7-1 Ionic Bonding
Ionic Bonds and Ionic
Compounds
Static Electrical attraction is the basis for ionic bonds.
In an ionic bond, positively charged ion (called a cation)
is attracted to a negatively charged ion (called an
anion).
Metals form positive cations by losing electrons
Non metals form negative anions by gaining electrons
Ionic compounds are electrically neutral.
The positive charges equal the negative charges
Salt
One of the more common ionic compounds that
you are familiar with is NaCl (salt).
Na is a soft silvery metal,
Cl is a poisonous gas.
When the two are combined a tremendous amount
of energy is released
Indicates that the compound NaCl is more stable than
either element alone.
http://www1.teachertube.com/viewVideo.php?video_id=
67110
Properties
Most exist as crystalline solids
High melting points
Indicates strong bonds
Tend to be brittle
Dissolve in water
Makes ionic solutions good conductors
As a solid- poor conductor
The Octet Rule
Chemical compounds tend to form so that each atom, by
gaining, losing or sharing electrons, has a full set of electrons in
its highest energy level.
How many electrons is that for most elements?
Who doesn’t follow that rule?
Which atoms have a full set of valence electrons?
By losing and gaining electrons atoms take a stable
configuration
Chlorine [Ne]3s23p5
Chlorine Ion [Ne]3s23p6 - same as Argon
Sodium
[Ne]3s1
Losing 1 electron has the configuration of [Ne], a
stable one.
Lewis Dot Diagrams
The focus for bonding is on the valence electrons.
The dots are placed on the imaginary four sides of
a “box” (right angles) around the element symbol.
Electrons are place singly first then paired after the
first four are place. (What rule are you following?)
Lewis dot diagrams are also used to represent
electron configuration during a reaction.
See figure 7-9 pg 230
Example
Nitrogen
Electron Configuration [He]2s22p3
Lewis Dot Structure
Practice
Draw the Lewis dot diagram for the following
atoms: C, Cl, Mg, K, Ne, Si, Ca, Kr, Al
Practice
Draw the Lewis dot diagram for the reaction
between Na + Cl
Types of Ions
Monoatomic ions
- one atom ions
Polyatomic ions – more than one atom
Monoatomic Cations – one atom positively
charged atoms
Some Transition metals have more than one ion and
are named with a roman numeral
Monoatomic Anoins
Nonmetals form anions most easily
When naming use the suffix “ide” in general
– Chlorine atom becomes a Chloride ion
Polyatomic ions
Sulfate – SO42As a unit form ionic bonds
Individual bonds are covalent
Monoatomic Cations
1+
2+
3+
H+ hydrogen
Mg2+ magnesium
Al3+ aluminum
Na+ sodium
Fe2+ iron(II)
Fe3+ iron(III)
K+ potassium
Co2+ cobalt(II)
Co3+ cobalt(III)
NH4+ ammonium
Ni2+ nickel(II)
Ni3+ nickel(III)
Li+ lithium
Ca2+ calcium
Ag+ silver
Zn2+ zinc
Cu+ copper
Cu2+ copper(II)
Polyatomic Ions
1-
2-
3-
F- fluoride
ClO3- chlorate
O2- oxide
N3- Nitride
Cl- Chloride
NO3- nitrate
S2- sulfide
P3- Phosphide
Br- Bromide
HCO3bicarbonate
C2H3O2Acetate
SO4 2- sulfate
PO43phosphate
I- Iodide
OH- hydroxide
CO3 2carbonate
Binary Ionic Compounds
Contain the ions of only two elements.
To name, write the cation followed by anion
Ex. Calcium and Fluorine form ionic bonds
– Called Calcium Fluoride
– MgO is magnesium oxide
– MgCl2- Even though there are 2Cl it's still the same type
of atom. Magnesium chloride.
– Name: BaS, Li2O
Binary Ionic Compounds
To denote the ratio of ions you use the empirical
formula.
Empirical formula-ratio of atoms
Use the element symbol and subscripts to denote
how many of each element you need to balance the
charge.
Always express the formula is it's lowest terms.
Binary Ionic Compounds
Use the crisscross method to determine the
empirical formula
1.
2.
Determine the charges of each atom when it forms
an ion
The number of the first element’s charge becomes
the subscript for the 2nd element in the formula.
Practice:
Write the formula for Sodium Chloride
1.
2.
Na1+ Cl1Na Cl = NaCl
Binary Ionic Compounds
Practice:
Write the formula for barium bromide
1. Ba2+ Br12. Ba Br = BaBr2
Binary Ionic Compounds
Practice:
Write the formula for potassium nitride
1. K1+ N32. K N = K3N
Binary Ionic Compounds
Practice:
Write the formula for magnesium oxide
1. Mg2+ O22. Mg O = Mg2O2 = MgO
Aluminum Oxide
Write the formula for aluminum oxide
Al203,
For every 2 aluminum ions there are 3
oxygen.
How can you determine the empirical formula?
Use the Charges
Al3+ O2criss-cross
Al203
How about a poly atomic?
Aluminum Ion
+ Sulfate Ion
Al3+ + SO42
Al (SO4)
= Al2(SO4)3
How about a poly atomic?
Aluminum Ion
+ Sulfate Ion
Al3+ + SO42- Al2(SO4)3
Don’t forget parenthesis
How about a poly atomic?
Write the formula for calcium carbonate
Ca2+
(CO3)2 Ca (CO3)
Ca2(CO3)2 Lowest terms?
Ca(CO3)
How about a poly atomic?
Write the formula for cobalt (II) phosphate
Co2+
(PO4)3 Co (PO4)
Co3(PO4)2 Lowest terms?
7-2 Covalent Bonding
Molecules and Their Formulas
Molecule – a neutral group of atoms that are
held together by covalent bonds
Molecular Compound – a chemical compound
whose simplest units are molecules.
Chemical Formula – indicates the relative
number of atoms of each kind in a chemical
compound
Ex – NaOH, Ca(OH)2
Molecular formula – indicates the relative
numbers of atoms of each kind in a molecule by
using atomic symbols and subscripts
H2O, O2, C6H12O6
Same thing as chemical formulas, but only with
NONmetals
Diatomic – molecule containing two elements
Cl2, O2, F2
The big seven
O2, Cl2, F2, H2, Br2, I2, N2,
Describing Covalent Bonds
The octet Rule describes covalent bonding.
Electrons are shared instead of transferred.
Lewis structures – reminder: draw dots around the
atom according to the number of valence electrons
This shows who is bonded to whom.
shared
shared
shared
Multiple Bonds
It is possible for atoms to share more than one
pair of electrons
Double bond shares two pairs
Triple bond shares three pairs
Is it possible to share 4 pairs? No.
If a molecule has multiple bonds, the strength of
the bond increases and the bond length
decreases.
Double Bond example
shared
Lines represent bonds
Unshared pair
Two shared pairs = double bond
Practice Drawing Bonds
Single
H2O
HCl
Br2
check: Is the octet rule satisfied for each atom?
Practice Drawing Bonds
Double
O2
CO2
H2CO (formaldehyde)
check: Is the octet rule satisfied for each atom?
Practice Drawing Bonds
Triple
N2
C2N2 (cyanogen)
check: Is the octet rule satisfied for each atom?
Exceptions to the Octet Rule
The octet rule cannot be satisfied in molecules whose total
number of valence electrons is an odd number.
There are also molecules in which an atom has fewer, or
more, than a complete octet of valence electrons.
Ex) NO2 , PCl5 , BF3 , SF6
Ways to remember exceptions:
Superman Sulfur – more
Boron the Moron – less
Phosphorus the prosperous – more
Nitrogen the ninny - less
Exceptions to the Octet Rule
Properties of Covalent Bonds
Bond length – the distance between two
bonded atoms at their minimum potential energy,
that is, the average distance between two
bonded atoms.
Bond energy – the energy required to break a
chemical bond and form neutral isolated atoms.
Atoms do not share electrons equally
Different atoms have different electronegativites
– Measure of atoms attraction for electrons
Electronegativity Dif. Bond Type
< 0.4
Non polar covalent
Between 0.5 and 1.9 Polar covalent
> 2.0
Ionic
If you have a polar covalent bond, the electron
density is concentrated on the more
electronegative atom.
If the electrons are pulled closer to the more
electronegative atom, it gains a partial negative
charge. The other atom in the bond has less
electron density, so it has a partial positive charge
Non polar covalent –
between atoms of similar
electronegativities.
Often between the same element (F2)
IONIC/COVALENT BONDS
Ionic
bonds
Metal to
nonmetal
Electrons are
transferred
Covalent
bonds
Nonmetal to
nonmetal
Electrons are
shared.
Can be
Nonpolar OR
polar
7-3 Naming Chemical
Compounds
Naming Molecular Compounds
mono
1
di
2
tri
3
tetra
4
penta
5
Mono is not added to the first element
hexa
Prefixes can be shortened (monooxide monoxide) hepta
Common names are allowed
Octa
6
Nona
9
Deca
10
Use the prefixes from the table
CO2 Carbon Dioxide
CCl4 Carbon Tetrachloride
Exceptions the the Rules
– Dihydrogen monoxide; H2O is Water
7
8
Practice Problems
IF5
Iodine pentafluoride
CO2
Carbon dioxide
PH3
Phosphorus trihydride
P4O10
tetraphosphorus decaoxide
Practice Problems
N2O4
Dinitrogen tetroxide
PCl5
Phosphorus pentachloride
NO2
Nitrogen dioxide
P2O5
Diphosphorus pentoxide
Naming Ionic Compounds
Cation is listed first.
Anion is second.
Examples:
MgCl2 - Magnesium Chloride
Cu(NO3)2 Copper (II) Nitrate
– Nitrate has a 1- Charge, and there are two of them.
– Therefore copper has a 2+ charge for the total charge to
be zero.
Sample Problem –
Name Fe(OH)3
Hydroxide has a 1- charge for a total of 3-.
Iron can be 2+ or 3+
For the charge to total zero, it must be iron 3+
Name = iron(III) hydroxide
Try the following
CuSO4 Copper(II) sulfate
Al2S3 Aluminum sulfide
Cu2O Copper(I) oxide
Hydrates
Ionic compounds that absorb water are called
Hydrates
Ionic compounds that do not absorb water are
called anhydrous substances
The properties of an ionic
substance depend if it is
hydrated.
The name must reflect this
hydration
Copper(II) sulfate when
hydrated becomes bright
blue.
Formula = CuSO4 • 5H2O
Given the name copper(II)
sulfate pentahydrate
mono
di
tri
tetra
penta
hexa
hepta
Octa
Nona
Deca
1
2
3
4
5
6
7
8
9
10
Prefixes
Acids
Many common items contain acids
Asprin, antiseptics, contact solution
An acid is a molecular substance that dissolves
in water to produce Hydrogen ions (H+)
Behave like an ionic compound
H+ is always the cation, anion depends on the acid
Name of the acid comes from the anion
Anion
Corresponding
Acid
Anion
Corresponding
Acid
F-, fluorine
HF, hydrofloric
acid
HCl, hydrocloric
acid
HBr,
hydrobromic
acid
HI, hydroiodic
acid
H2S,
hydrosulfuric
acid
NO3-, nitrate
HNO3, nitric
acid
H2CO3,
carbonic acid
H2SO4, sulfuric
acid
Cl-, chloride
Br-, bromide
I-, iodide
S2-, sulfide
CO32- ,
carbonate
SO42- , sulfate
PO43-,
phosphate
C2H3O2-,
acetate
H3PO4,
phosphoric acid
H C2H3O2,
acetic acid