Molecular Structure:

1
Chapter 10

• Molecular Structure
• Solids and Liquids

2
Homework

• Assigned Problems (odd numbers only)
• Questions and Problems 10.1 to 10.53 (begins
  on page 292)
• Additional Questions and Problems 10.65 to
  10.91 (page 325)
• Challenge Questions 10.93 to 10.99, (page
  327)

3
Electron Configuration of Ionic Compounds

• Metals Low ionization energy, form ions
  (cations)
• Nonmetals High ionization energy, accept
  electrons (anions)
• The representative elements lose or gain
  electrons to form ions with electron
  configuration of the nearest noble gas

4
Electron Configuration of Ions

• An ion An atom that is electrically charged from
  loss or gain of electrons
• Atoms are neutral due to the equal number of
  protons and electrons
• Loss or gain of an electron will leave a net
  charge on the atom

5
Electron Configuration of Ions

• Consider sodium
• Can attain the noble gas configuration (neon) if
  it loses 1 electron to obtain Na
• Noble gas core remains but it does not have the
  same chemical properties of neon

Na
Na
Loss of 1 e-
1s22s22p6
1s22s22p63s1
Electron configuration of neon
6
Electron Configurations

• We know for Representative Elements
• Group 1A metals form ions with 1 charge
• Group 2A metals form ions with 2 charge
• Group 7A nonmetals form ions with -1 charge
• Group 6A nonmetals form ions with -2 charge
• Group 8A nonmetals do not form ions, in fact they
  are extremely unreactive
• Transition Elements All of the elements of the d
  area of the periodic table
• Elements differ in the number of electrons in the
  d subshell
• Octet rule does not apply here
• Loss of electrons does not lead to the noble gas
  structure

7
Ions of the Metals of Groups IA, IIA, IIIA

• Metals form cations by losing enough electrons to
  get the same electron configuration as the
  previous noble gas
• A stable noble gas core attained in each case

8
Ions of the Nonmetals of Group VIA, VIIA

• Nonmetals form anions by gaining enough electrons
  to get the same electron configuration as the
  next noble gas
• A stable noble gas core attained in each case

9
Electron-Dot Formulas

• A 2-dimensional representation of how atoms are
  covalently bonded together
• Each covalent bond is represented by a pair of
  dots (bonding electrons)
• Must show all unshared pairs of (nonbonding)
  electrons
• All valence e- from every atom in a molecule must
  be accounted for in the form of bonds or
  nonbonding pairs

10
Drawing Electron-Dot Formulas

• Bonding involves ONLY valence electrons
• Sharing of the s and p valence electrons to
  achieve stability of the nearest noble gas
• Illustrates the sequence of atoms
• Shows atoms and their valence electrons
• How they are distributed in a molecule
• Use dot or X to represent an electron

11
Drawing Electron-Dot Formulas

• Determine the arrangement of atoms within a
  molecule
• If there are three or more atoms, the central
  atom (usually) appears only once in the formula
• Halogens are often terminal atoms (at the edges)
  unless it is combined with O as in oxyacids
• Hydrogen is always a terminal atom

12
Drawing Electron-Dot Formulas

• Determine the total number of valence electrons
• For main-group elements the group numbers equal
  the number of valence electrons for the element
  of that group
• If it is an anion, add one electron to the total
  for each negative charge
• If it is a cation, subtract one electron from the
  total for each positive charge

13
Drawing Electron-Dot Formulas

• Give the number of valence electrons for Mg, N,
  and Br. Draw the Lewis dot symbol for each of
  these elements
• Mg 2 valence electrons
• N 5 valence electrons
• Br 7 valence electrons

Mg
N
Br
14
Drawing Electron-Dot Formulas

• Attach the central atom to each bonded atom by a
  pair of electrons
• Subtract two electrons (from total valence) for
  each single bond drawn in the structure

15
Drawing Electron-Dot Formulas

• Distribute the remaining electrons
• Add electrons to each atom bonded to the central
  atom until each has eight electrons (complete
  octets), except hydrogen
• Any extra electrons should go to the central
  atoms

16
Drawing Electron-Dot Formulas

• Duet Rule
• Hydrogen wants two electrons to attain the noble
  gas configuration of helium
• Octet Rule
• All other main group elements want 8 electrons to
  achieve the noble gas configuration
• Filled valence shell is achieved by
  gaining/losing electrons or by sharing electrons

17
Constructing Lewis Structures

• If the central atom does not fulfill the octet
  rule, share one or more lone pairs between a
  terminal atom and the central atom (form multiple
  bonds)
• Double or triple bonds are formed ONLY when one
  or both of the atoms are C, N, O or S

18
Bonding Behavior of Elements

• Oxygen has 6 valence electrons and 2 octet
  vacancies
• Can complete its octet by forming two covalent
  bonds
• Nitrogen has 5 valence electrons and 3 octet
  vacancies
• Can complete its octet by forming three covalent
  bonds

Group 6A
Group 5A
19
Bonding Behavior of Elements

• Carbon has 4 valence electrons and 4 octet
  vacancies
• Can complete its octet by forming four covalent
  bonds
• Fluorine has 7 valence electrons and 1 octet
  vacancy
• Can complete its octet by forming one covalent
  bond

Group 4A
Group 7A
20
Bonding Behavior of Elements
Group 4A
Group 5A
Group 6A
Group 7A
21
Drawing Electron-Dot Formulas

• Create the electron-dot formula for F2
• Start with the atomic symbol for fluorine

• Bond pair Pair of electrons shared by the two
  atoms
• Lone pair Pair of electrons not involved in the
  bonding

22
Drawing Electron-Dot Formulas

• Draw Electron-dot formulas for the following
• NH4
• SO42-
• CO
• SCN-

23
NH4

• Determine the arrangement of the atoms
• N 5 valence electrons
• H 4x1(each) valence electron
• Charge Subtract one electron
• Total valence electrons 8

24
SO42-

• Determine the arrangement of the atoms
• S 6 electrons
• O 4 x 6 (each) electrons
• Charge Add two electrons
• Total electrons 32

25
Multiple Covalent Bonds

• Many molecules exist that need two or three pairs
  of electrons to provide a complete octet of
  electrons per atom
• Multiple covalent bonds Covalent bonds where two
  or three pairs of electrons are shared between
  the same two atoms

26
Bonding

• Single Bond
• Uses a single pair of electrons between two atoms
• Double Bond
• Uses two pairs of electrons between the same two
  atoms
• Triple Bond
• Uses three pairs of electrons between the same
  two atoms

27
CO

• Determine the arrangement of the atoms
• C 4 electrons
• O 6 electrons
• Total valence electrons 10
• If octets are not complete, form one or more
  multiple covalent bonds

28
SCN-

• Determine the arrangement of the atoms
• S 6 electrons
• C 4 electrons
• N 5 electrons
• Charge add one electron
• Total electrons 16
• If octets are not complete, form one or more
  multiple covalent bonds

29
Resonance

• Two or more e-dot structures for a molecule or
  ion that have the same arrangement of atoms
• Contain the same number of electrons
• Differ only in the location of the electrons

30
Resonance

• Occurs whenever it is possible to draw two or
  more electron-dot structures
• They differ only in the location of a double bond
  between the same two types of atoms
• Two of the four electrons in the double bond
  rapidly move between the two single bonds the
  true structure is the average of the individual
  structures

31
Resonance Example I

• Draw resonance structures for the nitrate ion.

32
Resonance Example I

• Draw resonance structures for ozone (O3)
• Ozone has 18 valence electrons

Resonance structures
Hybrid
33
Resonance Example II
34
Shapes of Molecules and Ions

• (Lewis) Electron-Dot structures describe the
  distribution of valence electrons among bonding
  pairs and nonbonding pairs
• They do not give info on the 3-D shape of the
  molecule

35
Shapes of Molecules and Ions

• A Lewis structure for water
• Are the atoms arranged in a straight line or do
  they form a v-shape?

36
Shapes of molecules and ionsVSEPR Theory

• The 3-D shapes of molecules and PA ions result
  from the orientation of atoms about the central
  atom
• VSEPR Valence Shell Electron-Pair Repulsion
  Theory focuses on the bonding and nonbonding
  electrons in the valence shell of the central
  atom
• The central atoms electrons play an important
  role in determining molecular shape

37
Shapes of molecules and ionsVSEPR Theory

• Bonding and nonbonding pairs of electrons have a
  natural electrostatic repulsion that pushes them
  as far apart from one another as possible
• This repulsion of electron groups (regions of
  negative charge) that causes a molecule to have a
  certain shape

38
Shapes of molecules and ionsVSEPR Theory

• Two electron groups
• 2 atoms attached
• Shape Linear
• Bond angle 180 apart
• Three electron groups
• 3 atoms attached
• Shape Trigonal Planar
• Bond angle 120 apart
• Four electron groups
• 4 atoms attached
• Shape Tetrahedral
• Bond angle 109.5 apart

39
Shapes of molecules and ionsVSEPR Theory

• Linear
• 2 atoms on opposite sides of central atom
• 180 bond angles
• Trigonal Planar
• 3 atoms form a triangle around the central atom
• Planar
• 120 bond angles

40
Shapes of molecules and ionsVSEPR Theory

• Tetrahedral
• 4 surrounding atoms form a tetrahedron around the
  central atom
• 109.5 bond angles

41
Shapes of molecules and ionsVSEPR Theory

• 2 atoms attached to the central atom
• Beryllium does not follow the octet rule
• Linear shape
• Bond angles are 180

42
Shapes of molecules and ionsVSEPR Theory

• 3 atoms attached to the central atom
• Boron does not follow the octet rule
• Trigonal planar shape
• Bond angles are 120

43
Shapes of molecules and ionsVSEPR Theory

• 4 atoms attached to the central atom
• Tetrahedral shape
• Bond angles are 109.5

44
VSEPR Theory Central atoms with Bonding pairs
and Lone pairs

• You cant actually see lone pairs, you can only
  see the atoms
• The electron-pair geometry around a central atom
  includes the spatial positions of all bond pairs
  and lone pairs
• Only the arrangement of atoms describes the
  molecular shape of the molecule, not the lone
  pairs of electrons

45
VSEPR Theory Central atoms with Bonding pairs
and Lone pairs

• 3 atoms attached to the central atom
• Four electron groups attached to the central atom
• A tetrahedral electron-pair geometry and a
  trigonal pyramidal molecular shape
• Bond angles are 109.5

46
VSEPR Theory Central atoms with Bonding pairs
and Lone pairs

• Two atoms attached to the central atom
• Four electron groups attached to the central atom
• A tetrahedral electron-pair geometry and a bent
  molecular shape
• Bond angles are 109.5

47
Electronegativity and Polarity

• The ability of an atom in a molecule to attract
  bonding electrons towards itself
• The higher the elements electronegativity, the
  greater its ability to attract electrons
• Fluorine (the reference element) is most
  electronegative, Francium is least electronegative

48
Electronegativity and Polarity

• Increases across period (left to right) on
  periodic table
• Decreases down group (top to bottom) on periodic
  table
• As electronegativity difference increases between
  two elements, bond polarity increases
• Nonmetals have higher electronegativity values
  than metals
• Metals tend to lose electrons and nonmetals tend
  to gain electrons when an ionic bond is formed

49
Electronegativity
50
Polarity of Bonds

• A covalent bond involves pairs of electrons
  shared equally between 2 atoms
• How they are shared (equally or unequally)
  depends on the electron donating and electron
  attracting nature of the atoms

51
Polarity of Bonds

• Nonpolar covalent bond
• Two identical atoms will share the bonding
  electrons equally
• Polar covalent bond
• Two different atoms will not share the bonding
  electrons equally
• One atom will have a greater attraction for the
  shared pair than the other atom

52
Electronegativity

• Increases across period (left to right) on
  Periodic Table
• Decreases down group (top to bottom) on Periodic
  Table
• Larger difference in electronegativities means
  more polar bond
• Negative end toward more electronegative atom

Molecule O-O C-O Mg-O
Electronegativity Values 3.5 and 3.5 2.5 and 3.5 1.2 and 3.5
Electronegativity Difference 0.0 1.0 2.3
Bond Type Pure Covalent Polar Covalent Ionic
53
Polarity of Bonds/Dipole Moments

• Polar covalent bond One of the two different
  elements will inevitably have a greater
  attraction for the shared pair than the other
• This unequal sharing causes the entire molecule
  to behave like an electric dipole
• Dipole A body with two poles, one partially
  negative and one positive

54
Polar Molecules
55
Polarity of Bonds/Dipole Moments

• In a water molecule, oxygen draws the shared pair
  of electrons closer to oxygen and partially
  withdrawn from hydrogen
• Since the dipoles do not cancel, water is a polar
  molecule

56
Attractive Forces in Compounds

• Intermolecular Forces
• Ionic compounds Ions ( and -) are held together
  by ionic bonds
• Neutral molecules One or more of these forces
  hold molecules together in liquids and solids
• Dipole-dipole
• Hydrogen-bonding
• Dispersion

57
Attractive Forces in Compounds

• Dipole-dipole Attractions between polar
  molecules
• The nonsymmetrical distribution of the charge
  causes the molecules to line up
• Positive end of one directed toward negative end
  of other

58
Attractive Forces in Compounds

• Hydrogen bonding A special type of
  dipole-dipole interaction
• Occurs between molecules that have a H atom
  bonded to F, O, or N
• The partially positive H and a lone pair of
  electrons on another N, O, or F atom

59
Attractive Forces in Compounds

• Dispersion (London) forces Short-lived dipoles
  caused by uneven shifts in electron density
• The uneven shift causes one end of the molecule
  to be slightly positive and one end slightly
  negative
• This induces the same electron shift in adjacent
  molecules which causes an attractive force
• Only intermolecular force possible in nonpolar
  substances

60
Matter and Changes of State

• Matter Anything that has mass and occupies space
• It is separated into three categories Solid,
  Liquid, and Gas

61
Matter and Changes of State

• A change in state is the most common type of
  physical change
• Melting/Freezing
• Vaporization/Condensation
• Sublimation/Deposition
• The composition of the substance does not change,
  only its appearance

62
Melting/Freezing

• A change of state requiring the input of heat
• Heat of Fusion
• Heat energy required to melt 1 g of a substance
• Heat energy that must be removed to freeze 1 g of
  a substance
• Heat energy (to melt) 1 g of ice
  (water at 0 C)

63
Change of State Problem

• Calculate the heat needed (in Joules) to melt 15
  g of ice at 0C, and to heat the water to 75 C
• Two parts to the problem
• Melt ice (use heat of fusion for water)
• Heat water (use specific heat of Water)

64
Change of State Problem

• Melt the ice
• Calculate the heat absorbed to melt the ice at 0
  C (no change in temperature)

65
Change of State Problem

• Heat the water
• Calculate the heat energy needed to warm the
  water from 0 C to 75 C

(75 C)
66
Change of State ProblemCombining Energy
Calculations

• Calculate the total heat
• Melting the ice (Q1)
• Heating the water (Q2)

67
Change of State

• Sublimation A phase change from solid to gas
  without going through the liquid state
• Requires the absorption of heat
• No temperature change occurs during process
• Deposition is the reverse process (heat is
  released)
• Evaporation A phase change from a liquid to a
  gas
• Requires the absorption of heat
• No temperature change occurs during process
• Condensation is the reverse process (heat is
  released)

68
Change of State

• Boiling A special form of evaporation where the
  liquid converts to vapor through bubble formation
• Boiling point Temp at which the vapor pressure
  of the liquid is the same as the atmospheric
  pressure
• This allows the bubbles at the liquid surface to
  escape in the atmosphere

69
Vaporization/Condensation

• Heat of Vaporization
• Heat energy required to vaporize 1 g of a
  substance
• Heat energy that must be removed to condense 1 g
  of a substance
• Heat energy (to vaporize) 1 g of water to vapor

70
Change of State Problem

• Calculate the heat needed (in Joules) to heat 15
  g of water from 75 C to 100 C, and to convert
  it to steam at 100 C
• Two parts to the problem
• Heat the water (use specific heat for water)
• Convert water to steam (use heat of vaporization
  for Water)

71
Change of State Problem

• Heat the water
• Calculate the heat energy needed to warm the
  water from 75 C to 100 C

25 C
72
Change of State Problem

• Calculate the heat absorbed to convert the liquid
  water to steam at 100 C (no change in
  temperature)

73
Change of State ProblemCombining Energy
Calculations

• Calculate the total heat
• Heat the water (Q1)
• Convert liquid water to steam (Q2)

74
Heats of Fusion and Vaporization
75
Heating Curve

• Illustrates the steps involved in changing a
  solid to a gas
• Heat added is shown on the x-axis
• Temperature is shown on the y-axis
• Energy required to undergo a series of phase
  changes depends on the (three) property values of
  the substance
• Specific Heat
• Heat of Fusion
• Heat of Vaporization

76
Heating Curve
No temp change
No temp change
77

• end