VALENCE ELECTRONS &
BONDING
Chapter 4
VALENCE ELECTRONS
• Valence electrons:
• found in the outermost shell of an atom
• determines the atom’s chemical properties.
• So how do we know the number of valence
electrons an element has?
• It based on their group on the periodic table.
1 ve
8 ve
3 ve
2 ve
5 ve
7 ve
4 ve
6 ve
VALENCE ELECTRONS
•
•
•
•
•
•
•
•
•
Lets find some valence electrons.
Si
Xe
K
Ba
I
O
Al
P
OXIDATION NUMBERS
• charge an atom would have if it were in a
compound composed of ions.
• AKA: charge
• change in oxidation number represents
number of electrons gained or lost in a
chemical reaction.
+1
+3
+2
-1
-3
+4
-2
OXIDATION NUMBERS
•
•
•
•
•
•
•
•
Lets find some oxidation numbers
In
F
Rb
Sn
Ca
N
O
Lewis Structures
• AKA: electron-dot structures or electron-dot
diagrams.
• Uses the number of valence electrons.
• indicated by dots placed around the element’s
symbol.
• diagrams that show the bonding between atoms
of a molecule.
• Always place the element with the largest
number of valence electrons first, next largest …
• They can form ionic & covalent bonds
Ionic Bonds
• TRANSFER of electrons from one
bonding atom to another
Ionic Bonds
• BaO
• FIRST WE SEE HOW MANY
ELECTRONS THEY ARE FOR EACH
ELEMENT.
• Ba = 2
• O=6
• DRAW LEWIS STRUCTURE FOR
OXYGEN
Ionic Bonds
Ba
o
Ionic Bonds
MgCl2
COVALENT BONDING
• chemical bond resulting from SHARING
of electrons between 2 bonding atoms.
COVALENT BONDING
• H2O
• FIRST WE SEE HOW MANY
ELECTRONS THEY ARE FOR EACH
ELEMENT.
• H=2
• OXYGEN = 6
• DRAW LEWIS STRUCTURE FOR
OXYGEN
COVALENT BONDING
H
H
O
BONDING
REVIEW
Directions: Each slide provides information in a question answer format.
Click once to see the question and then again for the answer.
Bonds Between Atoms
Polyatomic Ions
Ionic
Covalent
Molecular
Substance
Metallic
Network
Solids
Polar
Nonpolar
What are we going to learn about???
Coordinate
Covalent
See if you can define the following words before starting the
lesson…
• Anion- negative ion
• Cation-positive ion
• Octet Rule- rule that states that atoms tend to gain,
lose, or share electrons so that each atom has full
outermost energy level which is typically 8 electrons.
• Polyatomic Ion- charged group of covalently bound
atoms
• Monatomic Ion- ion formed from a single atom
• Molecule-neutral group of atoms united by covalent
bonds
• Alloy- homogeneous mixture of metals
• Unshared Pair- pair of electrons that is not involved
in bonding but instead is held exclusively by one
atom.
Marriage
Forming
of a bond
is like
marriage
•More stable
•exothermic
Divorce
•Less stable
•Endothermic
The
breaking
of a bond
relates to a
divorce.
Ionic Bonds
•What is an Ionic Bond?
- An Ionic Bond is a chemical bond
resulting from the TRANSFER of electrons
from one bonding atom to another
• When is an ionic bond formed?
- An ionic bond is formed when a cation
(positive ion) transfers electrons to an
anion (negative ion).
What are some characteristics of an ionic
bond?
1. Crystalline at room
temperatures
2. Have higher melting
points and boiling
points compared to
covalent compounds
3. Conduct electrical
current in molten or
solution state but not
in the solid state
4. Polar bonds
Covalent Bonds
•What is an Covalent Bond?
- A covalent bond is a chemical bond
resulting from SHARING of electrons
between 2 bonding atoms.
• What forms a covalent bond?
- A covalent bond is formed between two
nonmetals.
What are some characteristics of a covalent
bond?
1.
2.
3.
Covalent bonds have
definite and predicable
shapes.
Very strong
Low melting and boiling
points
Covalent Bonds can have multiple bonds, so you should be
familiar with the following…
Single Covalent Bondchemical bond resulting
from sharing of an electron
pair between two atoms.
Double Covalent Bondchemical bond resulting
from sharing of two electron
pairs between two atoms.
Triple Covalent Bondchemical bond resulting
from sharing of three
electron pairs between two
atoms.
There are five different categories associated with covalent
bonds. What are the 5 different categories?
Covalent
Molecular
Substance
Network
Solids
Polar
Nonpolar
Coordinate
Covalent
First, we are going to look at Polar Covalent…
What is polar covalent?
-Polar covalent is a description of a
c
bond that has an uneven
distribution
of charge due to an unequal sharing of
bonding electrons.
The boy is not equally
sharing with anyone
else but rather taking
all the food for himself.
Next, we are going to look at Non-Polar Covalent…
What is non-polar covalent?
-Non polar covalent is a covalent
bond that has an even distribution of
charge due to an equal sharing of
bonding electrons.
This couple is
non- polar
because they are
sharing the drink
equally between
them.
Next, we are going to look at Molecular Substances…
What is a molecular substance?
-A molecular substance is a substance
that has atoms held together by
covalent bonds.
Name 2 Characteristics of a
Molecular Substance.
1. Weak
2. Low melting and boiling
points
Next, we are going to look at Coordinate Covalent…
What is a Coordinate Covalent Bond?
-A coordinate covalent bond is a
bond formed when one atom donates
both electrons that are shared.
People donate their
blood to help others
just like atoms
“donate” electrons to
form stable octets.
*Think about the party analogy!
Now, we are going to look at Network Solids…
What is a Network Solid?
-A network solid is a solid that has
covalently bonded atoms linked in
one big network or one big
macromolecule.
Name 3 Characteristics of a
Network Solid.
1. Poor conductors of heat
and electricity
2. Hard / Strong
3. High melting and boiling
points
Diphenylglycoluril
assembles into a unique, twodimensional hydrogen
bonding network in the solid
state, while exhibiting a
twisted molecular structure.
Metallic Bonding
What is a Metallic Bond?
- A metallic bond occurs in
metals. A metal consists of
positive ions surrounded by a
“sea” of mobile electrons.
This
shows
what a
metallic
bond
might
look
like.
Name 4 Characteristics of a
Metallic Bond.
1. Good conductors of heat
and electricity
2. Great strength
3. Malleable and Ductile
4. Luster
Polyatomic Bonds
Polyatomic ions usually have a
charge because the collection of
atoms has either gained an extra
electron or else it has lost an
electron.
What is a Polyatomic Bond?
- A polyatomic bond is
charged group of covalently
bonded atoms. It is made up
of more than one atom.
Linear
Trigonal Planer
Trigonal Bipyramidal
Tetrahedral
Octahedral
Just as a summary to what each
bond looks like…
REVIEW
Now that you have completed the review provided you may wish to
practice for your test by answering the following regents questions.
The question is presented in the same way that
http://regentsprep.org/ does them. If you would like to get reasons to
why your answers were incorrect you can find the explanations
there. After you have gotten your answer, click the mouse and the
right answer will show up.
If you chose…
If you chose…
If you chose…
If you chose…
If you chose…
If you chose…
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
Li
18
He:



Be
B 


C


Na Mg


Al


N


 Si 


O



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
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 13 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.
Change in electrons for octet
1) lose 3e2) gain 3 e3) gain 5
eC.
Ionic charge of aluminum
1) 32) 5-
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.
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
Bonds in all
the
polyatomic
ions and
diatomics are
all covalent
NONPOLAR
COVALENT BONDS
when electrons
are shared
equally
H2 or Cl2
2. 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
Polar Covalent Bonds: Unevenly
matched, but willing to share.
- water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen.
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
CHEMICAL
BONDING
Chemistry I – Chapter
8
Chemistry I Honors –
Chapter 12
Cocaine
SAVE PAPER AND INK!!! When
you print out the notes on
PowerPoint, print "Handouts"
instead of "Slides" in the print
setup. Also, turn off the
backgrounds
(Tools>Options>Print>UNcheck
"Background Printing")!
Chemical
Bonding
Problems and
questions —
How is a molecule or
polyatomic ion held
together?
Why are atoms
distributed at strange
angles?
Why are molecules not
Review of Chemical
Bonds
Most bonds are
somewhere in
between ionic
and covalent.
• There are 3 forms of
bonding:
• _________—complete
transfer of 1 or more
electrons from one atom
to another (one loses,
the other gains) forming
oppositely charged ions
that attract one another
The type of bond can usually be
calculated by 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
Example: NaCl
–
0.3
to
1.7:
Polar
Na = 0.8, Cl = 3.0
Difference
is 2.2, so
Covalent
this is an ionic bond!
– 0.0 to 0.3: Non-Polar
Ionic Bonds
All those ionic compounds
were made from ionic
bonds. We’ve been
through this in great detail
already. Positive cations
and the negative anions
Therefore, ionic
are attracted to one compounds are usually
between metals and
another (remember the
nonmetals (opposite
Paula Abdul Principle ends
of of the periodic
Chemistry: Oppositestable).
Electron
Distribution
in Molecules
G. N. Lewis
1875 - 1946
• Electron distribution
is depicted with
Lewis
(electron dot)
structures
• This is how
you decide
how many
atoms will
Bond and Lone
Pairs
• Valence electrons are
distributed as shared or
BOND PAIRS and
H Cl or LONE PAIRS.
unshared
••
•
•
••
shared or
bond pair
lone pair (LP)
This is called a LEWIS
structure.
Bond Formation
A bond can result from an
overlap of atomic orbitals
on neighboring atoms.
••
H
+
•
•
Cl
••
H
••
•
•
Cl
••
Overlap of H (1s) and Cl (2p)
Note that each atom has a single,
unpaired electron.
Review of Valence
Electrons
• Remember from the electron
chapter that valence
electrons are the electrons in
the OUTERMOST energy
level… that’s why we did all
those electron configurations!
• B is 1s2 2s2 2p1; so the outer
energy level is 2, and there
are 2+1 = 3 electrons in level
2. These are the valence
electrons!
• Br is [Ar] 4s2 3d10 4p5
Review of Valence
Electrons
Number of valence electrons of a
main (A) group atom = Group
number
Steps for Building a Dot
Structure
Ammonia, NH3
1. Decide on the central atom; never H.
Why?
If there is a choice, the central atom is
atom of lowest affinity for electrons.
(Most of the time, this is the least
electronegative atom…in advanced
chemistry we use a thing called formal
charge to determine the central atom. But
that’s another story!)
Therefore, N is central on this one
Building a Dot Structure
3. Form a single bond
H
H
N
between the central
atom and each
H
surrounding atom
bond takes
2 form
4. (each
Remaining
electrons
electrons!)
LONE
PAIRS to complete the
H N H
••
octet as needed (or duet in the
case of H).
3 BOND PAIRS and 1 LONE
PAIR.that N has a share in 4 pairs (8
Note
electrons), while H shares 1 pair.
H
Building a Dot Structure
••
5.Check to make sure there H N H
are 8 electrons around
each atom except H. H
H
should only have 2
electrons. This includes
SHARED
6. Also,
check pairs.
the number of electrons in
your drawing with the number of
electrons from step 2. If you have more
electrons in the drawing than in step 2,
you must make double or triple bonds. If
you have less electrons in the drawing
than in step 2, you made a mistake!
Carbon Dioxide, CO2
C 4 e1. Central atom =
O 6 e- X 2 O’s = 12 e2. Valence electrons =Total: 16 valence
electrons
3. Form bonds.This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons
around it except for H, which can have 2.
Carbon Dioxide, CO2
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
How many are in the drawing?
6. There are too many electrons in our drawing.
We must form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond
shares 2 pairs. So one pair is taken away from
each atom and replaced with another bond.
even triple
bonds are
commonly
observed for
C, N, P, O,
and S
C2F4
H2CO
SO3
Now You Try One!
Draw Sulfur Dioxide, SO2
Violations of the Octet
Rule
(HonorsBonly)
Usually occurs with
and
elements of higher
periods. Common
Be: 4Be, B, P,
exceptions are:
S, and Xe. B: 6
P: 8 OR 10
S: 8, 10, OR 12
BF3
Xe: 8, 10, OR
12
SF4
MOLECULAR
GEOMETRY
MOLECULAR GEOMETRY
VSEPR
• Valence Shell Electron
Pair Repulsion theory.
• Most important factor
in determining
geometry is relative
repulsion between
Molecule
adopts the
shape that
minimizes the
electron pair
repulsions.
Some Common
Geometries
Linear
Trigonal Planar
Tetrahedr
al
VSEPR charts
• Use the Lewis structure to
determine the geometry of the
molecule
• Electron arrangement establishes
the bond angles
• Molecule takes the shape of that
portion of the electron arrangement
• Charts look at the CENTRAL atom
for all data!
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
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.
Bond Polarity
+d -d
••
••
H Cl
••
HCl is POLAR
because it has a
positive end and a
negative end.
Cl has a greater share
(difference
in
in bonding electrons
electronegativity)
than
does H.
Cl has slight negative charge (-d) and H
has slight positive charge (+ d)
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
Bond Polarity
• “Like Dissolves
Like”
–Polar dissolves
Polar
–Nonpolar
dissolves
Nonpolar
Ch. 11 - Chemical Bonds
II. Kinds of Chemical Bonds (p.304308)
 Ionic Bond
 Covalent Bond
 Comparison Chart
A. Ionic Bond
• Attraction between 2 oppositely
charged ions
– Ions - charged atoms
– formed by
transferring efrom a metal
to a nonmetal
A. Ionic Bond
– ions form a 3-D crystal lattice
NaCl
B. Covalent Bond
• Attraction between neutral atoms
– formed by sharing e- between two nonmetals
B. Covalent Bond
– covalent bonds result in discrete molecules
Cl2
NH3
H2O
B. Covalent Bond
• Nonpolar Covalent Bond
– e- are shared equally
– usually identical atoms
B. Covalent Bond
• Polar Covalent Bond
– e- are shared unequally between 2 different
atoms
– results in partial opposite charges
+
d
d
B. Covalent Bond
• Nonpolar
• Polar
• Ionic
View Bonding Animations.
C. Comparison Chart
IONIC
transferred from
metal to nonmetal
COVALENT
shared between
nonmetals
Melting
Point
high
low
Soluble in
Water
yes
usually not
Electrons
Conduct
Electricity
Other
Properties
yes
no
(solution or liquid)
crystal lattice of ions, molecules, odorous
liquids & gases
crystalline solids
Chemical Bondin
g
Valence Electrons
• Electrons in the outer
energy level are
called valence
electrons.
• It is these electrons
that determine the
formation of
chemical bonds.
Valence Electrons
• Certain numbers of
valance electrons
are more stable than
others.
• Atoms will gain or lose
electrons to become
more stable.
Chemical Bond Typ
es
• Ionic
• Covalent
• Metallic
Ionic Bond
Electrons are
transferred from
one atom to another.
Electron Transfe
r
Lithium
Neon
Ionic Bond
• When this happens, the
charges within each
atom are no longer
balanced.
• The atoms have become
ions - one with a
positive charge, the
other with a negative
charge.
Covalent Bond
Electrons are
shared between
atoms.
Covalent Bond
• The nucleus of both
atoms has an equal
attraction for the
electrons.
• The shared electrons
spend most of their
time between the
two atoms.
Polar covalent B o n d
Atoms share the electron unequally,
There is a slight difference in charge
between the two poles of the bond; water
is an example.
Hydrogen Bonding
In a hydrogen bond, an atom or a molecule interacts
weakly with a hydrogen atom already taking part in
a polar covalent bond.
Ionization Energy
• The energy needed to
remove an electron
from an atom.
• The attraction between
the negatively charged
electron and the
positively charged
nucleus must be
overcome.
Electron Affinity
• The tendency of an
atom to attract
electrons.
• Metals have a low
electron affinity.
• Nonmetals have a high
electron affinity.
Diatomic
Molecules
Covalent bonding
occurs between
atoms of the same
element.
Hydrogen
Nitrogen
Oxygen
Florine
Chlorine
Bromine
Iodine
7
Diatomic
Molecules
Polyatomic Ions
• A group of covalently
bonded atoms acting
like a single atom
when forming
compounds.
• This group of atoms is
not electrically
balanced and has an
overall charge.
Predicting Bond Type
s
• Will these pairs form
ionic or covalent bonds?
– sodium & chlorine
– calcium & oxygen
– carbon & oxygen
– aluminum & sulfur
Chemical Bondin
g
End
Ch. 11 - Chemical Bonds
IV. Naming Ionic Compounds
(p. 314-320)
 Oxidation Number
 Ionic Names
 Ionic Formulas
A. Oxidation Number
• The charge on an ion.
• Indicates the # of e- gained/lost to
become stable.
1+
2+
3+ 4+ 3- 2- 1-
0
B. Ionic Names
• Write the names of both elements,
cation first.
• Change the anion’s ending to -ide.
• Write the names of polyatomic ions.
• For ions with variable oxidation #’s,
write the ox. # in parentheses using
Roman numerals. Overall charge = 0.
B. Ionic Names
• NaBr
– sodium bromide
• Na2CO3
– sodium carbonate
• FeCl3
– iron(III) chloride
C. Ionic Formulas
• Write each ion. Put the cation first.
• Overall charge must equal zero.
– If charges cancel, just write the symbols.
– If not, crisscross the charges to find subscripts.
• Use parentheses when more than one
polyatomic ion is needed.
• Roman numerals indicate the oxidation #.
C. Ionic Formulas
• potassium chloride
– K+ Cl

KCl
• magnesium nitrate
– Mg2+ NO3 
Mg(NO3)2
• copper(II) chloride
– Cu2+ Cl

CuCl2
C. Ionic Formulas
• calcium oxide
– Ca2+ O2

CaO
• aluminum chlorate
– Al3+ ClO3

Al(ClO3)3
• iron(III) oxide
– Fe3+ O2

Fe2O3
Chemical Bonds
– Atoms gain or lose electrons through chemical
reactions to gain a filled outer shell and therefore a
lower energy level.
– A chemical reaction forms a chemical bond that is
an attractive force that holds atoms together in a
compound.
– Ionic bonds are formed when an atom transfers an
electron to another atom during a chemical reaction
• The opposite charges resulting forms an
electrostatic attraction between the ions that are
formed.
– Covalent bonds form when atoms share electrons
in a chemical bond.
– Metallic bonds form in metals.
– These new bonds form compounds which can be
described in several ways.
– Molecular orbital theory describes the electrons
as belonging to the whole molecule which gives the
orbital its own shape, orientation, and energy levels.
– Isolated atom description considers the electrons
around the atoms as being isolated from the rest of
the molecule.
• Ionic Bonds
– Ionic bonding occurs when one atom
transfers an electron to another atom
• The difference in electrical charge results
in an electrostatic attraction between unlike
electrical charges
• This occurs when a metal reacts with a
nonmetal
Na+1 + Cl-1  NaCl (table salt)
– Energy and Electrons in Ionic Bonding
• Example:
energy + Na+  Na+ + eCl + e-  Cl- + energy
Na+ + Cl-  NaCl + energy
• The energy that is released in steps 2 and
3 is greater that that absorbed in step one
and an ionic bond is formed.
–This energy is called the heat of
formation.
• Two rules for keeping track of electrons in ionic
bonding reactions.
–Ions are formed when atoms gain or lose
electrons to achieve a noble gas
configuration
–The number of electrons that are lost must
equal the number of electrons that are
gained.
– Ionic Compounds and Formulas
• The formula of a compound describes what
elements are in the compound and in what
proportions.
• Compounds that are held together by ionic bonds
are called ionic compounds.
• The elements in Group IA and IIA tend to lose
electrons for form positive ions
• The elements in Group VIA and VIIA tend to gain
electrons to form negative ions.
• Covalent Bonds
– A covalent bond is a chemical bond that is formed
when two atoms share a pair of electrons.
–H. + H.  H:H
– Covalent Compounds and Formulas
• Since a pair of electrons is shared in a covalent
bond, the electrons move throughout the entire
molecular orbital.
• In the above example, since both hydrogen share
the electron pair, each hydrogen has a filled
valence shell, since it has the electron
configuration of helium.
• Compounds that are held together by covalent
bonds are called covalent compounds.
• Covalent compounds form from atoms on the
right side of the periodic table
– Multiple Bonds.
• In electron dot notations, a pair of electrons
can be represented by a pair of dots : .
–This can be a bonding pair or a lone pair
(non-bonding pair).
• Bonding pairs can also be represented by
lines connecting atoms.
»H:H = H—H
• When one pair of electrons is shared, it is
called a single bond.
»H-H
• When two pairs of electrons are shared it
is called a double bond.
H
H
\
/
C C
/
\
H
H
• When three pairs of electrons are shared
it is called a triple bond.
H-C
C-H
• Electronegativity.
– Electronegativity is the ability of an atom to
attract bonding electrons.
– Elements with higher values have the
greatest attraction for bonding electrons.
– The difference in electronegativity can be
used to predict whether a bond will be ionic
or covalent.
• If the absolute difference is 0.5 or less, the
bond will be covalent.
• If the absolute difference is 1.7 or more the
bond will be ionic.
– When the absolute difference is between 0.5
and 1.7 a covalent bond is formed, but one in
which the electron pair is not shared
equally.
– This type of bond is called a polar covalent
bond.
– A polar covalent bond results in areas of
partial positive charge and areas of partial
negative charge since the electrons spend
more time around the more electronegative
atom.
• Naming chemical Compounds
• Ionic Compound Names
– Ionic compounds that are formed from metal
ions are named by naming the metal ion
(electropositive ion) first, followed by the
nonmetal (electronegative ion)
– The ending of the nonmetal is changed to
end in -ide
– When a metal can have various oxidation
states the oxidation state is give by roman
numerals in parenthesis after the name of the
metal.
Naming the ions
Names of main-group monatomic ions are straightforward.
A cation takes the name of the element plus the word "ion."
Na
Sr
Ba
Na+
Sr2+
Ba2+
sodium ion strontium ion barium ion
For anions, the element name has its
ending replaced with ide.
N
N3nitride ion
O
O2oxide ion
Cl
Clchloride ion
Some transition metals can form more than one
ion. Iron, for example, forms both Fe2+ and Fe3+.
To name such an ion unambiguously, we use the
name of the element, a Roman numeral in
parentheses to denote the charge, and the word
"ion." Fe2+ and Fe3+ would be iron(II) ion and
iron(III) ion, respectively.
• Ionic Compound Formulas
– Rules
• The symbol for the positive element is
written first, followed by the symbol of
the negative element
• Subscripts are used to indicate the
numbers of ions needed to produce an
electrically neutral compound.
Example calcium chloride
–calcium is Ca2+ and chlorine is Cl–in order to balance charges there needs to
be two negative charges to balance the 2+
on the calcium
–the formula is therefore CaCl2
Most of the common polyatomic ions are anions,
although a few are cations. It is important to know
the names, charges, and formulas of all of these ions.
Common polyatomic ion names
Formula
Name
NH4+
Ammonium ion
CO32-
Carbonate ion
PO43-
Phosphate ion
SO42-
Sulfate ion
OH-
Hydroxide ion
NO3-
Nitrate ion
Anions + Cations = ionic compound name
FFluoride ion
CO32- Carbonate ion
PO43- Phosphate ion
SO42- Sulfate ion
OHHydroxide ion
NO3- Nitrate ion
Cl-
Chloride ion
O2-
Oxide ion
S2-
Sulfide ion
NH4+
Na+
Ca2+
Fe3+
Ag+
Al3+
Ammonium ion
Sodium ion
Calcium ion
Iron(III) ion
Silver ion
Aluminum ion
Ammonium fluoride
Iron phosphate
Silver oxide
Note: Ag has a charge of +1 and oxide has a charge of –2 so
the chemical formula of silver oxide must be Ag2O
What is the correct name for the ionic
compound Na2SO4?
A. disodium sulfate
B. sodium sulfoxide
C. sodium sulfate
D. sodium sulfide
The procedures used for naming binary (two-element)
molecular compounds are similar to those used for
naming ionic compounds:
1. The name of the element farthest to the left in the
periodic table is usually written first.
2. If both elements are in the same group in the
periodic table, the lower one is named first.
3. The name of the second element is given an -ide ending.
4. Greek prefixes (Table 2.6) are used to indicate
the number of atoms of each element. The prefix
mono- is never used with the first element.
When the prefix ends in a or o and the name
of the anion begins with a vowel (such as oxide),
the a or o is often dropped.
CO2
Carbon Dioxide
H2O
Dihydrogen Monoxide
Carbon tetrachloride is
A. CCl4
B. CClO4
C. C2Cl4
D. CCl5
TETRA = 4
• Once you understand chemical names and formulas,
you can figure out what chemical compounds are
contained in different household products. For
example, (A) washing soda is sodium carbonate
(Na2CO3) and (B) oven cleaner is sodium hydroxide
Using Lewis Dot structure to help you
move electrons around
The number of valence electrons of any element is the same as
the group number of the element in the periodic table.
• For example, the Lewis symbol for oxygen, a member of
group 6A, shows 6 dots.
•A maximum of eight dots can be placed around a symbol,
where each dot represents a valence electron.
• Dots are placed above, below, to the left, and to the right of
the element symbol.
•Each position can accommodate two electrons, and electrons
are not "paired" until each of the four positions contains a
single electron.
The Lewis dot structure for Oxygen
O
Oxygen is in group VIA so it has 6 valence electrons
The Lewis dot structure for Chlorine
Cl
chlorine is in group VIIA so it has 7 valence electrons
The Lewis dot structure for calcium
Ca
calcium is in group IIA so it has 2 valence electrons
Making calcium chloride
Ca + Cl
Cl
Ca( Cl )2
Lewis dot structure of a compound
NH3
1) How many valence electrons does N have?
N is in group VA so it has 5 valence electrons
2) How many valence electrons does H have?
H is in group IA so each H has one valence electron
3) How many total valence electrons are there in this
molecule?
3x1+5=8
You know it had to be 8 because it
has NO CHARGE!
Lewis dot structure of a compound
NH3
H
N H
H
Lewis dot structure and making ammonium ion
NH4+
H
+
H
H N H
H
+
Forces.
• A force is viewed as a push or a pull, something
that changes the motion of an object.
• Forces can result from two kinds of interactions.
– Contact interactions.
– Interaction at a distance.
• The net force is the sum of all forces acting on
an object.
– When two forces act on an object the forces
are cumulative (the are added together.
– Net force is called a resultant and can be
calculated using geometry.
• Four important aspects to forces.
– The tail of a force arrow is placed on the
object that feels the force.
– The arrowhead points in the direction of the
applied force.
– The length of the arrow is proportional to the
magnitude of the applied force.
– The net force is the sum of all vector forces.
The rate of movement and the direction of movement of this
ship are determined by a combination of direction and
magnitude of force from each of the tugboats. A force is a
vector, since it has direction as well as magnitude. Which
direction are the two tugboats pushing? What evidence
would indicate that one tugboat is pushing with greater
magnitude of force? If the tugboat by the numbers is pushing
with a greater force and the back tugboat is keeping the ship
(A)When two parallel forces are acting on the cart in
the same direction, the net force is the two forces
added together.
• (B) When two forces are opposite and of equal
magnitude, the net force is zero.
• (C) When two parallel forces are not of equal
magnitude, the net force is the difference in the
direction of the larger force.
• You can find the result of adding two vector
forces that are not parallel by drawing thetwo
force vectors to scale, then moving one so the
tip of one is the tail of the other.
• A new arrow drawn to close the triangle will tell
you the sum of the two individual forces.
(A) This shows
the resultant of
two equal 200 N
acting at an
angle of 90O,
which gives a
single resultant
arrow
proportional to
a force of 280 N acting at 45O. (B) Two unequal
forces acting at an angle of 60O give a single
resultant of about 140 N.
Horizontal Motion on Land.
• It would appear as though Aristotle's theory of
motion was correct as objects do tend to stop
moving when the force is removed.
– Aristotle thought that the natural tendency of
objects was to be at rest.
– Objects remained at rest until a force acted
on it to make it move.
• Aristotle and Galileo differed in how they
viewed motion.
– Again, Aristotle thought that the natural
tendency of objects was to be at rest.
– Galileo thought that it was every bit as natural
for an object to be in motion.
• Inertia.
– Galileo explained the behavior of matter to
stay in motion by inertia.
– Inertia is the tendency of an object to remain
in motion in the absence of an unbalanced
force such as:
• friction
• gravity.
Galileo (left)
challenged the
Aristotelian view of
motion and focused
attention on the
concepts of distance,
time, velocity, and
acceleration.
Chemical Bonding
Physical Science
Chemical Bonds
• When two or more atoms attach to each other,
they form a chemical bond
• Compounds are any two elements chemically
bonded
–
–
–
–
Water
Sugar
Salt
And almost all other substances!!!!
• Electrons are responsible for the type, strength,
and size of a chemical bond
Lewis Structures
• Bohr-Rutherford diagrams are large and
difficult to show relationships between
multiple atoms
• Lewis diagrams are used to show multiple
atoms
• Lewis diagrams show only the valence
electrons
Lewis Structures
• Valence electrons form the charge of an
atom
• Electrons are always trying to get together
in groups of 8 (forget shells for a minute)
• Elements that have 8 valence electrons
have FULL outer groups
• We call these elements NOBLE or INERT
gases, they are found in group 8
Lewis Structures
• Label the Nobel (Inert) gases on your chart
Lewis Structures
• Elements with 1 valence electron are called the
Alkali metals (group 1) (Label)
Lewis Structures
• Elements with 2 valence electrons are called the
Alkaline Earth metals (group 2) (Label)
Lewis Structures
• Elements with 7 valence electrons are called the
Halogens (group 7) (Label)
Lewis Structures
• Consist of
– Element Symbol
– Electrons in each open spot
Sy
Lewis Structures
7
• Draw the element symbol
• Determine the # valence electrons
• Starting at the top, going clockwise, place one
electron in each spot around the element symbol
Cl
Lewis Structure
• Draw the Lewis Structure for Aluminum!
Al
Quick Recap!
• Draw the Lewis Structure for Lithium!
• Draw the Rutherford-Bohr Diagram for
Lithium!
-
Li
3P
3N
-
A note about charges…
• Writing a charge
– Valence electrons, Bohr-Rutherford, and Lewis
diagrams are used to determine charge
– Charges are a shortcut to determining bonding
properties
– RULES OF CHARGE
• IF the # of valence electrons is GREATER than 4, the charge
is negative (Mostly)
• IF the # of valence electrons is less than 4, the charge is
positive (Mostly)
• Charges are in reference to a full shell of 8
A note about charges…
• For example, Aluminum has 3 valance
electrons
– The possible charges are +3 OR -5
– It either has 3 OVER a full shell, or 5 LESS
than a full shell
– Because the number 3 is less than 4, we use
the charge of +3
Charges
• Any element with a
charge is called an
ION, the charge is an
ionic charge
• What are the ionic
charges of the
elements in the table?
Sodium?
+1
Nitrogen? -3
Oxygen?
-2
Argon?
0
Charges
Charges
• A few exceptions!
– Metals are always a positive charge!!
– Non metals are always negative!!
– Metalloids can go either way (you are not
responsible for choosing – I will tell you)
• Example: Boron
– According to rule of 4’s….its a +3 charge
– But since it’s a nonmetal, we use -5!
Rules of Bonding
1. All compounds must have neutral
charges
2. (That means the positive charges
(cations) and the negative charges
(anions) must equal
3. Subscript numbers are used to show the
number of ions
4. Coefficients are used to show the
number of molecules
Rules of Bonding
2H2O
Coefficient
Subscript
1 atom of O
“1’s” are
implied and
not written
Subscript
2 atoms of H
Rules of Bonding
H2O
H
H
O
Rules of Bonding
H2O
O
H
H
Rules of Bonding
• Try this one!
NaCl (table salt)
Na Cl
Rules of Bonding
• Last One!
Aluminum Bromide
Br Al Br
Br
Chemical Bonding
• Several Types including
– Covalent Bonds*
– Ionic Bonds*
– Metallic (only between metals)
Covalent Bonds
• Electrons are shared between two or more
atoms
• Covalent bonds can exist between atoms of the
same type…for example N-N (N2) or O-O (O2)
• Covalent bonds can form single, double, or triple
bonds
• Covalent bonds are strong and usually result in
stable molecules
• Carbon always forms covalent bonds and forms
the basic molecules for all life substances
Ionic Bonds
• Usually formed by members of the Alkali
group (ones with +1 electron)
• Electrons are donated to another molecule
• Between elements from opposite sides of
the chart
• Forms crystals (salts) & most dissolve in
water
Forming Compounds
1.
2.
3.
4.
Write ions with charges
Cross charges
Write subscripts (omit “1’s”)
Use parenthesis if needed
Forming Compounds
• What is the molecular formula of water?
+1
H
H2
-2
O
O1
H2O
O
H
H
Forming Compounds
• What is the molecular formula of carbon
dioxide?
+4
-2
C
O
C2
O4
C2O4
CO2
Yikes!
Reduce like a fraction to
lowest denominator
****note****
Forming Compounds
• What is the molecular formula of a
compound that has aluminum and
sulphur?
+3
-2
Al
S
Al2 S3
Al2S3
Any guesses
on the
name?