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20 Chemistry
Chapter 8 Notes
Ionic Compounds
There are two types of compounds
i) ionic
ii) molecular
Ionic compounds are formed by the attraction of positively charged atoms (called positive
ions) and negatively charged atoms (called negative ions). The force of attraction
between these two types of charge creates what is called an ionic bond. This bond is the
“glue” that holds ionic compounds together.
Recall that noble gases have their outer energy levels full of electrons. This makes noble
gases stable (thus they do not readily react).
Recall also that some elements have low ionization energies (and low electro negativities)
and others have high ionization energies (and high electro negativities). These tendencies
determine whether an atom is likely to loose electrons and become a positive ion or gain
electrons and become negative ions. In either case atoms react in ways that will fill their
outermost energy levels so that they are stable.
Formation of Positive Ions
Consider the electron configurations of Argon (a noble gas) and Potassium
Argon
Potassium
1s22s22p63s23p6
1s22s22p63s23p64s1
The only difference as far as electrons are concerned is potassium has one more electron.
And this one electron is considerably farther away from the positive nucleus than all the
others (it has a relatively low ionization energy – 419 kJ/mol). This means it will only
take a small amount of energy to remove this electron and make a much more stable
configuration like Argon’s. Put another way:
K  419kJ / mol  K   e
Group 1A elements have one valence electron and so tend to loose one electron and
obtain a 1+ charge. Group 2A elements have two valence electrons and tend to loose two
electrons and obtain a 2+ charge and so on up to about group 3A.
Group B elements
[noble gas]ns2(n-1)dx
like Vanadium
[Ar]4s23d3
Often loose their ns2 electrons becoming 2+, but sometimes loose a 3d electron as well.
Suffice it to say transition metals obtain charges of 2+ or 3+.
Positive ions are also “cations”.
Formation of Negative Ions
Recall that non-metals on the right hand side of the periodic table tend to have high
electro negativities and high ionization energies. For this reason they attract electrons in
an amount sufficient to fill their outer most energy levels and become stable. Compare
the electron configurations of Bromine to Krypton (a noble gas).
Bromine
Krypton
1s22s22p63s23p64s23d104p5
1s22s22p63s23p64s23d104p6
By gaining 1 electron, bromine becomes Br - and assumes the same electron
configuration as Krypton and is stable.
Negative ions are called “anions”.
Electron Configuration Representation
Although Argon and a Potassium ion have the same configuration, they are very
different. Argon still has a completely different nucleus and it doesn’t have a charge. To
tell the difference, we will represent ions as such:
Argon
Potassium
Calcium
- 1s22s22p63s23p6
or
- [1s22s22p63s23p6]+ or
- [1s22s22p63s23p6]2+ or
[Ar]
[Ar] +
[Ar] 2+ etc.
- [1s22s22p63s23p6]- or
- [1s22s22p63s23p6]2- or
- [1s22s22p63s23p6]3- or
[Ar] [Ar] 2[Ar] 3- etc.
And for negative ions:
Chlorine
Sulfur
Phosphorus
Try p. 214 # 2 – 5
Formation of Ionic Compounds:
Now that we have an idea how ions are formed we can explain how ionic compounds
form. We will use Sodium Chloride as a first example. The sodium atom looses an
electron to become positively charged. The chlorine atom gains an electron (from the
sodium atom) and becomes negatively charged. The two charged ions are attracted to
from Sodium Chloride. In this case the “exchange” is simple because sodium lost one
electron and chlorine gained one.
What about Calcium Chloride. Calcium must give up two electrons but a chlorine atom
will only take one electron. Therefore two chlorine atoms are needed to bond with each 1
calcium atom.
Ca2+
and
Cl-
Form
Fe3+
and
O2-
Form Fe2O3
CaCl2
Try p. 217 # 7 – 10
A common way of classifying ionic compounds:
i) metal oxides – form when a positively charged metal ion (cation) is
electrostatically attracted to a negatively charged oxygen ion
(anion)
ii) salts - form when a positively charged metal ion (cation) is electrostatically
attracted to any other kind of negative ion (anion)
Another way of classifying ionic compounds:
i) Binary Ionic Compounds – contain only two different elements i.e. NaCl
ii) Non-binary Ionic Compounds – ie. CuSO4
Properties of Ionic Compounds
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Crystalline – shape of crystals vary from one compound to the next. The crystal
lattice (arrangement of particles) in an ionic compound is such that each positive
ion is surrounded by negative ions and vice versa.
High melting and boiling points.
Hard, rigid brittle solids – because of the strong electrostatic attraction between
positive and negative ions
Conduct electricity when dissolved in water.
Note also that there are no molecules of an ionic compound.
Lattice Energy
Review exothermic and endothermic processes
The formation of an ionic compound is always exothermic. When ionic compounds are
formed, energy is released (positive energy). If energy is added (negative energy) to an
ionic compound, the ionic bonds will break. Lattice energy is the amount of energy
required to break the bonds of one mole of an ionic compound.
There are two main factors effecting the amount of lattice energy.
Lattice energy increases with decreasing ionic radii. This makes sense if you think about
it. After-all the smaller the ion, the closer the positive nucleus is to the valence electrons
responsible for bonding. So Magnesium compounds will have higher (negative) lattice
energies than Calcium compounds (all things being equal).
The higher the charges on the ions that make up an ionic compound, the greater the
(negative) lattice energy. Refer to table 8-3 on page 220 for some examples of lattice
energies.
Metallic Bonds and Properties of Metals
Consider a sample of a metal. Although the atoms that make up that sample have valence
electrons, they do not share these electrons (as is the case in molecular compounds) or
lose electrons (as in the case of ions). Instead, the valence electrons enter a “sea” of
electrons. The positively charged metallic cations float in this sea of delocalized
electrons. As such electrons are not held in position by any one specific atom. The
particles of this sample of a metal are held together because the positively charged
cations are attracted to the negatively charged “sea” of electrons. This attraction is a
metallic bond.
This “electron-sea” model helps explain some of the properties of electrons:






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Varied melting points.
Moderately high melting points
Very high boiling points
Malleable
Ductile
Conduct heat and electricity
Lustre
There seems to be a correlation between the number of valence electrons and such
properties as hardness.
Try
p. 231 # 42
p. 237 # 80 – 83