Honors Chemistry
Name: __________________________________________ Date: _____________ Mods: _________
Ch. 8: Concepts of Chemical Bonding
1) WHY and HOW do Atoms form Chemical Bonds?

WHY – Atoms have __________________ chemical potential energy when they are bonded to
other atoms than when they are independent particles and this is a GOOD thing!
o

Matter wants to be in its LOWEST POSSIBLE ENERGY STATE because they are
_________________ stable and _____________ reactive
HOW – Only the ________________ (outermost) electrons of atoms are involved in bonding (never
inner core electrons)
o
These electrons are the furthest from the nucleus and least attracted to the atom making it
easiest to remove/attract away from the atomic nucleus
o
In a chemical bond, as atoms get ______________ to one another, there are both attractive
and repulsive forces at work:

Attractive Forces: the nucleus of one atom is attracted to the __________________ of the
other atom (causes a decrease in total potential energy = good)

Repulsive Forces: the two sets of
_______________ and/or two sets of
electron clouds begin to repel one
another (causes an increase in total
potential energy = bad)

Bond Length: the ideal ________________ between bonded atoms in which attractive
forces are maximized and repulsive forces are minimized (this distance depends on the
identities of the atoms bonding)

Bond Strength: the more attracted two atoms are to one another, the greater the bond
strength (this can be quantifiably measured by calculating the bond _______________ (aka:
bond energy)


Bond Enthalpy: the energy required to ________ a chemical bond (unit: kJ/mol)
Relationship: As bond length ___________________, bond strength and bond
enthalpy __________________.
Chemical Bond
Cl-Cl
Br-Br
H-Cl
H-Br
Bond Length (picometers)
199
229
127
141
Bond Enthalpy (kJ/mol)
243
193
432
366
2) Ionic Bonding: nonmetal atom _________ electrons from
atom due to a _____________ difference in
electronegativity (∆EN > 1.7).
This complete ____________________ of electrons
in a cation and anion which are then electrostatically
attracted to one another (opposite charges) resulting in a
very strong bond between the ions.
a metal
results

Example:

Ionic compounds have a particular arrangement known as a crystal lattice structure; this structure
organizes the cations and anions so the surround one another giving the compound its lowest
potential energy (most stable arrangement)
o
Lattice Energy: the energy ________________when 1 mole of an ionic crystalline compound is
___________________ (unit (kJ/mol)

While bond enthalpy is energy required to break a chemical bond, lattice energy is energy is
energy released when 1 mole of ionic bonds are formed

The strong attraction of (+)/(-) ions is described by the lattice energies in the figure below. All
lattice energies are __________________ values (aka: energy is released)

Compound
NaCl
NaBr
CaF2
LiCl
Lattice Energy (kJ/mol)
-787.7
-751.4
-2634.7
-861.3
LiF
MgO
KCl
-1032
-3760
-715

Which ionic compound in the table
has the strongest ionic bonding?
________________

Why do you think that is?
The higher more negative) the lattice energy, the ___________________ the ionic bond,
and the more stable the compound.
3) Covalent Bonding: two nonmetal atoms ____________
electrons in order to achieve their octets since nonmetals
typically ___________ electronegativities
have

Polar Covalent Bonding
o

Unequal sharing of electrons between two nonmetal atoms due to an
___________________________ difference in electronegativity (∆EN = 0.4-1.69).

This means that the electrons in a polar bond are pulled closer to (and spend more time
near) the _____________ EN atom

Example:
Nonpolar Covalent Bonding
o
Equal sharing of electrons between two nonmetal atoms with _______________________ or
very ____________________ electronegativities (∆EN < 0.39).

This means that no one atom attracts electrons better than the other

Example:
4) Metallic Bonding: metal atoms have typically __________
electronegativities and tend to ___________ electrons
forming _________________. When these metals bond
together, there is an electrostatic attraction between these
lost/delocalized electrons which forms a ___________ of
randomly moving electrons

Example:
5) Strength of Chemical Bonds

______________ bonds are the STRONGEST (due to the presence of ions with ________ charges)

_________________ covalent bonds are STRONGER than ________________ bonds (due to the
_________________ charges in the bond that result in polar bonds from large ∆EN)
Partial Charges & Polarity of Chemical Bonds- Examples
1
H
2.1
3
Li
1.0
11
Na
1.0
19
K
0.9
37
Rb
0.9
55
Cs
0.8
87
Fr
0.8
4
Be
1.5
12
Mg
1.2
20
Ca
1.0
38
Sr
1.0
56
Ba
1.0
88
Ra
1.0
21
Sc
1.3
39
Y
1.2
57
La
1.1
89
Ac
1.1
23
V
1.5
41
Nb
1.5
73
Ta
1.4
24
Cr
1.6
42
Mo
1.6
74
W
1.5
25
Mn
1.6
43
Tc
1.7
75
Re
1.7
26
Fe
1.7
44
Ru
1.8
76
Os
1.9
27
Co
1.7
45
Rh
1.8
77
Ir
1.9
28
Ni
1.8
46
Pd
1.8
78
Pt
1.8
29
Cu
1.8
47
Ag
1.6
79
Au
1.9
30
Zn
1.6
48
Cd
1.6
80
Hg
1.7
•
The more EN element in a polar covalent bond will attract the
electrons closer to itself, resulting in a partial ________________
charge on that atom
•
The opposite end of the bond (less EN atom) is thus deprived of
electrons near it, so it forms a partial ________________ charge
•
1.
22
Ti
1.4
40
Zr
1.3
72
Hf
1.3
5
B
2.0
13
Al
1.5
31
Ga
1.7
49
In
1.6
81
Tl
1.6
6
C
2.5
14
Si
1.8
32
Ge
1.9
50
Sn
1.8
82
Pb
1.7
7
N
3.0
15
P
2.1
33
As
2.1
51
Sb
1.9
83
Bi
1.8
8
O
3.5
16
S
2.5
34
Se
2.4
52
Te
2.1
84
Po
1.9
9
F
4.0
17
Cl
3.0
35
Br
2.8
53
I
2.5
85
At
2.1
NOTE: only ionic bonds have full charges (transfer of
electrons) and nonpolar bonds have NO charges (equal
sharing)
For each problem below, use an arrow () to show the direction of polarity in each separate bond.
Then, circle the more polar bond in each problem
a)
C—O
and
N—C
c)
B—O
and
B—S
b)
P—Br
and
P—Cl
d)
B—F
and
I—B
e)
C—N
and
C—H
f)
and
S—O
N—S
2.
Calculate the difference in electronegativity and classify the chemical bonds below as either ionic
(I), polar covalent (P), or nonpolar covalent (NP). Once the bond type is known, indicate on the
bond either the full charges, partial charges, or no charges.
Ionic
Polar Covalent
Nonpolar Covalent
X > 1.7
1.69 > X > 0.4
X < 0.39
Indicate Charges (full,
partial, or none)
1
Na–Cl
2
C–B
3
Cs–F
4
Cu–O
5
Cl–Cl
6
Mg–H
7
Al–Cl
8
C–H
(symbolized as A+ and Z-) – full charges
(symbolized as A+ and Z-) – partial charges
(no charges)
∆EN
Bond Type
(I, P, NP)
Honors Chemistry
Name: __________________________________________ Date: _____________ Mods: _________
Ch. 8: Rules for Drawing Covalently Bonded Lewis Structures
1) Sum the valence electrons from all atoms in the molecule or ion

For an ANION  add one electron to the total for each negative charge

For a CATION  subtract one electron from the total for each positive charge
2) Determine the central atom in the molecule and attach all other atoms to it

The central atom is the least electronegative element in the compound excluding hydrogen –
hydrogen can NEVER be central! (Note: typically, the central atom is the one written first in
the molecular formula)

Write the symbols for all the other atoms around the central atom. Connect the atoms to the
central with a single bond (a dash). Keep track of the electrons being used. Each single bond
made uses 2 electrons.

Chemical formulas are often written in the order in which the atoms are connected in the
molecule (ex: HCN  carbon is the central atom)
3) Complete the octets around all the outer atoms bonded to the central using nonbonding
electrons ( : )

Note: Hydrogen atoms may only ever have 2 electrons associated with them at any given time

Keep track of the electrons being used to complete the octets
4) Place any leftover electrons on the central atom, even if doing so results in more than an
octet of electrons around the atom.

Note: Expanded octets occur when more than 8 electrons surround an atom. This exception
to the octet rule is allowed for any atom in the 3rd row of the periodic table and after!

Keep track of the electrons – make sure the total number of valance electrons available were
used in the Lewis structure
5) If there are not enough electrons to give the central atom an octet, multiple bonds are
needed.
Single Bond (
):
Double Bond (
Triple Bond (
):
):

Remove one of the nonbonding pairs of electrons on one of the outer atoms and draw a
double bond (2nd dash) connecting the outer atom to the central. If need be, a triple bond (3rd
dash) may be formed by removing another nonbonding pair from the same outer atom.

Keep track of the electrons – make sure the total number of valance electrons available were
used in the Lewis structure
Drawing Covalent Lewis Structures – Examples
Lewis Dot Structures:
1) Methane: CH4
4) Sulfite ion: SO32-
2) Hydrochloric Acid: HCl
5) Hydrocyanic Acid: HCN
3) Ammonia: NH3
6) Methanal (aka: formaldehyde): H2CO
Resonance Structures:
Multiple Lewis structures are used to describe molecules which have double or triples bonds that
can be moved to different sides of the central atom and still stay bonded to an identical outer atom.
7) Nitrite ion: NO2 –
Exceptions to the Octet Rule:
Less than an Octet  central
atoms which have FEWER than
8 electrons (seen with B only)
8) BF3
Expanded Octet  central atoms which have MORE
than 8 electrons (this is allowed for central atoms
found in the 3rd period and after)
9) SBr6
10) XeI4
Drawing Lewis Diagrams - Practice
# of valence e-‘s:________
1. SF4
3. PO43-
5. N2
2. CH2F2
# of valence e-‘s:________
# of valence e-‘s:________
7. CO32-
# of valence e-‘s:________
4. CBr4 # of valence e-‘s:________
6. CCl2O
# of valence e-‘s:________
# of valence e-‘s:________
8. ClO21–
10. H2S
12. NH4+
# of valence e-‘s:________
# of valence e-‘s:________
# of valence e-‘s:________
11. C2H4 # of valence e-‘s:________
(attach carbons together – make symmetrical)
# of valence e-‘s:________
14. NO31-
9. OF2
13. BCl3
# of valence e-‘s:________
# of valence e-‘s:________
15. IBr5
# of valence e-‘s:________
17. C2H2 # of valence e-‘s:________
(attach carbons together – make symmetrical)
19. PCl3 # of valence e-‘s:________
21. PF5
# of valence e-‘s:________
16. SO42-
# of valence e-‘s:________
18. CO # of valence e-‘s:________
20. SiO2
# of valence e-‘s:________
22. PH3 # of valence e-‘s:________
23. XeF4
# of valence e-‘s:________
25. CN- # of valence e-‘s:________
27. C2Br2 # of valence e-‘s:________
(attach carbons together – make symmetrical)
29. SeF4
# of valence e-‘s:________
24. SeH2
# of valence e-‘s:________
26. IF3 # of valence e-‘s:________
28. ClO1– # of valence e-‘s:________
30. AsF3 # of valence e-‘s:________