Presentation on theme: "1 September 2013 Inorganic Chemistry Part 2a Revised Adv Higher Unit 1 Based on Presentations produced by David P. White University of North Carolina,"— Presentation transcript:
1 September 2013 Inorganic Chemistry Part 2a Revised Adv Higher Unit 1 Based on Presentations produced by David P. White University of North Carolina, Wilmington Gordon Watson Chemistry Department, Kelso High School
2 September 2013Introduction Covalent Bonding Shapes This topic explores further aspects of Covalent Bonding, leading to an understanding of the Shapes of molecules & polyatomic ions
3 September 2013 Covalent Bonding Covalent Bond sharing A Covalent Bond will form when atoms can rearrange their electrons (by sharing ) to produce an arrangement of lower energy. Bond Enthalpy Bond Length
4 September 2013 Ionic Bonding Ionic Bond transfer An Ionic Bond will form when atoms can rearrange their electrons ( transfer ) to produce an arrangement of lower energy. ‘stable octet’ Ionic Lattice Despite the advantage gained by achieving a ‘stable octet’, the real driving force behind ionic bonding is the stability of the Ionic Lattice formed.
5 September 2013 Electronegativity 1 Ionic and Covalent are, in fact, rather arbitrary labels and it is no longer enough to simply look to see if a metal element is involved or not. Electronegativity Electronegativity values are a useful guide but properties will still need to be studied to provide confirmation.
6 September 2013 Electronegativity 2 bonding continuum Ionic and Covalent are just opposite ends of a bonding continuum Most bonds lie between these two extremes.
7 September 2013 Electronegativity 3
8 September 2013 Lewis Diagrams models Various methods ( models ) can be used to explain the properties associated with covalent bonding. Lewis Diagrams One useful model positions electrons ( both bonding and non- bonding) around atoms at ‘ four corners ’. These are called Lewis Diagrams
9 September 2013 Lewis Structures ‘stable octet’ In Ammonia the central nitrogen atom has 4 electron pairs, the ‘stable octet’ ‘bonding pairs’‘non- bonding pair’‘lone pair’ There are 3 ‘bonding pairs’ and 1 ‘non- bonding pair’ - often referred to as a ‘lone pair’ DativeCoordinate Sometimes atoms will use a lone pair as a bonding pair to form a Dative or Coordinate Covalent bond, as in the Ammonium ion, NH 4 +
10 September 2013 More Lewis Structures C HH HH HH Ethane CCl.. CCl 4 Carbon Tetrachloride.. CO CO Carbon Monoxide.. COO CO 2 Carbon Dioxide H O H.. H 2 O Water Hydrogen Oxide
11 September 2013 Unusual Lewis Structures ‘stable octet’ In both these molecules, the Be and B atoms do not achieve a ‘stable octet’. The resulting molecules are, predictably, very reactive (unstable). In both these molecules, the P and S atoms have in excess of a ‘ stable octet ’. d-orbital The extra orbitals needed, to cope with more than 8 electrons, come from the empty d-orbital set. hybridisation This ‘ mixing ’ of orbitals is called hybridisation
12 September 2013 Hybridisation - sp orbitals
13 September 2013 Hybridisation - sp 2 orbitals
14 September 2013 Hybridisation - sp 3 orbitals
15 September 2013 Hybridisation - sp 3 orbitals
16 September 2013 Hybridisation - sp 3 d orbitals
17 September 2013 Hybridisation - sp 3 d 2 orbitals
18 September 2013 Shapes of Molecules valence electrons In order to predict molecular shape, we consider that all outer shell electrons ( valence electrons ) of the central atom repel each other. Therefore, the molecule adopts whichever 3D geometry minimizes this repulsion. bondingnon-bonding Both bonding and non-bonding electron pairs must be considered. VSEPR We call this process the Valence Shell Electron Pair Repulsion ( VSEPR ) theory.
19 September 2013 Electron Pair Geometry 1
20 September 2013 Electron Pair Geometry 2
21 September 2013 Counting Electron Pairs number of atoms number of electron pairs It is important not to confuse the number of atoms with the number of electron pairs. For example, CF 4 and XeF 4 may, at first sight, appear likely to be the same shape. Electron pairs = electrons of central atom + no. of atoms attached divided by 2 CF 4 = XeF 4 = electron pairs = 4 electron pairs 6 electron pairs = 6 electron pairs
22 September 2013 Molecular Geometry 1
23 September 2013 Molecular Geometry 2
24 September 2013 Molecular Geometry 3
25 September 2013 Molecular Geometry 4
26 September 2013 Repulsive Forces 1 Bonding electrons, because they are attracted by two nuclei, do not repel as much as non-bonding electrons. This can cause ‘ distortions ’ in the shapes of molecules:
27 September 2013 Repulsive Forces 2 Electron pair repulsions decrease in strength in the order: non-bonding/non-bonding non-bonding/bonding bonding/bonding ???
28 September 2013 Repulsive Forces 3 Placing the non-bonding lone pairs at the Axial positions would appear to give least repulsion but they would only be 90° away from the 3 bonding pairs. To minimize e − − e − repulsion, lone pairs are always placed in equatorial positions, so …
29 September 2013 End of Part 2a Inorganic Chemistry Part 2a