Atom – the smallest unit of matter “indivisible”
Helium
atom
electron shells
a) Atomic number = number of Electrons
b) Electrons vary in the amount of energy
they possess, and they occur at certain
energy levels or electron shells.
c) Electron shells determine how an atom
behaves when it encounters other atoms
Electrons are placed in shells
according to rules:
1) The 1st shell can hold up to two electrons,
and each shell thereafter can hold up to 8
electrons.
Octet Rule = atoms tend to gain, lose or share electrons so
as to have 8 electrons
C would like to Gain 4 electrons
N would like to Gain 3 electrons
O would like to Gain 2 electrons
Why are electrons important?
1) Elements have different electron
configurations
 different electron configurations mean
different levels of bonding
Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell
electrons
1
2
13
14
15
16
17
H
18
He:

Li Be


B 


C


Na Mg


Al

N



O




 Si 
P
 S





: F  :Ne :




:Cl  :Ar :


Chemical bonds: an attempt to fill electron shells
1. Ionic bonds –
2. Covalent bonds –
3. Metallic bonds
Learning Check

A.
X would be the electron dot formula for
1) Na
B.

X

1) B
2) K
3) Al
would be the electron dot formula
2) N
3) P
IONIC BOND
bond formed between
two ions by the
transfer of electrons
When the electronegativities of two atoms are quite
different from each other:
One atom loses an electron (or electrons)
The other atom gains an electron (or electrons)
This results in an Ionic
Bond.
Formation of Ions from Metals
 Ionic compounds result when metals react with
nonmetals
 Metals lose electrons to match the number of valence
electrons of their nearest noble gas
 Positive ions form when the number of electrons are
less than the number of protons
•
Group 1 metals 
ion 1+
Group 2 metals 
ion 2+
Group 3 metals 
ion 3+
Formation of Sodium Ion
Sodium atom
Na 
2-8-1
11 p+
11 e0
– e
Sodium ion

Na +
2-8 ( = Ne)
11 p+
10 e1+
Formation of Magnesium Ion
Magnesium atom
Magnesium ion

Mg 
2-8-2
12 p+
12 e0
– 2e

Mg2+
2-8 (=Ne)
12 p+
10 e2+
Some Typical Ions with Positive
Charges (Cations)
Group 1
Group 2
Group 13
H+
Mg2+
Al3+
Li+
Ca2+
Na+
Sr2+
K+
Ba2+
Learning Check
A. Number of valence electrons in aluminum
1) 1 e2) 2 e3) 3 eB.
C.
Change in electrons for octet
1) lose 3e2) gain 3 e-
Ionic charge of aluminum
1) 32) 5-
3) gain 5 e-
3) 3+
Solution
A. Number of valence electrons in aluminum
3)
3 eB.
Change in electrons for octet
1)
lose 3e-
C.
Ionic charge of aluminum
3) 3+
Learning Check
Give the ionic charge for each of the following:
A. 12 p+ and 10 e1) 0
2) 2+
3) 2B. 50p+ and 46 e1) 2+
2) 4+
3) 4-
C. 15 p+ and 18e2) 3+
2) 3-
3) 5-
Ions from Nonmetal Ions
In ionic compounds, nonmetals in 15, 16, and 17
gain electrons from metals
Nonmetal add electrons to achieve the octet
arrangement
Nonmetal ionic charge:
3-, 2-, or 1-
Fluoride Ion
unpaired electron

:F

2-7
9 p+
9 e0
+ e
octet

1-
: F:

2-8 (= Ne)
9 p+
10 e1ionic charge
Ionic Bond
• Between atoms of metals and nonmetals
with very different electronegativity
• Bond formed by transfer of electrons
• Produce charged ions all states. Conductors
and have high melting point.
• Examples; NaCl, CaCl2, K2O
Ionic Bonds: One Big Greedy Thief Dog!
1). Ionic bond – electron from Na is transferred to Cl,
this causes a charge imbalance in each atom. The Na
becomes (Na+) and the Cl becomes (Cl-), charged
particles or ions.
Li
F
A Li Atom
An F Atom
+
Li
F
A Li+ Ion
An F- Ion
Be
F
An F Atom
F
An F Atom
A Be Atom
-
2+
F
Be
F
A Be2+ Ion
An F- Ion
An F- Ion
NaCl
Crystal Lattice
The melting
as follows:
NaF
KCl
LiCl
points of some Ionic Compounds are
993 oC
770 oC
605 oC
These high melting points are experimental evidence that
Ionic Bonds are VERY STRONG. (Hard to break just by
heating).
Task 1
• Describe step by step how the NaCl solid
occur from Na(s) and Cl2(g)
• What energy are involved?
COVALENT BOND
bond formed by the
sharing of electrons
Covalent Bond
• Between nonmetallic elements of similar
electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not
conductors at any state
• Examples; O2, CO2, C2H6, H2O, SiC
In Covalent bonds, electrons are
Shared
H
H
Covalent bonds in large networks (Network Bonding)
gives rise to substances with very high melting points.
diamond structure
Diamonds are “forever”!
Some melting points of
Network Solids:
Diamond (Carbon) 3550 oC
Silicon Carbide (SiC) 2700 oC
Boron Nitride (BN) 3000 oC
Covalent bonds
are very
strong!
Bonds in all the
polyatomic ions
and diatomics
are all covalent
bonds
NONPOLAR
COVALENT BONDS
when electrons are
shared equally
H2 or Cl2
Non Polar Covalent bonds-
Two atoms share one or more pairs of
outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)
POLAR COVALENT
BONDS
when electrons are
shared but shared
unequally
H2O
When electrons are shared unequally between two
atoms, the bond is called Polar Covalent. A type of
PC bond formed when “H” from one atom attracts
“O” or “N” from another atom is called Hydrogen
Bonding. polar covalent bonds
- water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen.
Hydrogen Bonding in Water
gives rise to the structure of
ice when water solidifies.
Hydrogen bonds
between the
“bases” hold the
two strands of
DNA together.
Polar Covalent Bonds: Unevenly
matched, but willing to share.
Bonds within molecules that hold the atoms of a molecule
together are called intramolecular
are strong covalent bonds.
Covalent Bonds
bonds. They
A dipole is a partial separation of charge which exists
when one end of a molecule has a slight positive charge
and the other end has a slight negative charge. Eg. A
water molecule has two dipoles.
The Greek letter
d
“delta” means “partial”
Just by pure chance, there are some times when
both electrons in helium are on the same side.
This forms temporary dipoles
ee-
e+2
e-
He
d
+2
He
d+
d
d+
The weak attractive forces between the
(+) side of one molecule and the (-) side
of another molecule are called London
Forces
The covalent
intramolecular bond
in I2 is very strong.
I
I
I
I
I
I
I
I
I
I
I
I
There are weaker
intermolecular forces which
hold covalent molecules
together in a molecular solid.
These are called London
Forces. Since they are
relatively weak, Iodine has a
low melting point.
Lewis Structures
(Electron-dot formulas) for Ionic Compounds.
Remember, in an ionic compound, the metal loses e-’s
and the non-metal gains. There is no sharing. Here is the
e-dot formula for sodium chloride (NaCl)
Na+
Cl
Here is the e-dot formula (Lewis Structure) for the ionic
compound MgF2 :
F
Mg2+
F
Notice, there is no sharing. The F atoms took both valence
e-’s from Mg, forming ions which do not share electrons.
The + and – charges on the ions cause them to attract each
other.
Electron-dot Formulas (Lewis Structures) for Covalent
Compounds.
When atoms form covalent bonds, they are trying to
achieve stable noble gas electron arrangements:
Hydrogen will share e-’s until it feels 2 e-’s like Helium.
Other elements share e-’s to achieve what is called a
“Stable Octet” (8 valence e-’s)
Electron-dot formula for Methane (CH4)
H
H
C
H
Here is a Carbon
atom (4 val e-’s) and
four Hydrogen atoms
(1 val e- each)
H
Electron-dot formula for Methane (CH4)
H
Now they have
formed a stable
molecule. Each
C atom “feels”
like it has a
stable octet.
H
C
H
H
Each H atom
“feels” like a
stable “He”
atom with 2e-s
Electron-dot formula for Ammonia (NH3)
H
N
H
Here is a Nitrogen
atom (5 val e-’s) and
three Hydrogen atoms
(1 val e- each)
H
Electron-dot formula for Ammonia (NH3)
“N” now feels
like it has a
stable octet
H
N
H
H
Each “H” feels
like it has 2 elike Helium.
Write the electron-dot formula for CF4
Because “F” is a
halogen, it has 7
valence e-s, so you
must show all 7 red
dots around each “F”
atom!
F
F
C
F
F
Write the electron-dot formula for H2S
S
H
H
The two H’s MUST
be at right angles to
each other!!
Write the Electron-Dot Formula for SeF2
Se
F
F
Because “F” is in Group 17,
they have 7 valence e-s, so
they must have 7 red dots
around them.
Hydrogen Bonding
(Shown in water)
d+ dH O
+
Hd
This hydrogen is bonded
covalently to: 1) the
highly negative oxygen,
and 2) a nearby unshared
pair.
Hydrogen bonding allows H2O to be a
liquid at room conditions.
H O
H
Attractions and properties
• Why are some chemicals gases,
some liquids, some solids?
–Depends on the type of bonding!
–Table 8.4, page 244
• Network solids – solids in which all
the atoms are covalently bonded to
each other
Attractions and properties
• Network solids melt at very high
temperatures, or not at all (decomposes)
– Diamond does not really melt, but
vaporizes to a gas at 3500 oC and
beyond
– SiC, used in grinding, has a melting
point of about 2700 oC
Multiple Bonds
• Sometimes atoms share more than one
pair of valence electrons.
• A double bond is when atoms share two
pairs of electrons (4 total)
• A triple bond is when atoms share three
pairs of electrons (6 total)
• Table 8.1, p.222 - Know these 7
elements as diatomic:
What’s the deal
with the oxygen
Br2 I2 N2 Cl2 H2 O2 F2 dot diagram?
Dot diagram for Carbon dioxide
C
O
• CO2 - Carbon is central
atom ( more metallic )
• Carbon has 4 valence
electrons
• Wants 4 more
• Oxygen has 6 valence
electrons
• Wants 2 more
Carbon dioxide
• Attaching 1 oxygen leaves the oxygen
1 short, and the carbon 3 short
CO
Carbon dioxide
 Attaching
the second oxygen leaves
both of the oxygen 1 short, and the
carbon 2 short
OC O
Carbon dioxide
 The
only solution is to share more
O CO
Carbon dioxide
 The
only solution is to share more
O CO
Carbon dioxide
 The
only solution is to share more
O CO
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
 Requires two double bonds
 Each atom can count all the electrons
in the bond
O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom can count all the electrons in the
bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom can count all the electrons in the
bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom can count all the electrons in the
bond
8 valence
electrons

O C O
How to draw them?
 Use the handout guidelines:
1) Add up all the valence electrons.
2) Count up the total number of electrons
to make all atoms happy.
3) Subtract; then Divide by 2
4) Tells you how many bonds to draw
5) Fill in the rest of the valence electrons
to fill atoms up.
Example
NH , which is ammonia
N
H
•
3
• N – central atom; has 5
valence electrons, wants 8
• H - has 1 (x3) valence
electrons, wants 2 (x3)
• NH3 has 5+3 = 8
• NH3 wants 8+6 = 14
• (14-8)/2= 3 bonds
• 4 atoms with 3 bonds
Examples
• Draw in the bonds; start with singles
• All 8 electrons are accounted for
• Everything is full – done with this one.
H
H NH
Example: HCN
•
•
•
•
•
•
•
•
HCN: C is central atom
N - has 5 valence electrons, wants 8
C - has 4 valence electrons, wants 8
H - has 1 valence electron, wants 2
HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds – this will require
HCN
• Put single bond between each atom
• Need to add 2 more bonds
• Must go between C and N (Hydrogen is full)
HC N
HCN
Put in single bonds
 Needs 2 more bonds
 Must go between C and N, not the H
 Uses 8 electrons – need 2 more to equal the
10 it has

HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on the N to fill its octet

HC N
Another way of indicating bonds
• Often use a line to indicate a bond
• Called a structural formula
• Each line is 2 valence electrons
HO H H O H
=
Other Structural Examples
H C N
H
C O
H
A Coordinate Covalent Bond...
• When one atom donates both electrons
in a covalent bond.
• Carbon monoxide (CO) is a good
example:
Both the carbon
and oxygen give
another single
electron to share
CO
Coordinate Covalent Bond
 When
one atom donates both electrons
in a covalent bond.
 Carbon monoxide (CO) is a good
example:
Oxygen
This carbon
electron
moves to
make a pair
with the other
single.
C O
gives both
of these
electrons,
since it has
no more
singles to
share.
Coordinate Covalent Bond
 When
one atom donates both electrons
in a covalent bond.
 Carbon monoxide (CO)
The coordinate
covalent bond is
shown with an
arrow as:
C
O
C O
Coordinate covalent bond
• Most polyatomic cations and anions
contain covalent and coordinate
covalent bonds
• Table 8.2, p.224
• Sample Problem 8.2, p.225
• The ammonium ion (NH41+) can be
shown as another example
Bond Dissociation Energies...
• The total energy required to break
the bond between 2 covalently
bonded atoms
• High dissociation energy usually
means the chemical is relatively
unreactive, because it takes a lot
of energy to break it down.
Resonance is...
• When more than one valid dot diagram
is possible.
• Consider the two ways to draw ozone
(O3)
• Which one is it? Does it go back and
forth?
• It is a hybrid of both, like a mule; and
shown by a double-headed arrow
• found in double-bond structures!
Resonance in Ozone
Note the different location of the double
bondstructure is correct, it is actually
Neither
hybrid of the two. To show it, draw all
varieties possible, and join them with a
double-headed arrow.
a
Resonance
Occurs when more than one valid Lewis structure
can be written for a particular molecule (due to
position of double bond)
•These are resonance structures of benzene.
•The actual structure is an average (or hybrid) of
these structures.
Polyatomic ions – note the different
positions of the double bond.
Resonance
in a
carbonate
ion (CO32-):
Resonance in
an acetate ion
(C2H3O21-):
The 3 Exceptions to Octet rule
• For some molecules, it is impossible to
satisfy the octet rule
#1. usually when there is an odd
number of valence electrons
– NO2 has 17 valence electrons,
because the N has 5, and each O
contributes 6. Note “N” page 228
• It is impossible to satisfy octet rule, yet
the stable molecule does exist
•
Exceptions
to
Octet
rule
Another exception: Boron
• Page 228 shows boron trifluoride, and
note that one of the fluorides might be
able to make a coordinate covalent bond
to fulfill the boron
• #2 -But fluorine has a high
electronegativity (it is greedy), so this
coordinate bond does not form
• #3 -Top page 229 examples exist because
they are in period 3 or beyond
Covalent Network Compounds
Some covalently bonded substances DO NOT
form discrete molecules.
Diamond, a network of
covalently bonded
carbon atoms
Graphite, a network of
covalently bonded
carbon atoms
METALLIC BOND
bond found in
metals; holds metal
atoms together
very strongly
Metallic Bond
• Formed between atoms of metallic elements
• Electron cloud around atoms
• Good conductors at all states, lustrous, very
high melting points
• Examples; Na, Fe, Al, Au, Co
Metallic Bonds: Mellow dogs with plenty
of bones to go around.
Ionic Bond, A Sea of Electrons
Metals Form Alloys
Metals do not combine with metals. They form
Alloys which is a solution of a metal in a metal.
Examples are steel, brass, bronze and pewter.
Formula Weights
• Formula weight is the sum of the atomic
masses.
• Example- CO2
• Mass, C + O + O
12.011 + 15.994 + 15.994
43.999
Practice
• Compute the mass of the following compounds
round to nearest tenth & state type of bond:
• NaCl;
• 23 + 35 = 58; Ionic Bond
• C2H6;
• 24 + 6 = 30; Covalent Bond
• Na(CO3)2;
• 23 + 2(12 + 3x16) = 123; Ionic & Covalent
Thank You