Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
NOTES - Unit 4: Bonding
PART 1: Ionic and Covalent Bonding
CHEMICAL BONDS
 A chemical bond is an attraction between 2 atoms or ions.
 Bonding occurs because it __________________________________________________________ of the system.
 Three broad classifications (general rule of thumb)
o Ionic ( ______________________ + _______________________ )
o Covalent ( ______________________ + _______________________ )
o Metallic ( ______________________ + _______________________ )
 ______________________________________: the power of an atom in a molecule to attract electrons to itself.
o The greater an atom’s electronegativity, the greater its ability to attract__________________________.
o In a compound, the element with the greater electronegativity will be the more _____________ species.
o The element with _______________________ electronegativity is sometimes referred to as the more
“_____________________________” element. This will become the more positive species in the bond.
 ______________________ bonds exist between atoms with low electronegativities & those with high
electronegativities. In general, between metals and nonmetals.
 _________________________________ bonds exist between dissimilar atoms with different electronegativities.
 ___________________________________________ bonds exist between identical nonmetallic atoms or
nonmetallic atoms with similar electronegativities.
 Atoms react with each other in chemical reactions in a quest to have complete outer electron energy levels like
the ____________________________________________.
o Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by ________ valence
electrons.
o Duet Rule: ___________________________ and _______________________________ want to have 2
electrons. Recall that the first energy level only holds two electrons.
IONIC BONDING
 Different texts report different “cutoff” values, but according to IB, _______________________ for ionic bonds.
 The greater the difference in electronegativity between the atoms, the more ____________________ the bond.
 Ionic bonds form between metals on the _____________________ of the periodic table and elements in groups
________________________________ (or polyatomic ions attracted to other ions)
o Note: group 4 elements (C and Si) do not form ionic bonds. They tend to form giant molecular
structures (network solids) or simple molecules.
 Example: NaCl
o Sodium (Na) has 11 electrons (but only one valence electron): _____________________________
o By losing this electron Na becomes Na+ : __________________________________or___________,
which has a full outer shell.
o Chlorine (Cl) has 17 electrons (seven valence electrons): __________________________________
1
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
By gaining an electron Cl becomes Cl-: ________________________________________or___________,
which has a filled outer shell.
o An ionic bond results: ______________________________________________
attractions between two oppositely charged ________________.
o These two ions do not exist in isolation. NaCl is simply the formula ____________,
a representation of the ration of cation:anion in the ionic lattice structure.
o In 3-dimentional cubic NaCl, each Na+ is surrounded by ___ Cl- ions, and vice versa.
Examples of Ions
o
Group #
1
2
3
4
5
6
7
Ex.
Number valence
e’s
Number e’s
transferred
e’s lost or gained
Charge of ion
formed
Type of element
Na
Ca
Al
C
P
O
Br
Polyatomic Ions: ions containing more than one element
 In ions formed from more than one element the charge is often spread (delocalized) over the whole ion.
 Cation example: ammonium, NH4+
o All four of the N-H bonds are identical; +1 charge is distributed evenly throughout the ion
 Polyatomic anions are sometimes known as acid radicals  formed when an acid loses one or more H+ ions.
 Anion examples:
Ion name
Formula
From…
Hydroxide
Water (H2O)
Nitrate
Nitric acid (HNO3)
Sulfate
Sulfuric acid (H2SO4)
Hydrogen sulfate (a.k.a. bisulfate)
Sulfuric acid (H2SO4)
Carbonate
Carbonic acid (H2CO3)
Hydrogen carbonate (a.k.a. bicarbonate)
Carbonic acid (H2CO3)
Ethanoate (a.k.a. acetate)
From ethanoic acid (vinegar), or acetic acid
(CH3COOH)
… and all the others on your “Things to Know and Love…” handout
2
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Formulas of Ionic Compounds: Lowest ratio of cation to anion that results in a net charge of zero.
 Examples:
o Copper (II) sulfate = __________________
o
Sodium oxide = _________________
o
Magnesium phosphate = ____________________
COVALENT BONDING: sharing of one or more pairs of electrons to achieve inert gas configuration.
 Single covalent bonds: two shared electrons (a single pair)
 Multiple covalent bonds:
o Double bonds: 4 shared electrons (two pairs)
o Triple bonds: 6 shared electrons (three pairs)
 Coordinate (dative) covalent bonds: formed when both electrons of the shared pair of electrons originate from
the same atom.
o Example: carbon monoxide (CO)
Lewis Structures: (a.k.a. electron dot structures) show all valence electrons
 Examples: All of the following are acceptable ways of representing the diatomic fluorine molecule.
Steps for Writing Lewis Structures:
1. Sum the valence electrons from all the atoms. Do not worry about keeping track of which electrons come from
which atoms (unless you are looking for dative bonds). It is the total number of electrons that is important.
2. Draw a skeletal structure. Use a pair of electrons to form a bond between each pair of bound atoms.
Hint #1: The atom with the smallest electronegativity is usually the central atom (H2O is a notable exception).
Hint #2: Polyatomic species are usually clumped and not spread out.
3. Arrange the remaining electrons to satisfy the duet rule for H and the octet rule for the second-row elements.
4. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d
orbitals (elements in Period 3 or beyond, often the central atom).
Lewis Structures: Comments about the Octet Rule
 The second-row elements C, N, O and F should ___________________________ obey the octet rule.
 The second-row elements B and Be often have _______________________________________________
electrons around them in their compounds. These electron-deficient compounds are very reactive.
 The second-row elements _________________________________________________________________,
since their valence orbitals (2s and 2p) can accommodate only eight electrons.
 Third-row and heavier elements often satisfy the octet rule but _______________________________________
___________________________________________________by using their empty valence d orbitals.
3
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Part 1 Practice Problems: Ionic and Covalent Bonding
Complete the following table of ionic compounds.
Name
1) barium hydroxide
Cation
Anion
Al3+
Cl-
Formula
2) sodium hydrogen sulfate
3)
4)
Sr(NO3)2
5) iron (II) oxide
6) iron (III) oxide
Draw Lewis structures for the following species:
7) O2
11) CH4
15) SF6
8) N2
12) C2H4 (ethane)
16) XeF4
9) CO2
13) C2H2 (ethyne)
17) PF6-
10) HCN
14) PCl5
Draw Lewis structures for the following species and indicate the coordinate (dative) covalent bond with a half-arrow
instead of just a line:
18) NH4+
19) H3O+
20) CO
4
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
PART 2: Shapes and Polarity
Bond Length and Bond Strength
• The strength of attraction that the two nuclei have for the shared electrons affects both the
_______________________ and ________________________ of the bond.
• Although there is a great deal of variation in the bond lengths and strengths of single bonds in different
compounds, double bonds are generally much ________________ and ________________ than single bonds.
• The strongest covalent bonds are shown by _____________________________ bonds.
Examples:
Bond
Cl – Cl
C–C
C=C
O=O
C≡C
N≡N
Bond Type
Length (nm)
Strength (kJ mol-1)
0.199
242
0.154
348
0.134
612
0.121
496
0.120
837
0.110
944
Note: while strength increases and length decreases from C – C  C = C  C ≡ C, the double bond is not twice as strong
as the single bond and the triple bond is not three times stronger (or shorter) than the single bond.
Valence Shell Electron Pair Repulsion (____________________) Theory
As the name implies, electron pairs in the outer energy level or valence shell of atoms repel each other and therefore
position themselves as far apart as possible.
The following points will help you apply the VSEPR Theory to predict molecular shapes:
• ___________________________________ applies to both bonding and non-bonding pairs of electrons.
• Double and triple bonded electron pairs are orientated together and so behave in terms of repulsion as a single
unit known as a negative charge center (or ___________________________________________).
• The total number of ____________________________________________________ around the central atom
determines the geometrical arrangement of the electrons.
• Non-bonding pairs of electrons (_________________________________) have a higher concentration of charge
than a bonding pair because they are not shared between two atoms and so they cause more repulsion than
bonding pairs. The repulsion _______________________________ in the following order:
o bonding pair – bonding pair < lone pair – bonding pair < lone pair – lone pair
Examples:
BF3
SO2
H2O
BF4-
NH3
5
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Polarity
Bond Polarity:
• ____________________________________ - equal sharing of electrons; atoms of identical electronegativity
o i.e. diatomic molecules such as H2 and Cl2
• ____________________________________ - unequal sharing of electrons; different atoms & the more
electronegative atom exerts greater attraction for the shared electrons.
o We label electron-rich and –poor atoms with ______________________________charges, - and +.
Molecular Polarity:
• When the bonds are arranged geometrically such that one side of the molecule is more electron-rich and the
other side is more electron-poor, the molecule is considered __________________________.
• If the entire molecule is polar (or there is a significant polar region), then label this resulting net dipole with an
_______________________________.
Examples:
HF
H2O
CO2
CCl4
CH3Cl
6
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Part 2 Practice Problems: Shapes & Polarity
Complete the chart below:
Formula
Lewis Dot Structure
Total eregions
# Bonding
e- regions
# Lone
e- pairs
VSEPR shape
(ball-n-stick drawing)
Name of Shape
Bond
angle(s)
Polarity
(P/NP)
1) BeH2
2) BCl3
3) CH4
4) PCl5
5) SF6
6) SnCl2
7
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Practice problems continued…
Formula
Lewis Dot Structure
Total eregions
# Bonding
e- regions
# Lone
e- pairs
VSEPR shape
(ball-n-stick drawing)
Name of Shape
Bond
angle(s)
Polarity
(P/NP)
7) NH3
8) H2O
9) SF4
10) ClF3
11)
XeF2
12)
BrF5
13)
XeF4
8
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
PART 3: Hybridization and Delocalization of Electrons
Hybridization: a modification of the localized electron model to account for the observation that atoms often seem to
use special atomic orbitals in forming molecules. This is part of both IB and AP curricula.
Examples:
BeF2
BF3
CH4
H2O
PF5
NH3
9
Unit 4: Bonding
IB Topics 4 & 14
Valence electron pair
geometry
Linear
# of
orbitals
2
Trigonal planar
3
Tetrahedral
4
Trigonal bipyramidal
5
Octahedral
6
Hybrid orbitals
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Electron density diagram
Examples
 and  bonds
 In Hybridization Theory there are two names for bonds, _____________ and _______________.
 Sigma bonds are the ___________________ bonds used to covalently attach atoms to each other.
 Pi bonds are used to provide the extra electrons needed to fulfill _______________requirements.
 Every pair of bonded atoms shares one or more pairs of electrons. In every bond at least one pair of electrons is
localized in the space between the atoms, in a sigma () bond.
 The electrons in a ___________________ bond are localized in the region between two bonded atoms and do
not make a significant contribution to the bonding between any other atoms.
 In almost all cases, _________________ bonds are sigma (____) bonds. A ______________ bond consists of one
sigma and one pi () bond, and a triple bond consists of ____ sigma and ____ pi bonds.
o Examples:

A Sigma bond is a bond formed by the overlap of two hybrid orbitals through areas of maximum electron
density. This corresponds to the orbitals combining at the tips of the lobes in the orbitals.
10
Unit 4: Bonding





IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
A Pi bond is a bond formed by the overlap of two unhybridized, parallel p orbitals through areas of low electron
density. This corresponds to the orbitals combining at the sides of the lobes and places stringent geometric
requirements on the arrangement of the atoms in space in order to establish the parallel qualities that are
essential for bonding.
Sigma bonds are ________________ than pi bonds.
A sigma plus a pi bond is stronger than a sigma bond. Thus, a ________________ bond is
__________________than a __________________ bond, but not twice as strong.
When atoms share more than one pair of electrons, the additional pairs are in pi () bonds. The centers of
charge density in a () is above and below (parallel to) the bond axis.
Molecules with two or more resonance structures can have bonds that extend over more than two bonded
atoms. Electrons in pi () bonds that extend over more than two atoms are said to be delocalized.
o Example: Benzene (C6H6)
Delocalization of Electrons
• Delocalization is a characteristic of electrons in pi bonds when there’s more than one possible position for a
double bond within the molecule.
• Example: ozone (O3)
o
o
These two drawn structures are known as resonance structures.
They are extreme forms of the true structure, which lies somewhere between the two.
 Evidence that this is true comes from bond lengths, as the bond lengths for oxygen atoms in
ozone are both the same and are an intermediates between an O=O double bond and an O-O
single bond.
• Resonance structures are usually drawn with a double headed arrow between them.
• Note that benzene (C6H6) has six delocalized electrons. Since the p-orbitals overlap (forming three pi bonds,
every-other-bond around the ring) all six electrons involved in pi bonding are free to move about the entire
carbon ring.
__________________________________________________________________________________________________
Formal Charge: A concept know as formal charge can help us choose the most plausible Lewis structure where there are
a number of possible structures. This is not part of the IB curriculum, but it is part of the AP curriculum. This theory
certainly has its critics; however, it has been included in this section of the course as it may help you in determining the
most likely structure.
Definition of formal charge: (# valence e’s on the free atom) – (# valence e’s assigned to the atom in the structure)
11
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Rules Governing Formal Charge
• To calculate the formal charge on an atom:
1. Take the sum of the lone pair electrons and one-half the shared electrons. This is the number of valence
electrons assigned to the atom in the molecule.
2. Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom
to obtain formal charge.
• The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that
species.
• If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any
negative formal charges on the most electronegative atoms are considered to best describe the bonding in the
molecule or ion.
Examples: CO2
NCO-
__________________________________________________________________________________________________
Part 3 Practice Problems: Hybridization & Delocalization of Electrons
#1-13) Go back to the Part 2 practice problems table of shapes and label the type of hybridization in each structure out
to the right of the table.
#14-18) Draw all resonance structures for the species below. State the number of sigma and pi bonds, as well as the
type of hybridization, in each of these structures.
14) NO317) CH3COO-
15) NO2-
18) C6H6 (benzene)
16) CO32-
12
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
PART 4: Intermolecular Forces
Intermolecular Forces
 Dipole-Dipole
 Hydrogen Bonding (special case of dipole-dipole)
 London Dispersion Forces
 Ionic


_____________________________________: the forces that act between different molecules are called
intermolecular forces.
These are the forces that make solids and liquids.
Solids
 In solids the intermolecular attractive forces are strong enough not only to hold molecules close together but to
virtually lock them in place.
 Solids, like liquids, are not very compressible because the molecules have little free space between them. Often the
molecules take up positions in a highly regular pattern. Solids that possess highly ordered structures are said to be
crystalline.
Liquids
 In liquids the intermolecular attractive forces are strong enough to hold molecules close together. Thus liquids are
much denser & far less compressible than gases.
 The attractive forces in liquids are not strong enough, however, to keep the molecules from moving past one
another. Thus liquids can be poured, and assume the shapes of their container.



Many properties of materials, including their boiling and melting points, reflect the strength of the intermolecular
forces.
A liquid boils when bubbles of its vapor form within the liquid. The molecules of a liquid must overcome their
attractive forces in order to separate and form a vapor.
o The ___________________ the attractive forces, the ___________________ the temperature at which the
liquid boils.
Similarly, the melting points of solids increase with an increase in the strength of the intermolecular forces.
Kinds of Solids
 Ionic solids
 Covalent-network solids
 Metallic solids
 Molecular solids solids and liquids held together primarily by one or more of the following forces…
 Dipole-dipole
 Hydrogen bonded
 London dispersion force
o Colletively these forces are called ___________________________________________________
13
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Dipole-Dipole Interactions - the electrical attractive forces that exist between polar molecules.
 The attractive forces are stronger than the repulsive forces, so there is an overall attraction between the molecules.
Hydrogen Bonding
 Special case of dipole-dipole interactions
 Seen among molecules where H is bonded to a highly electronegative atom, such as N, O or F.
 Example: water
Diagram:

Why is methane a gas at room temperatures, yet methanol is a liquid?
London dispersion forces
 Dispersion forces- attractions are electrical in nature. In a symmetrical molecule like hydrogen, however there
doesn’t seem to be any electrical distortion to produce positive or negative parts.
 Example: consider a small symmetrical molecule, such as H2 or Br2.

The even shading shows that on average there is no electrical distortion.

However, the electrons are mobile. At any one instant they might find themselves towards one end of the molecule,
making that end (-) and the other end (+). This is called an instantaneous dipole.

An instant later the electrons may well have moved to the other end, reversing the polarity of the molecule.

This constant “sloshing around” of the electrons in the molecule causes rapidly fluctuating dipoles even in the most
symmetrical molecules.
This “sloshing” even happens in monatomic atoms --- noble gases, like helium which consist of a single atom.

14
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7

If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer
properly covered by electrons for that instant.

Imagine a molecule which has a temporary polarity being approached by one which happens to be entirely nonpolar just at that moment.
 (This is actually pretty unlikely, but it makes the diagrams easier to draw. In reality, one of the molecules is likely
to have a greater polarity than the other at that time, and so will be the dominant one.)


As the molecule approaches, its electrons will tend to be attracted by the slightly positive end of the other
molecule.
This sets up an induced dipole in the molecule, and it too becomes polar (at least for the moment).

An instant later the electrons in the left-hand molecule may well have moved to the other end. In doing so, they
will repel the electrons in the right hand one.

The polarity of both molecules reverses, but you still have attraction. As long as the molecules stay close to each
other the polarities will continue to fluctuate in synchronization so that the attraction is always maintained.

There is no reason why this has to be restricted to two molecules. As long as the molecules are close together
this synchronized movement of the electrons can occur over huge numbers of molecules.
An instant later, of course, you would have to draw a quite different arrangement of the distribution of the
electrons as they shifted about—but always in synchronization.






It’s important to understand that dispersion forces act between all molecules.
They are usually only important when they are the only force acting.
POLARIZABILITY: The ease with which the charge distribution in a molecule can distorted by an external electric
field is called its polarizability.
You can think of the polarizability of a molecule as a measure of the “_________________________________”
of its electron cloud; the great greater the polarizability of a molecule, the more easily its electron cloud can be
distorted to give a momentary dipole, which leads to stronger London dispersion forces.
In general, larger molecules tend to have greater polarizabilities because they have a greater number of
_______________________________ and their electrons are farther from the nuclei.
15
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7

Dispersion forces tend to ___________________in strength with increasing
___________________________________.



The shapes of molecules can also play a role in the magnitudes of dispersion forces.
Typically, the greater the surface area of the molecule, the greater the dispersion forces.
Example: Which will have the higher BP, pentane (C5H12) or 2,2-dimethylpropane (C5H12)?
Rules of Thumb…
 When the molecules have comparable molecular weights and shapes, dispersion forces are approximately
equal.
 In this case, differences in the magnitudes of the attractive forces are due to differences in the strengths of
dipole-dipole attractions, with the most polar molecules having the strongest attractions.
Challenge example: benzene v. toluene v. phenol




The properties of molecular solids depend not only on the strength of the forces that operate between
molecules but also on the abilities on the molecules to pack efficiently in three dimensions
Benzene is a highly symmetrical planar molecule. It has a higher melting point than toluene. The lower
symmetry of toluene molecules prevents them from packing as efficiently as benzene molecules. As a result,
the intermolecular forces that depend on close contact are not as effective, and the melting point is lower.
In contrast, the boiling point of toluene is higher than that of benzene, indicating that the intermolecular
attractive forces are larger in liquid toluene than in liquid benzene. (greater molecular weight … greater
dispersion forces)
For phenol, both the melting and boiling points are higher than those of benzene because of the hydrogen
bonding ability of the OH group in phenol.
Relative strengths of forces:
16
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
PART 5: Giant Covalent Structures, Metallic Bonding & Physical Properties
GIANT COVALENT STRUCTURES
Allotropes of carbon: allotropes occur when an element can exist in different crystalline forms.
Three of carbon’s allotropes are described below:
Allotrope
Graphite
Diamond
Structural
diagram
Buckminsterfullerene (C60)
Hybridization
Density
Conductivity
Appearance
Uses
Lubrican; pencils
Jewelry; ornamentation; used in
tools and machinery for grinding
and cutting
Reacts w/ K to make
superconducting crystalline
material; related forms are used
to make nanotubes for the
electronics industry; catalysts
and lubricants
Silicon and Silica
Like carbon, silicon is a _______________________ element and so its atoms have four valence shell
electrons. In the elemental state, each silicon atom is covalently bonded to four others in a
tetrahedral arrangement. This results in a giant lattice structure much like ______________________.
Silicon dioxide (SiO2), commonly known as ________________ or ____________________, also forms a giant covalent
structure. This is a similar ____________________________________ bonded structure, but here the bonds
are between Si and O. Each Si atom is covalently bonded to __________ oxygen atoms, and each O to
____________ Si atoms.
17
Unit 4: Bonding
IB Topics 4 & 14
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
METALLIC BONDING






The valence electrons in metals become _____________________________ from the individual atoms so that
metals consist of a close packed lattice of positive ions in a sea of _____________________________ electrons.
A metallic bond is the attraction that two neighboring ____________________
ions have for the delocalized electrons between them.
Metals are ____________________________ --- they can be bent and reshaped
under pressure.
They are also ________________________ --- they can be drawn out into a wire.
Metals are malleable and ductile because the close-packed layers of positive ions
can slide over each other without breaking more bonds than are made.
_________________________________ added to the metal disturb the lattice and make the metal less
malleable and ductile. This is why alloys are harder than the pure metals from which they are made.
TYPE OF BONDING AND PHYSICAL PROPERTIES
Melting and boiling points
 When a liquid turns into a gas the attractive forces between the particles are completely broken, so boiling point
is a good indication of the strength of _________________________________________ forces.
 When solids melt the crystalline structure is broken down, but there are still some attractive forces between the
particles.
 Melting points are affected by impurities. These _______________________ the structure and result in lower
melting points.
 Covalent macromolecular structures have extremely ____________________ m.pts. and b.pts.
 Metals and ionic compounds also tend to have relatively __________________ b.pts. due to ionic attractions.
 H-bonds are in the order of _______________ the strength of a covalent bond.
 London dispersion forces are in the order of less than ____________________ of a covalent bond.

The weaker the attractive forces, the more _______________________________ the substance.

Intermolecular forces will increase with
o Increasing __________________________________________________
o
The extent of ____________________________ within the bonds of the structure
Example: diamond
Example: sodium chloride (NaCl)
18
Unit 4: Bonding
Example: predicting relative m.pts.
Compound
Lewis structure
IB Topics 4 & 14
Propane (C3H8)
AP Chapters (Zumdahl): 8 (sections 1-3;6-7;9-13); 9.1 & 9.5; 10.1-10.7
Ethanal (CH3COH)
Ethanol (C2H5OH)
Mr
Polarity
Intermolecular
bonding type
Melting Point (°C)
Solubility
 “_________________________________________________________”
o Polar substances tend to dissolve in __________________________ solvents, such as H2O.
o Nonpolar substances tend to dissolve in non-polar solvents, such as heptane or tetrachloromethane.
 Organic molecules often have a _____________________________ and a nonpolar carbon chain ____________.
 As the nonpolar carbon chain length increases in a homologous series, the molecules become ______________
soluble in water.
 Ethanol (C2H5OH) is a good solvent for other substances as it contains ______________ polar and nonpolar ends.
 Example: Put the following substances in order of decreasing solubility in water: methanol, butanol, propanol,
ethanol.
Conductivity
 Electricity = _________________________________________________________________________________.
 For conductivity to occur, the substance must possess electrons or ions that are _________________________
___________________________________.
 Metals (and graphite) contain delocalized electrons and are excellent __________________________________.
 Molten ionic salts also conduct electricity, but are chemically __________________________________ in the
process.
 When all electrons are held in fixed positions, such as in diamond and in simple molecules, _________________
_________________________________________________________.
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