chapter 9

Transcript chapter 9

Ch 9 Bonding and Molecular Structure: Fundamental Concepts

 Valence Electrons  Chemical Bond Formation  Bonding in Ionic Compounds  Covalent Bonding and Lewis Structures  Resonance  Exceptions to the Octet Rule  Charge Distribution in Covalent Molecules  Bond Properties  Molecular Shapes  Molecular Polarity

Valence Electrons

 There are

core electrons

and

valence electrons

 Valence electrons, outermost shell,determine chemical properties 

Main group elements



are the s and p,

 - valance electrons

Same as the group number



Transition elements

- valance electrons are the

and the

(n-1) d

orbitals.

Valence Electrons

 Lewis electron dot diagrams   

The element’s symbol represents the atomic nucleus and the core electrons.

Up to 4 valence electrons, dots, are placed around the symbol Any remaining electrons are paired with those 4

 A full set of 8 electrons is an

octet

and stable

Lewis electron dot diagrams

Covalent bonding involves sharing of valence electrons Between atoms.

Bonding in Ionic Compounds

Formation of these compounds is exothermic in part ionization Of Na and Ca is low.

Bonding in Ionic Compounds

Lattice Energy

F Cl Br I Li

+ 1036 853 807 757

+ 923 787 747 704

+ 821 715 682 649

+ 785 689 660 630

+ 740 659 631 604 The bond between ions of opposite charge is strongest when the ions are small.

Lattice Energy

The ionic bond should also become stronger as the charge on the ions becomes larger. The data in the table below show that the lattice energies for salts of the OH and O 2 ions increase rapidly as the charge on the ion becomes larger.

OH O 2-

Na +

900 2481

2+ 3006 3791

Al 3

+ 5627 15,916

Lattice Energy

   Why do NaCl 2 or Na 2 Cl not exist?

The energy for Na + to lose a 2nd electron is very  high and endothermic.

So the energy of the reaction would not be negative



The reaction would not be product favored

 Cl is not favorable to add a second electron

The amount of energy required is very large

Covalent Bonding and Lewis Structure

 Molecules or polyatomic ions made of nonmetal atoms bonding covalently.

Bonding Pair Shared Pair each 2 shared dots is one stick bond Nonbonding Pair, Unshared Pair, Lone Pair Double Bonds 2 shared pair 4 shared electrons 2 stick bonds Triple Bonds 3 shared pair 6 shared electrons 3 stick bonds

Drawing Lewis Structures

http://chemed.chem.purdue.edu/genchem/topicreview/ bp/ch8/lewis.html

misterguch.brinkster.net/lewisstructures.html

http://pages.towson.edu/ladon/lewis.html

Drawing Lewis Structures

Decide on the Central Atom

the one with the lowest electron affinity (or most bonds)

Determine the total number of valence electrons in the molecule or ion

-add up all electron dots and charges

Draw skeletal structure with single bonds Use remaining electrons to surround terminal atoms with 8 dots (except H) If central atom has fewer than 8 electrons, move lone pair to create multiple bonds and allow all atoms to have 8 electrons.

Drawing Lewis Structures

 Practice drawing structures for NH 4 + NO + CO SO 4 2-



Drawing Lewis Structures

Practice drawing structures for CH 4 C 2 H 4 NH 2 C 2 H 6 C 2 H 2 H 3 O +

Lewis Structures of Acids and their Anions

Isoelectronic Species

Resonance Structure

When two equivalent structures can be drawn for one molecule, the extra bond was believed to resonate (flip) between the two. If that were occurring,t he lengths of the bonds would be different on each side.

This difference in bond length is not observed.

Resonance theory was proposed by Linus Pauling as a composite. The resonance Hybrid will provide two equal bonds with bond lengths

Resonance Structure

Exceptions to the Octet Rule

 Atoms that have fewer than four pair of electrons on a central atom 

Generally Group 2A or 3A

 

Classic example BF 3 Non octet structures

 Atoms that have more than four pair of electrons on a central atom.



Generally Group 4A, 5A, 6A, 7A, or 8

 

(Yes, Group 8 can form bonds) Expanded Octet

Expanded Octet

Only elements of the third or higher periods in the periodic Table form compounds and ions in which an octet is exceeded.

Second period elements have only s and p orbitals Elements win the third and higher periods, the d orbitals in the Outer shell are traditionally included among valence orbitals For the elements.

The extra orbitals provide the elements with an opportunity to Accommodate up to 12 electrons.

Expanded Octet

Molecules with an Odd Number of Electrons

 There is a small group of molecules with an odd number of electrons  Molecules with an odd number of electrons do not obey the Octet  The single electron classifies the molecules as a

free radical



Free radicals

are more reactive than molecules with paired electrons

Charge Distribution in the Bonds &Molecules

 Lewis structures indicate that electrons are evenly distributed along bonds.

 Valence electrons are not distributed among the atoms as evenly as suggested.

 Some atoms may have a slight negative charge or a slight positive charge  Electrons are drawn more strongly to one atom in the bond.



Electron distribution

defines the uneven distribution of electrons.

Charge Distribution in the Covalent Bonds &Molecules

 Electrons in covalent bonds are not necessarily evenly distributed.



Charge distribution

defines the uneven distribution of electrons.

 

Polar diatomic molecules will line up by polarity Determines the site at which reactions occur (like the attachment of an H + ), will the H attach to the O or Cl end of a OCl -

Formal Charges on Atoms

Formal Charges on Atoms

Calculate formal charges for the atoms in NH 4 + CN CO 3 2 SO 3

Bond Polarity and Electronegativity

 Pure covalent equal sharing-few molecules  Polar covalent bonds - electron pair shared unequally molecule acts as a dipole  1930 Linus Pauling determined

Electronegativity



The ability of an atom in a molecule to attract electrons to itself (in a bond) F = 4.0



EN differences: Covalent < 1.67 < Ionic

Bond Polarity and Electronegativity

Assign a positive and negative pole to each of the following pair, decide which is ore polar

B - F Si - O C = O B - Cl P - P C = S

Combining Formal Charge and Bond Polarity

 Electroneutrality Principle electrons in a molecule are distributed in such a way that the charges on the atoms are as close to 0 as possible.



When a negative charge occurs, it is on the most electronegative element.



BF 4 formal charge gives B a (-) but actually the ( ) is distributed over the F

Combining Formal Charge and Bond Polarity

A is more satisfactory

Combining Formal Charge and Bond Polarity

Bond Properties- Bond Order

 The

Order of a Bond

is the number of bonding electron pair shared by two atoms.



Bond order = 1 single

H 2 CH 4 

Bond order = 2 double

H 2 C = C H 2 

Bond Order = 3 triple

H C C H

Bond Properties Bond Order for Resonance Structures

Bond Order = number of shared pair number of bonding sites

Bond Length and Bond Order

Bond Length is the distance Between the nuclei of two Bonded atoms. Single bonds are determined By the size of atoms Double bonds are shorter Triple bonds, shorter still.

Bond Length and Bond Order

Bond Length and Bond Order

Bond Energy: Bond Dissociation Energy

Bonds broken Bonds formed

C=C 610 kJ / mol and the H-H 1046 kJ / mol C-C 346 kJ and 2 C-H 2 x 413 kJ / mol D

H o rxn = 1046 kJ - 1172 kJ = -126 kJ

Molecular Shapes

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Molecular Shapes

 VSEPR valence shell electron-pair repulsion theory  

Lewis dot diagrams can predict the 3-d geometry of a molecule - which determines its properties Bond and lone electron pair in the valence shell of an element repel each other and seek to be as far apart as possible.

 Central Atoms Surrounded Only by Single -Bond Pairs 

Linear - trigonal planar -Tetrahedral



Trigonal-bipyrimadal -octahedral

Molecular Shapes single bond -Pair

Expanded Octet

Effect of Lone Pair on Bond Angles

Relative strengths of repulsion are in the order

Lone pair-lone pair < Lone pair-bond pair < bond pair-pond pair

Central Atoms with More Than Four Valence Electron Pairs

Multiple Bonds and Molecular Geometry

Bonding and Polarity

http://www.bsc2.ehb schweiz2.ch/Chemie/Simulationen%20Chemie/Bindu ng/Bindungstypen%20Animation.htm

Molecular Polarity

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Molecular Polarity Figure 1: A model of ethanol showing the volume that is occupied by its electrons.

The distribution of the electrons in an ethanol molecule is skewed relative to the protons to give a region having a partial negative charge, which is shown in red, and a corresponding region having a partial positive charge, which is shown in blue.

Molecular Polarity

Molecular Shapes and Molecular Polarity

Homework Ch 9

32. Which compound has the most negative energy of ion pair formation? Which has the least negative value? (a) NaCl (b) MgS (c) MgF 2 34. Place the following compounds in order of increasing lattice energy (from least negative to most negative): LiI, LiF, CaO, RbI. 37. Which compound in each of the following pairs should require the higher temperature to melt? (See Study Question 36.) (a) NaCI or RbCI (b) BaO or MgO (c) NaCI or MgS 39. Draw Lewis structures for the following molecules or ions: (a) CS 2 (b) NO 2 (c) BF 4 (d) Cl 2 SO

Homework Ch 9

41. Draw Lewis structures for each of the following molecules: (a) Methanol, CH30H (C is the central atom) (b) Vinyl chloride, H2C = CHCI, the molecule from which PVC plastics are made. Its structure resembles that of C 2 F 4 in Study Question 40. (c) Acrylonitrile, H 2 C= CHCN, the molecule from which materials such as Orlan are (a) S0 2 made 42. Show all possible resonance structures for each of the following molecules or ions: (a) BrF 5 (b) IF 3 (b) N0 (c)IBr 2 2 (d) BrF (c) SCN 2 + 45. Draw Lewis structures for each of the following molecules or ions:

Homework Ch 9

46. Determine the formal charge on each atom in each of the following molecules or ions: (a) N 2 H 4 (b) P0 4 3 (c) BH 4 (d) NH 2 0H 48. Determine the formal charge on each atom in the following molecules and ions: (a) N0 2 + (a) C-O and C-N (d) B-F andB-I (b) N0 2 (c) NF 3 (d) HN0 an arrow to show the direction of polarity in each bond. (b) P-Br andP-CI (c) B-O and B-S 3 51. For each pair of bonds, indicate the more polar bond and use

Homework Ch 9

53. Urea, (NH 2 ) 2 CO, is used in plastics and fertilizers. It is also the primary nitrogen-containing substance excreted by humans. Draw the molecule with C=O as the center (a) Which bonds in the molecule are polar and which are nonpolar? (b) Which is the most polar bond in the molecule? Which atom is the negative end of the bond dipole? 54. Considering both formal charges and bond polarities, predict on which atom or atoms the negative charge resides in the following anions: (a) BF 4 (b) BH 4 (c) OH (d) CH 3 C0 2 -

Homework Ch 9

57. Compare the electron dot structures of the carbonate (C0 3 and borate (B0 (a) Are these ions isoelectronic? (b) How many resonance structures does each ion have? (c) What are the formal atom charges in each ion? (d) If an H+ ion attaches to C0 3 3 ) ions. 3 2 to form the bicarbonate ion, HC0 (a) CN 3 , does it attach to an 0 atom or to the C atom? 61. Specify the number of bonds for each of the following molecules or ions. Give the bond order for each bond. (b) CH 3 CN (c) S0 3 (d) CH 3 CH=CH 2 2 )

Homework Ch 9

64. Consider the nitrogen-oxygen bond lengths in N0 2 N0 3 + , N0 2 , and . In which ion is the bond predicted to be longest? Which is predicted to be the shortest? Explain briefly. 69. Phosgene, Cl 2 CO, is a highly toxic gas that was used as a weapon in World War I. Using the bond energies of Table 9.9, estimate the enthalpy change for the reaction of carbon monoxide and chlorine to produce phosgene. (Hint: First draw the electron dot structures of the reactants and products so you know the types of bonds involved.) CO (g) + Cl 2(g) -> Cl 2 CO (g)

Homework Ch 9

74. The following molecules or ions all have two oxygen atoms attached to a central atom. Draw the Lewis structure for each one and then describe the electron-pair geometry and the molecular geometry. Comment on similarities and differences in the series. (a) CO 2 (b) N0 2 (c) 0 3 (d) CI0 2 75. The following molecules or ions all have three oxygen atoms attached to a central atom. Draw the Lewis structure for each one and then describe the electron-pair geometry and the molecular geometry. Comment on similarities and differences in the series. (a) C0 3 2 (b) N0 3 (c) S0 3 2 (d) CI0 3 -

Homework Ch 9

78. Give approximate values for the indicated bond angles. (a) O-S-O in S0 2 (b) F-B-F angle in BF 3 (b) (c) Cl- C - Cl angle in Cl CH3CN 2 CO (d) H -C- H (angle 1) and C-C=N (angle 2) in acetonitrile 83. Consider the following molecules: (a) CH 4 (b) NCl 3 (c) BF 3 (d) CS 2 (i) Which compound has bonds with the greatest degree of polarity? (ii) Which compounds in the list are not polar?

Homework Ch 9

85. Which of the following molecules is (are) polar? For each polar molecule indicate the direction of polarity, that is, which is the negative and which is the positive end of the molecule. (a) BeCl 2 (c) CH 3 Cl (b) HBF (d) S0 3 2