1
Bell Ringer
1) Pick up the practice sheet off the
front demo table and complete it.
2) Get out your notes from yesterday
Schedule
1. Bell Ringer
2. Periodic Trends
Practice
3. Go over Tests
4. Bonding Notes
I CAN……solve chemistry problems
by being an independent, creative
thinker.
You gain strength, courage, and
confidence by every experience in
which you really stop to look fear in
the face. You must do the thing
which you think you cannot do.
2
1
2
3
4
5
6
7
•
•
•
•
•
Alkali Metals
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
3
• Atomic Radius
–size of atom
First Ionization Energy
Energy required to remove one e- from a
neutral atom.
© 1998 LOGAL
Melting/Boiling Point
© 1998 LOGAL
4
Atomic Radius
• Atomic Radius
Increases as you move DOWN a family
Decreases as you move ACROSS a period.
1
2
3
4
5
6
7
5
Atomic Radius
• Why larger going down?
–Higher energy levels have larger orbitals
–Shielding - core e- block the attraction between the
nucleus (positive protons) and the valence
(negative) e• Why smaller to the right?
–Increased nuclear charge without additional
shielding pulls e- in tighter (adding electrons and
protons, but not adding energy levels)
6
Ionization Energy
• First Ionization Energy
Increases ACROSS a period
Decreases DOWN a group/family
1
2
3
4
5
6
7
7
Ionization Energy
• Why opposite of atomic radius?
–In small atoms, e- are close to the nucleus where
the attraction is stronger
• Why small jumps within each group?
–Stable e- configurations don’t want to lose e-
8
F. Melting/Boiling Point
• Melting/Boiling Point
Highest in the middle of a period.
1
2
3
4
5
6
7
9
Ionic Radius
• Ionic Radius
Cations (+)
lose esmaller
Anions (–)
gain elarger
© 2002 Prentice-Hall, Inc.
10
Electronegativity
• Measure of the ability of an atom in a chemical compound to
attract electrons
Increases ACROSS a period
Decreases DOWN a group/family
1
2
3
4
5
6
7
11
Examples
• Which atom has the larger atomic radius?
Be
or Ba
Ba
Ca
or Br
Ca
12
Examples
• Which atom has the higher 1st I.E.?
N or Bi
Ba or
N
Ne
Ne
13
Examples
• Which atom has the higher melting/boiling point?
Li
or C
C
Cr
or Kr
Cr
14
Examples
• Which particle has the larger ionic radius?
S or
Al
2S
or
2S
3+
Al
Al
15
Examples
• Which element has the higher electronegativity?
S or F
F
Al
Al
or In
16
Bell Ringer
1) Get out your notes from yesterday &
answer the following question:
1)
Draw the Lewis Dot Notation for the following
elements:
Na
Cl
B
C
N
O
I CAN……solve chemistry problems
by being an independent, creative
thinker.
Schedule
1. Bell Ringer
2. Go over Tests
3. Bonding Notes
You gain strength, courage, and
confidence by every experience in
which you really stop to look fear in
the face. You must do the thing
which you think you cannot do.
17
Chemical
Bonding
and
Molecular
Structure
Cartoon courtesy of NearingZero.net
18
Bonding
• Chemical Bond
–attractive force between atoms or ions
that binds them together as a unit
–bonds form in order to…
» decrease chemical potential energy (PE)
» increase stability
Intra VS Intermolecular Forces
• INTRA within
• INTER between, among
• 1. Intramolecular force: bonds between atoms or ions in molecules
Examples: metallic, covalent, & ionic bonds
– Hydrogen bonds to oxygen and forms water
• 2. Intermolecular force: attraction between molecules themselves
Intermolecular forces are only associated with covalently
bonded molecules (glucose & fructose hydrogen bridge to
make sucrose)
–3 different levels of strength:
» a) Hydrogen Bridge (used to be called H bond)
(strongest)
» b) Dipole-dipole forces
» c) London Dispersion forces
(weakest) (covalently bonded molecules have this)
– Water molecule attracted to another water molecule
19
Review of Chemical Bonds
• There are 2 main types of
bonding:
• IONIC—transfer of 1 or more
valence electrons
• COVALENT—sharing of valence
electrons
“Between ionic
and covalent
most bonds are”
20
The type of bond can usually be calculated by 21
finding the difference in electronegativity of
the two atoms that are going together.
Electronegativity Difference
• If the difference in electronegativities
is between:
– 1.7 to 4.0: Ionic
– 0.3 to 1.7: Polar Covalent
– 0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!
22
23
Ionic Bonds
Ions are positively or negatively
charged atoms due to the removal
or addition of an ELECTRON
1.
CATION (CAT – ION) is positively
charged – electron removed.
2.
ANION (AN – ION) is negatively
charged – electron gained.
In general
•
metals (Mg) lose electrons --->
cations (+)
•
nonmetals (F) gain electrons --->
anions (-)
Ionic Bonds
24
Positive cations and the negative
anions are attracted to one another
(remember: Opposites Attract!)
Therefore, ionic
compounds are usually
between metals and
nonmetals (opposite ends
of the periodic table).
Comparison
Ionic Compounds
• Crystalline solids (made of
ions)
• High melting and boiling
points
• Conduct electricity when
melted
• Many soluble in water but
not in nonpolar liquid
• Metal and nonmetals bond
25
Covalent Compounds
• Gases, liquids, or solids
(made of molecules)
• Low melting and boiling
points
• Poor electrical conductors
in all phases
• Many soluble in nonpolar
liquids but not in water
• Nonmetals bond with
other nonmetals
26
Diagram Ionic Bonding
NaCl
Al & Cl
• K&S
• Al & O
Get out the practice problems from
Friday and your notes on Bonding 
Paired Excercises:
• 1,5,7,9,11,13,15,23,25,
Happy April 2nd !!
27
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
Covalent Bond
45
A chemical bond in which two or more electrons are shared by two
atoms.
How should two atoms share electrons?
F
+
7e-
F
F F
7e-
8e- 8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
single covalent bond
lone pairs
F F
lone pairs
lone pairs
46
Lewis structure of water
H
+
O
+
H
single covalent bonds
H O H
or
H
O
2e-8e-2eDouble bond – two atoms share two pairs of electrons
O
C O
or
O
O
C
double bonds
- 8e8e- 8ebonds
double
Triple bond – two atoms share three pairs of electrons
N
N
triple
bond
8e-8e
or
N
N
triple bond
H
Double and
even triple
bonds are
commonly
observed for C,
N, P, O, and S
47
H2CO
SO3
C2F4
Wednesday 4/18/12
48
1. Get out your lab from
yesterday and turn it in
to your tray. Get out
paper for notes.
• 1. Draw the Lewis dot
structure for HCl.
• 2. If the penguin represents a
hydrogen atom and the polar
bear represents a chlorine
atom, what does the ice
cream represent in the
drawing? What do you think
the picture is trying to
illustrate?
• 3. Do you think HCl will be
attracted to a charged wand?
Explain your thinking.
Intraparticle vs. Intermolecular
Force
49
INTRAPARTICLE ATTRACTIONS
INTERMOLECULAR FORCES
(Types of Bonding WITHIN a substance;
atom to atom bonding)
(Types of forces BETWEEN neighboring
molecules; molecule to molecule)
Ionic
NA
Metallic
NA
Covalent
Dipole-Dipole
H-bridges
London Dispersion Forces (LDF)
Thursday 4/19/12
• Get out your notes from yesterday.
• Draw the Lewis structure of the following
molecules and identify if they are ionic, polar
covalent, or nonpolar covalent:
–1. NaCl
–2. NO2–3. CH4
50
Polarity
51
H F
+

A molecule, such as HF, that has a center of
positive charge and a center of negative
charge is said to be polar, or to have a dipole
moment.
Dipole Moment
• Direction of the polar bond in a molecule.
• Arrow points toward the more electronegative
atom.
+

H
Cl
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

52
53
Polar Covalent Bond
Electrons are shared unequally
A covalent bond with greater electron density around one of
the two atoms resulting in partial charges (dipole)
e- poor
electron poor
region
H
electron rich
region
F
H
+
e- rich
F
-
Bond Polarity
HCl is POLAR because it
has a positive end and a
negative end. (difference
in electronegativity)
+
-
••
••
H Cl
••
Cl has a greater share in
bonding electrons than
does H.
Cl has slight negative charge (-) and H has
slight positive charge (+ )
54
Bond Polarity
• This is why oil and water will not mix! Oil
is nonpolar, and water is polar.
• The two will repel each other, and so you
can not dissolve one in the other
55
Bond Polarity
• “Like Dissolves Like”
–Polar dissolves Polar
–Nonpolar dissolves
Nonpolar
56
57
Determining Molecular Polarity
• Nonpolar Molecules
-Electrons are shared equally
-Symmetrical electron density
-Often identical atoms
-Dipole moments are symmetrical and cancel out.
F
BF3
F
F
58
Polar Bonds
..
F
N
O
Cl
H
H
H
Polar
F
Polar
H
B
Polar
Cl
F
F
Cl
Polar
F
Cl
F
H
Cl
F
H
F
Nonpolar
H
C
Xe
F
F
Nonpolar
C
Cl
Cl
Nonpolar
H
H
Polar
A molecule has a zero dipole moment because their dipoles cancel one another.
Hydrogen Bridge
59
• Occurs with polar covalent structures only
• Includes a positive Hydrogen ion bonding
directly to a negative ion
• DNA, RNA, and enzymes have many Hydrogen
bonding sites
• Strongest intermolecular force
60
London Dispersion forces
•
•
•
•
Also known as van der Waals bonds
Occur only with nonpolar structures
Weakest intermolecular force
Electrostatic bonds between atoms or
molecules
• Temporary dipole moment
MOLECULAR
GEOMETRY
61
MOLECULAR GEOMETRY
VSEPR
• Valence Shell Electron Pair
Repulsion theory.
• Most important factor in
determining geometry is
relative repulsion between
electron pairs.
Molecule adopts
the shape that
minimizes the
electron pair
repulsions.
62
Some Common Geometries
Linear
Trigonal Planar
Tetrahedral
63
64
65
Other VSEPR charts
Structure Determination by VSEPR
Water, H2O
2 bond
pairs
2 lone
pairs
The molecular
geometry is
BENT.
The electron pair
geometry is
TETRAHEDRAL
66
67
Structure Determination by
VSEPR
Ammonia, NH3
The electron pair geometry is tetrahedral.
lone pair of electrons
in tetrahedral position
N
H
H
H
The MOLECULAR GEOMETRY — the
positions of the atoms — is TRIGONAL
PYRAMID.