BONDING
General Rule of Thumb:
metal + nonmetal = ionic
polyatomic ion + metal or polyatomic ion = ionic (both)
nonmetal + nonmetal(s) = covalent
Ionic Bonds
Isn’t it ionic that opposites attract?
Valence Electrons


Knowing electron configurations is
important because the number of
valence electrons determines the
chemical properties of an element.
Valence Electrons: The e- in the
highest occupied energy level of an
element’s atoms.
Valence Electrons


All elements in a particular group or family
have the same number of valence electrons
(and this number is equal to the group number
of that element)
Examples:



Group 1 elements (Na, K, Li, H): 1 valence electron.
Group 2 elements (Mg, Ca, Be): 2 valence electrons.
Group 17 elements (Cl, F, Br): 7 valence electrons.
Lewis Structures








Electron dot structures show the
valence electrons as dots around the
element’s symbol:
Li
B
Si
N
O
F
Ne
Lewis Structures








Electron dot structures show the
valence electrons as dots around the
element’s symbol:
Li
B
Si
N
O
F
Ne
Octet Rule


Noble gas atoms are very stable;
they have stable electron
configurations. In forming
compounds, atoms make
adjustments to achieve the lowest
possible (or most stable) energy.
Octet rule: atoms react by
changing the number of electrons so
as to acquire the stable electron
structure of a noble gas.
Octet Rule









Atoms of METALS obey this rule by losing
electrons.
Na:
Na+:
Atoms of NONMETALS obey this rule by
gaining electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a
relatively stable configuration with its 4d
sublevel filled (pseudo noble-gas)
Octet Rule









Atoms of METALS obey this rule by losing
electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining
electrons.
Cl:
Cl-:
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled
(pseudo noble-gas)
Octet Rule









Atoms of METALS obey this rule by losing
electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining
electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled
(pseudo noble-gas)
Octet Rule









Atoms of METALS obey this rule by losing
electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining
electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled
(pseudo noble-gas)
Octet Rule









Atoms of METALS obey this rule by losing
electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining
electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule.
Example: silver (Ag)
By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled
(pseudo noble-gas)
Ionic Bonds

Anions and cations have opposite
charges; they attract one another
by electrostatic forces
(IONIC BONDS)
Ionic Bonds

Ionic compounds are electrically
neutral groups of ions joined together
by electrostatic forces. (also known
as salts)


the positive charges of the cations
must equal the negative charges of
the anions.
use electron dot structures to predict the
ratios in which different cations and
anions will combine.
Ex. of Ionic Bonds: Determine charge & CRISS CROSS

Na
Cl
Na+Cl- = NaCl

Al
Br
Al3+Br- = AlBr3

K
O
K+O2- = K2O

Mg
N
Mg2+N3- = Mg3N2

K
P
K+P3- = K3P
DETERMINE THE CHARGE OF THE ION, CRISS CROSS CHARGES
so the compound is NEUTRAL!!! (compounds DO NOT have
charges)
Cation + always goes 1st, Anion – always goes last!!
Making Ionic Compounds
+ CATION always goes first in a
compound!!!
- ANION always goes last in a
compound!!!
Another Example: reducing to the
simplest formula
What is the ionic compound formula
for the calcium ion with the oxide
ion?
Ca
O
Ionic Compounds w/ polyatomic ions
POLYATOMIC:
DO NOT CHANGE THE FORMULA
OF A POLYATOMIC ION!!!
Determine the formula for:
Li+ and PO43NH4+ and N3-
More practice

Determine the charge on an
Aluminum ion, then pair it with the
sulfite ion, SO32-
More practice…

Determine the charge on a Calcium
ion, then pair it with the sulfate ion,
SO42-
Covalent Bonds
The joy of sharing!
Covalent Bonds


Covalent bonds: occur between
two or more nonmetals; electrons
are shared not transferred (as in
ionic bonds)
The result of sharing electrons is
that atoms attain a more stable
electron configuration.
Covalent Bonds

Most covalent bonds involve:



2 electrons (single covalent bond),
4 electrons (double covalent bond, or
6 electrons (triple covalent bond).
MOLECULES

Lewis structures (electron dot
structures) show the structure of
molecules. (Bonds can be shown
with dots for electrons, or with
dashes: 1 dash = 2 electrons)
Octet Rule

Octet Rule: The representative
elements achieve noble gas
configurations (8 electrons) by
sharing electrons.
THERE ARE A FEW
EXCEPTIONS!!!
Hydrogen can only have 1 bond
(2 electrons around it)

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO
Writing Lewis Formulas (for molecules)

How to:
1.
2.
3.
4.
5.
Add up all valence electrons for EACH atom in
the molecule
Attach atoms with a single bond (skeleton
drawing) (C = ALWAYS CENTRAL, H =
ALWAYS ON OUTSIDE OF STRUCTURE)
Subtract out 2 electrons for each single bond
you drew (EACH BOND = 2 electrons)
Distribute remaining electrons (in pairs)
around atoms to obtain octet rule (except H)
If there’s not enough electrons to satisfy the
octet rule, make MULTIPLE BONDS (double,
triple)
Now let’s get more complex…
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples
(remember your 5 steps):

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Electronegativity

We’ve learned how valence
electrons are shared to form
covalent bonds between elements.
So far, we have considered the
electrons to be shared equally.
However, in most cases, electrons
are not shared equally because of a
property called electronegativity.
Electronegativity


The ELECTRONEGATIVITY of an
element is: the tendency for an
atom to attract electrons to itself
when it is chemically combined with
another element.
The result: a “tug-of-war”
between the nuclei of the
atoms.
Electronegativity
Electronegativities are given numerical
values (the most electronegative element
has the highest value; the least
electronegative element has the lowest
value)
*** see Periodic Table, in lower part of
element box



Most electronegative element:
Fluorine (4.0)
Least electronegative elements:
Fr (0.70), Cs (0.79)
Electronegativity

Notice the periodic trend:



As we move from left to right across a
row, electronegativity increases (metals
have low values nonmetals have high
values – excluding noble gases)
As we move down a column,
electronegativity decreases.
The higher the electronegativity
value, the greater the ability to
attract electrons to itself.
Nonpolar Bonds


When the atoms in a molecule
are the same, the bonding
electrons are shared equally.
Result: a nonpolar covalent bond

Examples: O2, F2, H2, N2, Cl2
Polar Bonds



When 2 different atoms are joined by a
covalent bond, and the bonding electrons
are shared unequally, the bond is a polar
covalent bond, or POLAR BOND.
The atom with the stronger electron
attraction (the more electronegative
element) acquires a slightly negative
charge.
The less electronegative atom
acquires a slightly positive charge.
Polar Bonds

Example: HCl

Electronegativities:


H = 2.20
Cl = 3.16
+ H
Cl
-
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Predicting Bond Types

Electronegativities help us predict
the type of bond:
Electronegativity
Difference
0.00 – 0.40
0.41 – 1.00
1.01 – 2.00
2.01 or higher
Type of Bond
covalent
(nonpolar)
covalent
(slightly polar)
covalent
(very polar)
ionic
Example
H-H
H-Cl
H-F
Na+Cl-
Polar Molecules


A polar bond in a molecule can
make the entire molecule polar
A molecule that has 2 poles
(charged regions), like H-Cl, is
called a dipolar molecule, or dipole.
Polar Molecules


The effect of polar bonds on the polarity of
a molecule depends on the shape of the
molecule.
Example:
O=C=O
CO2
shape: linear
*The bond polarities cancel because they
are in opposite directions; CO2 is a
nonpolar molecule.
Polar Molecules


The effect of polar bonds on the
polarity of a molecule depends on the
shape of the molecule.
Water, H2O, also has 2 polar bonds:


But, the molecule is bent, so the bonds
do not cancel.
H2O is a polar molecule.
RESONANCE STRUCTURES
Resonance

A molecule or polyatomic ion for
which 2 or more dot formulas with
the same arrangement of atoms can
be drawn is said to exhibit
RESONANCE.
Resonance Example




CO32-
3 resonance structures can be drawn for CO32the relationship among them is indicated by
the double arrow.
the true structure is an average of the 3.
Resonance Example




CO32-
3 resonance structures can be drawn for CO32the relationship among them is indicated by
the double arrow.
the true structure is an average of the 3.
Resonance Example




CO32-
3 resonance structures can be drawn for CO32the relationship among them is indicated by
the double arrow.
the true structure is an average of the 3
Resonance Structures


Another way to represent this is by
delocalization of bonding electrons:
(the dashed lines indicate the 4
pairs of bonding electrons are
equally distributed among 3 C-O
bonds; unshared electron pairs are
not shown)
VSEPR
valence shell electron pair repulsion
Molecular Shape


Lewis structures (electron dot
structures) show the structure of
molecules…but only in 2 dimensions
(flat).
BUT, molecules are 3 dimensional!

for example, CH4 is:
Molecular Shape


Lewis structures (electron dot structures)
show the structure of molecules…but only in
2 dimensions (flat).
BUT, molecules are 3 dimensional!
 but in 3D it is:
a tetrahedron!
= coming out of page
= going into page
= flat on page
Why do molecules take on 3D shapes
instead of being flat?


Valence Shell Electron Pair
Repulsion theory
“because electron pairs repel one another,
molecules adjust their shapes so that the
valence electron pairs are as far apart from
another as possible.”
Why do molecules take on 3D shapes
instead of being flat?


Valence Shell Electron Pair Repulsion
theory
Remember: both shared and unshared
electron pairs will repel one another
(unshared electron pairs repel MORE than
shared electron pairs in bonds)
Non-Bonding
Pairs
H—N — H
H
Bonding
Pairs
5 Basic Molecule Shapes
1. tetrahedral

example: CH4
5 Basic Molecule Shapes
2. Pyramidal


Example: NH3
(note: unshared pair of electron
repels, but is not considered part of
overall shape; no atom there to
contribute to the shape)
5 Basic Molecule Shapes
3. Bent or angular


Example: H2O
Notice electron pair repulsion
5 Basic Molecule Shapes
4. Linear

Example: CO2
5 Basic Molecule Shapes
5. Trigonal planar or planar triangular

Example: BF3
CHEAT SHEET…
Let’s make this easy to remember shapes
Linear= only 2 regions of SHARED electron pairs
(no unshared electrons) around the CENTRAL
atom
Bent = 2 SHARED electron pairs, 2 UNSHARED
electron pairs around the CENTRAL atom
Tetrahedral = 4 SHARED electron pairs (meaning 4
bonds) around the CENTRAL atom
Trigonal Planar = 3 SHARED electron pairs (meaning 3
bonds) around the CENTRAL atom
Pyramidal = 3 SHARED electron pairs, 1 UNSHARED
pair of electrons around the CENTRAL atom
Geometry and polarity


Three shapes will cancel out
polarity.
Linear = CO2 = NONPOLAR
Geometry and polarity


Three shapes will cancel out polarity.
Trigonal Planar = BF3 = NONPOLAR
120º
Geometry and polarity
Three shapes will cancel out polarity.
Tetrahedral = CH4
= NONPOLAR

Geometry and polarity


Others don’t cancel
Bent = H2O = POLAR
Geometry and polarity


Others don’t cancel
Pyramidal = NH3 = POLAR