BONDING
Bond =
when two nuclei simultaneously attract the
same pair of electrons
*The electrons experience a force of
attraction from both nuclei.
*This_______________________
positive-negative-positive
attraction holds the two particles
together.
*This attraction is called ________________
a chemical bond.
* One Pair of Electrons = 1 bond
Bonding changes = Energy Changes
The ______
Lower the Energy, The More _______
Stable !!!!!!!
*Breaking Bonds= Endothermic
Endothermic= Absorbing Energy
A+B
absorbing Energy,
*By __________
Atoms become _____
less stable
when separated.
AB + Energy
A+B
AB
Remember… The
_________the
Energy, The More
Lower
_________
!!!!!!!
Stable
_____________
High Energy =____________
Less Stable
 Making Bonds=
Exothermic
Exothermic = Releasing Energy
(Exit, Exhale)
When bonds ______,
Form excess
energy is ________,
released resulting
in a more ________
stable
compound.
A+B
AB + Energy
A+B
AB
Chemical Energy
Potential Energy is associated with the changes in
* ______________
Energy
bonds between atoms. ( Stored
____________)
* Single, uncombined atoms have a _______________
high Potential
__________
Energy.
*When bonds form, some of the _________________
Potential Energy is
______________
released.
* The amount of Potential Energy associated with a
specific bond depends on: ______
Mass , ___________,
Bond Type
____________
Atomic Radius , etc.
Remember the Octet Rule…
*Atoms will _____
lose , _____,
gain or ______
share enough
electrons as to obtain a _____
full , stable octet.
A. Electrons can be ______________
Transferred
Lost or ______
_____
gained to form ions with a
______________
= ________________
stable octet.
IONIC BONDS
METALS TO ______________
__________
NON-METALS
B. Electrons can be ___________
Shared
When _______________
sharing electrons , both atoms
feel as if they have a __________.
full octet
Covalent Bonds
= __________________
NON-METAL TO ________________
NON-METAL
_____________
EXCEPTIONS TO THE OCTET RULE
Hydrogen and ______
Helium are Exceptions to
*_________
they can only
the octet rule because ___________
hold
a maximum of 2 electrons in their
________________________________
PEL.
________________________________
*They are stable and Happy :) with a full “ duet
____”
of electrons.
Are all bonds the same strength?
NO!!
A. Weak bonds
Require little energy to break,
and release little energy when
forming. (unstable)
B.Strong Bonds
Require lots of energy to break,
and release lots of energy when
forming (highly stable)
Practice:
1.For the following equation:
N(g) + N(g) --> N2(g) + Energy
a. Bond broken, energy absorbed
b. Bond broken, energy released
c. Bonds formed, energy absorbed
d.
Bond formed, energy is released
_____________________________
2. Which would be true of the reverse
reaction?
a. Bond broken, energy absorbed
More Practice
3. Which amount of energy would be given
off in the formation of a highly stable
compound?
a) -100 kJ/mol
b) -800 kJ/mol
___________________
c) +170 kJ/mol
d) +500 kJ/mol
4. Which amount of energy from above
would symbolize weak bond formed?
d) +500 kJ/mol
One way to predict bond type is by taking the
difference in atom’s electronegativities.
What is electronegativity?
-Ability to attract electrons
Who had the highest electronegativity? Lowest?
Highest = Non- Metals,upper right on table
Lowest = Metals, lower left corner of table
What did an atom’s electronegativity value mean
for its reactivity?
Higher electronegativity value,more attraction for
electrons, Higher reactivity. Lower electro.value,
less reactivity.
(More on Electronegativity and bonds ahead)
Bonds vs. Molecular
Attractions
BOND
Intramolecular Force of Attraction = ________
*When two nuclei simultaneously attract the same
pair(s) of electrons to form a covalent bond. (stronger)
Intermolecular Force of Attraction
*The attraction that exists b/w two separate
molecules that is responsible for holding them near
one another in space. (weak in comparison)
Bond vs Intermolecular forces of
attraction
Bond Types
1. IONIC
2. METALLIC
3. COVALENT
a. nonpolar covalent
b. polar covalent
c. coordinate covalent
Ionic bonds
*Form between a _____
nonmetal
metal and a _____________
+
-Metal loses
_____ 1 or more electrons and becomes __
gains those electrons and becomes __
-Nonmetal _____
cation
*Metal ______(+
ion) attracts
Nonmetal _____(a
negative ion) =
anion
Ionic Bonds
*The amount of electrons transferred depends on
how many each atom needs to lose or gain to
complete its octet.
Ionic bonds
NaCl =most common ionic
compound
Na
2-8-1
Cl
2-8-7
Valence
Electrons
1
7
Lewis Dot
Na
Cl
Electron
Configuration
New e- config.
With full shell
Ion Formed
Lewis Dot of Ion
[2-8]+1
[2-8-8]-1
Na+1
Cl-1
[ Na
]+1
* * * -1
[**Cl * ]
*
Ionic Bonding
Ionic Bonds can also be predicted
using electronegativity values
*If the difference in electronegativity is
1.7, the bond
greater than or equal to ______________
is ionic
___________
NaCl
Cl = 3.2
Na = 0.9
2.3= ionic
Remember: There are exceptions
to everything.
Using electronegativity values,
predict if the following compounds
are ionic.
1) LiBr
2) H2O
3) CaO
4) CO2
Properties of Ionic Substances
Non metals .
*Form when ________
react with __________
metals
*Electronegativity difference of ____________
1.7or higher .
*Have ______
high
high melting points and _______
boiling points.
solids
*All ionic compounds are _________
at room
temperature since the bonds are so _________.
strong
*Have very low vapor pressure.
-Atoms do not easily go from liquid to
gas because of strong bonds.
Properties of Ionic Substances
Non conductors of electricity in the solid
*Are ______________
state.
*Conducts electricity as a _______
or aqueous
______.
liquid
(This allows the electrons to be mobile,
move freely and conduct electricity)
Writing Lewis Dot Ionic
Compounds
Na11
Cl17
*In terms of electronegativity values, how can
you tell that NaCl is ionic?
Writing Lewis Dot Ionic
Compounds
Use:
Li and O
Use:
Ca and F
Binary Ionic Compound: One
metal and one nonmetal
NaCl, CaBr2, MgO, K2S
Some Ionic Compounds have more
than 2 elements, they usually contain
polyatomic ions
NaNO3
NH4Cl
CaSO4
LiCr2O7
AgNO2
Polyatomic Ions (E)
1) Grouping of 2 or more atoms that act
as a unit and carry a charge.
2) If a compound contains a polyatomic
ion, it has both ionic and covalents
bonds.
NaNO3
Na+1 and NO3-1
Metallic Bonds
Metal
Metal
Are formed between _________
and ____________
*Metals generally have:
-Few valence e-Low ionization energy
( Ionization =
How much energy is needed to
remove the outermost e-)
*Metallic bonds are relatively strong bonds:
-Moderately high melting point
-Moderately high boiling point
Metallic Bonds
*Kernel- Nucleus and non valence electrons
*Arranged in a fixed position of a crystal lattice
*Valence electrons move freely about the crystal
lattice and do not belong to
any given atom.
*Metals are good conductors because of these mobile
electrons
Metallic Bonds
*Metallic Bonds
Result from the force of attraction
of the mobile valence electrons
for an atoms positively charged
kernel.
****SEA OF MOBILE ELECTRONS !!!
Metal= crystal lattice
Red sphere = (+) kernel of metal
cations
Static = moving e-
Examples of Metallic Bonds
* The bonds found in copper wire
* The bonds found in a piece of magnesium strip
* The bonds found in a piece of gold.
Properties of metals
* Composed of Monoatomic metal atoms.
ex. Cu(s) Au(s) Hg(l)
* Moderate Bond Strength
* Moderate melting and boiling points
* Conducts electricity in the solid and liquid phase
* Malleable and ductile
*Within the metal crystal lattice, electrons belong to
the whole crystal rather than individual ions.
*Mobility of electrons distinguishes the metallic bonds
from all others.
Covalent Bonds = MOLECULES
*Electrons are shared between 2 or more
non metals to obtain a stable octet.
-Sharing can be equal or unequal
*Electronegativity difference is usually
below 1.7
H*
*H
H ** H
= H2
Outer shell is full for BOTH by SHARING.
Remember… Hydrogen can have only 2 in its outer shell.
By sharing 1 pair of
e- = 2 e-
Practice Covalent Bonds
Combine :
F and F
Show the formation of H2O
Double Bonds
*Double Bond- 2 Pairs of e- (4 e-) are shared
- Makes it more stable, Lower Potential Energy
*Each bond is usually 1 e- from each atom forming
the pair.
Example: Show the formation of the following:
CO2
Triple Bonds
*Triple Bonds- 3 Pairs of e- (6 e-) are shared
between 2 atoms.
* Very stable, low Potential Energy
Example: N2
How many e- are shared?
Single Bond: Cl-Cl, H-H
1
_____pairs
of e2
_____
e- total
Double Bond O=O
2
_____
pairs of e4
_____
e- total
Triple Bond
NΞN
3
_____
pairs of e6
_____
e- total
Polar or Non Polar Covalent??
*Each Covalent Bond is one of 2 types:
Polar or Non Polar
Non Polar Covalent Bonds:
-A bond formed between 2 atoms that have the
same attraction ( electronegativity) for the
shared pair of e-.
- Equal pull on bonding pair
- Symmetrical e- charge distribution
***Different non metals can have the same
electronegativity values
Examples of Non Polar Bonds
Examples:
Cl2
Cl - Cl
H2
CS2
H-H
C-S
Polar Covalent Bonds
*Polar Covalent Bonds:
-Bonds formed between 2 non metal
atoms which have different attractions for
shared e- (different electronegativities)
- Unequal pull on bonding pair of e- Asymmetrical e- charge distribution
- Leaves partial charges on atoms within
molecule
H – Cl
Polar Covalent
Examples:
H2O
HBr
NH3
e- cluster around
more electronegative
atom
Properties of Covalent Molecules
*Form when a non
_________
metal reacts with a __________
non metal
*Have ____
Low melting points and ____
low freezing points
*Molecular solids are soft
*Covalent Molecules are ______________
poor conductors of heat and
electricity
weak bonds
*Molecular substances are generally _____________
high vapor pressure
*Molecular substances have a ____
(goes into vapor easily because of weak bonds.)
Shapes of Molecules
*The shapes of molecules is
determined by _____________
the repelling
________________
of electrons
There are 5 Basic shapes:
*Tetrahedral
*Linear
*Planar
*Angular
*Pyramidal
Linear
When there are 2 atoms in a molecule
Any molecule in which the bonds cause
the atoms to fall in a straight line.
H2
C2H2
CO2
Angular
*V-shaped or Bent
H2O
SO2
H2S
Planar
*The atoms of a molecule fall in the same
plane.
C2H4
Pyramidal
*3 shared pairs of electrons will cause a
pyramid shape.
NH3
PCl3
Tetrahedral
*4 evenly separated bonds
CCl4
CH4
Polarity Of Molecules
*Molecules can either be polar or non polar.
(Do not confuse bond polarity with molecular polarity)
Bond Polarity- When checking between 2 atoms if there is
difference in electronegativity or not.
*Polar bond-difference in electronegativity between
2 atoms
*Non Polar Bond- No difference in electronegativity
between 2 atoms.
NonPolar Molecules
*Molecules containing only nonpolar covalent
bonds, are always nonpolar molecules!
All diatomics: Br2, I2,N2 ,Cl2, H2, O2, F2,
NonPolar Molecules
*Non polar molecules can also contain
polar bonds
*Bonding e- have to be symmetrically
distributed between atoms.
CH4
Tetrahedral
CCl4
CO2
Linear
Polar Molecules
*Molecules that contain polar covalent bonds can
also be polar molecules depending on overall
molecule symmetry (shape). (asymmetrical)
H 2O
HF
SO2
Leads to partial charges at different points in molecule,
even though overall molecule is neutral
***Use dipole arrows for each bond to determine the
overall symmetry of the molecule***
1) Draw the dipole arrow facing the more
electronegative atom.
a) The atom the arrow points towards
is - while where the arrow points
away from is +
2) If dipole arrows are equal in magnitude
but opposite in direction, they cancel out
and the molecule is nonpolar
What is the Polarity of the
following Molecules?
HCl
CO2
CH4
NH3
CCl4
H 20
S.N.A.P
Symmetrical
Non polar
Asymmetrical
Polar
Like Dissolves Like
Like Polarities Dissolve
Like Polarities
Polar dissolves Polar
Alcohol and water
Non Polar dissolves
Non Polar
Many organic molecules
Polar NOT Dissolving non Polar
Oil and Water
Intermolecular Forces Of
Attraction (IMF’s)
*The attractions between molecules that holds
them near one another in space.
Intermolecular Forces Of
Attraction (IMF’s)
4 types of Intermolecular forces of attraction:
1. Dipole-Dipole
2. Hydrogen Bonding
3. Van der Waals
4. Molecule-Ion or Dipole-Ion
IMF’s explain the change in phase going
down the halogen group (17)
F2
Cl2
Br2
I2
Generally IMF’s will INCREASE with:
1.Increasing #e2.Increasing mass of atoms
3.Increasing mass and complexity
(polarity) of a molecule.
1. Dipole-Dipole
1) When the partially δ + end of one dipole,
attracts the partially δ– end of a
neighboring dipole.
Show H2S:
Dipole Dipole
Dipole= Polar
*The temperature at which liquid boils is due
to the strength of the dipole attractions
*A higher attractive force= increased mp
and increased bp
*The larger the molecule, the stronger the
attractive forces
2. Hydrogen Bonds
1)
2)
3)
Not an actual bond, just a strong IMF
H-bonding is FON!
H-bonds are only possible if H is
covalently bonded to either F, O, or N.
a) those polar molecules then strongly
attracted each other which is the “hydrogen
bond”
b) Common molecules that can H-bond:
H2O, HF, and NH3
Hydrogen Bonds
*Hydrogen has such
a small share of the
e- pair, it is almost
like a bare proton
Van Der Waals Forces
*Weakest IMF between small, nonpolar
molecules
*Van der Waals strength ↑ with ↑ mass
and ↑ # e-
Molecule Ion attraction
1. Partially δ– end of polar molecule,
dipole, attracts + cation
2. Partially δ + end of dipole attracts - anion
Molecule Ion Attraction
Example: Saltwater, NaCl (aq)
(Dipole is usually water)
Order of IMF Strength
Molecule Ion
Hydrogen Bonding
Dipole Dipole
Van der Waals
Decreasing
Strength of IMF
What do IMF’s account for?
1)
2)
3)
4)
Boiling Point
Melting Point
Vapor Pressure
Phase
↑IMF ↑BP ↑MP
↓IMF ↓BP ↓MP
↑IMF↓VP, ↓IMF ↑VP
The ↑IMF, the more likely a
substance will be solid
↓IMF, the more likely the substance
will be a gas