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Chapter 4 Bonding: General Concepts. Chapter 8 Table of Contents 4.1 Types of Chemical Bonds 4.2 Electronegativity 4.3 Bond Polarity and Dipole Moments.

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Presentation on theme: "Chapter 4 Bonding: General Concepts. Chapter 8 Table of Contents 4.1 Types of Chemical Bonds 4.2 Electronegativity 4.3 Bond Polarity and Dipole Moments."— Presentation transcript:

1 Chapter 4 Bonding: General Concepts

2 Chapter 8 Table of Contents 4.1 Types of Chemical Bonds 4.2 Electronegativity 4.3 Bond Polarity and Dipole Moments 4.4 Ions: Electron Configurations and Sizes 4.5 Energy Effects in Binary Ionic Compounds 4.6Partial Ionic Character of Covalent Bonds 4.7The Covalent Chemical Bond: A Model 4.8Covalent Bond Energies and Chemical Reactions 4.9The Localized Electron Bonding Model 4.10Lewis Structures 4.11Exceptions to the Octet Rule 4.12Resonance

3 Chapter 4 Copyright © Cengage Learning. All rights reserved 3 Questions to Consider What is meant by the term “chemical bond”? No simple, and yet complete, way to define this.

4 Section 4.1 Types of Chemical Bonds Return to TOC molecular structure : atoms arranged within molecules molecular bonding : the forces that hold groups of atoms in a molecule together (1) Ionic bonding : transfer of e is nearly complete (s) ex. metal + nonmetal Na ‧ + ‧ Cl : ─→ [ Na + Cl - ].. charge

5 Section 4.1 Types of Chemical Bonds Return to TOC (2) Covalent bonding : atoms share es when they bind. (g, l, s) ex. H ‧ + ‧ H → H 2 nonmetal nonmetal H ‧ + ‧ F : ─→ H - F..  +  - (polar C.B.) e spend more time in F (nonpolar C.B.) es equally shared

6 Section 4.1 Types of Chemical Bonds Return to TOC Copyright © Cengage Learning. All rights reserved 6 The Interaction of Two Hydrogen Atoms

7 Section 4.2 Electronegativity Return to TOC (1)Linus Pauling (1930) The ability of an atom in a molecule to attract shared es to itself. (2)  ↑  attract e ↑  | E.A. | ↑  I.E. ↑ (3) (4) pure covalent bond : same  polar covalent bond : similar  ionic bond : widely different  ;    2  most  : halogen, oxygen, S

8 Section 4.2 Electronegativity Return to TOC (5)  & E.A.  tendency of an atom to attract e.  : relative number ; attraction of an atom in a chemical bond (with another atom) for the shared es. E.A. : refers to an isolated atom’s attraction for an addition e ; measurable quantity.

9 Section 4.2 Electronegativity Return to TOC Copyright © Cengage Learning. All rights reserved 9 The Pauling Electronegativity Values

10 Section 4.2 Electronegativity Return to TOC Copyright © Cengage Learning. All rights reserved 10 The Relationship Between Electronegativity and Bond Type

11 Section 4.3 Bond Polarity and Dipole Moments Return to TOC Copyright © Cengage Learning. All rights reserved 11 Dipole Moment Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge. Use an arrow to represent a dipole moment.  Point to the negative charge center with the tail of the arrow indicating the positive center of charge.

12 Section 4.3 Bond Polarity and Dipole Moments Return to TOC Copyright © Cengage Learning. All rights reserved 12 Dipole Moment

13 Section 4.3 Bond Polarity and Dipole Moments Return to TOC Copyright © Cengage Learning. All rights reserved 13 No Net Dipole Moment (Dipoles Cancel)

14 Section 4.4 Ions: Electron Configurations and Sizes Return to TOC Copyright © Cengage Learning. All rights reserved 14 Stable Compounds Atoms in stable compounds usually have a noble gas electron configuration. Why NaCl 2, NaNe not occur ?

15 Section 4.4 Ions: Electron Configurations and Sizes Return to TOC (1) Ionic size r cation < r atom r anion > r atom isoelectronic : same number of e O 2 -, F -, Na +, Mg 2 + r : O 2 - > F - > Na + > Mg 2 + ( 2) Lattice energy M + (g) + X - (g) → MX(s)  H  H : LiCl > NaCl > KCl ( ∵ r : Na + < Na + < K + )  H : MgO > NaCl ( ∵ Q : Mg 2+,O 2 - > Na +,Cl - ) always

16 Section 4.6 Partial Ionic Character of Covalent Bonds Return to TOC Copyright © Cengage Learning. All rights reserved 16 The relationship between the ionic character of a covalent bond and the electronegativity difference of the bonded atoms

17 Section 4.7 The Covalent Chemical Bond: A Model Return to TOC Copyright © Cengage Learning. All rights reserved 17 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

18 Section 4.7 The Covalent Chemical Bond: A Model Return to TOC Copyright © Cengage Learning. All rights reserved 18 Fundamental Properties of Models 1.A model does not equal reality. 2.Models are oversimplifications, and are therefore often wrong. 3.Models become more complicated and are modified as they age. 4.We must understand the underlying assumptions in a model so that we don’t misuse it. 5.When a model is wrong, we often learn much more than when it is right.

19 Section 4.8 Covalent Bond Energies and Chemical Reactions Return to TOC Copyright © Cengage Learning. All rights reserved 19 Bond Energies To break bonds, energy must be added to the system (endothermic). To form bonds, energy is released (exothermic).

20 Section 4.8 Covalent Bond Energies and Chemical Reactions Return to TOC (1) Bond order : the number of bonding e pairs shared by 2 atoms order 1 : single bondex. H-H order 2 : double bondex. O = C = O order 3 : triple bondex. : C  O : fractional bond order :ex. O

21 Section 4.8 Covalent Bond Energies and Chemical Reactions Return to TOC (2) Bond length : the distance between the nuclei of 2 bonded atoms a) C – N < C – C < C – P(size of atom) C = O < C = S b) C – O > C = O > C  O(bond order) (3) Bond energy (D) bond dissociation E always  molecule (g) molecular fragments (g) molecule (l, s)  H =  D (bonds broken) –  D (bonds formed)  H l → g s→ g E supplied,  H > 0 E released,  H < 0 E required E released

22 Section 4.8 Covalent Bond Energies and Chemical Reactions Return to TOC Copyright © Cengage Learning. All rights reserved 22 Exercise Predict  H for the following reaction: Given the following information: Bond Energy (kJ/mol) C–H 413 C–N 305 C–C  H = –42 kJ

23 Section 4.9 The Localized Electron Bonding Model Return to TOC Copyright © Cengage Learning. All rights reserved 23 Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

24 Section 4.9 The Localized Electron Bonding Model Return to TOC Copyright © Cengage Learning. All rights reserved 24 Localized Electron Model Electron pairs are assumed to be localized on a particular atom or in the space between two atoms:  Lone pairs – pairs of electrons localized on an atom  Bonding pairs – pairs of electrons found in the space between the atoms

25 Section 4.9 The Localized Electron Bonding Model Return to TOC Copyright © Cengage Learning. All rights reserved 25 Localized Electron Model 1.Description of valence electron arrangement (Lewis structure). 2.Prediction of geometry (VSEPR model). 3.Description of atomic orbital types used to share electrons or hold lone pairs.

26 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 26 Lewis Structure 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.

27 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 27 Duet Rule Hydrogen forms stable molecules where it shares two electrons.

28 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 28 Octet Rule Elements form stable molecules when surrounded by eight electrons.

29 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 29 Single Covalent Bond A covalent bond in which two atoms share one pair of electrons. H–H

30 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 30 Double Covalent Bond A covalent bond in which two atoms share two pairs of electrons. O=C=O

31 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 31 Triple Covalent Bond A covalent bond in which two atoms share three pairs of electrons.

32 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 32 Steps for Writing Lewis Structures 1.Sum the valence electrons from all the atoms. 2.Use a pair of electrons to form a bond between each pair of bound atoms. 3.Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

33 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 33 Steps for Writing Lewis Structures 1.Sum the valence electrons from all the atoms. (Use the periodic table.) Example: H 2 O 2 (1 e – ) + 6 e – = 8 e – total

34 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 34 Steps for Writing Lewis Structures 2.Use a pair of electrons to form a bond between each pair of bound atoms. Example: H 2 O

35 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 35 Steps for Writing Lewis Structures 3.Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Examples: H 2 O, PBr 3, and HCN

36 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 36 Concept Check Draw a Lewis structure for each of the following molecules: H2H2 F2F2 HF

37 Section 4.10 Lewis Structures Return to TOC Copyright © Cengage Learning. All rights reserved 37 Concept Check Draw a Lewis structure for each of the following molecules: NH 3 CO 2 CCl 4

38 Section 4.11 Exceptions to the Octet Rule Return to TOC Copyright © Cengage Learning. All rights reserved 38 Boron tends to form compounds in which the boron atom has fewer than eight electrons around it (it does not have a complete octet). BH 3 = 6e –

39 Section 4.11 Exceptions to the Octet Rule Return to TOC Copyright © Cengage Learning. All rights reserved 39 When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, place the extra electrons on the central atom. SF 4 = 34e – AsBr 5 = 40e –

40 Section 4.11 Exceptions to the Octet Rule Return to TOC Copyright © Cengage Learning. All rights reserved 40 Concept Check Draw a Lewis structure for each of the following molecules: BF 3 PCl 5 SF 6

41 Section 4.11 Exceptions to the Octet Rule Return to TOC Copyright © Cengage Learning. All rights reserved 41 Let’s Review C, N, O, and F should always be assumed to obey the octet rule. B and Be often have fewer than 8 electrons around them in their compounds. Second-row elements never exceed the octet rule. Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals.

42 Section 4.11 Exceptions to the Octet Rule Return to TOC Copyright © Cengage Learning. All rights reserved 42 Let’s Review When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond).

43 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 43 More than one valid Lewis structure can be written for a particular molecule. NO 3 – = 24e –

44 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 44 Actual structure is an average of the resonance structures. Electrons are really delocalized – they can move around the entire molecule.

45 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 45 Concept Check Draw a Lewis structure for each of the following molecules: COCO 2 CH 3 OHOCN –

46 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 46 Formal Charge Used to evaluate nonequivalent Lewis structures. Atoms in molecules try to achieve formal charges as close to zero as possible. Any negative formal charges are expected to reside on the most electronegative atoms.

47 Section 4.12 Resonance Return to TOC Formal charge (F.C.) F.C. = (# of v. es in free atom) – (# of es assigned to atom in Lewis structure) = group# – # of lone pair e – ½(# of bonding e) ex. O = C = O or O  C - O or C=O=O… favor unfavor  The best Lewis structure is usually the one with the lowest F.C. ex. CS 2,COCl 2,SO 4 2-

48 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 48 Concept Check Consider the Lewis structure for POCl 3. Assign the formal charge for each atom in the molecule. P: 5 – 4 = +1 O: 6 – 7 = –1 Cl: 7 – 7 = 0

49 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 49 Rules Governing Formal Charge The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species.

50 Section 4.12 Resonance Return to TOC Copyright © Cengage Learning. All rights reserved 50 Rules Governing Formal Charge If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion.


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