karen timberlake chapter 5 - pbworks
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1
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc. Chapter 3, Section 3
Electrons and Reactivity
Atoms contain
� a very small nucleus packed with neutrons and
positively charged protons.
� a large volume of space around the nucleus that
contains the negatively charged electrons.
It is the electrons that determine the physical and
chemical properties of atoms.
Elements react in order to achieve the same
electronic structure of the nearest noble gas
2
© 2013 Pearson Education, Inc. Chapter 3, Section 3
Electrons and Reactivity
Atoms contain
� a very small nucleus packed with neutrons and
positively charged protons.
� a large volume of space around the nucleus that
contains the negatively charged electrons.
It is the electrons that determine the physical and
chemical properties of atoms.
Elements react in order to achieve the same
electronic structure of the nearest noble gas
3 © 2013 Pearson Education, Inc. Chapter 5, Section 1 4
Ionic and Covalent Bonds
Atoms form octets
� to become more stable.
� by losing, gaining, or
sharing valence electrons.
� by forming ionic or
covalent bonds.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 5
Covalent Compounds
Covalent compounds are formed when nonmetals
share valence electrons forming a covalent bond.
Covalent compounds consist of molecules, which are
discrete groups of atoms, such as:
� water, H2O 2 H atoms + 1 O atom
� carbon dioxide, CO2 1 C atom + 2 O atoms
� glucose, C6H12O6 6 C atoms + 12 H atoms
+ 6 O atoms
© 2013 Pearson Education, Inc. Chapter 5, Section 1 6
Ionic Compounds
Ionic compounds are formed when metal atoms
transfer electrons to nonmetal atoms resulting in the
formation of ionic bonds.
Ionic compounds we use every day include:
� table salt NaCl
� baking soda NaHCO3
� milk of magnesia Mg(OH)2
2
© 2013 Pearson Education, Inc. Chapter 5, Section 1 7
Ionic Compounds, Minerals
Precious and semiprecious gemstones are also
examples of ionic compounds called minerals.
Sapphires and rubies are made of aluminum oxide
(Al2O3), which is an ionic compound. Their brilliant
colors come from the presence of other metals such as
� chromium, which make rubies red and
� iron and titanium, which make sapphires blue
© 2013 Pearson Education, Inc. Chapter 5, Section 1 8
The octet rule is the tendency for atoms to share or transfer electrons to obtain a stable configuration of 8 valence electrons.
Octet Rule
© 2013 Pearson Education, Inc. Chapter 5, Section 1 9
Octet Rule
The octet rule is associated with the stability of thenoble gases except for He, which is stable with 2valence electrons (duet).
Valence Electrons
He 1s2 2
Ne 1s22s22p6 8
Ar 1s22s22p63s23p6 8
Kr 1s22s22p63s23p64s23d104p6 8
© 2013 Pearson Education, Inc. Chapter 5, Section 1 10
Metals Form Positive Ions
Metals form positive ions called
cations
� by a loss of their valence electrons.
� with the electron configuration of
the nearest noble gas.
� that have fewer electrons than
protons.
Group 1A (1) metals ion 1+
Group 2A (2) metals ion 2+
Group 3A (3) metals ion 3+
© 2013 Pearson Education, Inc. Chapter 5, Section 1 11
Formation of a Sodium Ion, Na+
Sodium achieves an octet by losing its one valence
electron.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 12
Charge of Sodium Ion, Na+
With the loss of its valence
electron, a sodium ion has a 1+
charge.
Na atom Na+
ion
charge
3
© 2013 Pearson Education, Inc. Chapter 5, Section 1 13
Formation of Magnesium Ion, Mg2+
Magnesium achieves an octet by losing its two
valence electrons.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 14
Charge of Magnesium Ion, Mg2+
With the loss of two valence
electrons magnesium forms a
positive ion with a 2+ charge.
Mg atom Mg2+ ion
charge
© 2013 Pearson Education, Inc. Chapter 5, Section 1 15
Formation of Negative Ions
In ionic compounds, nonmetals
� achieve an octet arrangement.
� gain electrons.
� form negatively charged ions called anions with
3–, 2–, or 1– charges.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 16
Formation of a Chloride Ion, Cl–
Chlorine achieves an octet by adding one more
electron to its 7 valence electrons.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 17
Charge of a Chloride Ion, Cl–
By gaining one electron, the
chloride ion has a 1− charge.
charge
© 2013 Pearson Education, Inc. Chapter 5, Section 1 18
Ionic Charge from Group
Numbers
Ions achieve the electron configuration of their
nearest noble gas by forming
� positive ions with a charge equal to its Group
number.
Group 1A (1) = 1+
Group 2A (2) = 2+
Group 3A (3) = 3+
� negative ions with a charge obtained by subtracting
8 or 18 from its Group number.
4
© 2013 Pearson Education, Inc. Chapter 5, Section 1 19
Some Typical Ionic Charges
© 2013 Pearson Education, Inc. Chapter 5, Section 1 20
Group Numbers for Some Positive
and Negative Ions
© 2013 Pearson Education, Inc. Chapter 5, Section 1 21
5.2
Writing names and
formulas for ionic
compounds
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 22
Ionic compounds
� consist of positive and negative ions.
� have attractive forces between the positive and
negative ions called ionic bonds.
� have high melting and boiling points.
� are solid at room temperature.
Ionic Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 23
Salt Is an Ionic Compound
Sodium chloride, or “table salt,” is an example of an
ionic compound.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 24
Formulas of Ionic Compounds
The chemical formula of an ionic compound
represents the element symbols and subscripts, which
represent the lowest whole-number ratio of ions.
In the formula of an ionic compound,
� the sum of positively and negatively charged ions is
always zero Rule of Zero Charge
� the charges are balanced.
5
© 2013 Pearson Education, Inc. Chapter 5, Section 1 25
Charge Balance for NaCl, “Salt”
In NaCl,
� a Na atom loses its valence electron.
� a Cl atom gains an electron.
� the symbol of the metal (sodium) is written first,
followed by the symbol of the nonmetal (chlorine).
© 2013 Pearson Education, Inc. Chapter 5, Section 1 26
Charge Balance In MgCl2
In MgCl2,
� a Mg atom loses two valence electrons.
� two Cl atoms gain one electron each.
� subscripts indicate the number of ions needed to
give charge balance.
� the symbol of the metal (magnesium) is written
first, followed by the symbol of the nonmetal
(chlorine).
© 2013 Pearson Education, Inc. Chapter 5, Section 1 27
Writing Formulas For Binary
Ionic Compounds
� Binary ionic compounds are made from 2
elements, a metal and a nonmetal.
� Charge balance is used to write the formula for
sodium nitride, a compound containing Na+ and
N3−. The metal ion is written first, followed by
nonmetal, subscripts indicate how many of each
© 2013 Pearson Education, Inc. Chapter 5, Section 1 28
Charge Balance in Na2S
In sodium sulfide, Na2S,
� two Na atoms lose one valence electron each.
� one S atom gains two electrons.
� subscripts show the number of ions needed to give
charge balance.
The group of ions with the lowest ratio of ions in an
ionic compound is called a formula unit .
© 2013 Pearson Education, Inc. Chapter 5, Section 1 29
Formula from Ionic Charges
Write the ionic formula of the compound containing
Ba2+ and Cl−.
� Write the symbols of each ion.
Ba2+ Cl−
� Balance the charges.
Ba2+ Cl− (two Cl−−−− needed)
Cl−
� Write the metal first, followed by the nonmetal,
, using a subscript to represent the number of
. Cl− ions.
BaCl2
© 2013 Pearson Education, Inc. Chapter 5, Section 1 30
Writing Names For Binary Ionic
Compounds
The name of an Binary ionic compound
� is made up of 2 elements.
� has the name of the metal ion written first.
� has the name of the nonmetal ion written second
using the first syllable of its element name.
followed by ide.
� has a space separating the name of the metal and
nonmetal.
6
© 2013 Pearson Education, Inc. Chapter 5, Section 3 31
Charges of Representative
Elements
© 2013 Pearson Education, Inc. Chapter 5, Section 1 32
Naming Binary Ionic Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 33
Step 1 Identify the cation and anion.
The cation from Group 2A (2) is Mg2+, and
the anion from Group 5A (15) is N3−.
Step 2 Name the cation by its element name.
Cation: Mg2+ is magnesium
Step 3 Name the anion by using the first syllable of
its element name followed by ide.
Anion: nitrogen becomes nitride, N3−
Step 4 Write the name for the cation first and the
name for the anion second.
Mg3N2 is magnesium nitride.
Naming Mg3N2
© 2013 Pearson Education, Inc. Chapter 5, Section 1 34
What do you notice about these
names and formulas?Chemical Formula Written Form
Fe2O3 iron(III) oxide
FeCl2 iron(II) chloride
PbS lead(II) sulfide
AlP aluminum phosphide
CuF2 copper(II) fluoride
SnI4 tin(IV) iodide
KBr potassium bromide
34
© 2013 Pearson Education, Inc. Chapter 5, Section 1 35
Metals with Variable Ion Charge
Most transition metals (3-12) and Group 4A (14) metals
form 2 or more positive ions, except Zn2+, Ag+, and Cd2+,
which form only one ion.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 36
Metals with Variable Charge
The names of transition metals with two or more
positive ions (cations) use a Roman numeral after the
name of the metal to identify the ion charge.
7
© 2013 Pearson Education, Inc. Chapter 5, Section 1 37
Naming Ionic Compounds with
Variable Charge Metals
© 2013 Pearson Education, Inc. Chapter 5, Section 1 38
Naming FeCl2
Step 1 Determine the charge of the cation from the
anion.
Analyze the Problem.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 39
Naming FeCl2
Step 2 Name the cation by its element name and
use a Roman numeral in parentheses for the
charge.
Fe2+ = iron(II)
Step 3 Name the anion by using the first syllable of
its element name followed by ide .
Cl− = chloride
Step 4 Write the name for the cation first and the
name for the anion second.
iron(II) chloride
© 2013 Pearson Education, Inc. Chapter 5, Section 1 40
Examples of Names of Compounds
with Variable Charge Metals
© 2013 Pearson Education, Inc. Chapter 5, Section 1 41
What do you notice about these
names and formulas?Chemical Formula Written Form
(NH4)2S ammonium sulfide
CoSO4 cobalt (II) sulfate
Fe(OH)3 iron (III) hydroxide
Ca3(PO4)2 calcium phosphate
NH4NO3 ammonium nitrate
41© 2013 Pearson Education, Inc. Chapter 5, Section 1 42
A polyatomic ion
� is a group of atoms.
� has an overall ionic charge.
Examples:
NH4+ ammonium OH− hydroxide
NO3− nitrate NO2
− nitrite
CO32− carbonate PO4
3− phosphate
HCO3− hydrogen carbonate
(bicarbonate)
Polyatomic Ions
8
© 2013 Pearson Education, Inc. Chapter 5, Section 1 43
The names of common polyatomic anions
� end in ate.
NO3− nitrate PO4
3− phosphate
� with one oxygen less, end in ite
NO2− nitrite PO3
3− phosphite
� with hydrogen attached, use prefix hydrogen (or bi).
HCO3− hydrogen carbonate (bicarbonate)
HSO3− hydrogen sulfite (bisulfite)
Some Names of Polyatomic Ions
© 2013 Pearson Education, Inc. Chapter 5, Section 1 44
Compounds Containing
Polyatomic Ions
Polyatomic ions
� must be associated with an ion of opposite charge.
� form ionic bonds with ions of opposite charge to
achieve charge balance.
Example:
Ca2+ NO3−
calcium nitrate ion
charge balance:
Ca(NO3)2
calcium nitrate
© 2013 Pearson Education, Inc. Chapter 5, Section 1 45
Some Compounds with Polyatomic
Ions
© 2013 Pearson Education, Inc. Chapter 5, Section 1 46
Guide to Naming Compounds with
Polyatomic Ions
© 2013 Pearson Education, Inc. Chapter 5, Section 1 47
Step 1 Identify the cation and polyatomic ion(anion).
Cation: K+ Anion: SO42−
Step 2 Name the cation, using a Roman numeral if needed.
K+ = potassium ion
Step 3 Name the polyatomic ion.
SO42− = sulfate ion
Step 4 Write the name or the compound, cation first and the polyatomic ion second.
K2SO4 = potassium sulfate
Name K2SO4
© 2013 Pearson Education, Inc. Chapter 5, Section 1 48
Guide to Writing Formulas for
Compounds with Polyatomic Ions
9
© 2013 Pearson Education, Inc. Chapter 5, Section 1 49
Write the Formula for
Aluminium Hydroxide
Step 1 Identify the cation and polyatomic ion
(anion).
Al3+ and OH−
Step 2 Balance the charges.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 50
Write the Formula for
Aluminium Hydroxide
Step 3 Write the formula, cation first, using the
subscripts from charge balance.
Al(OH)3
© 2013 Pearson Education, Inc. Chapter 5, Section 1 51
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
5.5
Covalent Compounds:
Sharing Electrons
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.Lectures
© 2013 Pearson Education, Inc. Chapter 5, Section 1 52
Covalent bonds form when atoms
of nonmetals share electrons to
complete octets.
Valence electrons are not
transferred, but shared to achieve
stability.
Covalent Bonds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 53
Formation of H2
In the simplest covalent molecule, H2 , the H atoms
� increase attraction as they move closer.
� share electrons to achieve a stable configuration.
� form a covalent bond.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 54
Formation of H2
10
© 2013 Pearson Education, Inc. Chapter 5, Section 1 55
Electron-Dot Formulas of
Covalent Molecules
In a fluorine (F2) molecule, the F atoms
� share one of their valence electrons.
� acquire an octet.
� form a covalent bond.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 56
Elements That Exist as Diatomic
Molecules
These seven elements share
electrons to form diatomic,
covalent molecules.
BrINClHOF!
Remember it!
© 2013 Pearson Education, Inc. Chapter 5, Section 1 57
The number of covalent bonds a nonmetal forms is
usually equal to the number of electrons it needs to
acquire a stable electron configuration.
Typical bonding patterns for some nonmetals are
shown in the table below.
Bonding Patterns of Some
Nonmetals
© 2013 Pearson Education, Inc. Chapter 5, Section 1 58
� Biological molecules tend to be made out of H,
O,N, and C mostly.
� Bonding tendency rule for H, O, N, C
HONC 1234 Rule
Hydrogen tends to form 1 bond
Oxygen tends to form 2 bonds
Nitrogen tends to form 3 bonds
Carbon tends to form 4 bonds
Bonding Patterns of Some
Nonmetals
© 2013 Pearson Education, Inc. Chapter 5, Section 1 59
Electron-Dot Formulas for Some
Covalent Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 60
Guide to Drawing Electron-Dot
Formulas
11
© 2013 Pearson Education, Inc. Chapter 5, Section 1 61
Step 1 Determine the arrangement of atoms.
In NH3, N is the central atom and is bonded to
three H atoms.
Step 2 Determine the total number
of valence electrons.
Total valence electrons for NH3 = 8 e−
Draw the Electron-Dot Formula
for NH3
H N H
H
© 2013 Pearson Education, Inc. Chapter 5, Section 1 62
Step 3 Attach each bonded atom to the central
atom with a pair of electrons.
Draw the Electron-Dot Formula
for NH3
H N H
H
© 2013 Pearson Education, Inc. Chapter 5, Section 1 63
Step 4 Place the remaining electrons using single
or multiple bonds to complete the octets.
8 valence e− − 6 bonding e− = 2 e− remaining
Use the remaining 2 e− to complete the octet
around the N atom.
Draw the Electron-Dot Formula
for NH3
H N H
H
H N H
H
or
© 2013 Pearson Education, Inc. Chapter 5, Section 1 64
Exceptions to the Octet Rule
Not all atoms have octets.
� Some can have less than an octet, such as
H, which requires only 2 electrons,
B, which requires only 3 electrons, and
Be, which requires only 4 electrons.
� Some can have expanded octets, such as
P, which can have 10 electrons,
S, which can have 12 electrons, and
Cl, Br and I, which can have 14 electrons
© 2013 Pearson Education, Inc. Chapter 5, Section 1 65
Single and Multiple Bonds
In many covalent compounds, atoms share two or
three pairs of electrons to complete their octets.
� In a single bond, one pair of electrons is shared.
� In a double bond, two pairs of electrons are
shared.
� In a triple bond, three pairs of electrons are
shared.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 66
Step 1 Determine the arrangement of atoms.
In CS2, C is the central atom and is bonded to
two S atoms.
Draw the Electron-Dot Formula
for CS2
S C S
12
© 2013 Pearson Education, Inc. Chapter 5, Section 1 67
Step 2 Determine the total number
of valence electrons.
Total valence electrons for
Draw the Electron-Dot Formula
for CS2
© 2013 Pearson Education, Inc. Chapter 5, Section 1 68
Step 3 Attach each bonded atom to the central
atom with a pair of electrons.
A pair of bonding electrons (single bond) is
placed between each S atom and the central C
atom.
Draw the Electron-Dot Formula
for CS2
S C S
© 2013 Pearson Education, Inc. Chapter 5, Section 1 69
Step 4 Place the remaining electrons using single
or multiple bonds to complete the octets.
16 valence e− - 4 bonding e− = 12 e− remaining
The remaining 12 electrons are placed as six
lone pairs of electrons on both S atoms.
However, this does not complete the octet for
the C atom.
Draw the Electron-Dot Formula
for CS2
S C S
© 2013 Pearson Education, Inc. Chapter 5, Section 1 70
Step 4 Continued: Double and Triple Covalent
Bonds: To complete the octet for the C atom,
it needs to share an additional lone pair from
each of the S atoms, forming a double bond
with each S atom.
Draw the Electron-Dot Formula
for CS2
S C S
or
S C S
S C S
© 2013 Pearson Education, Inc. Chapter 5, Section 1 71
A Nitrogen Molecule has
a Triple Bond
In a nitrogen molecule, N2,
� each N atom shares 3 electrons,
� each N atom attains an octet, and
� the sharing of 3 sets of electrons is called a triple
bond.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 72
Resonance structures are
� two or more electron-dot formulas for the same
arrangement of atoms.
� related by a double-headed arrow ( ).
� written by changing the location of a double bond
between the central atom and a different attached
atom.
Resonance Structures
13
© 2013 Pearson Education, Inc. Chapter 5, Section 1 73
Sulfur dioxide has two resonance structures.
Step 1 Determine the arrangement of atoms. In SO2,
the S atom is the central atom.
O S O
Step 2 Determine the total number of valence
electrons.
Total valence electrons for
Writing Resonance Structures for
SO2
© 2013 Pearson Education, Inc. Chapter 5, Section 1 74
Step 3 Attach each bonded atom to the central
atom with a pair of electrons.
Writing Resonance Structures for
SO2
O S OO S O or
© 2013 Pearson Education, Inc. Chapter 5, Section 1 75
Step 4 Place the remaining electrons using single
or multiple bonds to complete the octets.
The remaining 14 electrons are drawn as lone
pairs of electrons to complete the octets of the
O atoms, but not the S atom.
Writing Resonance Structures for
SO2
O S O
O S O or
© 2013 Pearson Education, Inc. Chapter 5, Section 1 76
Step 4 Continued: To complete the octet for S, an
additional lone pair from one of the O atoms
is shared to form a double bond.
Because the shared lone pair of electrons can
come from either O atom, two resonance
structures can be drawn.
Writing Resonance Structures for
SO2
O S O
O S O
© 2013 Pearson Education, Inc. Chapter 5, Section 1 77
FNO2, a rocket propellant, has two resonance structures.
One is shown below. What is the other resonance
structure?
Learning Check
© 2013 Pearson Education, Inc. Chapter 5, Section 1 78
FNO2, a rocket propellant, has two resonance structures.
One is shown below. What is the other resonance
structure?
Solution
14
© 2013 Pearson Education, Inc. Chapter 5, Section 1 79
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
5.6
Naming and Writing
Covalent Formulas
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.Lectures
© 2013 Pearson Education, Inc. Chapter 5, Section 1 80
Names of Covalent Compounds
When naming a covalent compound,
� the first nonmetal in the formula is named by its
element name, and
� the second nonmetal is named using the first syllable
of its element name, followed by ide.
Prefixes are used
� to indicate the number of atoms present for each
element in the compound, and
� because two nonmetals can form two or more
different compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 81
Prefixes Used in Naming
Covalent Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 82
Some Covalent Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 83
Guide to Naming Covalent
Compounds with Two Nonmetals
© 2013 Pearson Education, Inc. Chapter 5, Section 1 84
Naming Covalent Compounds,
SO2
Analyze the Problem.
Step 1 Name the first nonmetal by its element
name. In SO2, the first nonmetal (S) is sulfur.
Symbols of
Elements
S O
Name sulfur oxygen
Subscripts 1 2
Prefix (none) understood di
15
© 2013 Pearson Education, Inc. Chapter 5, Section 1 85
Naming Covalent Compounds,
SO2
Step 2 Name the second nonmetal by using the first
syllable of its name followed by ide. The
second nonmetal (O) is named oxide.
Step 3 Add prefixes to indicate the number of
atoms (subscripts). Because there is
one sulfur atom, no prefix is needed. The
subscript two for the oxygen atoms is written as
the prefix di .
The name of SO2 is sulfur dioxide.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 86
Learning Check
Name the covalent compound P2O5.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 87
Solution
Name the covalent compound P2O5.
Analyze the Problem.
Step 1 Name the first nonmetal by its element
name. In P2O5, the first nonmetal (P) is
phosphorus.
Symbols of Elements P O
Name phosphorus oxygen
Subscripts 2 5
Prefix di penta
© 2013 Pearson Education, Inc. Chapter 5, Section 1 88
Solution
Step 2 Name the second nonmetal by using the first
syllable of its name followed by ide. The
second nonmetal (O) is named oxide.
Step 3 Add prefixes to indicate the number of
atoms (subscripts). The subscript for the two
P atoms is di; the subscript for the five oxygen
atoms is penta.
The name of P2O5 is phosphorus pentoxide.
When the vowels o and o, or a and o appear
together (pentaoxide), the first vowel is omitted.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 89
Guide to Writing Formulas for
Covalent Compounds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 90
Write the formula for sulfur hexafluoride.
Analyze the Problem.
Step 1 Write the symbols in order of the
elements in the name.
S F
Step 2 Write any prefixes as subscripts.
Prefix hexa = 6 Formula: SF6
Writing Formulas of Covalent
Compounds
Name sulfur hexafluoride
Symbols of Elements S F
Subscripts 1 6 (from hexa)
16
© 2013 Pearson Education, Inc. Chapter 5, Section 1 91
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
The Shape of Molecules
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.Lectures
© 2013 Pearson Education, Inc. Chapter 5, Section 1 92
Valence-Shell Electron-Pair Repulsion
Theory (VSEPR)
In the valence-shell electron-pair repulsion (VSEPR)
theory, the electron groups around a central atom
� are arranged as far apart from each other as possible.
� have the least amount of electron-electron repulsion.
� are used to predict the molecular shape.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 93
Molecular Shape – why does it
matter?
• Watson and Crick withtheir famous 3-D model of DNA
• How did they know? •Shape of molecules important as it influences their physical and chemical behavior
© 2013 Pearson Education, Inc. Chapter 5, Section 1 94
What Determines the 3-D Shape of
a Molecule?
� Need to understand electrostatic repulsion
(electron clouds around the central atom in a
molecule are going to repel each other)
� Need to draw Lewis structures
� Need to differentiate between lone pair
electrons and bonding pair electrons around
the central atom in a molecule
� Need to apply a theory based on all of the
above
© 2013 Pearson Education, Inc. Chapter 5, Section 1 95
Lewis Structure Lone pair e- (nonbonding e- domain)
Bonding pair e- (bonding domain)
• Electron pairs are also called domains
• Electron domains try to stay out of each
others way (to minimize repulsion)
© 2013 Pearson Education, Inc. Chapter 5, Section 1 96
Valence Shell Electron Pair
Repulsion Theory
� VSEPR, for short
� Proposed by English chemist Ron Gillespie in
the 1950s
� Based on Lewis structures of molecules and
electron repulsion
17
© 2013 Pearson Education, Inc. Chapter 5, Section 1 97
VSEPR Theory Rules
1. In a molecule, electron domains (pairs) will
orient themselves around the central atom
in an arrangement that minimizes the
repulsions among them.
2. The three dimensional orientation of the
outer atoms around the central atom
depends on the numbers and types of e-
domains
3. Lone pair electrons count as 1 electron
domain; single, double or triple bonding pair
electrons count as 1 electron domain.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 98
Shapes of Molecules
The three-dimensional shape of a molecule can be
predicted by
� the number of bonding groups (bonding domains) and
lone-pair electrons (nonbonding domains) around the
central atom in the molecule
� VSEPR theory (valence-shell-electron-pair repulsion).
© 2013 Pearson Education, Inc. Chapter 5, Section 1 99
Molecules with Two Electron Domains
In a molecule of BeCl2,
� there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule).
� to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other.
� the shape of the molecule is linear.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 100
Molecules With Three Electron
Domains
In a molecule of BF3,
� three electron groups are bonded to the central atom
B (B is an exception to the octet rule).
� repulsion is minimized with 3 electron groups at
angles of 120°.
� the shape is trigonal planar.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 101
Three Electron Domains with a
Lone Pair
In a molecule of SO2,
� S has 3 electron groups; 2 electron groups bonded to
O atoms and one lone pair.
� repulsion is minimized with the electron groups at
angles of 120°, a trigonal planar arrangement.
� the shape is determined by the two O atoms bonded
to S, giving SO2 a bent (~120°°°°) shape.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 102
Four Electron Domains
In a molecule of CH4,
� there are four electron groups around C.
� repulsion is minimized by placing four electron groups
at angles of 109°, which is a tetrahedral arrangement.
� the four bonded atoms form a tetrahedral shape.
18
© 2013 Pearson Education, Inc. Chapter 5, Section 1 103
Four Electron Groups with a Lone
Pair
In a molecule of NH3,
� three electron groups bond to H atoms, and the fourth
one is a lone (nonbonding) pair.
� repulsion is minimized with 4 electron groups in a
tetrahedral arrangement.
� the three bonded atoms form a pyramidal (~109°°°°)
shape.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 104
Four Electron Groups with Two
Lone Pairs
In a molecule of H2O,
� two electron groups are bonded to H atoms and two
are lone pairs (4 electron groups).
� four electron groups minimize repulsion in a
tetrahedral arrangement.
� the shape with two bonded atoms is bent (~109°°°°).
© 2013 Pearson Education, Inc. Chapter 5, Section 1 105
Molecular Shapes for 2 and 3
Bonded Atoms
© 2013 Pearson Education, Inc. Chapter 5, Section 1 106
Molecular Shapes for 4 Bonded
Atoms
© 2013 Pearson Education, Inc. Chapter 5, Section 1 107
Guide to Predicting Molecular
Shape (VSEPR Theory)
© 2013 Pearson Education, Inc. Chapter 5, Section 1 108
Predicting the Molecular Shape of
H2Se
Step 1 Draw the electron-dot formula. In the
electron-dot formula for H2Se, there are four
electron groups, including two lone pairs of
electrons around Se.
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© 2013 Pearson Education, Inc. Chapter 5, Section 1 109
Predicting the Molecular Shape of
H2Se
Step 2 Arrange the electron groups around the
central atom to minimize repulsion. The
four electron groups around Se would have a
tetrahedral arrangement.
Step 3 Use the atoms bonded to the central atom
to determine the molecular shape. Two
bonded atoms give H2Se a bent shape with a
bond angle of ~109°.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 110
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
A closer look at
bonding type – the
concept of
electronegativity
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 111
General Rules for Bond Type
Nonmetal
Nonmetal
Covalent
Bond
Metal
Ionic
Bond
Negative complex
ions
© 2013 Pearson Education, Inc. Chapter 5, Section 1 112
Ionic or Covalent?
� So far, we have been employing the following
generalization for the type of bonding holding
atoms together in a substance
Generalization
� metal + nonmetal = ionic bond
� nonmetal + nonmetal = covalent bond
Reality
� Bonding type is on a continuum, from 100% ionic
to 100% covalent!
100% covalent 100% ionic
© 2013 Pearson Education, Inc. Chapter 5, Section 1 113
Bonding Conundrum
� Arose in the 1930’s from American Chemist
Linus Pauling’s work on the nature of the
chemical bond.
� Pauling realized that there was some “in-
betweeness” with the bonding in many
substances
� Calculated the ability of elements to attract
electrons to themselves when bonding
� Called this ability the Electronegativity of the
element
113© 2013 Pearson Education, Inc. Chapter 5, Section 1 114
Determining Predominant Bond type
� Electronegativity is a relative measure of the
tendency for atoms of an element to attract
electrons to themselves in a chemical bond.
� Originated with American chemist Linus Pauling
(1901-1994), a 2x Nobel Prize winner who did
most of his work at Cal Poly Tech, but finished
his career at Stanford
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© 2013 Pearson Education, Inc. Chapter 5, Section 1 115 © 2013 Pearson Education, Inc. Chapter 5, Section 1 116
Electronegativity (EN) Values of the
Elements
Scale ranges from 0.7 to 4.0
© 2013 Pearson Education, Inc. Chapter 5, Section 1 117
Electronegativity
� is a measure of an atom’s ability to attract
electrons to itself in a chemical bond
� increases from left to right, going across a period
on the periodic table.
� decreases going down a group on the periodic
table.
� is high for the nonmetals, with fluorine as the
highest.
� is low for the metals and transition metals.
Electronegativity
© 2013 Pearson Education, Inc. Chapter 5, Section 1 118
Trends in Periodic table
© 2013 Pearson Education, Inc. Chapter 5, Section 1 119
Periodic trends in electronegativity
� Atomic radius increases, valence e- further
away from nucleus, decreases pulling power
of nucleus
Electronegativity decreases down a group
because:
Electronegativity increases from left to right across
a period because:
� Atomic radius decreases, valence e- closer to
the nucleus, increases pulling power of nucleus
© 2013 Pearson Education, Inc. Chapter 5, Section 1 120
Bonding and Electronegativity
� According to Pauling, the difference in the electronegativity
values of the two atoms involved in a chemical bond can
be used to predict the type of bond that forms.
� Pauling Classified chemical bonds into the following three
types, based on differences in electronegativity
� Ionic
� Polar Covalent
� Nonpolar Covalent
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© 2013 Pearson Education, Inc. Chapter 5, Section 1 121
Electronegativity and Bond Type
� It is the difference in the electronegativies (∆EN) of
two bonded atoms that determines the
predominant bond type
∆EN ≤ 0.4 � Nonpolar covalent bond
(e- shared equally between atoms)
∆EN 0.5 to 1.7 polar covalent bond,
(e- shared unequally)
∆EN > 1.7 ionic bond (e- transferred)
121© 2013 Pearson Education, Inc. Chapter 5, Section 1 122
Bond Character Summary
a) Nonpolar covalent bond �
electrons shared equally
b) Polar covalent bond � e- not
shared equally, more
electronegative atom has greater
share of e- cloud
c) Ionic bond – e- transferred from
one atom to other, creates a + and
– ion
© 2013 Pearson Education, Inc. Chapter 5, Section 1 123
A nonpolar covalent bond
� occurs between nonmetals.
� has an equal or almost equal sharing of electrons.
� has almost no electronegativity difference (0.0 to
0.4).
Examples:
Atoms Electronegativity Type of BondDifference
Nonpolar Covalent Bonds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 124
Example of Nonpolar Covalent
Bonds� Diatomic elements e.g. H2
� Bonds between P and H = 2.1-2.1=0
© 2013 Pearson Education, Inc. Chapter 5, Section 1 125
A polar covalent bond
� occurs between nonmetal atoms that do not share
electrons equally.
� has a moderate electronegativity difference
(0.5 to 1.7).
Examples:
Atoms Electronegativity Type of Bond
Difference
Polar Covalent Bonds
© 2013 Pearson Education, Inc. Chapter 5, Section 1 126
Comparing Nonpolar and Polar
Covalent Bonds
22
© 2013 Pearson Education, Inc. Chapter 5, Section 1 127
Result of differences in electronegativity:
• More electronegative element has greater
share of e- cloud, it is electron rich
•Less electronegative element is e- poor
•Creation of partial charges on the atoms in the
bond (bond dipoles) δδδδ+ and δδδδ -
© 2013 Pearson Education, Inc. Chapter 5, Section 1 128
H Cl
Example: Bond between H and Cl
• Electronegativity values H= 2.1, Cl = 3.0
• e- cloud pulled towards Cl atom = Polar Covalent
Bond
electron rich
region
electron poor
region
e- rich
e- poor
ClH
δ+ δ-
9.5
Direction e- cloud being pulled
2.1 3.0
© 2013 Pearson Education, Inc. Chapter 5, Section 1 129
Bond Polarity and Dipoles
Bonds become more polar as the difference in
electronegativity values of bonding atoms increases.
Polar covalent bonds have a separation of charges
called a dipole.
The positive and negative ends of the dipole are
indicated by the lowercase Greek letter delta with a
positive or negative sign, δ+ and δ-, or an arrow that
points from the positive to the negative charge.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 130
Ionic Bonds
An ionic bond
� occurs between metal and nonmetal ions.
� is a result of electron transfer.
� has a large electronegativity difference (1.8 or more).
Examples:
Atoms Electronegativity Type of Bond Difference
© 2013 Pearson Education, Inc. Chapter 5, Section 1 131
Determining Predominant Bond Type
� Look at the electronegativity values of atoms
involved in bond
� Calculate electronegativity difference (∆EN)
� If the bond is polar covalent, draw arrow in
direction that e- cloud is pulled (more
electronegative atom)
� Also use lower case Greek symbol for delta, δ,
along with + or - sign
© 2013 Pearson Education, Inc. Chapter 5, Section 1 132
Electronegativity and Bond Types
23
© 2013 Pearson Education, Inc. Chapter 5, Section 1 133
Predicting Bond Types
© 2013 Pearson Education, Inc. Chapter 5, Section 1 134
Use the electronegativity difference to identify the type
of bond between the following atoms as nonpolar
covalent (NP), polar covalent (P), or ionic (I).
A. K–N
B. N–O
C. Cl–Cl
D. B–Cl
Learning Check
© 2013 Pearson Education, Inc. Chapter 5, Section 1 135
Use the electronegativity difference to identify the type
of bond between the following atoms as nonpolar
covalent (NP), polar covalent (P), or ionic (I).
A. K–N 2.2 ionic (I)
B. N–O 0.5 polar covalent (P)
C. Cl–Cl 0.0 nonpolar covalent (NP)
D. B–Cl 1.0 polar covalent (P)
Solution
© 2013 Pearson Education, Inc. Chapter 5, Section 1 136
Polar Molecules
A polar molecule
� contains polar bonds.
� has a separation of positive and negative charge. called a dipole, indicated with δ+ and δ–.
� has dipoles that do not cancel.δ+ δ–
H–Cl NH3
dipole
Dipoles do not cancel.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 137
Nonpolar Molecules
A nonpolar molecule
� contains nonpolar bonds
Cl–Cl H–H
� or has a symmetrical arrangement of polar bonds.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 138
Guide to Determination of
Polarity
24
© 2013 Pearson Education, Inc. Chapter 5, Section 1 139
Molecular Polarity, H2O
Determine the polarity of the H2O molecule.
Step 1 Determine if the bonds are polar covalent or
nonpolar covalent. From the electronegativity
table, O 3.5 and H 2.1 gives a difference of 1.4,
which makes the O — H bonds, polar covalent.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 140
Molecular Polarity, H2O
Determine the polarity of the H2O molecule.
Step 2 If the bonds are polar covalent, draw the
electron-dot formula and determine if the
dipoles cancel or not. The four electron
groups of oxygen are bonded to two H atoms.
Thus, the H2O molecule has a net dipole, which
makes it a polar molecule.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 141
Polar Molecules
A polar molecule
� contains polar bonds.
� has a separation of positive and negative charge. called a dipole, indicated with δ+ and δ–.
� has dipoles that do not cancel.δ+ δ–
H–Cl NH3
dipole
Dipoles do not cancel.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 142
Learning Check
Determine the shape of each of the following molecules
and whether they are polar or nonpolar. Explain.
1. PBr3
2. HBr
3. Br2
4. SiBr4
© 2013 Pearson Education, Inc. Chapter 5, Section 1 143
Solution
Determine the shape of each of the following molecules
and whether they are polar or nonpolar. Explain.
1. PBr3 pyramidal; polar; dipoles don’t cancel
2. HBr linear; polar; one polar bond (dipole)
3. Br2 linear; nonpolar; nonpolar bond
4. SiBr4 tetrahedral; nonpolar; dipoles cancel
© 2013 Pearson Education, Inc. Chapter 5, Section 1 144
General, Organic, and
Biological ChemistryFourth Edition
Karen Timberlake
5.9
Attractive Forces in
Compounds
Chapter 5Compounds and
Their Bonds
© 2013 Pearson Education, Inc.Lectures
25
© 2013 Pearson Education, Inc. Chapter 5, Section 1 145
States of Matter
The fundamental difference between states of
matter is the distance between particles.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 146
States of Matter
Because in the solid and liquid states
particles are closer together, we refer to them
as condensed phases.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 147
The States of Matter
� The state a substance is
in at a particular
temperature and
pressure depends on
two antagonistic entities:
� The kinetic energy of the
particles
� The strength of the
attractions between the
particles
© 2013 Pearson Education, Inc. Chapter 5, Section 1 148
Ionic Compounds
In ionic compounds, ionic bonds
� require large amounts of energy to break.
� hold positive and negative ions together.
� explain their high melting points.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 149
What about covalent compounds?
The attractions between molecules are not
nearly as strong as the intramolecular
attractions that hold compounds together.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 150
Intermolecular Forces
They are, however, strong enough to control
physical properties such as the state a
molecular compound will be in at room
temperature, its melting and boiling point.
26
© 2013 Pearson Education, Inc. Chapter 5, Section 1 151
Intermolecular Forces
In addition, they are responsible for the
shapes of protein molecules, DNA, water’s
unique properties, the structure of the cell
membrane, etc
© 2013 Pearson Education, Inc. Chapter 5, Section 1 152
Covalent Compounds
In covalent compounds, the attractive forces between
solid and liquid molecules
� are weaker than ionic bonds.
� require less energy to break.
� explain why their melting points are lower than ionic
compounds.
These attractive forces include
� dipole-dipole attractions,
� dispersion forces, and
� hydrogen bonding.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 153
Intermolecular Forces
They are, however, strong enough to control
physical properties such as boiling and
melting points, vapor pressures, and
viscosities.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 154
Dipole-Dipole Attractions
In covalent compounds, polar molecules exert attractive
forces between molecules called dipole-dipole
attractions. dipole dipole forces
© 2013 Pearson Education, Inc. Chapter 5, Section 1 155
Dipole-Dipole Attractions,
Hydrogen Bonds
In covalent compounds, some polar molecules form
strong dipole attractions called hydrogen bonds, which
occur between the partially positive hydrogen atom of
one molecule and a lone pair of electrons on a nitrogen,
oxygen, or fluorine atom in another molecule.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 156
Dispersion Forces
Dispersion forces are
� weak attractions between nonpolar molecules.
� caused by temporary dipoles that develop when
electrons are not distributed equally.
Nonpolar molecules form attractions when they form temporary
dipoles.
27
© 2013 Pearson Education, Inc. Chapter 5, Section 1 157
London Dispersion Forces
While the electrons in the 1s orbital of helium
would repel each other (and, therefore, tend
to stay far away from each other), it does
happen that they occasionally wind up on the
same side of the atom.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 158
London Dispersion Forces
At that instant, then, the helium atom is polar,
with an excess of electrons on the left side
and a shortage on the right side.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 159
London Dispersion Forces
Another helium nearby, then, would have a
dipole induced in it, as the electrons on the
left side of helium atom 2 repel the electrons
in the cloud on helium atom 1.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 160
London Dispersion Forces
London dispersion forces, or dispersion
forces, are attractions between an
instantaneous dipole and an induced dipole.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 161
London Dispersion Forces
� These forces are present in all molecules, whether they are polar or nonpolar.
� The tendency of an electron cloud to distort in this way is called polarizability.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 162
Factors Affecting London Forces
� The shape of the molecule
affects the strength of dispersion
forces: long, skinny molecules
(like n-pentane tend to have
stronger dispersion forces than
short, fat ones (like neopentane).
� This is due to the increased
surface area in n-pentane.
28
© 2013 Pearson Education, Inc. Chapter 5, Section 1 163
Factors Affecting London Forces
� The strength of dispersion forces tends to
increase with increased molecular weight.
� Larger atoms have larger electron clouds,
which are easier to polarize.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 164
Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces?
� If two molecules are of comparable size and
shape, dipole-dipole interactions will likely be
the dominating force.
� If one molecule is much larger than another,
dispersion forces will likely determine its
physical properties.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 165
Comparison of Bonding and
Attractive Forces
© 2013 Pearson Education, Inc. Chapter 5, Section 1 166
Intermolecular Forces Affect Many
Physical Properties
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 167
Melting Points and Attractive
Forces
The stronger the attractive force between ions or
molecules, the higher the melting points.
Ionic compounds, have the strongest attractive force
and, therefore the highest melting points.
Covalent molecules have less attractive forces than ionic
compounds and, therefore lower melting points.
© 2013 Pearson Education, Inc. Chapter 5, Section 1 168
Melting Points and Attractive
Forces
The attractive forces between covalent molecules vary
in magnitude; the stronger the attractive force, the
higher its melting point.
� Hydrogen bonds are the strongest type of dipole–
dipole attractions, requiring the most energy to
break, followed by dipole–dipole forces.
� Dispersion forces are the weakest, requiring even
less energy to break them, and therefore have lower
melting points than hydrogen bonds and dipole–
dipole forces.
29
© 2013 Pearson Education, Inc. Chapter 5, Section 1 169
Melting Points of Selected
Substances
© 2013 Pearson Education, Inc. Chapter 5, Section 1 170
Learning Check
Identify the main type of attractive forces for each of the
following compounds: ionic bonds, dipole–dipole,
hydrogen bonds or dispersion.
1. NCl3
2. H2O
3. Br2
4. KCl
5. NH3
© 2013 Pearson Education, Inc. Chapter 5, Section 1 171
Solution
Identify the main type of attractive forces for each of the
following compounds: ionic bonds, dipole–dipole,
hydrogen bonds or dispersion.
1. NCl3 dipole–dipole
2. H2O hydrogen bonds
3. Br2 dispersion
4. KCl ionic bonds
5. NH3 hydrogen bonds