Ch_04_summary - CHAPTER 4 BONDING(IB TOPICS 4 AND 14...

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Unformatted text preview: CHAPTER 4 BONDING (IB TOPICS 4 AND 14) SUMMARY A chemical bond is the interaction between atoms within a molecule, between molecules or between ions of opposite charges. Bond formation is an exothermic process; it gives out energy and leads to a more energetically stable state. Bonding in liquids and solids 1. Covalent bonding results from electron sharing between non-metals or a non-metal and a metal of higher electronegativity. The electron pair is attracted by both nuclei leading to a bond that is directional in nature. (i) van der Waals’ Forces: For a non-polar molecule, only weak, temporary, instantaneous dipole-dipole interaction called van der Waals’ forces exist between molecules. These molecules are low melting point solids and low boiling point liquids and gases. Larger the molecule, stronger the van der Waals’ forces, higher its boiling and melting points. (ii) Dipole-dipole Interaction: In a polar molecule, besides weak van der Waals’ forces, the molecules experience stronger permanent dipole-dipole interaction. (iii) Hydrogen Bonding: Elements of high electronegativity (F, O, N), bonded to a (tiny) hydrogen atom give rise to a special case of dipole-dipole interaction called H-bonding. This is important in determining solubility, melting and boiling points, and stability of crystal structures. Hydrogen bonding also plays an important role in biological systems. Hydrogen bonded molecules experience stronger hydrogen bonding in addition to van der Waals’ forces and dipole-dipole interaction. Strength of bonding: van der Waals’ < dipole-dipole < H-bond << covalent bond ≈ ionic bond. However, van der Waals’ forces can become extensive depending on the size of the molecules, e.g., in polymers and long chain hydrocarbon molecules. 2. Ionic bonding occurs as a result of electron transfer from active metals of Groups 1, 2, 3 to active non-metals of Groups 6, 7, leading to electrostatic attraction between ions of opposite charges. Ionic crystals are regular repeating arrays of positive and negative ions, packed so that each positive ion is surrounded by negative ions and vice versa in a 3 dimensional lattice structure. Ionic crystals have high melting and boiling points, are water soluble (water solubility varies considerably) and conduct electricity when molten or in aqueous solution, but not in the solid state, as ions are not free to move about in a solid ionic crystal. 3. Metals are good conductors of electricity and heat, ductile and malleable. Metals have low ionization energies and vacant valence orbitals. Metals consist of regular, repeating arrangements of metal ions (cations) in which the bonding electrons are mobile or delocalized and can move rapidly through the metal, since an electron is much smaller than the spaces between the cations. This explains high conductivity of elements. Electrons can acquire large amounts of kinetic energy and rapidly transfer it to cooler parts of the metal: this explains thermal conductivity. The movement of cations between layers can take place with no change in metallic bonding. This explains ductility and malleability. 4. Allotropes of carbon: Diamond, Graphite and C60 Fullerene form covalent bonds and network solids. In diamond, each carbon is covalently bonded to four other carbon atoms in a three dimensional tetrahedral arrangement. Diamond is very hard, stable with a very high melting point and does not conduct electricity as all its valence electrons are involved in covalent bonding (sp3 hybridized). Graphite has a layered structure; each layer has carbon atoms covalently bonded to three other carbon atoms (sp2 hybridized) in a hexagonal arrangement but with only weak van der Waals’ forces between layers. Graphite is easy to break, and soft with lubricating qualities. Graphite is a good conductor due to mobility of electrons in pi orbitals that are delocalized. C60 fullerene contains 60 carbon atoms (20 hexagons and 12 pentagons). It is a highly symmetrical, closedcarbon-caged molecule. Within each molecule, each carbon atom is covalently bonded to three other carbon atoms (sp2 hybridized). C60 molecules contain two bond lengths – © IBID Press 2007 1 CHAPTER 4 BONDING (IB TOPICS 4 AND 14) SUMMARY the hexagons can be considered "double bonds" and are shorter; the longer bonds in the pentagonal rings suggest that electron delocalization is poor. • Lewis Electron-Dot Structures: Structures showing covalent bonds by using symbols of element(s) involved, and indicating all the valence electrons by using dots, crosses, a combination of dots and crosses or by using a line to represent a pair of electrons. • Valence Shell Electron Pair Repulsion (VSEPR) Theory: Electron pairs in the valence shell of the central atom are arranged as far apart as possible due to mutual repulsion. This minimizes the forces of repulsion between the electron pairs. The species therefore has minimum energy and maximum stability. SHAPES WITH ONLY BONDED ELECTRON PAIRS ON THE CENTRAL ATOM: #of Bonded e– Pairs 2 (and no lone e- pairs) 3 (and no lone e- pairs) 4 (and no lone e- pairs) 5 (and no lone e- pairs) 6 (and no lone e- pairs) Structure Linear Trigonal planar Regular Tetrahedral Trigonal bipyramidal Octahedral or square bipyramidal Angle 180o 120 o 109.5o 90o, 120 o, 180 o 90 o and 180 o Examples BeCl2, CO2, C2H2 BCl3, CO32–, NO3– CH4, BF4–, NH4+ PCl5 SF 6 SHAPES WITH 3 ELECTRON PAIRS (BONDED + LONE PAIRS) ON CENTRAL ATOM Number of bonded e– pairs 3 Number of lone e– pairs 0 2 1 Structure Angle Examples Trigonal planar Bent, V-Shaped, or angular 120o less than 120o (or ≈120o) BCl3, CO32–, NO3– O3, SO2, NO2– SHAPES WITH 4 ELECTRON PAIRS (BONDED + LONE PAIRS) ON CENTRAL ATOM Number of bonded e– pairs Number of lone e– pairs Structure Angle Examples 4 0 Regular Tetrahedral Regular Tetrahedral 109½o CH4, BF4–, NH4+ 3 1 2 ∠ ≈107o or less than 109o ∠ ≈105o or less than 109o NH3, PCl3, SO32– 2 Trigonal pyramidal Bent, V-shaped or Angular © IBID Press 2007 H2O 2 CHAPTER 4 BONDING (IB TOPICS 4 AND 14) SUMMARY SHAPES WITH 5 ELECTRON PAIRS (BONDED + LONE PAIRS) ON CENTRAL ATOM Number of bonded e– pairs Number of lone e– pairs 5 0 4 1 3 2 2 3 Structure Angle Examples Trigonal bipyramidal Unsymmetrical tetrahedral T-shaped Linear 90o, 120 o, 180 o PCl5 ≈ 90o, 120 o, 180 o SF4 90o, 180 o 180 o ClF3 I3 – SHAPES WITH 6 ELECTRON PAIRS (BONDED + LONE PAIRS) ON CENTRAL ATOM Number of bonded e– pairs Number of lone e– pairs 6 0 5 4 1 2 • • • • • Structure Angle Examples Octahedral or square bipyramidal Square pyramidal Square planar 90 o and 180 o SF6 ≈90 o and 180 o 90 o BrF5 XeF4 Note that it is important to specify the angle as both O3 and H2O are bent molecules, but with different angles around the central atoms. A lone e– pair is spread out over a larger volume than a shared e– pair, causing a distortion from regular geometry, and causing the angle to be slightly reduced. A Resonance structure is one of two or more Lewis structures for a species that cannot be described adequately by a single structure. Resonance structures have the same sigma bonds but differ in the arrangement of the pi bonds. The actual bonding is a mixture of the various possible arrangements, and will have a lower energy than any of the individual forms. The phenomenon is called delocalization because the valence electrons provided by individual atoms are no longer held in the vicinity of that atom, but are mobile and shared by a number of atoms. It is this spreading out of the electrons that gives the structure its lower potential energy than it would have if double and single bonds were arranged in such a way that orbital overlap could not take place. Thus, delocalization stabilizes a structure. Bonds have electron distribution with axial symmetry around the axis joining the two nuclei (from combination of two s orbitals or an s and a p or hybridized orbitals); π bonds result from the sideways overlap of parallel p orbitals with electron distribution above and below the nuclei. Hybridization involves mixing of atomic orbitals to give new hybrid orbitals of equivalent energy, i.e., it is the redistribution of energy in different orbitals (in the valence shell) of an atom. © IBID Press 2007 3 CHAPTER 4 BONDING (IB TOPICS 4 AND 14) SUMMARY Type of Hybridization Examples Bonds Shape Angle sp3 (combination of: one s and three p orbtals) CH4 NH3 H2O 4 σ bonds 3 σ; 1 lone e− pair 2 σ; 2 lone e− pairs Tetrahedral Trigonal pyramidal Bent/angular/”V” = 109.5° <109.5° (≈107°) <109.5° (≈105°) sp2(combination of: one s and two p orbtals) C2H4 Each C has 3 σ; 1 π C=C has 1 σ; 1 π B has 3 σ bonds Each C has 3 σ; 1 π Planar 120° Planar 120° Planar 120° Linear 180° Linear 180° BCl3 C6H6 sp(combination of: one s and one p orbtal) C2H2 Each C has 2 σ; 2 π C≡C has 1 σ; 2 π Be has 2 σ bonds BeCl2 (N.B. Shading indicates Topic 14 (AHL) material.) © IBID Press 2007 4 ...
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This note was uploaded on 09/23/2010 for the course CS 001 taught by Professor Jix during the Spring '10 term at Riverside Community College.

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