Tests are not graded yet
Turn in your project up front and work on
warm up:

Write the molecular formula for:
Trinitrogen hexoxide
Aluminum nitride
Copper (II) sulfate
Write the names for:
NO2
PCl3
CaI2
Chapter 8

Covalent compounds consist of what?
 Only nonmetals

When naming, we use …
 Prefixes: mono, di, tri, tetra…
 Prefix = number of atoms (subscript)


N2O7
SF6

Why are there no charges (like in ionic
compounds)?
 In ionic compounds, electrons are
_______________, so atoms gain or lose charge
 In covalent compounds, electrons are
_____________, so no charges are formed

What does the octet rule state?
 In order to be stable, an atom wants a full outer
shell (which generally means 8 valence electrons)
 Which nonmetal is the exception to this rule?
 Which group do all elements want to be like?

When neither atom wants to give up
their electrons, they will just share



Electronegativity
When 2 electrons are shared between
atoms, they form a single bond
When 2 or more atoms bond
covalently, this is called
a molecule


Consider ionization energy and
electronegativity– when 2 elements
are near each other on the periodic
table, these values will be very near
each other
Ionization energy


Energy required to remove an electron
Electronegativity

How well an element attracts electrons in a
bond

If both atoms have very similar strengths
(for holding on to their electrons) then….
 Neither one will be strong enough to take electrons
away from the other


Lewis structures – using electron dot diagrams,
shows the arrangement of the atoms in a
molecule
How many valence electrons does carbon
have? How many more electrons does it
need to be “happy”?





How many times do you think carbon will bond?
How about hydrogen? Oxygen?
Generally, the # of “missing” electrons will equal
how many times an element will bond
CH4
CCl4


Calculate the number of valence electrons
Arrange the atoms in the molecule
○ Generally, the atom you have one of will go in the
middle
○ Hydrogen only bonds once, bonds on the outside
○ How many times will carbon bond? Oxygen? (look
at their valence electrons)




Put pairs of electrons between the central
atom and all of the outer atoms
Put electrons to fill the central atom
Put remaining electrons around outer atoms
Check to see that every atom is “happy”

PH3

H2S

SiH4


When 2 electrons are shared between
atoms, you draw a line to show the bond
All other electrons that are not shared are
called lone pairs and are included in the
structure


Single covalent bonds are also called sigma
bonds
Orbitals – the area where you will most likely
find an electron


How many electrons per orbital?
When these orbitals overlap, they form a
sigma bond (σ)


Let’s try carbon dioxide…
Sometimes, atoms may share more than 2
electrons
 If 4 electrons are shared, how many bonds would
there be?
 This is called a double bond
 How many electrons would a triple bond share?

Double or triple bonds consist of sigma and
pi bonds (π)




Draw: O2
N2
F2
What do you notice about the bonds?
Bond length : the distance between two
bonding nuclei
Which of these 3 do you think would have
the shortest bond length?
Warm up:
Draw the Lewis structures for the following:
C2H6
C2H4
C2H2

Keep in mind how many times each
element wants to bond




As the number of bonds increases, the
bond length becomes shorter
Which bond would be the strongest?
Bond dissociation energy : energy required
to break a bond in a molecule
What is the relationship between bond
length and bond dissociation energy?

Shorter bonds = more energy


In chemical reactions, bonds are broken and
formed
Breaking bonds _____________ energy


Forming bonds _____________ energy


Gives off (Aladdin)
If more energy goes in, then it is _______________


Requires (breaking a stick)
Endothermic
If more energy is given off, then it is ___________

Exothermic

PO43- what is this called?
When an ion has a charge, that means it has lost or
gained ______________
What has phosphate done?
Start the lewis structure like we did for the
others – add up all valence electrons
Now we have 3 extra electrons

ClO4-

NH4+

CO32-

H3O+

sulfite

H2SO4

CH3OH

HCN
Warm up:
Name and draw the Lewis structures for the
following compounds
H3P
CS2
N2H2





H – 1 time
O – 2 times
N – 3 times
C – 4 times
Lowest electronegativity element goes in
the center

Look at the word…



Molecules that contain how many atoms?
My fish’s name will help you know these
In nature, when these elements are not
bonded to another element, they like to
exist with 2 of themselves. They are more
stable that way.

What does it mean when something
resonates?


To vibrate or sound, especially in response
to another vibration
Resonance structures are different
ways to draw Lewis structures for a
molecule or ion
 Only the arrangement of the electrons is
changed

Let’s draw the structure for NO3-

How many resonance structures do
each of these have?

O3

NO2-

SO2

CCl2O


Sometimes an atom may not obey the octet
rule
Odd number of valence electrons (NO2)
 Fulfill the octet of the “outer” atoms

Less than 8 electrons present around an atom
(BH3)
Compounds with Be or B
 Tend to be very reactive
 Coordinate covalent bond – when one atom
donates both electrons in a shared pair (BH3 + NH3)



Draw the Lewis structure for SO3 and draw
its resonance structures
Draw the Lewis structure for ClF3


Expanded octet: happens with elements in
period 3 and below – d orbital electrons
can hold more than 8
Generally, the central atom gets the extra
electrons
 PCl5
 SF6
 Let’s look at H2SO4 again
 The S-O bonds have been experimentally
determined shorter than single bonds

ClF5


ICl4 -1


More than an octet on iodine
BeH2


More than an octet on chlorine
Less than an octet - Beryllium and boron
generally follow the less than 8 exception
NO

Odd number of valence - Nitrogen generally
takes the odd number of electrons

Draw the Lewis structures for ammonia
(NH3) and the ammonium ion



The hypothetical charge on an atom in a
covalently bonded molecule
Helps to determine the best Lewis structure
Want to keep the formal charge low – most
stable structure
FC = (# valence e-) – [(# of bonds) + (# of unshared e-)]
In a molecule, the sum of the formal charges (for every
atom in the molecule) is zero
In a polyatomic ion, the sum is equal to the charge


Use the structures for NH3 and NH4+ from the
warm up
Determine the FC for each nitrogen and
hydrogen in both structures


Draw the structure for NOCl


Write the value next to the atom; if there is no
number, it is understood to be zero
There are 2 possibilities, one is more preferred
Draw the structure for sulfate
Draw the structures and determine the FC
for each atom
Cl2O

SO2
AsF3


Valence Shell Electron Pair Repulsion – used
to determine the shape of a molecule
What determines how a molecule will
arrange itself?



What part of the atom are we generally
concerned about?...
ELECTRONS
Something to keep in mind: lone pair
electrons occupy more space than bonded
electrons



On a separate sheet, draw the Lewis
structures for each of the compounds on
the handout
Let’s see how many bonded pairs there
are, and how many lone pairs on the
central atom there are
Don’t fill in the picture column or angle
column yet
Linear
180o
Bent
104.5o
Trigonal planar
120o
Tetrahedral
109.5o
Trigonal pyramidal
Trigonal bypramidal
107.3o
90o/ 120o
Octahedral
90o

If the bond is not lying in the plane, then
you use either dashes or wedges


When electrons are bonded, think of them
as “trapped” between the 2 atoms,
therefore occupying less space
Lone pairs occupy more space, therefore
causing the bonded electrons to repel
(and bend the molecule)








NCl3
OCl2
HOF
NHF2
CO2
H2Se
CH2O
NH4+1




Pick one of the VSEPR shapes and build a
molecule
Include: label the type, an example of a
specific molecule (none that are on the
table), the angle between the atoms,
represent lone pairs (if there are any)
Use anything you would like to build this –
no drawings, and the model must be an
accurate representation of the shape
Due next Wedn. Feb 10th




Hybrid – when 2 things combine and have
properties of both
When atoms bond, they want to arrange their
orbitals to have lowest energy possible
Hybridization – describes the arrangement of
the orbitals
Hybrid orbitals – combined orbitals;
intermediates between orbitals

between s and p lies the hybrid orbital sp


Draw the orbital diagram for Carbon
From this, it looks as if there are only 2
places for electrons from another atom to
pair up (in the p orbital), but how many
times does carbon like to bond?
sp3
Write the formulas for the following
compounds:





Aluminum sulfate
Iron (III) phosphide
Hydronitric acid
Nitrous acid
Dicarbon trisulfide
Regions of high edensity
VSEPR shape
Hybridization
2
Linear
sp
3
Trigonal planar
sp2
4
Tetrahedral
sp3
5
Trigonal bipyramidal
sp3d
6
octahedral
sp3d2
• When giving the hybridization, you are generally talking about the
hybridization for the central atom

Generally, the # of things you are bonded
to = the number of hybrid orbitals


Lone pairs(on the central atom) occupy
hybrid orbitals as well


Bonded to 2 things = sp
Ex: draw the Lewis structure for water
Those 2 lone pairs count towards the hybrid
orbitals, so water is sp3








NCl3
OCl2
HOF
NHF2
CO2
H2Se
CH2O
NH4+1


If something is polar, it means it has opposing
ends
Need to know electronegativity and shapes

Influenced by the electronegativities
of atoms in a molecule




What is electronegativity?
An atom’s attraction for electrons when in
a bond
What is the trend for electronegativity?
(remember shielding and nuclear strength)
Who has the highest electronegativity
value?


Ionic: Look at the electronegativities of
Na and Cl – who has more attraction
for the electrons?
Covalent: look at the values for the
nonmetals
Polar covalent – unequal sharing of the
electrons in a bond
 Nonpolar covalent – equal sharing of
electrons in a bond

Electronegativity Difference
Bond Type
Less than 0.4
Nonpolar covalent
0.5 to 1.9
Polar covalent
Greater than 2.0
Ionic



What kind of bond would carbon and
oxygen form?
Phosphorus and fluorine?
Chlorine and chlorine?

Draw the Lewis structure, determine the
shape and hybridization for the following:
BF3
SF4
PF6-

Draw the Lewis structure for water

What is water’s shape?
Who is stronger?
Who will the electrons be closer to?

This makes partial charges.




Draw carbon tetrachloride and label
the partial charges
Compare carbon tetrachloride’s
structure to water’s



Polar molecules are asymmetric, while
nonpolar are symmetrical
Which one of these would you consider
symmetrical?
You have to look at the polarity of
each bond, and look at the overall
molecule to determine if it is polar
Determine if the following molecules/ion
are polar:
NCl3
If the bonds are polar, it could be
H2S
polar or nonpolar, check the
structure
CS2
SF6


Solubility (what is this?) is determined by
polarity

What is the universal solvent?

Are most substances polar or nonpolar?
Determine the more polar molecule in each pair:
methyl chloride (CH3Cl) or methyl bromide (CH3Br)
water
or
hydrogen sulfide (H2S)
hydrochloric acid
or
hydroiodic acid
boron trihydride
OR
ammonia
silicon tetrabromide
OR
HCN

What were the properties of ionic
compounds in terms of conductivity,
melting point and solubility?


What are properties of covalent?


High melting point, conducts (when
dissociated), and soluble in water (meaning
ionic compounds are what?)
Many covalent compounds exist as liquid or gas
Which type is more strongly held together?

What are intermolecular forces? (interstate)

Forces that hold one molecule to another
3 Types:
Hydrogen bonding
Dipole-dipole
Dispersion/London forces

In the solid/liquid state (not concerned with
gaseous state – why?)


Dipole – contains oppositely charged
regions (partial charges)
Results from the attraction between the
partial positive end of one molecule and
partial negative end of another molecule

Also known as induced dipole forces



animation
Occur between nonpolar molecules with
no permanent dipoles
Result from a temporary shift of electrons,
and dipoles are instantaneously created

Ex. 2 chlorine molecules

Occur between hydrogen and O, N or F


Due to their high electronegativities it makes H
more partially positive
Causes these compounds to have higher
boiling points
What is the strongest intermolecular force present for
each of the following compounds?
1) water
2) carbon tetrachloride
3) ammonia
4) carbon dioxide
5) phosphorus trichloride
6) nitrogen
7) ethane (C2H6)
8) acetone (CH2O)
9) methanol (CH3OH)
10) borane (BH3)
1) water
2) carbon tetrachloride
3) ammonia
4) carbon dioxide
5) phosphorus trichloride
6) nitrogen
7) ethane (C2H6)
8) acetone (CH2O)
9) methanol (CH3OH)
10) borane (BH3)
hydrogen bonding
London dispersion forces
hydrogen bonding
London dispersion forces
dipole-dipole forces
London dispersion forces
London dispersion forces
dipole-dipole forces
hydrogen bonding
dipole-dipole forces
Grab a chemistry book, and work on the
following questions –
p. 274
83, 85, 89, 96, 98, 101, 108, 112, 114,
120, 127

Be sure to look through my powerpoints and
study guide on my website

Name the following compounds:
ZnCl2

KNO3

H2S

NF3


Name and draw the Lewis structures for the
following compounds:

CS2

PH3

CCl4







Write the formulas for the following
compounds:
Aluminum sulfate
Iron (III) phosphide
Hydronitric acid
Nitrous acid
Dicarbon trisulfide
Go ahead and take out the worksheet with
the PT with electronegativities from
yesterday

Name the following acids:
H3N
H3SO3
H2Se