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Bonding: General Concepts Chapter 8. Bonds Forces that hold groups of atoms together and make them function as a unit.

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Presentation on theme: "Bonding: General Concepts Chapter 8. Bonds Forces that hold groups of atoms together and make them function as a unit."— Presentation transcript:

1 Bonding: General Concepts Chapter 8

2 Bonds Forces that hold groups of atoms together and make them function as a unit.

3 Bond Energy -It is the energy required to break a bond. -It gives us information about the strength of a bonding interaction.

4 Ionic Bonds -Formed from electrostatic attractions of closely packed, oppositely charged ions. -Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

5 Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm) When E is positive (+), repulsion is indicated. When E is negative (-), attraction is indicated. This is a statement of Coulomb’s Law where:

6 Bond Length The distance where the system energy is a minimum.

7 Interaction of two hydrogen atoms.

8 Covalent Bonding - covalent bonds are formed by sharing electrons between nuclei. H. +. H ----> H-H - coordinate covalent bonds are bonds where both shared electrons originate on the same atom H 3 N: + H + ----> H 3 N-H +

9 Types of Covalent Bonds Polar covalent bond -- covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. A dipole moment exists. Nonpolar covalent bond -- covalent bond in which the electrons are shared equally between both atoms. No dipole moment exists.

10 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X) actual  (H  X) expected

11 Pauling Electronegativity Values

12 Percent Ionic Character where x A is the larger electronegativity and x B is the smaller value. Watch significant figures!!! Ionic Bond % IC > 50 % Polar Covalent % IC 5 - 50 % Nonpolar Covalent % IC < 5 %

13 Three Possible Types of Bonds Nonpolar Covalent (Electrons equally shared.) Polar Covalent (Electrons shared unequally.) Ionic (Electrons are transferred.)

14 Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

15 The Effect of an electric field on hydrogen fluoride molecules.

16 Dipole Moment for the water molecule.

17 Dipole moment for the ammonia molecule.

18 Nitrogen Trichloride Does nitrogen trichloride exhibit a dipole moment? Yes. It has three nonpolar bonds but, also, has a lone pair of electrons which makes it assymetrical and therefore, polar.

19 Nonpolar molecule--zero dipole moment.

20 Table 8.2 on page 357 in Zumdahl.

21 Achieving Noble Gas Electron Configurations (NGEC) Two nonmetals react: They share electrons to achieve NGEC. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC. See Table 8.3 on page 361 in Zumdahl.

22 Sizes of ion related to position on the periodic table.

23 Isoelectronic Ions Ions containing the same number of electrons (O 2 , F , Na +, Mg 2+, Al 3+ ) O 2  > F  > Na + > Mg 2+ > Al 3+ largest smallest

24 Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M + (g) + X  (g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

25 Formation of an Ionic Solid 1.Sublimation of the solid metal M(s)  M(g) [endothermic] 2.Ionization of the metal atoms M(g)  M + (g) + e  [endothermic] 3.Dissociation of the nonmetal 1 / 2 X 2 (g)  X(g) [endothermic]

26 Formation of an Ionic Solid (continued) 4.Formation of X  ions in the gas phase: X(g) + e   X  (g) [exothermic] 5.Formation of the solid MX M + (g) + X  (g)  MX(s) [quite exothermic]

27 Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions

28 .Comparison of the energy changes in the formation of sodium fluoride and magnesium oxide.

29 Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released Draw the Lewis Structure for each reactant and product before doing any calculations!

30 Single, Double, & Triple Bonds Single bonds -- one shared pair of electrons. Double bonds -- two shared pairs of electrons. Triple bonds -- three shared pairs of electrons. See bond energy Tables 8.4 & 8.5 on pages 373-374 in Zumdahl.

31 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

32 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 don’t misuse it.

33 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.

34 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.

35 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.

36 Lewis Structures Covalent Compounds Ionic Compounds In ionic compounds, electrons are transferred and ions are formed. In covalent compounds, electrons are shared to form a molecule.

37 Electron Deficient Molecules Beryllium chloride -- BeCl 2 -- is electron deficient with four electrons. It forms a linear molecule. Boron trifluoride -- BF 3 -- is electron deficient with six electrons. It forms a trigonal planar molecule. See page 381 for the reaction between boron trifluoride and ammonia.

38 Comments About the Octet Rule -2nd row elements C, N, O, F observe the octet rule. -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.

39 Rules for Writing Lewis Structures Sum the valence electrons from all the atoms. Use a pair of electrons to from a bond between each pair of bound atoms. Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row elements.

40 Lewis Structures NO + 5 e - + 6 e - - 1 e - = 10 e - [:N  O:] + Each atom has an octet and is satisfied.

41 Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures called a resonance hybrid. See the resonance structures for the nitrate ion on page 385 in Zumdahl.

42 Odd-Electron Molecules NO 2 contains 17 electrons. cannot satisfy the octet rule. a more sophisticated model is needed- the molecular orbital model.

43 Stereochemistry The study of the three- dimensional arrangement of atoms or groups of atoms within molecules and the properties which follow such arrangement.

44 VSEPR Model Valence Shell Electron Pair Repulsion -- The structure around a given atom is determined principally by minimizing electron pair repulsions.

45 Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Put pairs as far apart as possible. 3.Determine positions of atoms from the way electron pairs are shared.(Parent Geometry) 4.Determine the name of molecular structure from positions of the atoms.(Actual Geometry)

46 Molecular Geometry Parent Geometry is electron pair arrangement about the central atom. linear trigonal planar tetrahedral trigonal bipyramidal octahedral Actual Geometry is the arrangement of atoms about the central atom. linear bent trigonal pyramid seesaw T-shaped square pyramid square planar

47 Lone pair of electrons on the ammonia molecule.

48 08_143 (a)(b) H (c) Lone pair Bonding pair Bonding pair Lone pair H H O O H O Lone pairs on the water molecule.

49 Octahedral structure for phosphorus hexachloride.

50 Octahedral structure for xenon.

51 Parent and actual geometry for xenon tetrafluoride.

52 Three possible arrangements of the electron pairs in triiodide ion.


54 VSEPR Model Summary Determine the Lewis structure(s) for the molecule. For molecules with resonance structures, use any of the structures to predict the molecular structure. Sum the electron pairs around the central atom to determine the parent geometry. The arrangement of the pairs is determined by minimizing electron-pair repulsions.(Actual Geometry)

55 VSEPR Model Summary (Continued) Lone pairs require more space than bonding pairs since they are tightly attracted to only one nucleus. Lone pairs produce slight distortions of bond angles less than 120 o.

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