Name _________________________
Block _________
Unit 4
Bonding Notes

Atoms are generally found in nature in combination held together by _______________
____________ .

A chemical bond is a _______________________________ between the nuclei and
valence electrons of different atoms that binds the atoms together

There are three main types of bonding: __________, ___________ and __________.

Ionic Bonding occurs between a ______________ and a ______________.

Metallic bonding occurs between two ____________________.

Covalent bonding occurs between a ______________ and a ______________.

A positive ion is called a _____________________________.

A negative ion is called an_____________________________.
What determines the type of bond that forms?

The ______________________ of the two atoms involved are redistributed to the
___________________________________________.

The interaction and rearrangement of the _______________________ determines
which type of bond that forms.

Before bonding the atoms are at their highest possible _______________________.
There are __________ philosophies of atom to atom interaction.

One understanding of the formation of a chemical bond deals with balancing the
opposing forces of _____________________________________.
-

Repulsion occurs between the _________________ clouds of each atom.

Attraction occurs between the __________________ and the negative electron clouds.

When two atoms approach each other closely enough for their electron clouds to begin
to overlap

The electrons of one atom begin to
_______________ the electrons of the other atom
1

And repulsion occurs between the _______________ of the two atoms.

As the optimum distance is achieved that balances these forces, there is a release of
potential energy

The atoms _________________ within the window of maximum attraction/minimum
repulsion

The _______________ released the stronger the _______________ bond between the
atoms

Another understanding of the formation of a chemical bond between two atoms centers
on achieving the most __________________________ of the atoms’ valence electrons

By rearranging the electrons so that each atom achieves a ______________________
________________ ________________ creates a pair of stable atoms (____________
___________________________________)

Sometimes to establish this arrangement one or more valence electrons are ________
______________________________________________.

Basis for _________________ bonding

Sometimes valence electrons are _______________________________________

Basis for _____________________ bonding

A good predictor for which type of bonding will develop between a set of atoms is the
difference in their __________________________________________..

Remember, electro-negativity is a measure of the ______________________ an atom
has for e-s after developing a bond
2

The _________________________ the difference between the two atoms, __________
___________________________________________________ of electrons

Let’s consider the compound Cesium Fluoride, _______________.

The electro-negativity value (EV) for Cs is ____________; the EV for F is ___________.

The difference between the two is _____________, which falls within the scale of ionic
character.

When the electro-negativity difference between two atoms is greater than
______________________________________.
The take home lesson on electro-negativity and bonding is this:
3

-
-
The ________________________ are on the P.T., the more evenly their e interact, and
are therefore more likely to form a _______________________
-

The ________________________________ are on the P.T., the less evenly their e
interact, and are therefore more likely to form an ionic bond.

Forming ionic bonds can be represented as _____________________________.


Hint - metal w/nonmetal = ionic
nonmetal w/nonmetal = covalent
Covalent Bonding


In a Covalent bond:
The ___________________________ between the atoms involved is not extreme

So the interaction between the involved electrons is ____________________________

It may not be an equal sharing relationship, but at least the electrons are being
“___________________”.

Covalent bonds are between _____________________________________________.

Covalent bonds are formed when electrons are _____________________ between two
atoms.
If two atoms share 2 electrons, they form a _____________________________. If two

atoms share 4 electrons, they form a _____________________________.
If two atoms share 6 electrons, they form a ____________________________.

There are two types of covalent bonds: polar and non-polar. Polar bonds usually involve
__________________________________. Non-polar bonds usually involve
___________________________________.

In polar bonds, the electrons are shared ________________________________. In nonpolar bonds, the electrons are shared ________________________________.
4

Covalent compounds can exist in any state (solid, liquid or gas). They have
_____________ melting and boiling points.

A ____________________ is a diagram showing the arrangement of valence electrons
among the
atoms in a molecule.

We use ______________ to represent valence electrons involved in bonding.

______________________ are electrons present in the outermost energy level of an
atom.
Examples of Lewis Dot Symbol
5

Recall that the valence electrons for the elements can be determined based on the
elements position on the ________________.

Atoms can _________________ more than one electron pair.

They may double or triple up pairs of electrons to satisfy the _________________.

A _______________________ is the sharing of one pair of electrons between two
atoms.

A __________________________ is the sharing of ______ pairs of electrons between
two atoms.

A ____________________ is the sharing of _______ pairs of electrons between two
atoms.

_____________ element can form depends on the number of valence electrons.
Rules for Writing Lewis Structures

Determine whether the compound is __________________________. If covalent, treat
the entire molecule. If ionic, treat each ion separately.

Determine the __________ number of _______________ _________________ available
to the molecule or ion by summing the valence electrons of all the atoms in the unit.

Organize the atoms so there is _______________ ____________________ (usually the
least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central
atom.

So what is the bottom line? To be stable the two atoms involved in the covalent bond
share their electrons in order to achieve the arrangement of a ________________.
6
Ionic Bonds

In an Ionic bond:

The electro-negativity difference is ____________________,

So the atom with the _______________________ pull doesn’t really share the electron

Instead the electron is essentially ___________________________________________
_______________________________________________________________________

When a metal bonds with a nonmetal, an
bond is
formed.

An ionic bond always involves the TRANSFER of electrons from the
to the
.

The cation and anion are held together by ___________________________________.

An ionic compound does not consist of individual molecules. Instead, there is a huge
network of positive and negative ions that are packed together in a
.

Because their bonds are so strong, ionic compounds tend to have very
_________________ points.

Ionic compounds are _____________________________, which means they can
conduct electricity.

When forming ionic compounds the positive and negative charges must
_____________.

_______________________________________________________________________
_______________________________________________________________________ .

An electron is _______________________ from the sodium atom to the chlorine atom.

The bottom line is, to be stable the two atoms involved in the ionic bond will either
___________________ or _________________ their valence electrons in order to
achieve a stable __________________ arrangement of electrons.
7
Metallic Bonding

Metallic bonds consist of positively charged metallic cations that donate electrons to the
__________________.

The “sea” of electrons are shared by all atoms and can move throughout the structure.

Properties:
o
o
o
o
Thermal Conductivity
Electrical Conductivity
Malleability- the ability to be hammered down into thin sheets.
Ductility- the ability to be drawn into a wire.

In a metallic bond:

The resulting bond is a cross between ____________ and _____________ bonding

Valence electrons are transferred from one metal atom to the __________________ metal
atoms

But none of the involved metal atoms want the electrons from the original atom, nor their
own so they ___________________________________

What results is a ________________________________ of valence electrons that none
of the atoms in the collection ________________ the valence electrons

It resembles collection of positive ions floating around in a sea of electrons
Determining the Type of Bond:
Determine the type of bond (Ionic, Covalent or Metallic) in the following compounds:
Compound
NaCl
CO
FeNi
SiS2
Bond Type
Compound
NCl3
PF3
CaCl2
Fe2O3
Bond Type
End of PPT 1
8
Draw the Lewis structures for the following compounds:
1) PBr3
4)
NO2-1
2) N2H2
5)
C2H4
3) CH3OH
5)
HBr
Determine if the elements in the following compounds are metals or non-metals. Describe
the type of bonding that occurs in the compound.
Compound
NO2
Element 1
(metal or non-metal?)
N = non-metal
Element 2
(metal or non-metal?)
O = non-metal
Bond Type
covalent
NaCl
SO2
PO43MgBr2
CaO
H2O
Cu-Zn alloy
9
Rules of writing formulas:




· positive ion is written first … this is usually a metal
· negative ion is written second … this is usually a nonmetal
· subscripts are used to show how many ions of each part are in the compound. They
are used to balance the charge of the ions.
· criss-cross method:
Examples:
1.
sodium oxide
sodium is the positive ion = +1
oxide is the negative ion = -2
therefore … it takes 2 sodium ions to balance the charge of the oxide
Formula = Na2O
2.
calcium nitrate
calcium is the positive ion = +2
nitrate is the negative ion = -1
therefore … it takes 2 nitrates to balance the charge of calcium
Formula = Ca(NO3)2
3.
aluminum sulfide
aluminum is the positive ion = +3
sulfide is the negative ion = -2
therefore … it takes 2 aluminum ions and 3 sulfide to balance the charge
Formula = Al2S3
10
Covalent Naming
 Binary covalent compounds are characterized by having two nonmetals.
Naming these compounds involves the use of numerical prefixes:
Prefix
Number
Prefix
Number
1
6
2
7
3
8
4
9
5
10
 If there is only ONE atom of the first element, you DON’T need a prefix. The
FIRST element is named as a normal element. The SECOND element has an
–IDE ending.

o N2O4
o XeF4
Diarsenic pentoxide
o Phosphorous pentabromide
o Carbon tetraiodide
o N2O5
o Trisilicon tetranitride
o CO
o Tetraphosphorous decoxide
o CBr4
11
Write the correct formulas for each covalent compound:
Compound Name
Oxidation States
water
O (-2)
H (+1)
Carbon Dioxide
C (+4)
O (-2)
Chlorine (Diatomic Element)
Cl (-1)
Methane (5 total atoms)
C (-4)
H (+1)
Ammonia (4 total atoms)
N (-3)
H (+1)
Carbon tetrabromide (5 total
atoms)
C (+4)
Br (-1)
Phosphorous trichloride (4 total
atoms)
P (-3)
Cl (-1)
Diphosphorous trioxide (5 total
atoms)
P (-3)
O (-2)
Covalent Formula
Balancing Charges
Criss-Cross rule
1. Write out symbols and charge of elements
2. Criss-Cross charges as subscripts (Swap and Drop)
3. Combine as a formula unit
Equation Form of Balancing Charges
(Number of Cations)x(Cation Charge) + (Number of Anions)x(Anion Charge) = 0
(1)(+3) + (X)(-1) = 0, x = 3
o EX: Aluminum and Oxygen
EX: Barium and Oxygen
12

Balancing Charges Practice:
o Lithium Iodide
o Strontium Chloride
o Sodium Sulfide
In each box, write the formula of the ionic compound consisting of the positive ion to the left of
the box and the negative ion above the box.
-2
-3
-3
Cl
Mg
S
F
N
O
P
+2
Cs+
Cr+3
Na
Zn+2
Al+3
K
I have provided you with charges on some but not all the elements above.
13
Valence Shell Electron Pair Repulsion Theory (VSEPR Theory).
This is the way that we predict the geometry shape of molecules, A model was developed a
qualitative model known as Valence Shell Electron Pair Repulsion Theory (__________________
Theory).
The basic assumptions of this theory are summarized below.
1) The electron pairs in the valence shell around the central atom of a molecule repel each other
and tend to orient in space so as to minimize the repulsions and maximize the distance between
them.
2) There are two types of valence shell electron pairs:
______________ pairs and
________________ pairs
Bond pairs are ______________ _________________ by two atoms and are attracted by two
nuclei. Hence they occupy less space and cause less repulsion.
Lone pairs ________________ ______________ involved in bond formation and are in
attraction with only one nucleus. Hence they occupy more space. As a result, the lone pairs
cause more repulsion.
Note: The bond pairs are usually represented by a ___________ _______________, whereas
the lone pairs are represented by a lobe with two electrons.
3) In VSEPR theory, the ___________________ bonds are treated as if they were single bonds. The
electron pairs in multiple bonds are treated collectively as a single super pair.
4) The shape of a molecule can be predicted from the number and type of valence shell electron
pairs around the central atom.
When the valence shell of central atom contains only bond pairs, the molecule assumes
symmetrical geometry due to even repulsions between them.
14
Bonds
VSPER- Valence Shell Electron Pair Repulsion
Lone Pairs
Shape
Linear
Bent
Trigonal Planar
Trigonal Pyramidal
Tetrahedral
Steric number is the total number of ______________________________ AND
______________________.
15
VSEPR Practice
Complete the table with the requested information.
Molecule
Structural Diagram
Oxidation State of
each element
Molecular
Geometry
CClF3
SF2
BF3
SiBr4
NH3
16
Bonding Notes Activity:
Bond Type
Formula
(Ionic/Covalent)
Lewis
Structure
Shared
Lone
Total
Electron Electron Electron Shape
Pairs
Pairs
Pairs
Structural
Formula
NaCl
NH3
SCl2
CH4
PCl3
MgCl2
CO2
SiBr4
H2O
BCl3
17
Polarity
Bond Polarity
Electronegativity
Ionic bonds have an electronegativity difference that is greater than 1.7. Covalent bonds have
an electronegativity difference less than (or equal to) 1.7. Electronegativity differences
between 0 and 0.4 indicate non-polar covalent bonds. Electronegativity differences between
0.4 and 1.7 indicate polar covalent bonds.
Polar Covalent Bond- a covalent bond in which the electrons are not shared equally because
one atom attracts them more strongly than the other.
Non-polar Covalent Bond- a covalent bond in which the electrons are shared equally.
Use the periodic table of electronegativities to answer the questions on electronegativity
differences.
H
2.1
Li Be
1.0 1.5
ELECTRONEGATIVITY
(electron attraction!)
Na Mg
0.9 1.2
B C N O F
2.0 2.5 3.0 3.5 4.0
Al Si P S Cl
1.5 1.8 2.1 2.5 3.0
K Ca
0.8 1.0
Sc
1.3
Ti V Cr
1.5 1.6 1.6
Mn
1.5
Fe Co Ni Cu Zn Ga Ge As Se Br
1.8 1.9 1.9 1.9 1.6 1.6 1.8 2.0 2.4 2.8
Rb Sr
0.8 1.0
Y
1.2
Zr Nb Mo
1.4 1.6 1.8
Tc
1.9
Ru Rh Pd Ag Cd In Sn Sb Te I
2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5
Cs Ba La-Lu Hf Ta W
0.7 0.9 1.0-1.2 1.3 1.5 1.7
Re
1.9
Os Ir Pt Au Hg Tl Pb Bi Po At
2.2 2.2 2.2 2.4 1.9 1.8 1.9 1.9 2.0 2.2
Fr Ra
0.7 0.9
Ac
1.1
Th Pa U Np-No
1.3 1.4 1.4 1.4-1.3
Determine the type of bond that would form between the following two elements using
differences in electronegativity.
Example: Mg – O
O is 3.5 and Mg is 1.2, therefore, the difference is 3.5 – 1.2 = 2.3 IONIC
Example: Cl – Cl
Cl is 3.0. The difference is 3.0 – 3.0 = 0
NON-POLAR COVALENT
18
Bond
Electronegativity Difference
Bond Type
Bond
Electronegativity Difference
Bond Type
1. C – N
2. Li – F
3. N – Cl
4. Na - Cl
5. O – F
6. B – H
7. Ba – F
8. C – H
Molecular Polarity
Dipole moment- a property of a molecule whereby the charge distribution can be represented
by a center of positive charge and a center of negative charge.
Polar Molecule- a molecule that has a permanent dipole moment.
Determining if a molecule is polar.
1. If ALL of the bonds are non-polar, then the molecule is non-polar.
2. If some or all of the bonds are polar, you can consider the vectors. Vectors are arrows
that point in the direction of the negative charge (the direction the electrons are
pulled). Examples:
O
POLAR
H
H
a. If all the arrows point toward the central atom AND the central atom has lone
pairs, then the atom is polar.
H
NON-POLAR
H
C
H
H
19
b. If all the arrows point toward the central atom AND the central atom has no lone
pairs, then the atom is non-polar.
H
F
POLAR
c. If the arrow points toward one atom in a linear molecule, it is polar.
d. If the arrows all point away from the central atom AND there are no lone pairs
on the central atom, then the atom is non-polar.
e. If the arrows all point away from the central atom AND there are lone pairs on
the central atom, then the atom is polar.
3. Another way to determine if a molecular is polar or not, is to look at symmetry.
H
NON-POLAR
H
C
H
O
C
O
H
a. If the molecule is symmetrical, then the molecule is non-polar.
b.
H
H
C
H
F
POLAR
H
Cl
c. If the molecule is not symmetrical, then the molecule is polar.
d. If the molecule is bent or trigonal pyrimidal, then the molecule is polar (lone
pairs on the central atom mean that it is polar).
20
Draw the structural formula for each molecule, and determine if it is polar or non-polar.
Formula
Structural Formula
Polar/Non-polar
NH3
SCl2
CF4
PCl3
H2 S
C2H2
21
Bonding and Shape Test Review
Determine the formula for the compound formed by the two atoms and indicate
if it is an ionic or covalent compound
1. Calcium and Oxygen
2. Nitrogen and Fluorine
3. Sodium and Chlorine
4. Carbon and Oxygen
Draw the dot diagram for each of the IONIC compounds below
5. CaO
6. Na2S
7. SrF2
10. KI
Review; http://www.youtube.com/watch?v=9DKId82_rUc
Complete the table below.
Formula
Electron Dot
Diagram
Bonding
Orbitals
Shape
Structural Formula
Polar?
NCl3
CO2
H2O
22
CH3F
VSPER Worksheet:
1)
What is the main idea behind VSEPR theory?
2)
For each of the following compounds, draw the Lewis diagram, molecular
shapes, and structural diagram for all atoms:
a)
carbon tetrachloride
b)
BH3
c)
silicon disulfide
d)
PF3
e) carbon dioxide
f) SF2
23
Unit 4 Test Review
1. Calcium and Oxygen
2. Carbon and Fluorine
Complete the table below
Formula
Electron Dot Diagram
Shape
Structural Formula
Polar?
CH4
H2O
BH3
NH3
CO2
Draw the Lewis dot diagram for the compounds below
6. AlN
7. HgCl2
8. SCl2
9. NaF
10. BF3
11. SiCl4
Write the formula for the following compounds
12. Aluminum Bromide
13. Dinitrogen Tetroxide
14. Ammonium Fluoride
15. Iron (II) Sulfate
16. Copper (I) Chloride
17. Carbon Dioxide
24