Honors Chemistry
Chapters 7 and 8 Notes
(Student edition)
Chapter 7 problem set:
Chapter 8 problem set:
53, 60, 67, 73, 83
40, 43, 45, 47, 54, 56, 58, 61, 63, 65-68, 74, 79, 80
Useful diagrams:
This book has excellent figures and tables – most of them have
something to offer you intellectually.
NIB – An introduction:
-
Attachment Between Atoms
in nature. They are “
Most elements are not found
” to other atoms.
Chemical Bond –
-
8.4
Types of chemical bonds:
Ionic Covalent Metallic -
Polar Bonds and Polar Molecules (part 1)
-
Polar Nonpolar –
sharing of electrons
sharing of electrons
-
There are two ways to predict polar vs nonpolar ( and covalent vs ionic)
#1 Use electronegativity difference
nonpolar covalent
0.4 - 1.7 - polar covalent
greater than 1.7 - ionic
Examples:
NaCl
HCl
Cl2
#2 - There is an easier way to predict
Ionic =
+
Polar Covalent =
Nonpolar Covalent =
or
+
1
7.1
and 7.2 Ions and Ionic bonds and Ionic Compounds
Ionic compound - a substance composed of positive and neg. ions
So that the charges are
. It involves a
electrons.
(Draw crystal lattice below)
of
Ca+2 with Cl–1 will form the compound CaCl2. It takes
chlorine ions to cancel out the the +2 charge on the calcium ion.
Formula unit Metals Nonmetals -
electrons - why?
They form cations.
electrons - why?
Metals lose electrons until they become like a
Nonmetals gain electrons until they do the same.
They form anions.
. (usually 8 valence electrons)
Both go to s2p6 - 8 valence electrons - called a
The tendency to arrange electrons so each atom has 8 is called the
.
The formation of an ionic bond:
an easier way.....
The ionic bonding picture looks like this....
Other examples of ionic bonding pictures:
2
NIB
-
Ions make up ionic substances - not atoms
-
Crystal Lattice - regular repeating pattern of
-
Ions are held in fixed position by
-
A large amount of energy is needed to break this structure. Because of this, ionic
compounds have a high
.
-
Ionics can’t conduct electricity as
-
Property summary of ionic compounds - hard, shatter (not
when in
, high melting point, no smell (low
in an ionic substance
.
, but they can when dissolved in
.
), conduct
)
(although this covered somewhat in section 8.2)
-
Energy is involved in all chemical reactions.
Na + Cl
-
yields
NaCl
+ 769 kJ
Lattice energy - energy released when an
NaCl = - 769 kJ/mole
smaller ions have
forms.
NaF = - 922 kJ/mole
KCl = -718 kJ/mole
LE
Exothermic reactions are those where heat is released. We use a (-) (negative) sign to show that
heat is given off when a value is reported outside of a chemical equation. Exothermic reactions
have their heat recorded as a positive value on the right hand side of the arrow when written in a
chemical equation.
3
7.3
Bonding in Metals
-
“Sea of electrons theory” Momma Fedell and her big wooden spoon…..
Cu+
Cu+
Cu+
Cu+
2
2
2
2
Cu+
Cu+
Cu+
Cu+
2
2
2
2
Properties of metals alloy - metallic solution
amalgam - mercury solution
alloys are important because they have properties that are better than the properties of the
constituent elements (an example is steel)
8.1 and 8.2
Molecular Compounds and The Nature of Covalent Bonding
Molecule - smallest quantity of matter that can exist by itself and still retain the properties of that
substance. Usually used when describing a covalently bonded substance.
Listed below are different types of elements and the types of molecules these elements form.
These terms do not apply to compounds.
-
monatomic molecules -
-
diatomic molecules -
-
polyatomic molecules -
-
chemical formulas show the relative #’s of atoms in a chemical compound
ex.
C6H12O6
Pb(NO3)2
(NH4)2Cr2O7
C=
Pb =
N=
, H=
, N=
, H=
, O=
, O=
, Cr =
, O=
4
-
The formation of a covalent bond:
→
Bond Length vs. Bond Energy :
Relationship (generally) Bond length =
-
Bond Energy
Diatomic Molecules and Orbital Notation (Orbital overlap or notation diagrams):
H2
1s
1s
O2
1s
2s
2p
1s
2s
2p
1s
2s
2p
1s
2s
2p
N2
- Atoms lose, gain, or share electrons in order to have 8 electrons in their outer shell.
-
Electron cloud representations:
F2
HCl
- HF - orbital notation diagram
H
1s
F
1s
2s
2p
5
Dot Diagrams of molecules and polyatomic ions
-
Lewis Structures for covalent compounds
-
Basic rules (try to use as many as the rules as possible)
1.
2.
3.
4.
Each atom wants 8 electrons (except H wants 2).
Each atom goes for close to the right # of bonds.
The least electronegative atoms goes in the middle.
The atom that makes the most bonds goes in the middle. H always on
the outside.
5. The single atom (the atom that does not have a subscript after it) goes
in the middle.
6. Attempt to make the structures as symmetrical as possible. Place the
atoms in order (left, right, bottom, and top) around a central atom.
-
Examples: Draw the following Lewis structures
H2O, PCl3, SiH2F2, CS2, C2H6, C2H4, C2H2, CH2O, HCN, FON
6
Drawing polyatomic ions
count electrons - if the charge is - 3, add 3 electrons
example PO4-3
p-5, O- 6x4=24, total = 29 but need to add 3 = 32
less bonds than atoms want = negative charge
more bonds than atoms want = positive charge
Ckbe - check bonds, check electrons
P wants 3 bonds, has 4 - + 1 charge Each O wants 2, has 1 - so each O = -1 Total = - 3
Draw NH4+
- 2 shared electrons in a
bond are donated by 1 atom
Draw OH-1, sulfate, nitrate, nitrite, carbonate, bicarbonate, H2SO4, H3PO4
7
Exceptions to the Rule of 8
resonance - refers to bonding in molecules that can’t be represented by 1 Lewis structure
ex.
SO2
originally it was thought that electrons in the second bond of a double bond “resonate”
back and forth between on both sides of the sulfur atom. Experiments show that each bond is
equal so the electrons must be evenly shared
other examples: O3, SO3
Exceptions to the rule of 8
Group 13 - Boron is small - can only accommodate 3 pairs of electrons
BH3 violates the octet rule
Groups 15, 16, 17 can find themselves in “expanded” valence shells
expanded means that the shells hold more than 8 electrons
examples:
SF6
PCl5
XeF4
Some molecules have odd #’s of electrons
NO - 15 total electrons - 11 valence electrons
NO2 - 17 valence electrons
related idea - Paramagnetism - atoms that interact weakly with a magnetic field
This results from unpaired electrons - magnetic fields cancel out from opposing spin
electrons. D block elements tend to be paramagnetic. These unpaired electrons also
cause the coloring of these compounds.
8
Hybridization
Carbon’s full electron configuration is: 1s22s22px12py1
carbon’s orbital notation:
1s
2s
2p
It looks like it wants to make 2 bonds. However, experiments show it makes 4 identical
bonds.
Hybridization - Rearrangement of electrons within the valence orbitals of atoms in a
chemical reaction.
Carbon will hybridize and have an electron configuration of:
1s22s12px12py12pz1
These orbitals have different character than just “s” or just “p”.
They form a new type of orbital called the sp3 hybrid. Carbon has four sp3 orbitals.
Carbon’s hybridized orbital notation:
1s
2sp3
also happens in Be
example - BeH2 - violates the octet rule - sp hybrid
also happens in B
example - BF3 - violates the octet rule - sp2 hybrid
9
8.3 and 8.4 Bonding Theories and Polar Bonds and Molecules
Shapes of Molecules - VSEPR Model - Valence shell electron pair repulsion theory (vesper) electron pairs get as far away from each other as possible
Example
Shared
pairs
Lone pairs
Shape
Bond Angles
Drawing/example
AB
AB2
AB2
AB2
AB3
AB3
AB4
AB5
AB6
to predict - structure must be drawn first - examples: CCl4, H2S, HBr, SO2, ClO4 -1, BH3, PF5, etc.
10
Now that we know about the shapes of molecules, we can learn more about why they have
certain physical properties….
Intramolecular forces –
- happens within a molecule or compound
bonds (ionic, covalent, metallic)
- always
Intermolecular forces (IMF)- forces that hold molecules together
- happens between
compounds
- can be
TYPES OF INTERMOLECULAR FORCES:
1. dipole-dipole
dipole
-
when electrons are
distributed
predict the IMF that occurs with HCl
predict the IMF that occurs with H2S
predict the IMF that occurs with CO2
2. Hydrogen Bonding: H-bonding is a “super” dipole-dipole
predict the IMF that occurs with HF
Are we having FON? - H-bonding happens any time H is bonded to
Why? A large difference in
one end of the molecule being very
.
between F, O, or N and H results in
, while the other end is very
.
Why N and not Cl? They have identical electronegativities
Well, N is so much
than Cl so the negative charge is spread over a
smaller area which exerts more force.
Effect of H-bonds on physical properties:
H-bonding tends to cause the following in substances:
Boiling Point
Heat of Vaporization
Also…
Vapor Pressure
Melting Point
H-bonds causes water to expand when it freezes (swimming pool bags, density,
polar bears, hockey players). H-bonding is also responsible for the shapes of
proteins (curly hair vs straight hair). H-bonding graph (boiling pt. vs. increasing
molecular mass)
11
Molecular Substances
Previously, we learned polar molecules have dipole-dipole IMF holding them together.
Previously we learned about H-bonding or “super-duper” dipole-dipole IMF.
These two types of IMF usually result in substances being
at room temp.
Most nonpolar covalent substances are gases at room temp. as the forces holding them together
are not strong enough to keep the molecules attracted - hence they are gases.
O2, H2, N2 CO2 - has dipoles, but nonpolar due to its
now, a third type of IMF:
3. Van der Waals Forces (London Forces) - Temporarily induced dipoles caused by
the
.
He
CO2
more electrons =
Thus, bigger atoms have stronger Van der Waals forces
Network Solids
-
They are covalent
or covalent
-
Examples:
-
They form do not form separate, distinct, molecules. It is one continuous
-
Properties:
-
They make good
and
.
.
.
12
NIB
Summary
- see, chemistry makes sense after all!
IMF
Molecule Type
Noble Gas
Noble Gas
Nonpolar
Nonpolar
Nonpolar
Polar
Polar
Ionic
Ionic
Metallic
Metallic
Molecule
He
Ar
H2
O2
Cl2
HF
ICl
NaCl
MgF2
Cu
Fe
Boiling Pt. (Co)
-269
-186
-253
-183
-34
19.5
97
1413
2237
2567
2750
Bond Energy
- basic idea - what is the strength of chemical bonds?
-
bond energy - energy needed to
-
bond strength and stability:
a bond - measured in kJ/mole
stronger bond - more stable -needs more energy to break the bond
weaker bond - takes little energy to break the bond so the chemical is
-
NIB
chemical changes favor
favored
energy states -
.
reactions are
Bond Strength
- which is stronger? - single, double, or triple bond? uh.....
-
which is shortest bond length? s, d, or t? - the answer is
-
which is stronger, short or long bonds? uh...
13