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Basic Concepts of Chemical Bonding The ability of an atom in a molecule to attract shared electrons to itself. Electronegativity: The ability of an atom.

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Presentation on theme: "Basic Concepts of Chemical Bonding The ability of an atom in a molecule to attract shared electrons to itself. Electronegativity: The ability of an atom."— Presentation transcript:


2 Basic Concepts of Chemical Bonding

3 The ability of an atom in a molecule to attract shared electrons to itself. Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself.

4 Ionic Bonds Electrons are transferred Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!

5 Determination of Ionic Character Compounds are ionic if they conduct electricity in their molten state Electronegativity difference is not the final determination of ionic character

6 Coulombs Law The energy of interaction between a pair of ions is proportional to the product of their charges, divided by the distance between their centers

7 Table of Ion Sizes

8 Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

9 Estimate H f for Sodium Chloride Na(s) + ½ Cl 2 (g) NaCl(s) Lattice Energy-786 kJ/mol Ionization Energy for Na495 kJ/mol Electron Affinity for Cl-349 kJ/mol Bond energy of Cl 2 239 kJ/mol Enthalpy of sublimation for Na109 kJ/mol Na(s) Na(g) + 109 kJ Na(g) Na + (g) + e - + 495 kJ ½ Cl 2 (g) Cl(g) + ½(239 kJ) Cl(g) + e - Cl - (g) - 349 kJ Na + (g) + Cl - (g) NaCl(s) -786 kJ Na(s) + ½ Cl 2 (g) NaCl(s) -412 kJ/mol

10 Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds Electrons are unequally shared Electronegativity difference between.3 and 1.7 Electrons are equally shared Electronegativity difference of 0 to 0.3

11 Covalent Bonding Forces Electron – electron Electron – electron repulsive forces repulsive forces Proton – proton repulsive forces Electron – proton attractive forces

12 Bond Length Diagram

13 Bond Length and Energy BondBond type Bond length (pm) Bond Energy (kJ/mol) C - CSingle154347 C = CDouble134614 C Triple120839 C - OSingle143358 C = ODouble123745 C - NSingle143305 C = NDouble138615 C NTriple116891 Bonds between elements become shorter and stronger as multiplicity increases.

14 Bond Energy and Enthalpy D = Bond energy per mole of bonds Energy requiredEnergy released Breaking bonds always requires energy Breaking = endothermic Forming bonds always releases energy Forming = exothermic

15 The Octet Rule Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

16 Formation of Water by the Octet Rule

17 Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule (HONC rule as well). 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

18 Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

19 C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make carbon the central atom..

20 Multiple Covalent Bonds: Double bonds Two pairs of shared electrons Ethene

21 Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons Ethyne

22 Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The actual structure is an average of the resonance structures. Benzene, C 6 H 6 The bond lengths in the ring are identical, and between those of single and double bonds.

23 Resonance Bond Length and Bond Energy Resonance bonds are shorter and stronger than single bonds. Resonance bonds are longer and weaker than double bonds.

24 Resonance in Ozone, O 3 Neither structure is correct. Oxygen bond lengths are identical, and intermediate to single and double bonds

25 Resonance in a carbonate ion: Resonance in an acetate ion: Resonance in Polyatomic Ions

26 Localized Electron Model Lewis structures are an application of theLocalized Electron Model L.E.M. says: Electron pairs can be thought of as belonging to pairs of atoms when bonding Resonance points out a weakness in the Localized Electron Model.

27 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

28 Fundamental Properties of Models A model does not equal reality. Models are oversimplifications, and are therefore often wrong. Models become more complicated as they age. We must understand the underlying assumptions in a model so that we dont misuse it.

29 VSEPR – Valence Shell Electron Pair Repulsion X + E Overall Structure Forms 2 LinearAX 2 3 Trigonal PlanarAX 3, AX 2 E 4 TetrahedralAX 4, AX 3 E, AX 2 E 2 5 Trigonal bipyramidalAX 5, AX 4 E, AX 3 E 2, AX 2 E 3 6 OctahedralAX 6, AX 5 E, AX 4 E 2 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

30 VSEPR: Linear AX 2 CO 2

31 VSEPR: Trigonal Planar AX 3 AX 2 E BF 3 SnCl 2

32 VSEPR: Tetrahedral AX 4 AX 3 E AX 2 E 2 CCl 4 PCl 3 Cl 2 O

33 VSEPR: Trigonal Bi-pyramidal AX 5 AX 4 E AX 3 E 2 AX 2 E 3 PCl 5 SF 4 ClF 3 I3-I3-I3-I3-

34 VSEPR: Octahedral AX 6 AX 5 E AX 4 E 2 SF 6 ICl 4 - BrF 5

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