Transcript of Chemical Bonding (Predicting Bond Types) Lewis (Electron) Dot Diagrams Binary Molecular Nomenclature...
- Slide 1
- Chemical Bonding (Predicting Bond Types) Lewis (Electron) Dot
Diagrams Binary Molecular Nomenclature Exceptions to the Octet Rule
Coordinate Covalent Bonding Resonance Structures Molecular Shapes
and Polarity Intermolecular Forces of Attraction
- Slide 2
- What is a chemical bond? A chemical bond is a strong attractive
force between atoms or ions in a chemical compound. Back to main
menu
- Slide 3
- Why do elements form chemical bonds? 1. Uncombined elements
have relatively high potential energy. 2. Atoms will gain, lose or
share valence electrons in order to chemically combine with other
atoms. 3. By combining with other atoms, atoms decrease potential
energy and create more stable arrangements. Back to main menu
- Slide 4
- What two factors determine whether or not a chemical bond will
form? 1. the electron configurations of the atoms involved 2. the
attraction the atoms have for electrons Back to main menu
- Slide 5
- How is the type of chemical bond formed between two atoms
determined? The type of chemical bond formed depends upon the
degree to which the valence electrons are shared between the atoms.
Back to main menu
- Slide 6
- Covalent Bonding In a covalent bond, valence electrons are
shared by the atoms. Covalent bonds can be nonpolar or polar. Back
to main menu
- Slide 7
- Nonpolar vs. Polar In a nonpolar covalent bond, electrons are
shared equally. Bonding which occurs between two atoms of the same
element is an example of nonpolar covalent bonding. Examples: H 2,
Br 2, O 2, N 2, Cl 2, I 2, F 2 In a polar covalent bond, electrons
are shared unequally. Back to main menu
- Slide 8
- Ionic Bonding In an ionic bond, valence electrons are
transferred between atoms. One atom gains electrons to form a
negative ion (anion) and the other atom loses electrons to form a
positive ion (cation). Back to main menu
- Slide 9
- Ionic Bonding Which category of elements tends to gain
electrons and form negative ions (anions)? nonmetals Which category
of elements tends to lose electrons and form positive ions
(cations)? metals Back to main menu
- Slide 10
- Using differences in electronegativity to determine bond type
Electronegativity is a measure of an atoms ability to attract
electrons when chemically combining with another element. The
higher the electronegativity value, the stronger the attraction the
atom has for another atoms electrons. The degree to which bonding
between atoms of two elements is ionic or covalent can be estimated
by calculating the difference in the elements electronegativities
(EN). Back to main menu
- Slide 11
- Using differences in electronegativity to determine bond type
Back to main menu Type of Bond EN (Difference in Electronegativity)
Nonpolar Covalent 0.2 Polar Covalent 0.2 to 1.7 Ionic1.7
- Slide 12
- 1 H 2.1 Periodic Table of Electronegativities 2 He - 3 Li 1.0 4
Be 1.5 5 B 2.0 6 C 2.5 7 N 3.0 8 O 3.5 9 F 4.0 10 Ne - 11 Na 0.9 12
Mg 1.2 13 Al 1.5 14 Si 1.8 15 P 2.1 16 S 2.5 17 Cl 3.0 18 Ar - 19 K
0.8 20 Ca 1.0 21 Sc 1.3 22 Ti 1.5 23 V 1.6 24 Cr 1.6 25 Mn 1.5 26
Fe 1.8 27 Co 1.9 28 Ni 1.8 29 Cu 1.9 30 Zn 1.6 31 Ga 1.6 32 Ge 1.8
33 As 2.0 34 Se 2.4 35 Br 2.8 36 Kr 3.0 37 Rb 0.8 38 Sr 1.0 39 Y
1.2 40 Zr 1.4 41 Nb 1.6 42 Mo 1.8 43 Tc 1.9 44 Ru 2.2 45 Rh 2.2 46
Pd 2.2 47 Ag 1.9 48 Cd 1.7 49 In 1.7 50 Sn 1.8 51 Sb 1.9 52 Te 2.1
53 I 2.5 54 Xe 2.6 55 Cs 0.7 56 Ba 0.9 57 La 1.1 72 Hf 1.3 73 Ta
1.4 74 W 1.7 75 Re 1.9 76 Os 2.2 77 Ir 2.2 78 Pt 2.2 79 Au 2.4 80
Hg 1.9 81 Tl 1.8 82 Pb 1.8 83 Bi 1.9 84 Po 2.0 85 At 2.2 86 Rn 2.4
87 Fr 0.7 88 Ra 0.9 89 Ac 1.1 104 Rf - 105 Db - 106 Sg - 107 Bh -
108 Hs - 109 Mt - 110 Uun - 111 Uuu - 112 Uub - 113114 Uuq - 115116
Uuh - 117118 Uuo - Back to main menu
- Slide 13
- Example 1: What type of bond would form between an atom of
nitrogen and an atom of chlorine? a.Nitrogen has an
electronegativity value of 3.0. b.Chlorine has an electronegativity
value of 3.0. c.The difference in the electronegativity values for
nitrogen and chlorine is EN = - = d.Therefore the type of bond
formed would be nonpolar covalent. The electrons would be shared
equally.. Back to main menu 3.0 0.0
- Slide 14
- Example 2: What type of bond would form between an atom of
hydrogen and an atom of chlorine? a.Hydrogen has an
electronegativity value of 2.1. b.Chlorine has an electronegativity
value of 3.0. c.The difference in the electronegativity values for
hydrogen and chlorine is EN = - = d.Therefore the type of bond
formed would be polar covalent. The electrons would be shared
unequally.. Back to main menu 3.02.10.9
- Slide 15
- Dipole A bond formed between atoms which are not shared equally
is called a dipole. a) In the bond formed between hydrogen and
chlorine, the chlorine would form the negative dipole (symbolized
by - ) because it has the higher electronegativity value. b) The
hydrogen would form the positive dipole (symbolized by + ) because
it has the lower electronegativity value. Back to main menu
- Slide 16
- Example 3: What type of bond would form between an atom of
lithium and an atom of chlorine? a.Lithium has an electronegativity
value of 1.0. b.Chlorine has an electronegativity value of 3.0.
c.The difference in the electronegativity values for lithium and
chlorine is EN = - = d.Therefore the type of bond formed would be
ionic. The electrons would be transferred between atoms. Back to
main menu 3.01.02.0
- Slide 17
- Example 3: What type of bond would form between an atom of
lithium and an atom of chlorine? The lithium atom would lose
electrons and form a positive ion, also known as a cation. The
chlorine atom would gain electrons and form a negative ion, also
known as an anion. Back to main menu
- Slide 18
- You Try It 1. Complete the following table.
CompoundElementsElectronegativityENBond Type KF K F O2O2 O O ICl I
Cl 0.8 4.0 3.2 Ionic 3.5 0.0 Nonpolar Covalent 2.5 3.0 0.5 Polar
Covalent Back to main menu
- Slide 19
- You Try It 2. For each of the bonds in question 1 that were
polar covalent, identify the negative dipole ( - ) and the positive
dipole ( + ). ICl Iodine is the positive dipole and chlorine is the
negative dipole. Back to main menu
- Slide 20
- You Try It Nonpolar covalent 3. Elements that exist as two
atoms chemically bonded together are called diatomic elements. The
diatomic elements are hydrogen, bromine, oxygen, nitrogen,
chlorine, iodine, and fluorine. (You need to memorize the diatomic
elements.) What type of chemical bond exists between the diatomic
elements? Back to main menu
- Slide 21
- You Try It 4. Using the three classifications of bonds
discussed, predict the type of bond that is most likely to be
present in compounds made from elements of groups 1 (1A) and 17
(7A). Ionic Back to main menu
- Slide 22
- You Try It 5. Using the three classifications of bonds
discussed, predict the type of bond that is most likely to be
present in compounds made from elements of groups 16 (6A) and 17
(7A). Polar Covalent Back to main menu
- Slide 23
- You Try It 6. Arrange the following chemical bonds in order of
least covalent to most covalent: H-H, H-Cl, H-Br, Li-Cl Li-Cl,
H-Cl, H-Br, H-H Back to main menu
- Slide 24
- Drawing Lewis Dot Diagrams for Atoms The electrons that play
the most important role in determining whether or not a chemical
bond will form are the valence electrons. In a Lewis dot diagram,
dots are placed around the chemical symbol of an element to
illustrate the valence electrons. The chemical symbol represents
the nucleus of the atom. Back to main menu
- Slide 25
- Drawing Lewis Dot Diagrams for Atoms Examples Back to main menu
Group 1Group 2Group 13Group 14Group 15Group 16Group 17Group 18 HHe
LiBeBCNOFNe
- Slide 26
- Drawing Lewis Structures for Covalent Compounds Types of
Covalent Bonds Single Covalent Bond one pair of valence electrons
is shared. Double Covalent Bond - two pairs of valence electrons
are shared. Triple Covalent Bond - three pairs of valence electrons
are shared. Back to main menu
- Slide 27
- Example 1: H 2 Back to main menu H H HH The two hydrogen atoms
will form a single, nonpolar covalent bond.
- Slide 28
- Example 2: O 2 Back to main menu O O OO The two oxygen atoms
will form a double, nonpolar covalent bond.
- Slide 29
- Example 3: N 2 Back to main menu N N NN The two nitrogen atoms
will form a triple, nonpolar covalent bond.
- Slide 30
- Example 4: HCl Back to main menu H Cl H
- Slide 31
- Example 5: NH 3 Back to main menu H H NH
- Slide 32
- You Try It Back to main menu HBrCCl 4 H2OH2O C 2 H 5 ClC2H4C2H4
C2H2C2H2 H2O2H2O2 HCNCO 2
- Slide 33
- Structural Formulas Structural formulas can also be used to
show the arrangement of atoms in molecules. In a structural
formula, dashes are used to represent shared pairs of electrons.
Back to main menu
- Slide 34
- Structural Formulas Example: H 2 Back to main menu H H HH
- Slide 35
- Structural Formulas Example: H 2 S Back to main menu H S S H H
H
- Slide 36
- Structural Formulas Example: N 2 Back to main menu NN NN
- Slide 37
- Binary Molecular Nomenclature Compounds formed when atoms
covalently bond are called molecular compounds. Binary molecular
compounds are generally composed of two nonmetallic elements. When
two nonmetallic elements combine, they often do so in more than one
way. For example carbon can combine with oxygen to form carbon
dioxide, CO 2 and carbon monoxide, CO. Back to main menu
- Slide 38
- Naming Binary Molecular Compounds Prefixes are used to show how
many atoms of each element are present in each molecule of the
compound. Back to main menu mono-1 di-2 tri-3 tetra-4 penta-5
hexa-6 hepta-7 octa- 8 nona- 9 deca-10
- Slide 39
- Naming Binary Molecular Compounds The names of molecular
compounds have this form: (prefix + element name) (prefix + element
root + ide) Back to main menu
- Slide 40
- Naming Binary Molecular Compounds Back to main menu The prefix
mono is usually omitted if there is just a single atom of the first
element. Example: CO 2 is carbon dioxide not monocarbon dioxide. If
the vowel combinations o-o or a-o appear next to each other in the
name, the first of the pair is omitted to simplify the name.
Example: N 2 O is dinitrogen monoxide not dinitrogen
monooxide.
- Slide 41
- You Try It Back to main menu Name the following compounds.
a.CBr 4 b.Cl 2 O 7 c.N 2 O 5 d.BCl 3 e.PCl 5 f.NO a.Carbon
tetrabromide b.Dichlorine heptoxide c.Dinitrogen pentoxide d.Boron
trichloride e.Phosphorus pentachloride f.nitrogen monoxide
- Slide 42
- Writing Formulas for Binary Molecular Compounds To write the
formula for a binary molecular compound you simply write down the
number of atoms of each element indicated by the name. Back to main
menu Example: Carbon tetrachloride CCl 4
- Slide 43
- You Try It Write formulas for the following binary molecular
compounds. Back to main menu a.dinitrogen tetrahydride b.carbon
disulfide c.iodine heptafluoride d.sulfur dioxide N2H4N2H4 CS 2 IF
7 SO 2
- Slide 44
- Writing Formulas for Binary Molecular Compounds A few molecular
compounds have common names that all scientists use in place of
formal names. CH 4 is methane H 2 O is water NH 3 is ammonia You
need to memorize these. Back to main menu
- Slide 45
- Exceptions to the Octet Rule Some molecules are stable even
though the atoms do not all obtain an octet. There are three common
exceptions to the octet rule. Back to main menu
- Slide 46
- Exception #1 In some molecules the central atom has less than
eight valence electrons. This is called an incomplete octet.
Incomplete octets are common in covalent compounds in which the
central atom is beryllium, boron or aluminum. Back to main
menu
- Slide 47
- Exception #1 Example: BeH 2 Back to main menu Beryllium has
only four electrons around it.
- Slide 48
- Exception #2 Molecules almost always have an even number of
electrons, allowing electrons to be paired, but there are some
exception in which there are an odd number of electrons. These
exceptions usually involve nitrogen. Back to main menu
- Slide 49
- Exception #2 Example: NO Back to main menu You will not be
expected to draw exceptions with odd numbers of electrons in this
course.
- Slide 50
- Exception #3 In some molecules the central atom has more than
eight valence electrons. This is called an expanded octet. Some
common central elements that have expanded octets are sulfur,
chlorine, bromine, iodine, xenon, phosphorus, and arsenic. Back to
main menu
- Slide 51
- Exception #3 Example: SF 6 Back to main menu Sulfur has twelve
electrons around it.
- Slide 52
- You Try It Back to main menu BF 3 AsH 5 BeI 2 PCl 5 ClF 5 XeF
4
- Slide 53
- Coordinate Covalent Bonding Objectives 1.Define coordinate
covalent bonding and give an example. Back to main menu
- Slide 54
- Coordinate Covalent Bonding A coordinate covalent bond is
formed when one atom contributes both electrons in a shared pair.
Example: CO Back to main menu
- Slide 55
- Coordinate Covalent Bonding and Polyatomic Ions Polyatomic ions
form coordinate covalent bonds. A polyatomic ion is covalently
bonded within itself, but is ionically bonded to another atom or
polyatomic ion to form a neutral compound. Back to main menu
- Slide 56
- Coordinate Covalent Bonding and Polyatomic Ions Example: NH 4 +
Back to main menu Example: SO 4 2-
- Slide 57
- You Try It Back to main menu OH - PO 4 3- ClO 3 -
- Slide 58
- Resonance Structures Objectives 1.Define resonance and draw
resonance structures for molecules. Back to main menu
- Slide 59
- Resonance Structures Resonance occurs when more than one valid
Lewis structure can be written for a particular molecule. The
different Lewis structures possible for a molecule are referred to
as resonance structures. Back to main menu
- Slide 60
- Resonance Structures Lets look at the Lewis structure for the
ozone, O 3, molecule. Another possible structure for the ozone
molecule is as follows: Back to main menu O=O-O: :O-O=O.
- Slide 61
- Resonance Structures Notice that each structure indicates that
the ozone molecule has two types of O-O bonds, one single bond and
one double bond. Back to main menu
- Slide 62
- Resonance Structures Based on what we just learned about bond
length, you would expect the bond lengths between the atoms to be
different. Back to main menu
- Slide 63
- Resonance Structures Scientists, however, have experimentally
determined that the bond lengths between the oxygen atoms are
identical. Back to main menu
- Slide 64
- Resonance Structures No one structure correctly describes the
ozone molecule. Scientists have determined that the structure for
ozone is the average of the two structures. A double-headed arrow
is used to indicate resonance. Back to main menu O=O-O:
:O-O=O.
- Slide 65
- You Try It Back to main menu Resonance structures can often be
written for polyatomic ions. Draw the possible resonance structures
for NO 2 -. O=N-O::O-N=O.
- Slide 66
- Molecular Shapes and Polarity Objectives 1.Define VSEPR and
given a chemical formula of a simple molecule, identify its
geometric shape as linear, trigonal planar, angular, tetrahedral,
trigonal pyramidal, trigonal bypyramidal, or octahedral. 2.Using
the shape of a molecule and electronegativites of its atoms,
determine the polarity of the molecule. Back to main menu
- Slide 67
- Molecular Shapes and Polarity The valence shell electron pair
repulsion (VSEPR) theory can be used to predict the three
dimensional shapes of a molecule. The main idea behind VSEPR theory
is that electron pairs (bonding and nonbonding) will orient
themselves so that repulsions between electron pairs are minimized.
Back to main menu
- Slide 68
- Linear Back to main menu Formula Lewis Structure Drawing of
Model Bond Angle HI H I 180
- Slide 69
- Linear Back to main menu Formula Lewis Structure Drawing of
Model Bond Angle HCN HCN is a 3-atom linear molecule. Which atom is
the central atom in the HCN molecule? How many pairs of nonbonding
electrons on the central atom of the HCN molecule? Carbon None 180
H C N
- Slide 70
- Bent (also called angular) Back to main menu Formula Lewis
Structure Drawing of Model Bond Angle H2OH2O 104.5 Which atom of
the water molecule is the central atom? How many pairs of
nonbonding electrons on the central atom of the H 2 O molecule?
Oxygen Two
- Slide 71
- How can you differentiate between a linear molecule and a bent
molecule in terms of nonbonding electron pairs on the central atom?
Linear molecules do not have nonbonding electrons on the central
atom. Bent molecules have nonbonding electrons on the central atom.
Back to main menu
- Slide 72
- Trigonal Planar Back to main menu Formula Lewis Structure
Drawing of Model Bond Angle H 2 CO 120 Which atom is the central
atom? How many atoms are bonded to the central atom? Carbon Three
How many nonbonded electrons are there on the central atom?
Zero
- Slide 73
- Trigonal Pyramidal Back to main menu Formula Lewis Structure
Drawing of Model Bond Angle NI 3 107 Which atom is the central
atom? How many atoms are bonded to the central atom? Nitrogen Three
How many nonbonded electrons are there on the central atom?
one
- Slide 74
- How can you differentiate between a trigonal planar molecule
and a trigonal pyramidal molecule in terms of nonbonding electron
pairs on the central atom? Trigonal planar molecules do not have
nonbonding electrons on the central atom. Trigonal pyramidal
molecules have a pair of nonbonding electrons on the central atom.
Back to main menu
- Slide 75
- Tetrahedral Back to main menu Formula Lewis Structure Drawing
of Model Bond Angle CH 4 109.5 Which atom is the central atom? How
many atoms are bonded to the central atom? Carbon Four How many
nonbonded electrons are there on the central atom? zero
- Slide 76
- Trigonal Bipyramidal Back to main menu Formula Lewis Structure
Drawing of Model Bond Angle PH 5 120 90 Which atom is the central
atom? How many atoms are bonded to the central atom? Phosphorus
Five How many nonbonded electrons are there on the central atom?
zero
- Slide 77
- Octahedral Back to main menu Formula Lewis Structure Drawing of
Model Bond Angle SH 6 120 90 Which atom is the central atom? How
many atoms are bonded to the central atom? Sulfur Six How many
nonbonded electrons are there on the central atom? zero
- Slide 78
- Summary of Molecular Shapes Back to main menu NameDrawing of
Shape Number of Atoms Bonded to Central Atom Number of Lone Pairs
of Electrons Bond Angle 2-Atom Linearnot applicable 180 3-Atom
Linear 180 Bent 104.5 Trigonal Planar 120 Trigonal Pyramidal 107
Tetrahedral 109.5 2 0 2 2 3 0 31 40
- Slide 79
- Summary of Molecular Shapes Back to main menu NameDrawing of
Shape Number of Atoms Bonded to Central Atom Number of Lone Pairs
of Electrons Bond Angle Trigonal Bypryamidal 180 Octahedral 180 50
60
- Slide 80
- Determining Molecular Polarity The polarity of each bond, along
with the geometry of the molecule, determines the polarity of the
molecule. A nonpolar molecule has an even distribution of molecular
charge. A polar molecule has an uneven distribution of molecular
charge. Back to main menu
- Slide 81
- Steps in Determining Molecular Polarity First look at the
geometric shape of the molecule. Molecules with nonbonding pairs of
electrons on the central atom are polar. Which two shapes are
always polar? bent and trigonal pyramidal Back to main menu
- Slide 82
- Steps in Determining Molecular Polarity If the molecule does
contain nonbonding pairs of electrons on the central atom, the
polarity is determined by the atoms surrounding the central atom.
If all of the atoms surrounding the central atom are the same, the
molecule is nonpolar. This is because the bond dipoles will cancel
out. If all of the atoms surrounding the central atom are not
alike, the molecule is polar. The bond dipoles will not cancel out.
Back to main menu
- Slide 83
- Steps in Determining Molecular Polarity H-C N: This molecule is
polar. H-Be-H This molecule is nonpolar. Back to main menu
- Slide 84
- You Try It Determine the polarity of each of the following
molecules. Back to main menu a.HI b.H 2 O c.H 2 CO d.NI 3 e.CH 4
polar nonpolar
- Slide 85
- Intermolecular Forces of Attraction Objectives 13.Define van
der Waals forces, dipole- dipole forces, hydrogen bonds, and London
forces. 14.Given a molecule, identify the dominant type of
intermolecular force of attraction. 15.Given chemical formulas for
two substances, identify which type of intermolecular forces they
exhibit and compare their boiling and freezing points. Back to main
menu
- Slide 86
- Intramolecular vs. Intermolecular Intramolecular forces forces
within a molecule that hold atoms together, that is, covalent
bonds. Intermolecular forces forces between molecules that hold
molecules to each other. These intermolecular forces are
collectively referred to as Van der Waals Forces. They are much
weaker than covalent bonds. Back to main menu
- Slide 87
- Importance of Intermolecular Forces The strength of the
intermolecular forces can be used to determine whether a covalent
compound exists as a solid, liquid, or gas under standard
conditions. Solids have the strongest intermolecular forces of
attraction between their particles. The intermolecular forces of
attraction between the molecules of liquids are not as strong as
those found between the particles of a solid. Gases have the
weakest intermolecular forces of attraction between their
particles. Back to main menu
- Slide 88
- Importance of Intermolecular Forces The strength of the
intermolecular forces can also be used to compare melting and
boiling points. The more strongly the molecules are attracted to
each other, the higher the boiling and melting points. Back to main
menu
- Slide 89
- Types of Intermolecular Forces London Dispersion Forces London
dispersion forces exist in all covalent molecules, however; they
are the most noticeable between nonpolar molecules and the
nonbonding atoms of noble gases. Back to main menu
- Slide 90
- Types of Intermolecular Forces London Dispersion Forces London
dispersion forces arise from the motion of valence electrons. From
the probability distributions of orbitals, it is concluded that the
electrons are evenly distributed around the nucleus. However, at
any one instant, the electron cloud may become distorted as the
electrons shift to an unequal distribution. Back to main menu
- Slide 91
- Types of Intermolecular Forces London Dispersion Forces It is
during this instant that a molecule develops a temporary dipole.
This temporary dipole introduces a similar response in neighboring
molecules, thus producing a short-lived attraction between
molecules. In general the larger the electron cloud, the more
likely the molecule is to form temporary dipoles. Back to main
menu
- Slide 92
- Types of Intermolecular Forces London Dispersion Forces London
forces are the weakest type of intermolecular forces of attraction.
Examples: CO 2, H 2, Ar Back to main menu
- Slide 93
- Types of Intermolecular Forces Dipole-Dipole Forces
Dipole-dipole forces of attraction exist between polar molecules.
Polar molecules contain uneven distributions of charge. The
negative dipole of one molecule is attracted to the positive dipole
of another molecule. Back to main menu
- Slide 94
- Example of Dipole-Dipole Forces HCl HCl is a polar molecule.
The hydrogen end of the molecule forms the positive dipole because
it has the lower electronegativity. The chloride end of the
molecule forms the negative dipole because it has the higher
electronegativity. The chloride end of the molecule is attracted to
the hydrogen end of a neighboring molecule. Back to main menu HCl
HCl Dipole-dipole forces ClH + -
- Slide 95
- Types of Intermolecular Forces Dipole-Dipole Forces
Dipole-dipole forces of attraction are stronger than London
dispersion forces. Back to main menu
- Slide 96
- Types of Intermolecular Forces Hydrogen Bonding Hydrogen
Bonding is a special type of dipole-dipole force. Since no
electrons are shared or transferred, hydrogen bonding is not a
chemical bond. Hydrogen bonding exists between where the very
electronegative elements of nitrogen, oxygen and fluorine are
covalently bonded to hydrogen. Hydrogen bonding occurs between
hydrogen and the unbonded electron pairs of nearby N, O, or F
molecules Back to main menu
- Slide 97
- Examples of Hydrogen Bonding Hydrogen bonding occurs in pure
substances. The hydrogen bonding is represented by a dotted line.
Back to main menu
- Slide 98
- Examples of Hydrogen Bonding Hydrogen bonding can also occur in
mixtures. Back to main menu
- Slide 99
- Examples of Hydrogen Bonding Hydrogen bonding occurs in pure
substances. The hydrogen bonding is represented by a dotted line.
Back to main menu
- Slide 100
- Types of Intermolecular Forces Hydrogen Bonding Hydrogen
bonding is about ten times stronger than ordinary dipole-dipole
forces. Back to main menu
- Slide 101
- Identifying the Types of Intermolecular Forces of Attractions
The chart below can help you identify the types of intermolecular
forces of attraction exhibited by a substance. Reminder: London
Dispersion Forces are exhibited by all covalent molecules. Back to
main menu
- Slide 102
- You Try It 1.List the intermolecular forces of attraction in
order of increasing strength. London dispersion forces,
dipole-dipole forces, hydrogen bonding Back to main menu
- Slide 103
- You Try It 2.What type of intermolecular forces of attraction
would be exhibited by each of the following substances? Justify
your answer. The first one has been done for you. (Hint: Draw the
Lewis Structure for the molecule in order to help you determine the
polarity of the molecule.) a.NH 3 Back to main menu London
dispersion forces, dipole-dipole, hydrogen bonding. NH 3 exhibits
London dispersion forces because all covalent molecules exhibit
London dispersion forces. NH 3 exhibits dipole-dipole forces
because its a polar molecule. NH 3 exhibits hydrogen bonding
because its a polar molecule in which hydrogen is bonded to a
nitrogen, oxygen, or fluorine atom. In this case, hydrogen is
bonded to nitrogen.
- Slide 104
- You Try It 2.What type of intermolecular forces of attraction
would be exhibited by each of the following substances? Justify
your answer. The first one has been done for you. (Hint: Draw the
Lewis Structure for the molecule in order to help you determine the
polarity of the molecule.) b.CO 2 Back to main menu London
dispersion forces only CO 2 is a nonpolar molecule. Nonpolar
molecules only exhibit London dispersion forces.
- Slide 105
- You Try It 3.What type of intermolecular forces of attraction
would be exhibited by each of the following substances? Justify
your answer. The first one has been done for you. (Hint: Draw the
Lewis Structure for the molecule in order to help you determine the
polarity of the molecule.) c.HI Back to main menu London dispersion
forces and dipole-dipole forces HI exhibits London dispersion
forces because all covalent molecules exhibit London dispersion
forces. HI also exhibits dipole-dipole forces because its a polar
molecule. It does not exhibit hydrogen bonding because since H is
not bonded to O, N or F.
- Slide 106
- You Try It 2.What type of intermolecular forces of attraction
would be exhibited by each of the following substances? Justify
your answer. The first one has been done for you. (Hint: Draw the
Lewis Structure for the molecule in order to help you determine the
polarity of the molecule.) d.NH 3 Back to main menu London
dispersion forces only BeH 2 is a nonpolar molecule. Nonpolar
molecules only exhibit London dispersion forces.
- Slide 107
- Comparing Boiling Points Two factors that affect boiling point
are the mass of the compound (molar mass) and the strength of the
intermolecular forces of attraction. The stronger the
intermolecular forces of attraction the higher the boiling point.
Back to main menu
- Slide 108
- Comparing Boiling Points Examine the table below. Back to main
menu Boiling Points of Halogens NameFormula Physical State at Room
Temperature Molar Mass (g/mol) Boiling Point (K, at 1 atm)
fluorineF2F2 gas 38.085.0 chlorineCl 2 gas 70.9239.1 bromineBr 2
liquid 159.8331.9 iodineI2I2 solid 253.8457.4 1.What relationship
exists between the mass of the halogens and the boiling point? The
larger the molar mass, the higher the boiling point.
- Slide 109
- Comparing Boiling Points Examine the table below. Back to main
menu Boiling Points of Halogens NameFormula Physical State at Room
Temperature Molar Mass (g/mol) Boiling Point (K, at 1 atm)
fluorineF2F2 gas 38.085.0 chlorineCl 2 gas 70.9239.1 bromineBr 2
liquid 159.8331.9 iodineI2I2 solid 253.8457.4 2.Arrange the
halogens in order of increasing intermolecular strength of
attraction. Justify your answer. F 2, Cl 2, Br 2, I 2 The stronger
the intermolecular forces of attraction, the greater the boiling
points.
- Slide 110
- Comparing Boiling Points 3.The graph below is a plot of the
boiling points of the hydrogen compounds in the groups headed by
fluorine (HF, HCl, HBr, and HI), oxygen (H 2 O, H 2 S, H 2 Se, H 2
Te), nitrogen (NH 3, PH 3, AsH 3, SbH 3 ), and carbon (CH 4, SiH 4,
GeH 4, SnH 4 ). Use the graph below to answer the following
questions. Back to main menu a.Which group of elements has the
lowest boiling points for each period? Why do they have the lowest
boiling points for each period? The group headed by carbon has the
lowest boiling points for each period. They are all nonpolar
molecules. Nonpolar molecules exhibit weaker London dispersion
forces.
- Slide 111
- Comparing Boiling Points 3.The graph below is a plot of the
boiling points of the hydrogen compounds in the groups headed by
fluorine (HF, HCl, HBr, and HI), oxygen (H 2 O, H 2 S, H 2 Se, H 2
Te), nitrogen (NH 3, PH 3, AsH 3, SbH 3 ), and carbon (CH 4, SiH 4,
GeH 4, SnH 4 ). Use the graph below to answer the following
questions. Back to main menu b.Notice in each of the other three
groups that the first compound (H 2 O, NH 3, and HF) in each group
has a significantly higher boiling point than the other elements in
their groups. What accounts for this phenomenon? H 2 O, NH 3, and
HF all exhibit hydrogen bonding. The other substances in the groups
exhibit dipole-dipole forces of attraction which are not as strong
as hydrogen bonding. Since H 2 O, NH 3, and HF all exhibit hydrogen
bonding they have higher than expected boiling points.
- Slide 112
- Comparing Boiling Points 3.The graph below is a plot of the
boiling points of the hydrogen compounds in the groups headed by
fluorine (HF, HCl, HBr, and HI), oxygen (H 2 O, H 2 S, H 2 Se, H 2
Te), nitrogen (NH 3, PH 3, AsH 3, SbH 3 ), and carbon (CH 4, SiH 4,
GeH 4, SnH 4 ). Use the graph below to answer the following
questions. Back to main menu c.With the exception of H 2 O, NH 3,
and HF, why do the boiling points generally increase within a
group? The boiling points increase because the molar mass of the
compounds increases.
- Slide 113
- You Try It 1.Determine whether each of the following would more
likely be formed by polar or nonpolar molecules. a.a solid at room
temperature b.a liquid with a high boiling point c.a gas at room
temperature d.a liquid with a low-boiling point Back to main menu
polar nonpolar
- Slide 114
- You Try It 2.Considering what you have learned about forces
between atoms and molecules, why do you think all of the elements
in group 18 exist as gases at room temperature? Back to main menu
The noble gases exhibit London dispersion forces. London dispersion
forces are the weakest of the intermolecular forces of attraction.
Substances with weak intermolecular forces of attraction tend to
have lower boiling points.
- Slide 115
- You Try It 3. Arrange the following according to increasing
boiling point: H 2 O, H 2 S, CO 2. Justify your ranking. Back to
main menu CO 2 < H 2 S < H 2 O CO 2 has only London
dispersion forces. H 2 S has dipole-dipole forces. H 2 O has
hydrogen bonding.
- Slide 116
- You Try It 4. Arrange the following according to increasing
boiling point: CH 4, CI 4, CF 4. Justify your ranking. Back to main
menu CH 4 < CF 4 < CI 4 All three molecules are nonpolar and
thus only have London dispersion forces between them. The bigger
the molecule, the more electrons and thus the larger the temporary
dipole. The larger the temporary dipole, the stronger the
intermolecular force and thus the higher the melting point.
- Slide 117
- You Try It 5.NH 3 is a gas at room temperature and H 2 O is a
liquid at room temperature. However, they both exhibit hydrogen
bonding. What does that tell you about the strength of the hydrogen
bonding in H 2 O as compared to NH 3 ? Back to main menu The
hydrogen bonding in H 2 O is stronger than the hydrogen bonding
which occurs in NH 3. H 2 O has two H atoms that can potentially
form four hydrogen bonds with surround water molecules. There are
exactly the right number of hydrogens and lone so pairs that every
one of them can be involved in hydrogen bonding. In the case of
ammonia, the amount of hydrogen bonding is limited by the fact that
each nitrogen has only one lone pair.