Electron Dot Notation (aka “Lewis Dot Structure”): Sometimes, it’s easier to
understand how certain compounds form when considering the structure of the electrons
in the pertinent valence shells of each atom. Except for Helium, which has only two
electrons in its valence shell, all noble gases have 8 electrons in their valence shells. We
therefore surround these atoms with 8 dots (2 above, 2 to each side and 2 below) to
represent this. Halogens have 7 electrons in their valence shells, so they’re represented
with 7 dots surrounding each of them (it doesn’t matter which side has the lone dot).
Alkalis have only 1 valence electron and alkali earths have only 2 valence electrons, so
they have 1 and 2 dots, respectively. More of these structures can be seen in Figure 10.18
of your textbook. Knowing that atoms would like to have their valence shell consist of 8
electrons (or 2, if it’s mimicking helium) will be useful when considering the molecular
structure of such important compounds as carbon monoxide and ozone. It will also help
in understanding why some molecules, such as water, adopt asymmetric shapes, while
others, such as carbon dioxide, don’t.
Metals and Alloys: Solids bond together in basically three different ways. The most
different of these three ways is the metallic bond. Remember that in the case of most
metals, the outermost electrons are only loosely held to the parent atom; they are easily
dislodged. So if a group of metal atoms got together, there would be a sea of electrons
among all the atoms. Since the atoms have effectively lost electrons, they are positively
charged ions. There is thus a sea of negatively-charged electrons surrounding a bunch of
positively-charged ions. This is how most solid metals stay together.
Alloys are combinations of metals that are bonded together in this fashion. They are
solid solutions. Common alloys are white gold (gold and palladium in varying degrees),
sterling silver (92.5% silver, 7.5% copper), dental fillings (70% silver, 18% tin, 10%
copper, 2% mercury), ordinary steel (97% iron, 2% manganese, 1% carbon) and stainless
steel (71% iron, 18% chromium, 10% nickel, 1% carbon); naturally these percentages can
be varied somewhat from one sample to another.
Ionic Bonds: Of the three primary types of chemical bonds, the ionic bond is the most
binding--it is harder to break the ionic bond than any of the other types of bonds. Ionic
bonds occur between elements that are no more than 2 electrons away from having a
noble gas structure. Thus, ionic bonds form between alkalis and halogens, alkalis and
group 16 elements, alkali earths and halogens, and alkali earths and group 16 elements. It
is possible to have ionic bonds involving greater degrees of ionization. For instance,
Al2O3 is an ionic bond even though the aluminum atoms are ionized to +3. But the
predominant trend is to have ionic compounds from among the four possibilities listed
above.
Ionic compounds share several features. One is that they always involve a metal and a
non-metal. Thus, lead and sodium will never form an ionic bond with each other.
Second, when the ionic molecules group together into a large-scale solid, the molecules
are very regularly spaced throughout the solid in a crystalline arrangement. When ionic
compounds come in a large-scale size, there is significant sharing between adjacent
atoms. For instance, table salt comes in a plain cubic structure in which along any axis
you have a sodium atom alternating with a chlorine atom. However, there are 6 sodium
atoms in close contact with each chlorine atom and 6 chlorine atoms in contact with each
sodium atom. Thus, in a bulk solid, any chlorine atom is sharing only 1/6 of its bonding
with any particular adjacent sodium atom. However, since there are 6 adjacent sodium
atoms, everything works out fine. The situation is a little more difficult for an ionic solid
like cesium chloride, CsCl. Here, we still have one chlorine atom for every cesium atom.
But it turns out that the structure isn't strictly cubic like sodium chloride. Instead, in
cesium chloride, the chlorine atoms form the outline of a cube, requiring 8 chlorine
atoms. The cesium atoms lie in the center of the cube; one-eighth of each chlorine atom
is effectively bonding to the cesium atom in the center, so the total attraction balances
out. The ratio of chlorine to cesium looks suspiciously far from 1 to 1. However, if we
start examining large numbers of such cube-like structures strung together in 3
dimensions, we find that the ratio of chlorine to cesium soon approaches 1 to 1,
especially when taking into consideration the large number of atoms typically involved in
an everyday-sized object. Things get even more complicated for such substances as
calcium fluoride (aka fluorite), CaF2. In this configuration, fluorine atoms lie at the
vertices of a cube; in addition, there are fluorine atoms in the center of each face of a
cube and at the midpoints between fluorine atoms at the vertices of the cube and another
fluorine atom at the very center of the cube. This gives a total of 27 fluorine atoms in one
cube. Within the cube, there are 4 calcium atoms. Over a large repetition of this pattern,
there winds up being a ratio of 2 fluorine atoms per calcium atom.
Here we have a representation of NaCl. A unit cell is represented by a cube with the
chlorine atoms (the large circles) at the vertices of the cubes. Visualizing this, we see
that we will also have a chlorine atom at the center of each face of the cube. Each
chlorine atom at a vertex contributes one-eighth of a chlorine atom while each chlorine
atom at a face contributes half an atom. The total number of whole chlorine “atoms” in
the cube is then 4. We also see sodium atoms (the small circles). Each cube as defined
above will have a sodium atom midway between each chlorine atom. There is a sodium
atom at each of the 12 edges of the cube plus one whole sodium atom in the center. Since
each sodium atom on the edge is only ¼ in the cube, the total number of whole sodium
“atoms” is 12  ¼ + 1 = 4. There is therefore a 1 : 1 ratio between chlorine atoms and
sodium atoms, exactly what we expect from the chemical formula NaCl.
Covalent Bonds: Whereas an ionic bond basically involves one atom giving up one or
more electrons to another atom to form a strong electrical bond, the covalent bond
involves atoms sharing electrons with each other. The prototypical example of a covalent
bond is the hydrogen molecule. A single hydrogen atom can gain a single electron to
assume a stable helium-like structure. Thus, when two hydrogen atoms approach each
other, each will share its electron with the other; in effect, each hydrogen atom gains an
electron without giving up its own electron. Technically, it is possible for all atoms-even the noble gases--to bond covalently, provided each atom feels like it has a stable,
noble gas-like electronic structure.
Some atoms require more than one electron to assume a stable structure. Oxygen is an
example--it needs two electrons to become stable. Thus, when two oxygen atoms come
together, each shares two electrons with the other; this is called a double bond. Nitrogen
needs three atoms, so when two nitrogen atoms come together, each shares three
electrons with the other; this is a triple bond. The atoms can also be satisfied by bonding
covalently with other elements. For instance, two hydrogen atoms can each share their
electrons with an oxygen atom and the oxygen atom can share two of its electrons, one
with each hydrogen atom. The result is H2O, water. Or three hydrogen atoms can share
their electrons with a nitrogen atom and the nitrogen atom can share three of its electrons
with the hydrogens, giving us NH3, ammonia. Variations can be played upon this theme.
A special case is the carbon atom. Carbon needs four electrons to approach stability.
However, when two carbon atoms come together, they will never share 4 electrons each
with each other; this is due to the subshell structure. However, other combinations are
possible. One carbon atom can share four electrons, split evenly between two oxygen
atoms, which in turn are each sharing two of their electrons with the carbon. The result is
CO2, carbon dioxide. Two molecules of special interest in this discussion are carbon
monoxide, CO, and ozone, O3. They actually have some unoccupied bonds. However,
they still have quasi-stability and can live for a reasonably long time (though they react
with other substances fairly readily).
Naming Compounds: In this chapter, we have examined how different atoms bond
together to form molecules. They do so based upon how far away they are from having
the same electronic structure as a noble gas. For instance, the alkalis need to lose one
electron, the alkali earths need to lose two electrons; similarly, the halogens need to gain
one electron and the elements in the column next to the halogens, topped by oxygen, need
to gain two electrons. Due to the existence of subshells, this simple method does not
apply uniformly to the transition metals. For those atoms, we can have varying degrees
of ionization. In the first part of today's lecture, we will discuss the nomenclature
involved in these situations. We see that different levels of ionization lead to different
names for the ion in question. Some of these ions are listed below.
Symbols and Charges for Monatomic Ions: a monatomic ion is simply one atom that has
had an appropriate number of electrons either removed (left column) or added (right
column) in order to achieve a noble gas-like valence shell.
This table contains those ions that don't have a variable charge allowed to them.
Symbol
H+
Li+
Na+
K+
Rb+
Cs+
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Ra2+
Ag+
Zn2+
Name
hydrogen ion
lithium ion
sodium ion
potassium ion
rubidium ion
cesium ion
beryllium ion
magnesium ion
calcium ion
strontium ion
barium ion
radium ion
silver ion
zinc ion
Symbol
Name
H
hydride
Ffluoride
Clchloride
Brbromide
I
iodide
O2oxide
S2sulfide
2Se
selenide
Te2telluride
3N
nitride
P3phosphide
As3arsenide
Al3+
aluminum ion
Note that the letters in an ion's name before the -ide ending is the stem. For example, the
stem for bromide is brom-.
We have more monotomic ions here; however, these are such that they can have a
variable number of electrons removed to achieve relative stability in their shell structure.
Systematic name Common
Symbol (Stock system)
name
Cu+
copper(I)
cuprous
Cu2+
copper(II)
cupric
Fe2+
iron(II)
ferrous
Fe3+
iron(III)
ferric
2+
Sn
tin(II)
stannous
Sn4+
tin(IV)
stannic
2+
Cr
chromium(II)
chromous
Cr3+ chromium(III)
chromic
2+
Mn
manganese(II)
manganous
Mn3+ manganese(III)
manganic
Systematic name Common
Symbol (Stock system) name
Hg22+
mercury(I)
mercurous
2+
Hg
mercury(II)
mercuric
Pb2+
lead(II)
plumbous
Pb4+
lead(IV)
plumbic
2+
Co
cobalt(II)
cobaltous
Co3+
cobalt(III)
cobaltic
2+
Ni
nickel(II)
nickelous
Ni4+
nickel(IV)
nickelic
+
Au
gold(I)
aurous
Au3+
gold(III)
auric
In addition for single atoms to be ionized, it is also very common for a group of atoms,
called polyatoms, to behave very similarly to the montomic ions. They are listed below.
Symbols and Charges for Polyatomic Ions
Formula
NO3NO2CrO42Cr2O72CNMnO4OHO22NH2CO32SO42SO32C2O42PO43PO33S2O32AsO43SeO42SiO32C4H4O62C2H3O2-
Name
nitrate
nitrite
chromate
dichromate
cyanide
permanganate
hydroxide
peroxide
amide
carbonate
sulfate
sulfite
oxalate
phosphate
phosphite
thiosulfate
arsenate
selenate
silicate
tartrate
Formula
Name
ClO4perchlorate
ClO3chlorate
ClO2
chlorite
ClOhypochlorite
IO4periodate
IO3iodate
IO
hypoiodite
BrO3bromate
BrO
hypobromite
HCO3- hydrogen carbonate (bicarbonate)
HSO4- hydrogen sulfate (bisulfate)
HSO3- hydrogen sulfite (bisulfite)
HC2O4- hydrogen oxalate (binoxalate)
HPO42- hydrogen phosphate
H2PO4- dihydrogen phosphate
HShydrogen sulfide
BO33borate
2B4O7
tetraborate
SiF62- hexafluorosilicate
acetate (an alternate way to write acetate is CH3COO-)
There is one positive polyatomic ion. It is NH4+ and is called the ammonium ion.
Note: Writing just the plus sign or minus sign for ions with +1 or -1 charges is
acceptable.
The trick in putting the positively-charged ions with the negatively-charged ions is to
keep the resulting molecule with a total charge of zero. Thus, if we want to combine
calcium with borate, we have to recognize that the calcium ion has two positive charges
and the borate ion has three negative charges. In order to have the molecule electrically
neutral, we will need to have 3 calciums and two borates: 3 · (+2) + 2 · (-3) = 0. This
give us the molecule Ca3(BO3)2, calcium borate. We can also work the other way. Let's
supposed we have one of our ions be one with a variable ionization. Let's look at FeO
and Fe2O3. In the first case, we know that the negatively-charge ion, oxygen, must have a
total ionization of -2. Therefore, the one iron ion must have a total ionization of +2. This
makes FeO ferrous oxide. In the second case, the three oxygen ions have a total charge
of 3 · (-2) = -6. Therefore, the two irons must have a total charge of +6; this means that
each iron atom must have an ionization of 6/2 = +3. We now see that Fe2O3 is ferric
oxide.