Forces that hold atoms together
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There are several major types of bonds.
Ionic, covalent and metallic bonds are
the three most common types of bonds.
Covalent bonds – electrons are shared
between atoms.
Ionic bonds – electrons are transferred
between atoms, creating cations and
anions.
Metallic bonds – two or more metals
bonded together.
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There are two different types of covalent
bonds, polar covalent and nonpolar covalent.
◦ polar covalent – electrons are not shared
equally between the two bonded atoms.
The electrons are pulled toward the more
electronegative of the elements.
◦ nonpolar covalent – electrons are shared
equally between the two bonded atoms.
Electronegativities
9_12
IA
IIA
Li
1.0
Be
1.5
Na
0.9
Mg
1.2
K
0.8
H
2.1
VIIIB
IIIB
IVB
VB
VIB
VIIB
Ca
1.0
Sc
1.3
Ti
1.5
V
1.6
Cr
1.6
Mn
1.5
Fe
1.8
Co
1.8
Rb
0.8
Sr
1.0
Y
1.2
Zr
1.4
Nb
1.6
Mo
1.8
Tc
1.9
Ru
2.2
Cs
0.7
Ba
0.9
La–Lu
1.1–1.2
Hf
1.3
Ta
1.5
W
1.7
Re
1.9
Os
2.2
Fr
0.7
Ra
0.9
Ac–No
1.1–1.7
IIIA
IVA
VA
VIA
VIIA
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
IB
IIB
Ni
1.8
Cu
1.9
Zn
1.6
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Rh
2.2
Pd
2.2
Ag
1.9
Cd
1.7
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Ir
2.2
Pt
2.2
Au
2.4
Hg
1.9
Tl
1.8
Pb
1.8
Bi
1.9
Po
2.0
At
2.2
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Ionic bonds are formed when there is an
electronegativity difference (DEN) greater than
2.0.
Polar covalent bonds form when there is a
DEN between 0.5 and 1.7.
Nonpolar covalent bonds form when there is
a DEN between 0 and 0.49.
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If the DEN is
between 1.7 and
2.0, an ionic bond
will form if a metal
is one of the
elements, and a
polar covalent bond
will form if only
nonmetals or
metalloids are
present.
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What type of bond
is formed between
the following
elements?
N and O
K and F
Mg and Cl
P and F
C and H
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Metals tend to lose their valence electrons,
leaving a complete octet in their next-lowest
energy level.
Sodium – (1 valence electron) loses 1 electron
and becomes Na+1.
Na ([Ne]3s1) → 1e- + Na+1([Ne])
Calcium – (2 valence electrons) loses 2
electrons and becomes Ca+2.
Ca ([Ar]4s2)
→ 2e- + Ca+2([Ar])
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Nonmetals tend to gain or share valence
electrons to complete an octet in their
highest energy level.
Oxygen – (6 valence electrons) gains two
electrons to become O-2 .
O ([He]2s22p4) + 2e- → O-2 ([He] 2s22p6)
Phosphorus – (5 valence electrons) gains
three electrons to become P-3.
P ([Ne]3s23p3) + 3e- → P-3 ([Ne] 3s23p6)
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Ionic bonds – forces of attraction that bind
cations and anions together.
Ionic compound – consists of electrically
neutral group of ions joined by electrostatic
forces.
Example: Sodium chloride
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At room temperature, most ionic compounds
are crystalline solids, where ions are arranged
in various 3-D patterns.
Because of the large attractive forces of the
ions to each other the compounds become
very stable and have high melting points.
Source: ©Clyde
H. Smith/Peter
Arnold, Inc.
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Scientists have learned that all of the
elements within each group behave similarly
because they have the same number of
valence electrons.
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Valence electrons - # of electrons in the
highest occupied energy level of an atom.
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The number of valence electrons is related to
the group numbers on the periodic table.
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Group 1 elements = 1 valence electron.
Group 2 elements = 2 valence electrons.
Groups 3-12 elements = 2 valence electrons.
Group 13 elements = 3 valence electrons.
Group 14 elements = 4 valence electrons.
Group 15 elements = 5 valence electrons.
Group 16 elements = 6 valence electrons.
Group 17 elements = 7 valence electrons.
Group 18 elements = 8 valence electrons.
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1. Multiply the number of valence electrons
by the number of moles of each element.
2. Add up all the electrons for each of the
elements.
3. If there is a charge and it is negative, add
that number of electrons to the total.
4. If there is a charge and it is positive,
subtract that number of electrons from the
total.
Total # of electrons should always be an even
number!
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Determine the number of valence electrons in
each of the following compounds and ions:

NH4+1

CH2ClBr
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PO4-3
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Valence electrons are the only electrons
involved in bonding, and are the only ones
written when drawing electron dot structures.
In forming compounds, atoms tend to
achieve the electron configuration of a noble
gas, having 8 valence electrons which as
known as having a stable octet (octet for 8
valence electrons).
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Also known as electron dot symbols
Use symbol of element to represent nucleus and inner
electrons
Use dots around the symbol to represent valence
electrons
Elements in the same group have the same Lewis
symbol
◦ Because they have the same number of valence
electrons
Cations have Lewis symbols without valence electrons
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Anions have Lewis symbols with 8 valence electrons
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Structural formula – chemical formulas that
show the arrangement of atoms in molecules
and in polyatomic ions.
Octet rule – atoms gain or lose electrons to
acquire the stable electron configuration of a
noble gas, usually having 8 valence electrons.
Exceptions to the octet rule:
◦ H needs 2 electrons to be stable
◦ Be needs 4 electrons to be stable
◦ B needs 6 electrons to be stable
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You can represent the formation of the
covalent bond in H2 as follows:
– This uses the Lewis dot symbols for the
hydrogen atom and represents the covalent
bond by a pair of dots.
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The shared electrons in H2 spend part of the
time in the region around each atom.
:
H H
– In this sense, each atom in H2 has a helium
configuration.
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The formation of a bond between H and Cl to
give an HCl molecule can be represented in a
similar way.
– Thus, hydrogen has two valence electrons
about it (as in He) and Cl has eight valence
electrons about it (as in Ar).
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Formulas such as these are referred to as
Lewis electron-dot formulas or Lewis
structures.
bonding pair
lone pair
– An electron pair is either a bonding pair
(shared between two atoms) or a lone pair
(an electron pair that is not shared).
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Single covalent bond – a bond in which two
atoms share a pair of electrons.
Double covalent bond – a bond in which two
atoms share two pairs of electrons.
Triple covalent bond – a bond in which two
atoms share three pairs of electrons.
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1.
2.
Steps for Drawing Lewis-dot structures
Determine the number of valence electrons
in the molecule.
- When drawing determining valence electrons
for an ion, add electrons if it an anion, and
subtract electrons if it is a cation.
The first element in the compound will be
the central atom. Exception: hydrogen will
never be the central atom.
Steps for Drawing Lewis-dot Structures
3. Use one pair of electrons to bond each outer
or terminal atom to the central atom.
4. Make all outer or terminal atoms stable
using the valence electrons (8 total dots
except for Hydrogen which only needs 2).
5. Put any remaining electrons around the
central atom as lone pairs.

Draw the Lewis structure for:
NH3

PO43-
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CHFClBr
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PF5
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Single covalent bond – a bond in which two
atoms share a pair of electrons.
Double covalent bond – a bond in which two
atoms share two pairs of electrons.
Triple covalent bond – a bond in which two
atoms share three pairs of electrons.
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Single bonds are longer (length between the
atoms) than double and triple bonds.
Double bonds are longer than triple bonds.
Single bonds are not as strong as double
bonds, and can be broken much easier than
double bonds.
Triple bonds are stronger than double bonds.
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If you have used up all of the valence
electrons and you still need two more
electrons to make the central atom stable,
you must have one double bond.
If you still need four more electrons to
make the central atom stable, you must
have either one triple bond or two double
bonds.
Double and triple bonds exist most
commonly between C, N, O, and S atoms.
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Draw Lewis structures for:
NOCl
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CO2
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N2
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SiO3-2
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The valence-shell electron pair repulsion
(VSEPR) model predicts the shapes of
molecules and ions by assuming that the
valence shell electron pairs are arranged as
far from one another as possible.
– To predict the relative positions of atoms
around a given atom using the VSEPR model,
you first note the arrangement of the electron
pairs around that central atom.
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The following rules and figures will help
discern electron pair arrangements.
1. Draw the Lewis structure
2. Determine how many bonding pairs are
around the central atom. Count a multiple
bond as one pair.
3. Determine how many lone pairs, if any, are
around the central atom.
All diatomic molecules have a linear shape.
2 pairs
Linear
3 pairs
Trigonal planar
5 pairs
Trigonal bipyramidal
4 pairs
Tetrahedral
6 pairs
Octahedral
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NH3

NOCl

PO43-

N2

CHFClBr
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H2S
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CO2

SiO3-2
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Nonpolar covalent bond – equal sharing of
electrons between two atoms.
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Polar covalent bond – unequal sharing of
electrons between two atoms.
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In polar covalent bonds the electrons are
pulled closer to the atom with the larger
electronegativity value.
Polar bonds can create polar or nonpolar
molecules and ions.
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The easiest way to determine if a molecule or
ion is polar or nonpolar is to look at the
central atom.
If the central atom has lone pairs of electrons,
the molecule or ion is polar.
If the central atom does not have any lone
pairs of electrons, the molecule or ion is
nonpolar.
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SO2
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PO4-3
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N2
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BrO2-1
Attractions Between Molecules
 Molecules are attracted to one another by a
variety of forces.
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These intermolecular forces are weaker than
ionic or covalent bonds.
These forces are responsible for whether or
not a molecular compound is a solid, liquid,
or a gas.
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van der Waals forces – consist of dispersion
forces and dipole interactions (dipole-dipole
moments).
Dispersion forces – weakest of all
intermolecular forces. They are caused by
the motion of electrons. The strength of
dispersion forces increases with the
increasing number of electrons in a molecule.
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All molecules contain dispersion forces.
As molar mass and the number of electrons
increase, dispersion forces increase.
Halogens are the most common molecules to
have dispersion forces. Fluorine is a gas,
Bromine is a liquid and Iodine is a solid.
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Dipole interactions – occur when polar
molecules or ions are attracted to one
another. This occurs when a partial positive
charge and a partial negative charge come
close to each other.
Dipole interactions are very similar to, but
much weaker than ionic bonds.
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Hydrogen bonds – force exerted between a
hydrogen atom bonded to an F, O, or N atom
in one molecule and an unshared pair on
another F, O, or N atom in a nearby molecule.
Hydrogen bonds can have a great effect on
the boiling point of a substance.