonding 1.1 Formation of Metallic Bonds Definition: Metallic bonding is the electrostatic attraction between a lattice of positive ions and delocalised electrons.  Metal atoms have low electronegativities and low ionisation energies.  Hence, they lose electron s readily to form positive ions which are arranged in a lattice structure.  The valence electrons lost form a mobile sea of electrons which are free to move throughout the three-dimensional structure. i.e. delocalised. They occupy the space between the metal cations. Note: (+) cations should be regularly arranaged and should not touch. (a) show understanding that all chemical bonds are electrostatic in nature and describe: (iii) metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons (m) describe, in simple terms, the lattice structure of a crystalline solid which is: (v) metallic, as in copper e– e–

St Andrew’s Junior College JC1 H1 Chemistry 2021 4 1.2 Strength of Metallic Bond (a) Metallic bond is strong. (b) The stronger the metallic bond, the higher its melting and boiling point. The strength of metallic bond is dependent on: (i) Number of valence electrons contributed to the sea of delocalised electrons / Charge of the cation : Generally, the more valence electrons the metal contribute s (the higher the charge of the metal cation), the stronger the electrostatic forces of attraction between metal cation and sea of delocalised electrons. Hence, the metallic bonding will be stronger. (ii) Size of the cation: Given the same type of packing, the smaller the size of the cations, the closer the delocalised electrons are to the positively charged nuclei. Hence, the metallic bonding will be stronger. E.g. m.p. of A l > Mg > Na due to increasing ionic size (i.e. A l3+ < Mg 2+ < Na +) and decreasing number of delocalised valence electrons. E.g. Transition metals have very high m.p. (> 1000 oC) due to their very small ionic size and large number of delocalised electrons (the d electrons and the s electrons in the valence shell can both delocalise). Format of answering boiling/melting point questions: Structure Interaction Amount of Energy

St Andrew’s Junior College JC1 H1 Chemistry 2021 5 Exercise 1 [Nov 2013 / II / 4b(i)] (a) Describe the structure and bonding in calcium with the aid of a labelled diagram. (b) Suggest why the melting point of calcium is higher that potassium. Answer: (a) Calcium has a giant metallic lattice structure with strong electrostatic forces of attraction between the lattice of Ca2+ and the delocalised electrons. (b) Both calcium and potassium have giant metallic lattice structure but calcium has a greater number of delocalised valence electrons and a smaller cationic size than potassium and stronger electrostatic forces of attraction between the cations and the valence ele ctrons. Hence, more energy is required to overcome the metallic bonds in calcium and hence higher melting point. Calcium Potassium Structure Giant Metallic lattice Giant Metallic lattice Interaction Stronger metallic bonds Weaker metallic bonds Amount of Energy More energy Less energy Thinking Proc