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Chapter 6.3
“Ionic Bonding”
Introduction to Bonding
• Chemical bond: an interaction between
atoms or ions that results in a reduction
of the potential energy of the system,
thereby becoming more stable
• Three types of bonds: ionic, metallic, and
covalent
• The bond type depends on the atoms’
electronegativities
More
• Subtract the electronegativities to
determine the nature of the bond
• If the difference is greater than 1.7, the
bond is ionic
• If the difference is from 0 to 0.3, the
bond is non-polar covalent
• If the difference is from 0.3 to 1.7, the
bond is polar covalent
Summary
• If the atoms have very different
electronegativities, then ionic bonding
occurs
• If they both have high electronegativities,
then covalent bonding occurs
• If they both have low electronegativities,
then metallic bonding occurs
Practice: What kind of bond?
•
•
•
•
•
•
•
Na and Cl
Sr and O
C and O
Ni and Fe
N and O
H and B
Ti and Cr
•
•
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•
•
•
ionic
ionic
polar covalent
metallic
polar covalent
non-polar covalent
metallic
Valence Electrons are…?
• The electrons responsible for the
chemical properties of atoms, and are
those in the outer energy level.
• Valence electrons - The s and p electrons
in the outer energy level
– the highest occupied energy level
• Core electrons – are those in the energy
levels below.
Keeping Track of Electrons
• Atoms in the same group have the same
outer electronic structure and therefore the
same number of valence electrons.
• The number of valence electrons is easily
determined. It is the group number for
groups 1 and 2
• Group 1: H, Li, Na, K, etc. have 1 valence e• Group 2: Be, Mg, Ca, etc. have 2 valence e-
More about Keeping Track
• For elements in groups 13-17:
• Subtract 10 from the group number
• This is the number of valence electrons
Electron Dot diagrams are…
• A way of showing & keeping
track of valence electrons.
• How to write them?
• Write the symbol - it represents
the nucleus and inner (core)
electrons
• Put one dot for each valence
electron (8 maximum)
• They don’t pair up until they
have to (Hund’s rule)
X
The Electron Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons to show.
l First we write the symbol.
lThen add 1 electron at a
time to each side.
lNow they are forced to pair up.
lWe have now written the electron dot
diagram for nitrogen.
l
N
The Octet Rule
l
l
l
The noble gases are unreactive in
chemical reactions
In 1916, Gilbert Lewis used this fact to
explain why atoms form certain kinds of
ions and molecules
The Octet Rule: in forming compounds,
atoms tend to achieve a noble gas
configuration; 8 in the outer level is stable
l Each noble gas (except He, which has
2) has 8 electrons in the outer level
Formation of Cations
• Metals lose electrons to attain a noble gas
configuration.
• They make positive ions (cations)
• If we look at the electron configuration, it
makes sense to lose electrons:
• Na 1s22s22p63s1 1 valence electron
• Na1+ 1s22s22p6 This is a noble gas
configuration with 8 electrons in the outer
level.
Electron Dots For Cations
• Metals will have few valence electrons
(usually 3 or less); calcium has only 2
valence electrons
Ca
Electron Dots For Cations
• Metals will have few valence electrons
• Metals will lose the valence electrons
Ca
Electron Dots For Cations
• Metals will have few valence electrons
• Metals will lose the valence electrons
• Forming positive ions
2+
Ca
This is named the
“calcium ion”.
No dots are now shown for the cation.
Electron Dots For Cations
• Let’s do scandium, #21
• The electron configuration is:
1s22s22p63s23p64s23d1
• Thus, it can lose 2e- (making it
2+), or lose 3e- (making 3+)
Sc
2+
= Sc
scandium (II) ion
Sc =
3+
Sc
scandium (III) ion
Electron Configurations: Anions
• Nonmetals gain electrons to attain noble
gas configuration.
• They make negative ions (anions)
• S = 1s22s22p63s23p4 = 6 valence electrons
• S2- = 1s22s22p63s23p6 = noble gas
configuration.
• Halide ions are ions from chlorine or
other halogens that gain electrons
Electron Dots For Anions
• Nonmetals will have many valence electrons
(usually 5 or more)
• They will gain electrons to fill outer shell.
P
3(This is called the “phosphide
ion”, and should show dots)
Stable Electron Configurations
• All atoms react to try and achieve a noble gas
configuration.
• Noble gases have 2 s and 6 p electrons.
• 8 valence electrons = already stable!
• This is the octet rule (8 in the outer level is
particularly stable).
Ar
Ionic Bonding
• Anions and cations are held together by
opposite charges (+ and -)
• Ionic compounds are called salts.
• Simplest ratio of elements in an ionic
compound is called the formula unit.
• The bond is formed through the transfer
of electrons (lose and gain)
• Electrons are transferred to achieve noble
gas configuration.
Ionic Compounds
1) Also called salts
2) Made from: a cation with an anion (or
literally from a metal combining with a
nonmetal)
Ionic Bonding
Na Cl
The metal (sodium) tends to lose its one
electron from the outer level.
The nonmetal (chlorine) needs to gain one
more to fill its outer level, and will accept the
one electron that sodium is going to lose.
Ionic Bonding
+
Na
Cl
1-
Note: Remember that NO DOTS
are now shown for the cation!
Ionic Bond
• Negative charges are attracted to positive
charges.
• Negative anions are attracted to positive
cations.
• The result is an ionic bond.
• A three-dimensional crystal lattice of anions
and cations is formed.
Preserve Electroneutrality
• When ions combine, electroneutrality
must be preserved.
• In the formation of magnesium chloride,
• 2 Cl- ions must balance a Mg2+ ion:
• Mg2+ + 2 Cl- →
MgCl2
Ionic Bonding
Let’s do an example by combining
calcium and phosphorus:
Ca
P
• All the electrons must be accounted for, and
each atom will have a noble gas
configuration (which is stable).
Ionic Bonding
Ca
P
Ionic Bonding
2+
Ca
P
Ionic Bonding
2+
Ca
Ca
P
Ionic Bonding
2+
Ca
Ca
P
3-
Ionic Bonding
2+
Ca
P
Ca
P
3-
Ionic Bonding
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
Ca
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
Ca
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
2+
Ca
2+
Ca
2+
Ca
P
P
33-
Ionic Bonding
= Ca3P2
Formula Unit
This is a chemical formula, which shows
the kinds and numbers of atoms in the
smallest representative particle of the
substance.
For an ionic compound, the simplest ratio
of the ions is called a formula unit
Properties of Ionic Compounds
1. Crystalline solids - a regular repeating
arrangement of ions in the solid:
–
–
Ions are strongly bonded together.
Structure is a rigid crystal lattice.
2. High melting points
• Coordination number- number of ions of
opposite charge surrounding it
- Page 198
Coordination Numbers:
NaCl
Both the sodium
and chlorine have 6
CsCl
Both the cesium
and chlorine have 8
TiO2
Each titanium has
6, and each oxygen
has 3
Do they Conduct?
•
Conducting electricity means allowing
charges to move.
• In a solid, the ions are locked in place.
• Ionic solids are insulators.
• When melted, the ions can move around.
3. Melted ionic compounds conduct.
–
–
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to
move in aqueous solutions)
- Page 198
The ions are free to move when they are
molten (or in aqueous solution), and thus
they are able to conduct the electric current.
Atoms and Ions
• Atoms are electrically neutral.
– Because there is the same number of
protons (+) and electrons (-).
• Ions are atoms, or groups of atoms, with
a charge (positive or negative)
– They have different numbers of protons
and electrons.
• Only electrons can move, and ions are made
by gaining or losing electrons.
An Anion is…
• A negative ion.
• Has gained electrons.
• Nonmetals can gain electrons.
• Charge is written as a superscript on the right.
1F
Has gained one electron (-ide
is new ending = fluoride)
2O
Gained two electrons (oxide)
A Cation is…
A positive ion.
Formed by losing electrons.
More protons than electrons.
Metals can lose electrons
+
K
2+
Ca
Has lost one electron (no
name change for positive ions)
Has lost two electrons
Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions
H+
Li+
Na+
K+
Rb+
Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 13: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Neither! Group 14
elements rarely form
ions (they tend to share)
Group 14: Do they
lose 4 electrons or
gain 4 electrons?
Predicting Ionic Charges
N3-
nitride
P3-
phosphide
As3- arsenide
Group 15: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2-
oxide
S2-
sulfide
Se2- selenide
Group 16: Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F- fluoride
Cl- chloride
Group 17: Gains
Br- bromide 1 electron to form
1- ions
I- iodide
Predicting Ionic Charges
Group 18: Stable
noble gases do not
form ions!
Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Note the use of Roman
iron (II) = Fe2+
numerals to show charges
iron (III) = Fe3+
Naming cations
•
Two methods can clarify when more than
one charge is possible:
1) Stock system – uses roman numerals in
parenthesis to indicate the numerical value
2) Classical method – uses root word with suffixes
(-ous, -ic)
• Does not give true value
Naming cations
• We will use the Stock system.
• Cation - if the charge is always the same (like
in the main group of metals) just write the
name of the metal.
• Transition metals can have more than one
type of charge.
– Indicate their charge as a roman numeral in
parenthesis after the name of the metal
Predicting Ionic Charges
Some of the post-transition elements also
have more than one possible oxidation state.
tin (II) = Sn2+
lead (II) = Pb2+
tin (IV) = Sn4+
lead (IV) = Pb 4+
Predicting Ionic Charges
Some transition elements have only one
possible oxidation state, such as these three:
(memorize these)
silver = Ag+
zinc = Zn2+
cadmium = Cd2+
Exceptions:
• Some of the transition metals have
only one ionic charge:
–Do not need to use roman
numerals for these:
–silver is always 1+ (Ag+)
–cadmium and zinc are always 2+
(Cd2+ and Zn2+)
Practice by naming these:
•
•
•
•
•
•
•
Na+
Ca2+
Al3+
Fe3+
Fe2+
Pb2+
Li+
Write symbols for these:
• potassium ion
• magnesium ion
• copper (II) ion
• chromium (VI) ion
• barium ion
• mercury (II) ion
Naming Anions
• Anions are always the same
charge
• Change the monatomic
element ending to – ide
• F- a fluorine atom will
become a fluoride ion.
Practice by naming these:
• Cl• N3• Br
2•O
• Ga3+
Write symbols for these:
• sulfide ion
• iodide ion
• phosphide ion
• strontium ion
Polyatomic ions are…
• Groups of atoms that stay together and
have an overall charge, and one name.
• Usually end in –ate, -ite, or -ide
• acetate: C2H3O2• nitrate: NO3• nitrite: NO2• hydroxide: OH• cyanide: CN-
Common Polyatomic Ions
•
•
•
•
•
2-
sulfate: SO4
carbonate: CO32chromate: CrO42chlorate: ClO3chlorite: ClO2-
• phosphate: PO43
+
• ammonium: NH4
(One of the few positive
polyatomic ions)
If the polyatomic ion begins with H, then combine the
word hydrogen with the other polyatomic ion
present:
H+ + CO32- →
HCO3hydrogen + carbonate → hydrogen carbonate ion
Writing Ionic Compound Formulas
Example: iron (III) chloride (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
Fe3+ Cl-
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
3
Not balanced!
Now balanced.
= FeCl3
Writing Ionic Compound Formulas
Example: aluminum sulfide (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
3+
Al
2
2S
3
Not balanced!
Now balanced.
= Al2S3
Writing Ionic Compound Formulas
Example: magnesium carbonate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
Mg2+ CO32They are balanced!
= MgCO3
Writing Ionic Compound Formulas
Example: barium nitrate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2+
Ba ( NO3 )2
2. Check to see if charges are
balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance subscripts.
Now balanced.
Not balanced!
= Ba(NO3)2
Writing Ionic Compound Formulas
Example: ammonium sulfate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
( NH4+ ) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
2
Now balanced.
Not balanced!
= (NH4)2SO4
Writing Ionic Compound Formulas
Example: zinc hydroxide (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2+
Zn
2. Check to see if charges are
balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
( OH- )2
Not balanced!
Now balanced.
= Zn(OH)2
Writing Ionic Compound Formulas
Example: aluminum phosphate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges are
balanced.
3+
Al
PO4
3-
They ARE balanced!
= AlPO4
Naming Ionic Compounds
• 1. Name the cation first, then anion
• 2. Monoatomic cation = name of the
element
Ca2+ = calcium ion
• 3. Monoatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride
Naming Ionic Compounds
(Metals with multiple oxidation states)
• some metals can form more than one charge
(usually the transition metals)
• use a Roman numeral in their name:
PbCl2 – use the anion to find the charge on the
cation (chloride is always 1-)
Pb2+ is the lead (II) cation
PbCl2 = lead (II) chloride
Things to look for:
1) If cations have ( ), the number in
parenthesis is their charge.
2) If anions end in -ide they are
probably off the periodic table
(Monoatomic)
3) If anion ends in -ate or –ite, then
it is polyatomic
Practice by writing the formula or
name as required…
•
•
•
•
•
iron (II) phosphate
potassium sulfide
ammonium chromate
MgSO4
FeCl3
Practice by writing the formula for
the following:
• magnesium hydroxide
• iron (III) hydroxide
• zinc hydroxide
Hydrates
• Some compounds contain H2O in their structure. These compounds are called hydrates.
• The H2O can usually be removed if heated.
• A dot separates water: e.g. CuSO4•5H2O is
copper(II) sulfate pentahydrate.
• A Greek prefix indicates the # of H2O groups.
Na2SO4•10H2O sodium sulfate decahydrate
nickel(II) sulfate hexahydrate
NiSO4•6H2O
sodium carbonate monohydrate Na2CO3•H2O
BaCl2•2H2O
barium chloride dihydrate
Prefixes
•
•
•
•
•
•
•
•
•
•
1
2
3
4
5
6
7
8
9
10
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
Hydrates
• Examples:
I. Give the name of the following:
1. CuSO4  5H2O
2. MgCl2  6H2O
3. Na2SO4  10H2O
II. Write the formula for:
1. zinc chloride hexahydrate
2. calcium phosphate dihydrate
3. copper (I) chloride pentahydrate
Helpful to remember...
1. In an ionic compound, the net ionic
charge is zero (criss-cross method)
2. An -ide ending generally indicates a
binary compound
3. An -ite or -ate ending means there is a
polyatomic ion that has oxygen
4. Prefixes generally mean molecular; they
show the number of each atom
Helpful to remember...
5. A Roman numeral after the name of
a cation is the ionic charge of the
cation