AP Chemistry 10504

1
AP Chemistry 1/05/04

• Welcome Back! Happy New Year
• New Grades
• Grading Rubric
• Website is UP!!!!!!
  (drschrempp.com)
• Final Exam

2
Grading Rubric (labs, etc)
3
Grading, continued

• Expectations are that you will provide
  college-level work
• Expectations are that quality of work will
  steadily improve from now through 2nd semester
• Expectations are that most of your work is
  completed outside of the classroom

4
Grading, continued (continued)
If you feel there is a discrepancy in your grade,
come see me. A few points I grade a lot of
papers (170 students). I make mistakes. You
dont need to come barter for just a point or
two, especially on regular and lab assignments.
This wont really affect your grade I usually
use lower scores to point out areas for
improvement Tests are the most important scores
with the greatest impact on your grade. STUDY
FOR THEM!
5
AP Chemistry Website
http//drschrempp.com
6

• Chemical Bonding
• The chapters covered are
• Chapter 8 Basic Concepts
• Chapter 9 Molecular Geometry

7

• Structure of Matter (20)
• B. Chemical bonding
• 1. Binding forces
• a. Types ionic, covalent, metallic, hydrogen
  bonding, van der Waals (including London
  dispersion forces)
• b. Relationships to states, structure, and
  properties of matter
• c. Polarity of bonds, electronegativities
• 2. Molecular models
• a. Lewis structures
• b. Valence bond hybridization of orbitals,
  resonance, sigma and pi bonds
• c. VSEPR
• 3. Geometry of molecules and ions, structural
  isomerism of simple organic molecules and
  coordination complexes dipole moments of
  molecules relation of properties to structure

8
8.1 Chemical Bonds
3 General Types of Chemical Bonds Ionic Bond
Electrostatic forced existing between ions of
opposite charge Covalent bond results form the
sharing of electrons between two atoms Metallic
Bonds each atom is bonded to several neighboring
atoms
9
Covalent Bonds Most familiar examples are in the
interaction of nonmetallic elements with one
another
10
Metallic Bonds Bonding electrons are relatively
free to move throughout the 3-dimensional
structure of the metal, and give rise to metallic
properties such as high conductivity and luster
11
Lewis Symbols

• The electrons involved in chemical bonding are
  the valence electrons.
• G.N. Lewis (1875-1946) suggested illustrating
  this by using the chemical symbol of the element
  plus a dot for each valence electron.
• Example Sulfur
• Electron configuration Ne3s23p4
• S

.
.
.
.
.
.
12
Electron-dot (Lewis) Structures for Selected Main
Group Elements
13

• Atoms often gain, lose, or share electrons to
  achieve the same number of electrons as the noble
  gas closest to them in the periodic table.
• Nobel gases have very stable electron
  arrangements (shown by high ionization energies,
  low electron affinity, and lack of chemical
  reactivity
• Octet Rule atoms tend to gain, lose, or share
  electrons until they are surrounded by eight
  valence electrons

14
Exceptions to the Octet Rule

• The octet rule is NOT always obeyed.
• Molecules such as ClO2, NO, and NO2 have an odd
  number of electrons and therefore each atom
  cannot have a octet.
• Compounds of boron and beryllium sometimes have
  fewer than eight electrons.
• It is rather common for atoms to have an expanded
  octet, more than eight electrons.
• Example PCl5

15
8.2 Ionic Bonding Lewis Structures Multiple
Bonds
16
The electron dot (Lewis) of Na and Cl in the
process of transferring electrons
This process is very exothermic ?Hfo -410.9 kJ
17

• 8.2 Ionic Bonding
• The formation of Na from Na and Cl- from Cl
  indicates that an electron has been lost by Na
  and one gained by Cl.
• This is a typical example of an ionic compound,
  as it is composed of a metal with low ionization
  energy and a nonmetal with high electron affinity
• Lewis-dot symbol

18
Energetics of Ionic Bond Formation
Enthalpy A thermodynamic function accounting for
heat flow in chemical changes ?H E PV
Enthalpy
Internal energy
Product of pressure and volume of the system
19
Enthalpy change, or heat of formation ?Hfo Remem
ber, a -?Hfo signifies that the reaction has lost
energy, making it exothermic
20

• Enthalpy change, or heat of formation for ionic
  bonds
• Remember from Chapter 5 that loss of an electron
  from an atom is an endothermic process (removing
  an electron from a Na atom requires 496 kJ/mol)
• Remember also that gaining an electron by a
  nonmetal is an exothermic process (adding an e-
  to Cl releases 349 kJ/mol)
• If this were all that was involved, most ionic
  bonding processes would be endothermic (496 349
  147kJ/mol) however, this assumes the Na and
  Cl atoms are infinitely far apart

21

• The main reason ionic compounds are stable is the
  attraction between ions of unlike charge. This
  draws the ions together, releasing energy and
  causing the ions to form a solid lattice. This
  stabilization is measured by
• Lattice energy the energy required to
  completely separate a mole of a solid ionic
  compound into its gaseous ions

22

• The magnitude of the lattice energy of a solid
  depends on the charges of the ions, their sizes,
  and their arrangement in the solid

Potential energy of 2 interacting charged
particles
Charges on the particles
Q1Q2
Eel
K
d
Constant 8.99 X 109 J-m/C2
Distance between their centers
Attraction between ions increases as the
magnitudes of their charges increase and as the
distance between them decreases
23
Ionic Bonding
Consider the reaction between sodium and
chlorine Na(s) ½Cl2(g) ? NaCl(s) DHºf -410.9
kJ
24
Ionic Bonding

• The reaction is violently exothermic.
• We infer that the NaCl is more stable than its
  constituent elements. Why?
• Na has lost an electron to become Na and
  chlorine has gained the electron to become Cl-.
  Note Na has an Ne electron configuration and
  Cl- has an Ar configuration.
• That is, both Na and Cl- have an octet of
  electrons surrounding the central ion.

25
Ionic Bonding
26
Ionic Bonding

• Energetics of Ionic Bond Formation
• The formation of Na(g) and Cl-(g) from Na(g) and
  Cl(g) is endothermic.
• Why is the formation of Na(s) exothermic?
• The reaction NaCl(s) ? Na(g) Cl-(g) is
  endothermic (?H 788 kJ/mol).
• The formation of a crystal lattice from the ions
  in the gas phase is exothermic
• Na(g) Cl-(g) ? NaCl(s) ?H -788 kJ/mol

27
Ionic Bonding

• Energetics of Ionic Bond Formation
• Lattice energy the energy required to completely
  separate an ionic solid into its gaseous ions.
• Lattice energy depends on the charges on the ions
  and the sizes of the ions
• k is a constant (8.99 x 10 9 Jm/C2), Q1 and Q2
  are the charges on the ions, and d is the
  distance between ions.

28
Ionic Bonding

• Energetics of Ionic Bond Formation
• Lattice energy increases as
• The charges on the ions increase
• The distance between the ions decreases.

29
(No Transcript)
30
Ionic Bonding

• Electron Configurations of Ions of the
  Representative Elements
• These are derived from the electron configuration
  of elements with the required number of electrons
  added or removed from the most accessible
  orbital.
• Electron configurations can predict stable ion
  formation
• Mg Ne3s2
• Mg Ne3s1 not stable
• Mg2 Ne stable
• Cl Ne3s23p5
• Cl- Ne3s23p6 Ar stable

31
Ionic Bonding

• Transition Metal Ions
• Lattice energies compensate for the loss of up to
  three electrons.
• In general, electrons are removed from orbitals
  in order of decreasing n (i.e. electrons are
  removed from 4s before the 3d).
• Polyatomic Ions
• Polyatomic ions are formed when there is an
  overall charge on a compound containing covalent
  bonds. E.g. SO42-, NO3-.

32
Covalent Bonding

• When two similar atoms bond, none of them wants
  to lose or gain an electron to form an octet.
• When similar atoms bond, they share pairs of
  electrons to each obtain an octet.
• Each pair of shared electrons constitutes one
  chemical bond.
• Example H H ? H2 has electrons on a line
  connecting the two H nuclei.

33
Covalent Bonding
34
Covalent Bonding

• Lewis Structures
• Covalent bonds can be represented by the Lewis
  symbols of the elements
• In Lewis structures, each pair of electrons in a
  bond is represented by a single line

35
AP Chemistry 1/07/04

• Daily Homework AP Practice Exam Question, read
  sections 8.5, 8.6
• Rubric for grading homework (5 points)
• Sections 8.3 8.4 Notes

36
Grading Rubric - homework
37
Covalent Bonding

• Multiple Bonds
• It is possible for more than one pair of
  electrons to be shared between two atoms
  (multiple bonds)
• One shared pair of electrons single bond (e.g.
  H2)
• Two shared pairs of electrons double bond (e.g.
  O2)
• Three shared pairs of electrons triple bond
  (e.g. N2).
• Generally, bond distances decrease as we move
  from single through double to triple bonds.

38
Bond Polarity and Electronegativity

• In a covalent bond, electrons are shared.
• Sharing of electrons to form a covalent bond does
  not imply equal sharing of those electrons.
• There are some covalent bonds in which the
  electrons are located closer to one atom than the
  other.
• Unequal sharing of electrons results in polar
  bonds.

39
Bond Polarity and Electronegativity

• Bond polarity helps describe the sharing of
  electrons between atoms.
• Nonpolar covalent bond one in which the
  electrons are shared equally between two atoms
• Polar covalent bond one of the atoms exerts a
  greater attraction for the bonding electrons than
  the other. If the difference in ability to
  attract electrons is large enough, an ionic bond
  is formed.

40
Bond Polarity and Electronegativity

• Electronegativity
• Electronegativity The ability of one atoms in a
  molecule to attract electrons to itself. i.e. the
  greater the electronegativity, the greater the
  ability to attract electrons to itself

41
Bond Polarity and Electronegativity
Ionization Energy Measures how strongly an atom
holds on to its electrons (section 7.4) Electron
affinity A measure of how strongly an atom
attracts additional electrons (section 7.5) (the
greater the attraction between a given atom and
an added electron, the more negative the atoms
electron affinity will be the greater the
affinity)
42
Bond Polarity and Electronegativity
Electron affinities
43
Bond Polarity and Electronegativity

• Electronegativity
• Electronegativity The ability of one atoms in a
  molecule to attract electrons to itself. i.e. the
  greater the electronegativity, the greater the
  ability to attract electrons to itself
• Pauling set electronegativities on a scale from
  0.7 (Cs) to 4.0 (F).
• Electronegativity increases
• across a period and
• down a group.

44
Bond Polarity and Electronegativity
Electronegativity
45
Bond Polarity and Electronegativity

• Electronegativity and Bond Polarity
• Difference in electronegativity is a gauge of
  bond polarity
• electronegativity differences around 0 result in
  non-polar covalent bonds (equal or almost equal
  sharing of electrons)
• electronegativity differences around 2 result in
  polar covalent bonds (unequal sharing of
  electrons)
• electronegativity differences around 3 result in
  ionic bonds (transfer of electrons).

46
Bond Polarity and Electronegativity

• Electronegativity and Bond Polarity
• There is no sharp distinction between bonding
  types.
• The positive end (or pole) in a polar bond is
  represented ? and the negative pole ?-.

47
AP Chemistry 1/08/04

• Daily Homework AP Practice Exam Question
  read
  sections 8.7, 8.8
• Chapter 8 Exam on Tuesday
• Sections 8.5 8.6 Notes

48
Homework Problem

• All the isotopes have 34 protons but a different
  number of neutrons in the nucleus
• 1s22s22p63s23p64s23d104p4
• 2 unpaired electrons 4S ?? 4p ?? ? ?
• Hunds rule indicates that each of the orbitals
  will be filled with with a single electron before
  it gets paired

49
Homework Problem

• In Se, the single paired 4p electrons have 1
  electron easily removed to create the 3 unpaired
  4p orbitals which is energetically favorable in
  bromine the romoval of 1 electron still leaves a
  paired 4p orbital
• The shielding effect is stronger in Te and makes
  it easier to remove an electron (lower ionization
  energy)

50
Homework Problem
(d) Because F is very electronegative and the
molecule is asymmetric with respect to the
fluorines, this molecule is polar
51
Bond Polarity and Electronegativity

• Dipole Moments
• Consider HF
• The difference in electronegativity leads to a
  polar bond.
• There is more electron density on F than on H.
• Since there are two different ends of the
  molecule, we call HF a dipole.
• Dipole moment, m, is the magnitude of the dipole
• where Q is the magnitude of the charges.
• Dipole moments are measured in debyes, D.

52
Bond Polarity and Electronegativity

• Bond Types and Nomenclature
• In general, the least electronegative element is
  named first.
• The name of the more electronegative element ends
  in ide.
• Ionic compounds are named according to their
  ions, including the charge on the cation if it is
  variable.
• Molecular compounds are named with prefixes.

53
(No Transcript)
54
Bond Polarity and Electronegativity
Bond Types and Nomenclature
55
Drawing Lewis Structures

• Add the valence electrons.
• Write symbols for the atoms and show which atoms
  are connected to which.
• Complete the octet for the central atom the
  complete the octets of the other atoms.
• Place leftover electrons on the central atom.
• If there are not enough electrons to give the
  central atom an octet, try multiple bonds.

56
Drawing Lewis Structures

• Formal Charge
• It is possible to draw more than one Lewis
  structure with the octet rule obeyed for all the
  atoms.
• To determine which structure is most reasonable,
  we use formal charge.
• Formal charge is the charge on an atom that it
  would have if all the atoms had the same
  electronegativity.

57
Drawing Lewis Structures

• Formal Charge
• To calculate formal charge
• All nonbonding electrons are assigned to the atom
  on which they are found.
• Half the bonding electrons are assigned to each
  atom in a bond.
• Formal charge is
• valence electrons - number of bonds - lone pair
  electrons

58
Drawing Lewis Structures

• Formal Charge
• Consider
• For C
• There are 4 valence electrons (from periodic
  table).
• In the Lewis structure there are 2 nonbonding
  electrons and 3 from the triple bond. There are
  5 electrons from the Lewis structure.
• Formal charge 4 - 5 -1.

59
Drawing Lewis Structures

• Formal Charge
• Consider
• For N
• There are 5 valence electrons.
• In the Lewis structure there are 2 nonbonding
  electrons and 3 from the triple bond. There are
  5 electrons from the Lewis structure.
• Formal charge 5 - 5 0.
• We write

60
Drawing Lewis Structures

• Formal Charge
• The most stable structure has
• the lowest formal charge on each atom,
• the most negative formal charge on the most
  electronegative atoms.
• Resonance Structures
• Some molecules are not well described by Lewis
  Structures.
• Typically, structures with multiple bonds can
  have similar structures with the multiple bonds
  between different pairs of atoms

61
Drawing Lewis Structures

• Resonance Structures
• Example experimentally, ozone has two identical
  bonds whereas the Lewis Structure requires one
  single (longer) and one double bond (shorter).

62
Drawing Lewis Structures
Resonance Structures
63
Drawing Lewis Structures

• Resonance Structures
• Resonance structures are attempts to represent a
  real structure that is a mix between several
  extreme possibilities.

64
Drawing Lewis Structures

• Resonance Structures
• Example in ozone the extreme possibilities have
  one double and one single bond. The resonance
  structure has two identical bonds of intermediate
  character.
• Common examples O3, NO3-, SO42-, NO2, and
  benzene.

65
Drawing Lewis Structures

• Resonance in Benzene
• Benzene consists of 6 carbon atoms in a hexagon.
  Each C atom is attached to two other C atoms and
  one hydrogen atom.
• There are alternating double and single bonds
  between the C atoms.
• Experimentally, the C-C bonds in benzene are all
  the same length.
• Experimentally, benzene is planar.

66
Drawing Lewis Structures

• Resonance in Benzene
• We write resonance structures for benzene in
  which there are single bonds between each pair of
  C atoms and the 6 additional electrons are
  delocalized over the entire ring
• Benzene belongs to a category of organic
  molecules called aromatic compounds (due to their
  odor).

67
Exceptions to the Octet Rule

• There are three classes of exceptions to the
  octet rule
• Molecules with an odd number of electrons
• Molecules in which one atom has less than an
  octet
• Molecules in which one atom has more than an
  octet.
• Odd Number of Electrons
• Few examples. Generally molecules such as ClO2,
  NO, and NO2 have an odd number of electrons.

68
Exceptions to the Octet Rule

• Less than an Octet
• Relatively rare.
• Molecules with less than an octet are typical for
  compounds of Groups 1A, 2A, and 3A.
• Most typical example is BF3.
• Formal charges indicate that the Lewis structure
  with an incomplete octet is more important than
  the ones with double bonds.

69
Exceptions to the Octet Rule

• More than an Octet
• This is the largest class of exceptions.
• Atoms from the 3rd period onwards can accommodate
  more than an octet.
• Beyond the third period, the d-orbitals are low
  enough in energy to participate in bonding and
  accept the extra electron density.

70
Strengths of Covalent Bonds

• The energy required to break a covalent bond is
  called the bond dissociation enthalpy, D. That
  is, for the Cl2 molecule, D(Cl-Cl) is given by ?H
  for the reaction
• Cl2(g) ? 2Cl(g).
• When more than one bond is broken
• CH4(g) ? C(g) 4H(g) ?H 1660 kJ
• the bond enthalpy is a fraction of ?H for the
  atomization reaction
• D(C-H) ¼?H ¼(1660 kJ) 415 kJ.
• Bond enthalpies can either be positive or
  negative.

71
(No Transcript)
72
Strengths of Covalent Bonds

• Bond Enthalpies and the Enthalpies of Reactions
• We can use bond enthalpies to calculate the
  enthalpy for a chemical reaction.
• We recognize that in any chemical reaction bonds
  need to be broken and then new bonds get formed.
• The enthalpy of the reaction is given by the sum
  of bond enthalpies for bonds broken less the sum
  of bond enthalpies for bonds formed.

73
Strengths of Covalent Bonds

• Bond Enthalpies and the Enthalpies of Reactions
• Mathematically, if ?Hrxn is the enthalpy for a
  reaction, then
• We illustrate the concept with the reaction
  between methane, CH4, and chlorine
• CH4(g) Cl2(g) ? CH3Cl(g) HCl(g) ?Hrxn ?

74
Strengths of Covalent Bonds
75
Strengths of Covalent Bonds

• Bond Enthalpies and the Enthalpies of Reactions
• In this reaction one C-H bond and one Cl-Cl bond
  gets broken while one C-Cl bond and one H-Cl bond
  gets formed.
• The overall reaction is exothermic which means
  than the bonds formed are stronger than the bonds
  broken.
• The above result is consistent with Hesss law.

76
Strengths of Covalent Bonds

• Bond Enthalpy and Bond Length
• We know that multiple bonds are shorter than
  single bonds.
• We can show that multiple bonds are stronger than
  single bonds.
• As the number of bonds between atoms increases,
  the atoms are held closer and more tightly
  together.

77
Strengths of Covalent Bonds
78
End of Chapter 8Basic Concepts of Chemical
Bonding
79
Drawing Lewis Structures

• Sum the valence electrons from all atoms.
• For an anion, add an electron for each negative
  charge.
• For a cation, subtract an electron for each
  positive charge.
• Dont worry about keeping track of which
  electrons come from which atoms.
• Write the symbols for the atoms to show which
  atoms are attached to which, and connect them
  with a single bond.
• Chemical formulas are often written in order in
  which the atoms are connected. (ex. HCN)
• When a central atom has a group of other atoms
  bonded to it, the central atom is usually written
  first. (ex. CH4, SF6)
• Sometimes youll need more information or some
  chemical intuition.
• Fill in the electrons by first completing the
  octets of the atoms bonded to the central atom,
  then to the central atom.
• Each bond represents two electrons.
• If there are not enough electrons to give the
  central atom an octet, try multiple bonds.

Example PCl3
80
Resonance

• Some species, such as CO2, SO2, and O3, are
  described by more than one Lewis strucuture.
• In this case, we write both Lewis structures and
  indicate that the real molecule is described by a
  resonance structure which is an average of the
  equivalent Lewis structures.

81
8.4 Bond Polarity Electronegativity
82
Representation of the polar hydrogen chloride
molecule
83
Chlorine hogs the electron blanket,
leaving hydrogen partially, but positively,
exposed
84
Electron Configurations of Ions
85
8.3 Covalent Bonding
86
Polyatomic ions have both covalent bonds (dashes)
and ionic charges
87
VSEPR

• Lewis structures alone tell use nothing about the
  geometry of a molecule.
• BUT
• Molecules will naturally adopt their
  lowest-energy arrangement.
• This is done by minimizing the valence shell
  electron pair repulsion (VSEPR)

88
VSEPR

• Repulsion between two electron pairs is the
  greatest.
• Repulsion between between a lone electron pair
  and a bonding electron pair is less.
• Repulsion between two bonding pairs is the least.

89
VSEPR Example
90
Reactions This topic makes up 30-40 of the AP
Exam The chapters covered are
91
III. Reactions (35-40) A. Reaction types 1.
Acid-base reactions concepts of Arrhenius,
Brønsted-Lowry, and Lewis coordination
complexes amphoterism 2. Precipitation
reactions 3. Oxidation-reduction
reactions a. Oxidation number b. The role
of the electron in oxidation- reduction c.
Electrochemistry electrolytic and galvanic
cells Faraday's laws standard half-cell
potentials Nernst equation prediction of
the direction of redox reactions
92
AP Exam Topics
B. Stoichiometry 1. Ionic and molecular species
present in chemical systems net ionic
equations2. Balancing of equations including
those for redox reactions3. Mass and volume
relations with emphasis on the mole concept,
including empirical formulas and limiting
reactants
93
AP Exam Topics
C. Equilibrium 1. Concept of dynamic equilibrium,
physical and chemical Le Chatelier's principle
equilibrium constants2. Quantitative
treatment a. Equilibrium constants for gaseous
reactions Kp , Kcb. Equilibrium constants for
reactions in solution (1) Constants for acids and
bases pK pH(2) Solubility product constants
and their application to precipitation and the
dissolution of slightly soluble compounds(3)
Common ion effect buffers hydrolysis
94
AP Exam Topics
D. Kinetics 1. Concept of rate of reaction2. Use
of experimental data and graphical analysis to
determine reactant order, rate constants, and
reaction rate laws3. Effect of temperature
change on rates4. Energy of activation the role
of catalysts5. The relationship between the
rate-determining step and a mechanism
95
AP Exam Topics
E. Thermodynamics 1. State functions2. First
law change in enthalpy heat of formation heat
of reaction Hess's law heats of vaporization
and fusion calorimetry3. Second law entropy
free energy of formation free energy of
reaction dependence of change in free energy on
enthalpy and entropy changes4. Relationship of
change in free energy to equilibrium constants
and electrode potentials