9/30/2011
Ionic vs. Covalent Compounds

Ionic compounds – contain a metal


formula units
Covalent compounds – only non-metals

molecules
UNIT 3: BONDING
Covalent & Ionic
Covalent or Ionic?

H2O




ionic
HgSO4



covalent
NaCl


Ionic Compounds

ionic
PF4


covalent

monoatomic cations – name of element


chloride ion



NO PREFIXES


creates Na+
endothermic process
349 kJ is released when adding one mole of
electrons to one mole of Cl (electron affinity)

check memory work!
compound name = cation name + anion name
495 kJ is required to remove one mole of electrons
from one mole of Na (ionization energy)

polyatomic ions – see chart


Cl-

sodium ion
monoatomic anions – root of element name +
ide ending


Na+
charges for entire compound must add up to zero
Ionic Bonding – NaCl example
Ionic Nomenclature

EN difference is high (> 2.1)
electrons are transferred – NOT shared
cation – atom (or group of atoms) that loses
electrons
anion – atom (or group of atoms) that gains
electrons
#electrons lost = # electrons gained
creates Clexothermic process
This would be an overall endothermic process,
but…
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Ionic Bonding – NaCl example

Lattice energy – the amount of energy
required to separate ionic compounds
Naming Acids

Watch ending of anion

energy is required to separate the ions
 energy is released when ions get together




1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
-ate
becomes
root-ic acid
HNO3
 nitric acid

-ite
becomes
root-ous acid
HNO2
 nitrous acid

Net energy change is negative, more stable
salt name prefix-hydrate
CuClO3 5 H2O
copper chlorate pentahydrate
hydro-root-ic acid

Net energy = Ionization Energy – Electron
Affinity – Lattice Energy
Naming Hydrates
becomes
HCl
 hydrochloric acid


-ide


Valence Electrons


outer shell electrons involved in bonding
Remember “A” group numbers
6 – hexa
7 – hepta
8 – octa
9 – nona
10 - deca
Covalent Nomenclature

Ionic Bonds – transfer electrons
one or more electrons leave on atom and join
another
 a cation (positive) and an anion (negative) form
 cation sticks to anion like a magnet


Covalent Bonds – share electrons

MEMORIZE THIS!
1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – hepta
8 – octa
9 – nona
10 - deca
neither atom loses or gains electrons
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Naming
First element:
prefix (except mono) + name of element
Second element:
prefix + root of element name + ide ending
Name the following

SeI4

N2O3

CSe2

P2Cl5




Write formulas for the following

phophorus pentabromide

PBr5

disulfur trioxide

aresenic trifluoride

dichlorine monoxide



S2O3

dinitrogen trioxide
carbon diselenide
diphosphorus pentachloride
Diatomic Elements


elements not found alone in nature
memorize these! (think 7)
H2
N2
O2
F2
Cl2
Br2
I2
AsF3
Cl2O
Lewis Dot Structures

selenium tetraiodide
Drawings that show how electrons are shared
between atoms in covalent bonds
Use only valence electrons
Lewis Structures for Atoms
Na
C
B
S
I
Ne
H
He
Special
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Octet Rule



All atoms want 8 valence electrons
Atoms will bond to satisfy this rule
H and He only want two electrons!
 B is satisfied with 6 electrons
 Can have expanded octets
Drawing Lewis Structure for
Molecules
2.
3.
4.
5.
6.
7.
Find the total number of valence electrons
Determine the central atom
Place single bonds between atoms
Place lone pairs
Check octet rule & total number of electrons
If needed add double or triple bonds one at a
time
The best arrangement has the lowest formal
charge or the negative formal charge on the most
EN element
Example


Single bond

Double bond

Triple bond

Exceptions!

1.
Covalent Bonds
CH3OH vs CH2OH2


one pair of shared electrons
two pair of shared electrons
three pair of shared electrons
Determining Formal Charge





Cf = Ev – (Eu + ½ Eb)
Cf = formal charge
Ev = valence electrons
Eu = unshared electrons
Eb = bonding electrons
Practice Drawing Lewis Structures
CH4
NH3
H2O
CSe2
CO
HCN
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Polyatomic Ions

Resonance
Add or subtract the appropriate number of
electrons
NH4+

When more than one Lewis Structure is
possible
SO4-2
O3
Free Radical







Molecular Geometries
describe the 3-D shape of a molecule
VSEPR – valence shell electron pair repulsion
Draw Lewis Structure
Count bonds (doesn’t matter what type)
Count lone pairs (unshared pairs of electrons)
Molecular Geometries for Expanded Octets
Bonds
on central
atom
2
Lone Pairs
on central
atom
0
Angle
Hybridization
Geometry
180
sp
Trigonal
Planar
Bent
3
0
120
sp2
Trigonal
Bipyramidal
Seesaw
2
1
Tetrahedral
4
0
Linear
Trigonal
Pyramidal
Bent
3
2
CO3-2
Molecular Geometries
Sometimes there is an odd number of electrons
NO
Geometry
NO3-
1
2
109.5
sp3
Bonds
on central
atom
5
Lone Pairs
on central
atom
0
4
1
T-shaped
3
2
Linear
2
3
Octahedral
6
0
Square
Pyramidal
Square
Planar
5
1
4
2
Angle
Hybridization
120 , 90
sp3d
90
sp3d2
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Drawing 3-D Lewis Structures
_______
----------
(solid line)
(dotted line)
(wedge)
CO2
SO2
NH3
in the plane of the page
away from the reader
toward reader
CH2O


Depend on the amount of repulsion

Lone pairs > triple bond > double > single
H2O
Electronegativity
Molecules are polar if they are asymmetrical
Look for


CH4
Polarity of Molecules

Bond Angles
Lone pairs on the central atom
Differing atoms attached to the central atom
Polarity of Bonds
EN Difference


the ability of an atom to attract electrons
involved in bonding with another atom
values predict the type of bonding that will
occur
What type of bond is formed?
Type of Bond

KCl

< 0. 5
non-polar covalent
(even sharing)


NH

0.5 – 2.1
> 2.1
polar covalent
(uneven sharing)
ionic
(electrons transfer)


3.2 – 0.8 = 2.4
ionic bond
3.0 – 2.2 = 0.8
polar covalent bond
PS


2.6 – 2.2 = 0.4
non-polar covalent
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Dipole

A polar bond
Draw Lewis Structure & Indicate Net Dipole



Strength is called the dipole moment

Ex. H2O



A molecule can have polar bonds but be
symmetrical (non-polar) overall (CCl4)
Valence Bond Theory



BrF5
ICl4ClF3
AsF5
NH3
Sigma ( ) Bond
Combines Lewis structures and atomic orbital
theory
Orbitals from two different atoms overlap
Pi ( ) Bond

Overlapping two p orbitals



single bond = bond
double bond = + bond
Triple bond = 2 bond + bond
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Hybrid Orbitals




sp hybridization
Created from existing orbitals forming a new
“sublevel”
E.g. HCl
H = 1s1
Cl = [Ne]3s23p5
Unfilled s and unfilled p orbitals overlap
forming 2 sp orbitals ( bond)
sp2 hybridization
sp3 hybridization (no lone pairs)
sp3 hybridization (one lone pair)
sp3d hybridization
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9/30/2011
sp3d2 hybridization
Hybrid Orbitals determine Geometry
Hybrid
Orbital
sp
Geometry
sp2
trigonal planar
linear
sp3
tetrahedral
sp3d
trigonal bipyramidal
sp3d2
octahedral
Intermolecular Forces





attractions between molecules that hold them
together forming a liquid or a solid
ion-dipole: occurs between an ionic compound
and polar molecules
H-bonding (not real bonding) – occurs between
the H ( +) on one molecule and the lone pairs on
the F, O, or N of another molecule
dipole-dipole: occurs between two polar
molecules
London forces: occur because of a disturbance
creating temporary dipoles (occur in ALL
molecules)
Strength of IM forces

increase as polarizability increases



Melting and boiling points

viscosity – resistance to flow
left to right on previous chart
increases with increasing molecular weight

Intermolecular Forces affect Properties

larger atoms have larger electron clouds and are
easier to polarize



as IM forces increase, MP & BP increase
as IM forces increase, viscosity increases
as temperature increases, viscosity decreases
surface tension – a measure of the inward
forces that must be overcome to expand the
surface area

as IM forces increase
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9/30/2011
List the following in order of
increasing boiling points.





barium chloride
neon
hydrogen
carbon monoxide
hydrogen fluoride
H2<Ne<CO<HF<BaCl2


HF – hydrogen bonding
CO – dipole-dipole
 Ne – nonpolar



Solid





fixed volume and shape
incompressible
does not flow
diffusion within a solid is very, very slow
molecules move but do not leave their location
Gas





BaCl2 – ionic – highest
Ne, CO, HF – all have similar molecular weights
H2 – non-polar & lowest molecular weight –
lowest
Liquid





shape of container, fixed volume
non-compressible
fluid
other substance diffuse slowly
molecules move around each other
State depends on energy & IM forces
shape & volume of container
compressible
fluid
other substance diffuse quickly within a gas
molecules move in roughly a straight line until
they run into something
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9/30/2011
Phase Changes
Heat = Enthalpy = Energy

Hfus = heat (enthalpy) of fusion


Hvap = heat (enthalpy) of vaporization

Heating Curve

Notice:

the amount of energy required to change a solid into a
liquid at its melting point
the amount of energy required to change a liquid into
a gas at its boiling point

Hvap > Hfus

Hsub = Hvap + Hfus
Calculating H for Temperature/Phase
Changes

energy can
be used to
raise
temperature
or change
state, but not
both at the
same time
Calculate the enthalpy change upon
converting 1.00 mole of ice at -25 C to water
vapor (steam) at 125 C under a constant
pressure of 1 atm.

specific heats
ice = 2.09 J/g-K
water = 4.18 J/g-K
 ice = 1.84 J/g-K



For water



Hfus = 6.01 kJ/mol
Hvap = 40.67 kJ/mol
answer: H = 56.0 kJ
Vapor Pressure

Critical temperature


the highest temperature at which a liquid can
exist


Critical pressure

the pressure required to bring about liquefication
at the critical temperature


molecules on the top layer of a liquid gain enough
energy from the environment to escape
these gas molecules “sitting” on top of the liquid
exert a pressure on the surface of the liquid
as the amount of vapor increases, the likelihood
that a gas molecule will hit the surface of the
liquid increases
dynamic equilibrium will be established

liquid molecules escaping to become gas at the same
rate gas molecules condense back to liquid
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Vapor Pressure affects BP


boiling occurs when the vapor pressure of a
liquid is equal to the external (atmospheric)
pressure
as external pressure decreases, liquid surface
experiences less pressure allowing more
molecules to turn into gas form
Triple Point

the temperature and pressure at which a
substance is in equilibrium among all three
phases (solid, liquid, gas)
Ionic Solids

Phase Diagram
Pack themselves to maximize ionic
attractions & minimize repulsions
Solids

Amorphous – no order

Crystalline – highly ordered
Metals


Metal atoms bond
because electrons are
easily delocalized
forms an electron “sea”
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Covalent Network Solids

Diamonds (left) are examples of solids where
molecules are covalently bonded
Covalent Network Solids

Graphite is an example of a solid where
molecules are held together by van der Waals
forces
Types of Bonding in Crystalline Solids
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