bonding_powerpoint.ppt

Ch. 6
The Structure of Matter
The Importance of BONDING

Important Terms
• Element = a pure substance that cannot
be separated or broken down into simpler
substances by chemical means
• Atom = the smallest unit of an element
that maintains the chemical properties of
that element
• Compound = a substance made up of
atoms of two or more different elements
joined by chemical bonds

Bonding
• Atoms with unfilled valence shells are
considered unstable.
• Atoms will try to fill their outer shells by
bonding with other atoms.
• Chemical bond = the attractive force
that holds atoms or ions together in a
compound

Atomic Bonds
• Atoms form atomic
bonds to become
more stable.
– Atoms become more
stable by filling their
valence shell or at
least meeting the octet
rule by getting 8
valence electrons.
Exception to Octet Rule
of 8 valence electrons:
Helium—which only has
1 energy level and holds
a max. of 2 electrons

Atomic Bonds
• There are three main types of chemical
bonds used by atoms to fill their valence
shell:
– Covalent
– Metallic
– Ionic
“Bond,
Chemical Bond”

Chemical Formulas
• A chemical formula tells us:
–the type of atoms present
–the number of atoms present
–the type of compound

Chemical Formulas
• Example: table
salt: Sodium
Chloride
• Chemical formula:
– NaCl
• Count the atoms
present:
– 1 Na atom
– 1 Cl atom

Chemical Formulas
• Sometimes there are subscripts present.
– A subscript is a small number that is in a
chemical formula. If no subscript is present
assume that it is 1.
– Example - water: H2O
• 2 H atoms
• 1 O atom
Subscript

Chemical Formulas
• Sometimes there are parentheses with a
subscript. The subscript only applies to the
atoms within the parentheses.
• Example - calcium hydroxide (kidney stones):
Ca(OH)2.
– 1 Ca atom
– 2 O atoms
– 2 H atoms

Chemical Formulas
• Sometimes there are subscripts in the
parentheses. Multiply the subscript outside the
parentheses by the subscript of each element
within the parentheses.
• Example - calcium nitrate: Ca(NO3)2
– 1 Ca atom
– 2 N atoms
– 6 O atoms (3 oxygens x 2 = 6)

Covalent Bonds
• Covalent bonds form between two non-
metals. Groups 14-17 on the Periodic Table
• Covalent bonds are formed when atoms
SHARE electrons.
– Both atoms need to gain electrons to become
stable, so they share the electrons they have.
• Atoms can share more than one pair of
electrons to create double and triple bonds.

Properties of Covalent
Compounds
Results in a NEUTRAL molecule
Weak bonds
 Physical State usually liquids or
gases
 Low Melting and Boiling Points
 Poor conductors of electricity
(no free electrons to move around)

Covalent Bonds
Each chlorine atom wants to gain one
electron to achieve an octet.
Use Lewis structures to draw valence electrons for
each atom in the covalent pair.

Covalent Bonds
The octet is achieved by each atom
sharing the electron pair in the middle.
Now, each Chlorine atom has 8 valence
electrons because it is sharing one pair.

Chlorine Molecule
It is a single bonding pair so it is called a
single covalent bond. The compound
is now called a molecule.
Cl Cl Cl2

Covalent Bonds
How will oxygen bond?

Covalent Bonds
Two bonding pairs, making a double bond.
The double bond can be shown as two dashes
O O
O2

Covalent Bonds
• Elements can share up to three pairs of
electrons. (6 total electrons).
Single Bond
(2e)
Double Bond
(4e)
Triple Bond
(6e)

Covalent Bonds
• Atoms can share their electrons equally or unequally.
• When atoms share electrons equally, it is called a non-polar
covalent bond.
– Non-polar covalent bonds form between atoms of the same type.
Ex: H2, Cl2,
• When atoms share electrons unequally it is called a polar
covalent bond.
– One atom pulls the electrons closer to itself.
– The atom that pulls the electrons more gets a slightly negative
charge.
– The other atom gets a slightly positive charge.
• Ex: Water molecule
Bonding Animation

Covalent Bonds Nomenclature
• Naming binary covalent
compounds:
– Two nonmetals
– Name each element
– Change the ending of
the 2nd element to
–ide
– Use prefixes to
indicate the # of atoms of
each element
– Do not use “mono” with
the first element
# of Atoms Prefix
1 mono-
2 di-
3 tri-
4 tetra-
5 penta-
6 hexa-
7 hepta-
8 octa-
9 nona-
10 deca-

• CO
– carbon monoxide
• CO2
– carbon dioxide
• PCl3
– phosphorus trichloride
• CCl4
– carbon tetrachloride
• N2O
– dinitrogen monoxide
# of Atoms Prefix
1 mono-
2 di-
3 tri-
4 tetra-
5 penta-
6 hexa-
7 hepta-
8 octa-
9 nona-
10 deca-

Given the following covalent compounds,
WRITE the correct chemical formula.
Name Chemical Formula
Hydrogen Disulfide
Diphosphorus pentoxide
Trinitrogen hexafluoride
HS2
P2O5
N3F6

Practice: Drawing Covalent Bonds
• We can illustrate covalent bonding using
Lewis structures.
• 1 – Draw a Lewis structure for each element.
– Ex: C H
• 2 - Continue adding atoms until all atoms have a full valence
H
H C H
CH4
carbon tetrahydride
H

Ions
• Ions are formed when atoms gain or lose
electrons.
• Ions are charged atoms (positive or negative).
• Positive ions are called cations.
– Formed when the atom loses electrons.
– Lose negative charge, becomes positive ION
– Metals
• Negative ions are call anions.
– Formed when the atom gains electrons.
– Gain negative charge, become negative ION
– Non-metals

Ionic Bonds
• Ionic bonds are formed between metals
and non-metals.
• Ionic bonds are formed between
oppositely charged atoms (ions).
• Ionic bonds are formed by the transfer of
electrons.
– One atom loses (gives away) electrons.
– One atom gains (receives) electrons.

Ionic Bonds
• Use the number of valence electrons to
determine the # of electrons that are lost or
needing to be gained.
• The transfer of electrons create a positive ion
and a negative ion. The opposite charges attract
one another, causing a chemical bond to form.
Bonding Animation

Atoms with 4 or less valence
electrons want to LOSE (give
away) their valence electrons.
[Groups 1, 2, 13, 14]
Atoms with 4 or more valence
electrons want to GAIN (receive)
more electrons to satisfy their
octet. [Groups 14, 15, 16, 17]

Ionic Bonds
• The normal charge of an
ion can be quickly
determined using the
oxidation number of an
element.
– The oxidation number of
an atom is the charge that
atom would have if the
compound was composed
of ions.

Ionic Bonds
• To find the oxidation
number :
Look at Group #
Determine # of valence
electrons
If 4 or less, atom will
lose (give away)
valence electrons (ion is
positive)
If 4 or more, atom will
gain the needed # to fill
valence shell. (ion is
negative)

Ionic Bonds
• Example:
– Beryllium is in Group 2
– Be has 2 e-
– Wants to achieve octet
– Loses the 2 e-
– Oxidation #/Ion charge of
+2
• Example:
– Nitrogen is in Group 15
– N has 5 e-
– Needs 3 more for octet
– Gains 3 e-
– Oxidation #/Ion charge of
-3

Practice: Determining Oxidation
Numbers
Atom Group Valence
Electrons
Oxidation
Number
Oxygen
Calcium
Fluorine
Phosphorus
Sodium
16 6 -2
2 2 +2
17 7 -1
15 5 -3
1 1 +1

Ionic Bonding Nomenclature
To name Binary Ionic Compounds:
 2 elements—one METAL and one NON-METAL
 Cation is always written first [Metal]
 Cation name stays the same
 Anion is written second [Non-metal]
 Change the non-metal’s ending to “-ide”.
 NO PREFIXES ARE USED FOR IONIC COMPOUND
NAMING

Examples
NaCl
Name the metal ion
Sodium
Name the nonmetal
ion, changing the
suffix to –ide.
Chloride
CaO
Calcium Oxide
Al2S3
Aluminum Sulfide
MgI2
Magnesium Iodide
BaNa2 You should recognize a problem with this one
This is two metals – not a binary ionic
compound
The name of this is Banana (JOKE – haha)

Drawing Ionic Bonds
• 1 – Draw the Lewis structure for each
element.
– Ex: Na Cl
• 2 – Draw arrows to show the TRANSFER
(gain/loss) of electrons [draw extra atoms
if needed]

Drawing Ionic Bonds (continued)
• 3 – Draw ion Lewis diagrams showing the
new charge for each ion.
– Ex:
• 4- Write the chemical formula for the
compound formed represents the ratio of
negative ions to positive ions.
– Ex: NaCl – for every 1 sodium ion, there is
also 1 chlorine ion.
Chemical Formula = NaCl

Practice Drawing Ionic Bonds
Elements Lewis Transfer Formula
Diagram
Calcium
Fluorine
Sodium
Oxygen

“Swap & Drop” Method
Given the name of an Ionic Compound, you can determine the chemical
formula using the “swap and drop” method:
1. Write the symbols for each ion.
2. Determine the oxidation number of each ion.
3. Swap and Drop
4. Reduce (if necessary).
5. Rewrite

Ionic vs. Covalent Bonds in
Binary Compounds
Ionic Bonds
• Form when electrons
are transferred
between atoms.
• Form between a
metal and a non-
metal.
Covalent Bonds
• Form when electrons
are shared between
atoms.
• Form between two
non-metals.
Both types of bonds result in all atoms
having a full outer energy level.

Ionic vs. Covalent Bonds in
Binary Compounds
Other comparisons between Ionic and Covalent Compounds:
Ionic Compounds
• Results in a
Neutral Compound
• Crystalline Solid
• Strong Bonds
• High Melting
Point
Covalent Compounds
• Results in a Neutral
Molecule
• Mostly results in
gases or liquids
• Weak Bonds
• Low Melting Points

Polyatomic Ions
• A polyatomic ion is a group of
covalently bonded atoms that have lost
or gained an electron. (Example: Nitrate
NO3
- and Ammonium NH4
+).
– Oppositely charged polyatomic ions can
form compounds. (Example: Ammonium
nitrate NH4NO3).

Polyatomic Ions
• Naming of these
compounds follows
the same rules as
binary ionic
compounds.
– The most important
part is recognizing
there is a polyatomic
ion present.
Common Polyatomic Ions
ammonium NH4
+
carbonate CO3
2-
bicarbonate HCO3
-
hydroxide OH-
nitrate NO3
-
nitrite NO2
-
phosphate PO4
3-
sulfate SO4
2-
sulfite SO3
2-
acetate C2H3O2
-

Practice: Polyatomic Ions
To go from the formula to
the name:
1. Name the cation.
2. Name the anion.

Polyatomic Ions
To go from
name to formula:
1. Write the symbols for
each ion.
2. Determine the
oxidation number of
each ion.
3. Swap and Drop
4. Reduce (if necessary).
5. Put parentheses
around the polyatomic
ion if receives a
subscript greater than
one.
6. Rewrite
O2-
(NH4)2O
** Remember charges CANCEL
out each other!!

Practice: Polyatomic Ions
Compound Name Oxidation #s Chemical Formula
Calcium phosphate
Sodium hydroxide
Ammonium sulfate
Ca2+ PO4
3-
Ca3(PO4)2
Na1+ OH1- NaOH
NH4
1+ SO4
2- (NH4)2SO4

Metallic Bonds
• Metallic bonds are metal to
metal bonds formed by the
attraction between positively
charged metal ions and the
electrons around them.
– Atoms are packed tightly together
to the point where outermost
energy levels overlap.
• This allows electrons to move
freely from one atom to the next
making them great conductors of
electricity.

• Transition metals are cations that
have variable charges that makes
them hard to name.
– We use Roman numerals to indicate the
charge of a transition metal.
• Example:
– copper (II) oxide – charge of copper for this
compound is +2
– titanium (IV) sulfide – charge of titanium for this
compound is +4
Transition Metals--Ionic Compounds

• To go from formula to name
you need to determine the
Roman numeral for your
transition metal.
1. If there are no subscripts,
simply give the transition
metal the equal and opposite
charge to the nonmetal.
2. Now use normal ionic
bonding rules putting your
new number in Roman
numerals to the right of your
transition metal ONLY.
Transition Metal Ionic
Compounds

• To go from formula to name
you need to determine the
Roman numeral for your
transition metal:
1. If there are subscripts present
use the reverse “Swap and
Drop.”
2. Now use normal ionic
bonding rules putting your
new number in Roman
numerals to the right of your
transition metal ONLY.
Transition Metal Ionic
Compounds

bonding_powerpoint.ppt

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