Chapter 8 - Chapter 8 Bonding: General Concepts 8.1 Types...

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Chapter 8 Bonding: General Concepts 8.1 Types of Chemical Bonds Atoms or ions combine with each others to form compounds to achieve the lowest possible energy state. 1. Ionic bonding : e.g : (Na + + Cl - NaCl) The electrostatic attractions are the driving force . Ionic compound results when a metal reacts with a nonmetal . The energy of interaction between a pair of ions can be calculated using Coulomb’s law: Χ = - r Q Q nm J E 2 1 19 . 10 31 . 2 where, r is the distance between the ion centers and Q 1 and Q 2 are numerical ion charges . e.g : In solid NaCl r = 2.76A o = 0.276 nm J nm nm J E 19 19 10 37 . 8 276 . 0 ) 1 )( 1 ( . 10 31 . 2 - - Χ - = - + Χ = Negative sign indicates an attractive force(Q 1 areQ 2 are different) , while Positive sign indicates a repulsive force(Q 1 areQ 2 are the same)
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Bond energy : It is the energy required to break the bond. Bond length : It is the equilibrium distance where the energy of the system is minimum. 2. Covalent bonding : e.g : (H + H H 2 ) The mutual attraction of the nuclei for the shared electrons is the driving force . When bonded atoms are different , they form a polar covalent bond: e.g : H δ + F δ - and H δ + O δ - 8.2 Electronegativity It is the ability of an atom in a molecule to attract shared electrons to itself. Pauling electronegativity : The relative elcetronegativities of the H and X atoms are determined by comparing the measured H - X bond energy which is an average of the H H and X - X bond energies (b.e): H - X bond energy = [( H - H b.e ) + ( X - X b.e )]/2 2
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(In Table 8.3 at page 354 are the assigned electronegativity values for all elements ) In general : Electronegativity and (hence polarity ) increases across a period and decreases down a group The range of electronegativity values is from 4.0 for fluorine to 0.7 for cesium . Example 8.1 : Order the following bonds according to polarity : H-H < S-H < Cl-H < O-H < F-H (2.1) (2.1) (2.5) (2.1) (3.0) (2.1) (3.5) (2.1) (4.0) (2.1) 0 0.4 0.9 1.4 1.9 Electronegativity difference and polarity increases Electronegativity difference Bond type Zero Covalent Intermediate Polar covalent Large Ionic 3
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8.3 Bond Polarity and Dipole Moments: The bond H δ + F δ - is said to be dipolar and to have a dipole moment that points to the negative charge 1. Molecules with polar bonds and have net dipole moment : e.g: HX, H 2 O and NH 3 2. Molecules with polar bonds but no net dipole moment: a. Linear molecule : e.g : CO 2 (O=C=O) b. Planar molecules: e.g: BF 3 and SO 3 c. Tetrahedral molecules: e.g: CH 4 4
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8.4 Ions: Electron Configurations and sizes: Atoms in stable compounds usually have a noble gas electron configuration. When we speak in this text of the stability of an ionic
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This note was uploaded on 10/25/2010 for the course CHEMISTRY 2010 taught by Professor Thwapiah during the Fall '10 term at Hashemite University.

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Chapter 8 - Chapter 8 Bonding: General Concepts 8.1 Types...

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