Transcript PowerPoint Notes

IV

Lesson Essential Questions:
› Why do atoms form chemical bonds?
› How is the type of chemical bond
determined?
Vocabulary: chemical bond, ionic bonding,
covalent bonding, nonpolar-covalent
bonding, polar, polar-covalent bonding

Chemical Bond
› attractive force between atoms or ions that
binds them together as a unit
› bonds form in order to…
 decrease potential energy (PE)
 increase stability
ION
1 atom
2 or more atoms
Monatomic
Ion
Polyatomic
Ion
+
Na
NO3
-
IONIC
COVALENT
Bond
Formation
e- are transferred from
metal to nonmetal
e- are shared between
two nonmetals
Type of
Structure
crystal lattice
true molecules
Physical
State
solid
Solid, liquid, or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Other
Properties
odorous
METALLIC
Bond
Formation
e- are delocalized
among metal atoms
Type of
Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water
no
Electrical
Conductivity
yes
(any form)
Other
Properties
malleable, ductile,
lustrous
Ionic Bonding - Crystal Lattice
RETURN
Covalent Bonding - True Molecules
Diatomic
Molecule
RETURN
Metallic Bonding - “Electron Sea”
RETURN

Most bonds are a
blend of ionic and
covalent
characteristics.

Difference in
electronegativity
determines bond
type.

Electronegativity
› Attraction an atom has for a shared pair of
electrons.
› higher e-neg atom  › lower e-neg atom +

Electronegativity Trend
› Increases up and to the right.

Nonpolar Covalent Bond
› e- are shared equally
› symmetrical e- density
› usually identical atoms

Polar Covalent Bond
› e- are shared unequally
› asymmetrical e- density
› results in partial charges (dipole)
+


Nonpolar
Polar
Ionic
View Bonding Animations.
Examples:

Cl2

HCl

NaCl
3.0-3.0=0.0
Nonpolar
3.0-2.1=0.9
Polar
3.0-0.9=2.1
Ionic

What type of bonding would be
expected between the following atoms?
› Li and Cl
› Ca and Ga
› I and Cl
› K and Na

Lesson Essential Questions:
› How is a molecular compound formed?
› What are some of the characteristics of a
covalent bond?
Vocabulary: molecule, chemical formula,
molecular formula, bond energy, electron-dot,
Lewis structure, structural formula, single bond,
multiple bonds, resonance
Covalent bond – bond that is created by
the sharing of electrons
 Molecule – neutral group of atoms held
together by covalent bonds
 Molecular compound – chemical
compound made of molecules

CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2

Potential Energy
› based on position of an object
› low PE =
high stability

Potential Energy Diagram
attraction vs. repulsion
no interaction
increased
attraction

Potential Energy Diagram
attraction vs. repulsion
increased
repulsion
balanced attraction
& repulsion

Bond Energy
› Energy required to break a bond
Bond
Energy
Bond
Length

Bond Energy
› Short bond = high bond energy

Electron Dot Diagrams
1. Pick the central atom
2. Count the valence electrons (they are what
electron dot diagrams show)
3. Place electrons around the atom

Octet Rule
› Most atoms form bonds in order to obtain 8
valence e› Full energy level stability ~ Noble Gases
Ne

Nonpolar Covalent - no charges

Polar Covalent - partial charges
+
+

On page 186 in your text book do practice
problems #1-4
1.
Draw the Lewis structure of ammonia, NH3
Draw the Lewis structure for hydrogen
sulfide, H2S
Draw the Lewis structure for silane, SiH4
Draw the Lewis structure for phosphorus
trifluoride, PF3
2.
3.
4.

Some elements can share more than
one electron pair.
› Double bond (two pairs of electrons are
shared)
› Triple bond (three pairs of electrons are
shared)

Draw Lewis structures for each of the
following molecules:
› O2
› CO2
› N3
› N2
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.
Note the different location of the
double bond
Neither structure is correct, it is actually a
hybrid of the two. To show it, draw all varieties
possible, and join them with a double-headed
arrow.
Resonance
in a carbonate
ion (CO32-):
Resonance
in an acetate
ion (C2H3O21-):

Prefix System (binary compounds)
1. Less e-neg atom
comes first.
2. Add prefixes to indicate # of atoms. Omit
mono- prefix on first element.
3. Change the ending of the
second element to -ide.
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
CCl4
carbon tetrachloride
N2O
dinitrogen monoxide
SF6
sulfur hexafluoride
arsenic trichloride
AsCl3
dinitrogen pentoxide
N2O5
tetraphosphorus decoxide
P4O10

Lesson Essential Questions:
› How is an ionic bond formed?
› What are some of the characteristics of an
ionic bond?
Vocabulary: ionic compound, formula unit,
lattice energy, polyatomic ion

Ionic compound – composed of positive
and negative ions that are combined so
that the charges are equal.
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2
Electron dot notation is used to note
changes.
 Form to create an atmosphere of
stability


Covalent – show sharing of e-

Ionic – show transfer of e-
Ions minimize potential energy in crystals by
forming a crystal lattice.
 Distance between all ions represent a
balance of attraction between oppositely
charged particles and repulsion between
like charged particles


Lattice Energy
› Energy released when one mole
of an ionic crystalline compound
is formed from gaseous ions





Ionic
High melting
temperature
High boiling point
Hard
Brittle, because slight
shift of crystal can
cause it to break
Conduct electricity
when dissolved in
water




Covalent
Low melting
temperature
Low boiling point
Do not conduct
electricity
Not as brittle
Ionic Formulas

Write each ion, cation first. Don’t show
charges in the final formula.

Overall charge must equal zero.

Use parentheses to show more than one
polyatomic ion.

Stock System - Roman numerals indicate
the ion’s charge.
› If charges cancel, just write symbols.
› If not, use subscripts to balance charges.
Ionic Names

Write the names of both ions, cation first.

Change ending of monatomic ions to -ide.

Polyatomic ions have special names.

Stock System - Use Roman numerals to
show the ion’s charge if more than one is
possible. Overall charge must equal zero.

Consider the following:
› Does it contain a polyatomic ion?
 -ide, 2 elements  no
 -ate, -ite, 3+ elements  yes
› Does it contain a Roman numeral?
 Check the table for metals not in Groups 1 or
2.
› No prefixes!
Common Ion Charges
1+
0
2+
3+ NA 3- 2- 1-
potassium chloride
K+ Cl-

KCl
magnesium nitrate
Mg2+ NO3-

Mg(NO3)2
copper(II) chloride
Cu2+ Cl-

CuCl2
NaBr
sodium bromide
Na2CO3
sodium carbonate
FeCl3
iron(III) chloride

Lesson Essential Questions:
› How is a metallic bond formed?
› What are some of the characteristics of a
metallic bond?
Vocabulary: metallic bond, alloy
Metal ions held together by attraction to
free floating electrons. (Sea of electrons)
 Good conductors of electricity – Why?

Malleable
 Ductile
 Bond strength – related to enthalpy of
vaporization

› The more energy required to vaporize, the
stronger the bond.
› See table on page 196.
A mixture of two or more substances,
one of which must be a metal.
 Common alloys include steel, 14K gold,
18K gold, cast iron, sterling silver, and
bronze.
 Within different alloys, there can be
different types of mixtures – ex. Steel
 Where do we find alloys?


Use the 3 circle Venn diagram to compare
and contrast ionic, metallic, and covalent
bonding.
Venn Diagram on Chemical Bonding
Name______________________________________
Ionic
Covalent
Metallic

Lesson Essential Questions:
› How is the VSEPR Theory useful?
› What are the different forces present in
bonding?
Vocabulary:
VSEPR theory, hybridization, dipole, hydrogen
bonding, London dispersion forces

Valence Shell Electron Pair Repulsion
Theory

Electron pairs orient themselves in
order to minimize repulsive forces.

Types of e- Pairs
› Bonding pairs - form bonds
› Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!

Lone pairs reduce the bond angle
between atoms.
Bond
Angle

Draw the Lewis Diagram.

Tally up e- pairs on central atom.
› double/triple bonds = ONE pair

Shape is determined by the # of
bonding pairs and lone pairs.
Know the 8 common shapes &
their bond angles!
2 total
2 bond
0 lone
BeH2
LINEAR
180°
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
3 total
2 bond
1 lone
SO2
BENT
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
4 total
2 bond
2 lone
H2O
BENT
5 total
5 bond
0 lone
PCl5
TRIGONAL
BIPYRAMIDAL120°/90°
6 total
6 bond
0 lone
SF6
OCTAHEDRAL
 PF3
4 total
3 bond
1 lone
TRIGONAL
PYRAMIDAL
107°
 CO2
2 total
2 bond
0 lone
LINEAR
180°

Identify the molecular geometry for the
following molecules:
› HI
› CBr4
› CH2Cl2
Intermolecular forces = forces between
molecules.
› The boiling point of a liquid is a good
measure of the intermolecular forces
between its molecules: the higher the
boiling point, the stronger the forces
between the molecules.
 Types of intermolecular forces
› Dipole-dipole forces
› Hydrogen bonding
› London dispersion forces


Dipole – created by equal but opposite charges
that are separated by a short distance.
A dipole is represented by an arrow with its head
pointing toward the negative pole and a crossed
tail at the positive pole. The dipole created by a
hydrogen chloride molecule is indicated as follows:
H Cl

Dipole-dipole forces are the forces of attraction
between polar molecules.
The negative region in one polar molecule
attracts the positive region in adjacent
molecules. So the molecules all attract each
other from opposite sides.
Dipole-dipole forces act at short range, only
between nearby molecules.

Hydrogen bonding = intermolecular
force in which a hydrogen atom that is
bonded to a highly electronegative
atom is attracted to an unshared pair of
electrons in a nearby molecule.

London Dispersion Forces =
intermolecular attractions resulting from
the constant motion of electrons and the
creation of instantaneous dipoles.
http://itl.chem.ufl.edu/2045/matter/FG11_005.GIF
Modern Chemistry Textbook
 www.nclark.net
 http://mrsj.exofire.net/chem/
 http://cottonchemistry.bizland.com/che
m/chemnotes1.htm
 http://www.unit5.org/chemistry/
