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Chemical Bonding
We will consider two types: ionic (ionic bonds) and molecular (covalent bonds)
We can tell them apart because ionic substances are made of ions,
- when separated in the dissolved or molten state, free ions allow the substance to conduct electricity
- substances that exhibit this feature are termed electrolytes
Our job in describing the bonding of a substance is to describe:
1) what atoms are connected to what other atoms (e.g., is H2O H-H-O or H-O-H)
2) what types of bonds hold these atoms together (ionic or covalent)
Why do atoms form bonds?
This is a topic that we will expand on greatly as we progress through grade 11 and 12 chemistry.
For our purposes, we can understand bonding as a consequence of the maximum stability that atoms have
when their valence electron shell is filled, as is the case for noble gases.
Atoms have three means to fill their valence shells:
1) transfer all valence electrons to other atom(s): the next shell down becomes the outer shell and is full,
the atom becomes positively charged
2) take electrons from another atom(s) to fill up their valence shell,
the atom becomes negatively charged
3) share valence electrons with other atoms.
Ionic Bonding
An ionic bond is the electrostatic attraction between two ions of opposite charge.
The actual trading of electrons to form ions is not actually part of the ionic bonding process.
Example:
Ionization (e- trading): Na + Energy —> Na+ + eCl + e- —> Cl- + Energy
Ionic Bond Formation: Na+ + Cl- —> NaCl
Ionic bonds can happen when metals and non-metals (especially halogens, Oxygen, or Sulfur) react:
2Na(s) + Cl2(g) —> 2NaCl(s)
[the subscripts (s) and (g) indicate solid and gas states respectively]
They also can form when the ions already exist, such as when a solution evaporates:
[the subscript (aq) indicates a substance that is dissolved in water]
Na(aq)+ + Cl(aq)- —> 2NaCl(s)
Once formed, ions are attracted to those with opposite charges.
They join together as a result of these attractions, much like
items of laundry in your clothes dryer. This joining of ions
forms large structures called crystal lattices. When dissolved,
the water molecules break this lattice apart and surround the
ions, holding them in solution as hydrated ions.
These crystal lattices are electrically neutral. Thus, the ions must be present in the lattice in a ratio that
causes the negative and positive charges to balance out. This ratio is what the formula for an ionic substance
shows. In other words:
the formula
NaCl tells you that a salt crystal has 1 sodium ion for each chloride ion
a salt crystal could have a few hundred of each ion, or a 1020 of each ion
MgCl2 tells you that these salt crystals have 1 magnesium ion for each two chloride ions
In general, for any pair of positive and negative ions there is only one ratio and, thus, one formula that
results in overall electrical neutrality.
The trading of electrons between elemental metals and
non-metals does demonstrate one method to arrive at
a viable formula for an ionic compound. Each atom
must either lose (metals) or gain (non-metals) enough
electrons to fill or empty its valence shell.
This determines the formula, because an particular
number of atoms of each element are needed to allow
each atom to achieve a full or empty valence shell.
A simpler method to arrive at the correct formula is
to use the criss-cross method (the absolute value of an ion’s charge is used as the subscript for
the other ion) followed by a reduction to the smallest whole-number ratio. The ionic formula
gives only the simplest ratio of the ions involved because the absolute number of ions of each
type is not important: ionic crystal lattices can be of any size. Only the ratio is important in
determining the compound type.
So, for Mg2+ and O2- ions, the formula Mg2O2 must be reduced to MgO, since this is the simplest ratio.
Molecular Bonding
Just as for ion formation, atoms form covalent bonds in order to fill their valence shells.
They do this by sharing valence electrons. Thus occurs because the atoms involved (non-metals in this case)
all want to gain electrons, The number of bonds that must form for each type of atoms depends on the
number of electrons needed to complete the valence shell.
Molecules are held together by
covalent bonds.
A covalent bond is a pair of electrons
that is shared by two atoms. They are
sharing to maintain a full valence shell.
They must remain close together to
continue sharing. Thus, the two atoms
are held together.
This process produces a small structure
called a molecule. Molecular structures
are constructed by circling unpaired
electrons between two atoms.
The actual Lewis structure must show bonded electron pairs as a short line.
Unlike ionic substances, most molecular substances do not form crystal lattices. Once each atom in your
structure has shared a sufficient number of electrons to achieve a full valence
shell, no further atoms need be attached to the molecule.
The other way in which molecular bonding differs from ionic bonding is that most
combinations of elements can result in a wide variety of different plausible
molecules. For example, both N2O2 and N2O3 are perfectly reasonable molecules.
For a molecule, the actual number of atoms is important, not just the ratio.
For example, H2O2 is not the same as HO, even though the ratio is the same. HO
does not result in a molecule in which the Oxygen atom has a full valence shell.