Notes
Electrons in Atoms
Electrons form the outer shell of the atom. It is these outer shells which interact with the world and are responsible for the
chemical and physical properties of atoms. Electrons are located in orbitals, regions which correspond to specific energy
levels (see the Bohr and Quantum theories of the atom). These energy levels are very easily seen on the periodic table (see
handout).
There are 4 types of orbitals:
i)
ii)
iii)
iv)
s orbital - can hold 2 electrons
p orbital - has three sub-orbitals, each of which can hold 2 electrons for a total of 6
d orbital - has five sub-orbitals, each of which can hold 2 electrons for a total of 10.
f orbital - has 7 sub-orbitals, each of which can hold 2 electrons for a total of 14.
Electrons tend to occupy the lowest available orbital. The simplest atom, hydrogen has 1 electron. It its lowest, or ground
state, this electron will occupy the 1s orbital, the lowest energy orbital available (see chart, page 297) The next element,
helium, has two electrons, both of which will occupy the 1s orbital. Element three, lithium, has three electrons. The first
two will fill the 1s orbital while the third must move up to the next energy level, 2s. Thus the electron configuration of an
atom is the arrangement of the electrons from the lowest energy level to the highest.
Electron configurations of period two elements
Representative
Element
Group
Electron configuration
1
lithium
1s22s1
2
beryllium
1s22s2
13
boron
1s22s22p1
14
carbon
1s22s22p2
15
nitrogen
1s22s22p3
16
oxygen
1s22s22p4
17
fluorine
1s22s22p5
18
neon
1s22s22p6
Note that the order of the orbitals does not always follow numerically with the energy level or the period. For instance, the
3d level comes between the 4s and the 4p. The reason for this has to do with quantum theory and is not important here. It
is important, however, that the order be observed when giving the electron configuration of an element. For instance, the
electron configuration of selenium (34 protons, 34 electrons) is:
1s22s22p63s23p64s23d104p4
Assignment:
Write the electron configuration for the following elements:
N, P, As, Br, Kr, Na1+, F1-, Ne
Orbital Diagrams are another way to illustrate the position of electrons. They take more room to draw, but can give
information concerning bonding. They are best learned by comparision with electron configuration:
eg.
Na (11 protons, 11 electrons)
electron configuration:
1s22s22p63s1
orbital diagram
1s
2s
2p
↑↓
↑↓
↑↓ ↑↓ ↑↓
:
3s
↑
Each line represents an orbital. Note the three lines under the 2p orbital to represent the three sub-orbitals. An arrow above
a space represents an electron occupying an orbital. It can be either up or down, to represent the spin of the electron. Two
arrows, one up, the other down represents a situation where the orbital is full. Following are the electron configurations
and orbital representations for the period two elements:
Electron arrangements and orbital diagrams of period two elements
Representative
Element
Electron configuration
Orbital Diagram
Group
1s
2s
2p
1
lithium
1s22s1
↑↓
↑
2
beryllium
1s22s2
↑↓
↑↓
13
boron
1s22s22p1
↑↓
↑↓
↑
14
carbon
1s22s22p2
↑↓
↑↓
↑ ↑
15
nitrogen
1s22s22p3
↑↓
↑↓
↑ ↑ ↑
16
oxygen
1s22s22p4
↑↓
↑↓
↑↓ ↑ ↑
17
fluorine
1s22s22p5
↑↓
↑↓
↑↓ ↑↓ ↑
18
neon
1s22s22p6
↑↓
↑↓
↑↓ ↑↓ ↑↓
Assignment:
Repeat the last assignment, giving the orbital representations for the elements.
Electron dot (Lewis) diagrams form a third way to illustrate the position of the electrons in an atom or ion. It is more
compact than either the electron configuration or the orbital representation, because it gives information only concerning
the valence electrons. Valence electrons are the electrons on the outside of an atom; they are the electrons responsible for
bonding and are also the electrons gained or lost when an atom ionizes. Valence electrons are electrons in the s and p
orbitals of the highest energy level reached by the electrons of an atom. In this class when valence electrons are mentioned,
the only elements concerned are those in groups 1, 2, and 13 through 18. The elements in period two are illustrative of the
groups:
Electron arrangements, orbital diagrams and valence electrons of period two elements
Representative
Electron
Valence
Group
Element
configuration
Orbital Diagram
electrons
1s
2s
2p
1
lithium
1s22s1
↑↓
↑
1
2
beryllium
1s22s2
↑↓
↑↓
2
13
boron
1s22s22p1
↑↓
↑↓
↑
3
14
carbon
1s22s22p2
↑↓
↑↓
↑ ↑
4
15
nitrogen
1s22s22p3
↑↓
↑↓
↑ ↑ ↑
5
16
oxygen
1s22s22p4
↑↓
↑↓
↑↓ ↑ ↑
6
17
fluorine
1s22s22p5
↑↓
↑↓
↑↓ ↑↓ ↑
7
18
neon
1s22s22p6
↑↓
↑↓
↑↓ ↑↓ ↑↓
8
2
Again note that the 1s electrons are not considered valence electrons since they are in a lower energy level. Only the
electrons in the second energy level (in this case) are valence electrons.
An electron-dot diagram (or Lewis diagram) begins with the symbol of the element. Dots are drawn around the symbol to
represent the valence electrons; one for each electron. The placement of the electrons is important; they must be placed in
four pairs, left, right, top, and bottom. Pay close attention to the placement of the electrons in the table below:
Electron arrangements, orbital diagrams and Lewis diagrams of period two elements
Group
Representative
Element
Electron
configuration
Valence
electrons
Orbital Diagram
1s
2s
2p
1
lithium
1s 2s
↑↓
↑
1
2
beryllium
1s22s2
↑↓
↑↓
2
13
boron
1s22s22p1
↑↓
↑↓
↑
3
14
carbon
1s22s22p2
↑↓
↑↓
↑ ↑
4
15
nitrogen
1s22s22p3
↑↓
↑↓
↑ ↑ ↑
5
16
oxygen
1s22s22p4
↑↓
↑↓
↑↓ ↑ ↑
6
17
fluorine
1s22s22p5
↑↓
↑↓
↑↓ ↑↓ ↑
7
18
neon
1s22s22p6
↑↓
↑↓
↑↓ ↑↓ ↑↓
8
Assignment:
2
1
Lewis
diagram
Li
Be
B
C
N
O
F
Ne
For the same elements as the last two assignments, give the correct Lewis diagram.
On the table on the previous page the Lewis diagram gives an indication of the bonding capabilities of each atom. Where
two dots are together, these represent paired valence electrons which fill an orbital and cannot be used for bonding. These
are called lone pair electrons. Where a single dot is found, this represents an unpaired electron in a half-filled orbital. This
electron can be used for bonding and is thus called a bonding electron.
If one looks at the pattern of bonding capabilities in the table above, the following pattern is evident:
1s
Orbital and Lewis diagrams with Lone Pair and Bonding Electrons
Valence
Electron-dot
Lone Bonding
Orbital Diagram
electrons
Diagram
Pairs Electrons
2s
2p
lithium
↑↓
↑
1
beryllium
↑↓
↑↓
2
boron
↑↓
↑↓
↑
3
carbon
↑↓
↑↓
↑ ↑
4
nitrogen
↑↓
↑↓
↑ ↑ ↑
5
oxygen
↑↓
↑↓
↑↓ ↑ ↑
6
fluorine
↑↓
↑↓
↑↓ ↑↓ ↑
7
neon
↑↓
↑↓
↑↓ ↑↓ ↑↓
8
Element
Li
Be
B
C
N
O
F
Ne
3
Bond
Capacity
0
1
1
0
2
2
0
3
3
0
4
4
1
3
3
2
2
2
3
1
1
4
0
0
You will note that for beryllium, boron and carbon that there is a discrepency between the orbital representation and the
electron-dot diagram. On an orbital representation crossed lines indicates a lone pair, while a single slash indicates a
bonding electron. On the orbital representations above, beryllium has one lone pair in the 2s orbital and no bonding
electrons, while the Lewis diagram shows two bonding electrons and no lone pairs. A similar situation exists for boron and
carbon.
The Lewis diagram is right. The situation can be explained using the principle of electron promotion, or hybridization. If
you look at figure 10-16 on page 297 you can see the energy differences between electrons in different orbitals. The
valence electrons in the s and p orbitals for each energy level have a very small difference in energy, so it is possible for an
electron in an s orbital to receive a small input of energy and move up into the adjacent p orbital. This gives the atom a
half-full s orbital and an extra half-full p orbital for bonding purposes.
Why electron promotion happens can be explained in terms of energy. All atoms desire to achieve the lowest possible total
energy. When an atom bonds with another atom it moves to a lower energy level, thus explaining why atoms bond. When
an atom promotes an electron, it gains energy, but the extra bonding capacity achieved allows it to get to a lower total
energy than it could without electron promotion. See the graph below:
│
│
│
│
2
Total
│

Energy
│ 1
│ 
│
│
│
3
│

└───────────────────────────────────
Point 1 represents the energy of an unhybridized atom. Point 2 represents the energy increase needed to promote an
electron. Point 3 represents the net decrease in energy which occurs when a hybridized atom forms bonds with other atoms.
Electron promotion happens only when it increases the bonding capacity and only happens between the s and p orbitals of
the same energy level. Following these criteria, hybridization occurs only in atoms of groups 2, 13, and 14:
Orbital Diagrams of Hybridized Atoms
Unhybridized
Orbital Diagram
Element
1s
2s
beryllium
↑↓
↑↓
boron
↑↓
↑↓
carbon
↑↓
↑↓
Hybridized
Orbital Diagram
2p
1s
2s
2p
↑↓
↑
↑
↑
↑↓
↑
↑ ↑
↑ ↑
↑↓
↑
↑ ↑ ↑
Note: For atoms of groups 2, 13, and 14, hybridization is the rule. It happens all the time for these elements, so it must be
taken into account when using these elements. Thus the complete picture is like this:
4
Element
Valence
electrons
1s
Orbital Diagram
2s
2p
lithium
↑↓
↑
beryllium
↑↓
↑
↑
2
boron
↑↓
↑
↑ ↑
3
carbon
↑↓
↑
↑ ↑ ↑
4
nitrogen
↑↓
↑↓
↑ ↑ ↑
5
oxygen
↑↓
↑↓
↑↓ ↑ ↑
6
fluorine
↑↓
↑↓
↑↓ ↑↓ ↑
7
neon
↑↓
↑↓
↑↓ ↑↓ ↑↓
8
Electron-dot
Diagram
Li
1
Lone Bonding
Pairs Electrons
0
Be
B
C
N
O
F
Ne
Bond
Capacity
1
1
0
2
2
0
3
3
0
4
4
1
3
3
2
2
2
3
1
1
4
0
0
Valence level expansion
Some compounds occur which cannot be easily explained with the information given so far:
PF5, SF6, ClF7, ArF8
Normally these elements have a bonding capacity much lower than is indicated by these formulas. What is happening is
that every valence electron becomes a bonding electron. The number of valence orbitals expand allowing a large increase
in bonding capacity. The reason for this is the same as for hybridization; to achieve a lower total energy.
Elements which expand their valence level include elements of groups 15 to 18, from period 3 down. The reason for this is
that expansion is into the unused d orbital of the same energy level; periods 1 and 2 do not have a d orbital, so there is no
valence level expansion for elements of these periods.
Chemical Bonding
When two atoms are joined together chemically, there is a chemical bond between them. Bonds form for reasons of energy.
When atoms combine to form molecules they achieve a lower energy state than they could as individual atoms (see figure
6-5, page 165). There are two main reasons why elements will combine:
1)
To share electrons - All elements want to have the same number of valence electrons as the noble gases, and
will either give electrons away or steal electrons in order to achieve this. Sometimes, however, this is not
possible. The next best thing is to get close to another atom with the same problem. (see figure 6-7, page 167).
For example, oxygen has 6 valence electrons. It would like to have 8 in order to have the same number of
valence electrons as Neon. If it cannot steal two electrons it will get close to another oxygen atom. Each atom
will put forward its 2 bonding electrons and their electron clouds will mesh, forming the compound O2. This
type of bonding is called covalent bonding.
2)
To balance charge - When atoms gain or lose electrons they become ions. These ions tend to be attracted to
ions of opposite charge and repelled by ions of the same charge, much like the poles of a magnet. These ions
will combine in proportions which completely balance the charges and form a compound which is electrically
neutral You used this principle when you studied chemical nomenclature. This type of bonding is called ionic
bonding (see figure 6-17, page 179)
Now the question remains; how can we determine whether two atoms will form covalent bonds or ionic bonds? This
question is easily answered using the principle of electronegativity. Electronegativity is a measure of how strongly an atom
is holding on to its valence electrons. It is related to ionization energy (chapter 5). If an atom loses an electron fairly easily
5
it has a low electronegativity (and tends to be a cation). If an atom tends not to lose electrons, but tends to steal them from
other atoms (and become an anion) it has a high electronegativity.
You have been given a table of electronegativities of the elements. To determine what type of bond exists between two
atoms you subtract their respective electronegativities:
- if the electronegativity difference is 0.2 or less, the bond is covalent
- if the electronegativity difference is 1.7 or greater the bond is ionic.
There exists a grey area where the electronegativity difference is greater than 0.2, but less than 1.7. When the
electronegativity difference is in this range the atom with the greater electronegativity is strong enough to pull the bonding
electrons so that they spend more time around it than the other atom, but is not strong enough to pull the bonding electrons
away completely and form ions. These bonding arrangements form electric dipoles where one end has a slightly positive
charge and the other end has a slightly negative charge. The greater the electronegativity difference, the greater the dipole:
The bond made in this electronegativity range is called a polar covalent bond.
Electronegativity
Electronegativity
Element bonded to F
Fluorine
difference
Compound
Bond
FF
FF
F
4.0
4.0
0.0
F2O
OF
O
3.5
4.0
0.5
NF3
NF
N
3.0
4.0
1.0
CF4
CF
C
2.5
4.0
1.5
BF3
BF
B
2.0
4.0
2.0
BeF2
BeF
Be
1.5
4.0
2.5
LiF
LiF
Li
1.0
4.0
3.0
Bond type
COVALENT
│
↑
increasingly ionic │
│
│


│
│


│
│


│
│
│ increasingly covalent
↓
│
IONIC
Molecular Diagrams
Both orbital representations and electron-dot diagrams may be used to show how atoms bond to form molecules.
It is important to remember that only the valence electrons may be used for bonding, and of those, only the bonding
electrons may be shared. Thus an element like chlorine with 3 lone pairs and 1 bonding electron may form only 1 bond.
To show a bond formed using these two methods is quite simple. Diagrams may be made of single, double, or triple bonds.
Draw the examples used in class.
6
Making diagrams of covalently-bonded molecules
Any covalently-bonded molecule is held together by shared valence electrons. This is different from ionic bonds, which
will be discussed later. The types of molecules this method covers includes molecular compounds (non-metal bonded to
non-metal) and organic compounds (containing C and H). We will also include some group 1 and 2 elements.
The key for making diagrams of these molecules is to remember some simple facts:
i)
Atoms form covalent bonds in order get close to a complete valence set of electrons; an atom will continue to
form bonds until it is surrounded by a complete set of valence electrons.
ii)
When you are given a chemical formula, the assumption is that each atom in the molecule has a complete set of
valence electrons surrounding it; you should be able to draw a structure which satisfies the bonding capacities of
each atom in the molecule.
iii)
When given a molecule for diagramming which has more than 2 atoms, look for bonding capacity. The atom
(or atoms) with the greatest bonding capacity will be the central atom; the atom at the middle of the molecule
which has all the other atoms attached to it.
There are three types of diagrams you should be able to draw; orbital diagrams, lewis diagrams and structural diagrams.
Draw the examples from class.
Coordinate Covalent Bond
Explains structures where conventional covalent bonding does not work. The best example is the ammonium ion, NH41+. It is a
stable, covalently bonded polyatomic ion. It is the preferred form of ammonia (over NH 3). The question is, where does the extra
hydrogen go? A coordinate covalent bond is one where one atom provides both the bonding electrons; it uses a lone pair to bond
with an atom that has no bonding electrons.
Write down the drawings in class.
Resonance
Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. Sometimes there
is a structure that includes both a double and a single bond. If the two bonds could reverse positions without changing any other
bonds in the molecule, they can and do reverse positions. Write down the examples in class of ozone and benzene. The main result
of resonance is in bond energy. Bond energy refers to the amount of energy that must be applied to a chemical bond to break it.
Single bonds take less energy to break than double bonds. Thus a molecule held together with single bonds is easier to take apart
than one held together with double bonds. When resonance occurs the bonding electrons responsible for the double bond change
position so often that the bond energy of a resonant structure is approximately midway between that of a single bond and a double
bond. This can lead to these structures being very stable, like benzene. Sometimes resonance does not result in a stable molecule;
the example is ozone.
7
Molecular Shape
The shape of molecules is determined by the nature of the central atom. The central atom is the atom in a molecule with the
greatest bonding capacity and can be an atom with an ability to make 2 or more bonds. This limits this discussion to
elements from groups 2, 13, 14, 15 and 16.
Molecular shape is guided by one principle; the tendency of electrons to repel one another. This includes the electrons from
individual atoms in a chemical compound. The electron cloud from each atom in a compound will arrange itself so as to
put the maximum distance between it and all the other electron clouds from the other atoms in the compound. Because the
electrons in question are the valence electrons, the theory which describes molecular shape is called Valence Shell Electron
Pair Repulsion (VSEPR) Theory.
Picture an element like beryllium; it has two bonding electrons and no lone pair electrons. It thus can bond with two other
atoms to create a compound like BeF2. Each bond consists of two pairs of electrons, each forming its own electron cloud.
These electron clouds repel one another and try to get as far from each other as possible. This compound forms the
following structure about the central atom:
F  Be  F
This is called a linear structure, meaning the three atoms arrange themselves in a line, so that the electron clouds from the
two fluorines are the maximum distance apart.
Look at figure 6-5 on page 186 of your text. It summarizes the electron configuration about the central atoms from groups
2, 13 to 16, in addition to two arrangements which result from valence level expansion.
Two shapes; pyramidal and angular, result when the bonding electrons (and the atoms bonded to them) must also deal with
repulsion from the lone pair electrons on the central atom. Group 15 elements, like nitrogen, have 3 bonding electrons and
1 lone pair electron. The three bonds are forced out of the planar position by the lone pair, resulting in a pyramidal
structure. Group 16 elements, like oxygen, have 2 bonding electrons and 2 lone pairs. The two bonds made are forced out
of a linear shape by repulsion from the two lone pairs, resulting in an angular shape.
Polarity
Polarity is concerned with whether a molecule has a positively charged end and a negatively charged end, or not. If this is
the case the molecule is polar and the chemical and physical properties of such molecules will be altered. The importance
of this will be discussed in the next unit.
Whether a molecule is polar or not depends on two things:
i)
ii)
the existence of bond dipoles
molecular symmetry
Bond dipoles exist when the electronegativity difference between the members of a bond is sufficiently high (between 0.2
and 1.7). If a molecule contains bond dipoles it has the potential to be polar.
Molecular symmetry is concerned with how the atoms in a molecule are distributed about a central atom. Molecules may
be either symmetrical or asymmetrical. Of the groups on the periodic table we have studied, this applies to groups 2, 13,
14, 15 and 16. Of these groups, 2, 13 and 14 form symmetrical molecules while groups 15 and 16 form asymmetrical
molecules. Very simply, polarity and symmetry are related as follows:
a)
A symmetrical molecule in which the central atom is bonded to atoms of the same element will be nonpolar in all cases, even when bond dipoles exist. e.g. CCl4, BCl3 and BeCl2
b)
An asymmetrical molecule with no bond dipoles will be non-polar. eg. NBr3
c)
An asymmetrical molecule which contains bond dipoles will be polar in every case. The direction of the
polarity will be in the direction of the atom(s) with the greatest electronegativity; that atom will be the
8
negative end of the molecule, the atom(s) with the least electronegativity will be the positive end. e.g.
H2O, NH3
d)
A symmetrical molecule which contains bond dipoles may be polar if the central atom is bonded to two or
more different types of atoms with different electronegativities. Again, the direction of polarity is in the
direction of least electronegativity to greatest. e.g. CHCl3.
Properties of Molecules
Once we have determined whether a molecule is polar or not we can make certain predictions about properties. The
properties of a substance are determined by how the molecules interact with each other. How they interact is in turn
determined by the forces of attraction between molecules. These are called intermolecular forces; the forces which act
between molecules, to draw them together, forming the various phases of matter.
To facilitate this discussion an understanding of energy and its relationship to matter is necessary. Kinetic energy is the energy of
motion. All matter is in motion; the nature of the motion depends upon the phase of the substance, whether it is solid, liquid or
gas.
The motion of the particles of matter is connected to temperature. The higher the temperature, the faster the motion. The lower
the temperature, the slower the motion. At a temperature of 0 Kelvin (-273°C), all motion stops. This is called absolute zero, and
is the point where matter has no kinetic energy. As energy is put into matter, temperature goes up. Thus temperature is a rough
measure of the kinetic energy in a substance.
A solid has its atoms held relatively rigidly in place; the only motion allowed the atoms is for them to vibrate in place. A liquid
allows its molecules to slide by each other (which gives a liquid its properties), but the molecules are still very close to one
another and interact strongly. A gas gives the molecules free rein; they move about very rapidly and interact with each other only
weakly.
Whether a particular substance is a particular phase depends both upon the kinetic energy of the substance and the forces which
bind the molecules together. In general, the weaker the intermolecular forces, the less energy which is required for the substance
to go from solid to liquid and from liquid to gas. Thus the physical properties of melting point and boiling point gives a rough
measure of the intermolecular forces which bind the molecules to each other. The higher the melting and boiling points, the
greater the strength of the intermolecular forces which bind the molecules together.
Following is a list of the intermolecular forces which we are concerned with:
1.
Van der Waals forces
a)
London dispersion forces
b)
Dipole-dipole attraction
c)
Hydrogen bonding
2.
Ionic bonding
3.
Metallic bonding
4.
Network covalent bonding
Van der Waals forces are weak forces which bind all matter together (in many substances, this is the only attractive force).
London Dispersion Forces are the dominant forces between covalently bonded, non-polar molecules
(ex. N2, O2, CH4). They result from the formation of "instantaneous dipoles" in electrically neutral matter. They are weak and act
only over a short distance Force increases with increasing number of electrons; the more electrons, the stronger the force. In very
large molecules it can be a very large force. (Analogy: Gravity)
Small molecules generally have a low melting point and boiling point (if molecules have a molar mass of less than 50, they are
generally gases at room temperature).
9
The force varies with shape of the molecule. A long molecule has a large number of electrons exposed; this increases the force, so
it has a relatively high boiling point. This same structure is very flexible and does not stack well, so it has a low melting point. A
more compact molecule exposes fewer electrons, so its boiling point is lower, but it stacks better, so its melting point is higher:
Normal Pentane (C5H12)
Neopentane (C5H12)
m.p. -130C,
m.p. -20°C,
b.p. 36°C
b.p. 9°C
Dipole-Dipole Attraction is a force which acts between polar molecules (ex. H2S). It results from the attraction of the opposite
poles of the permanent molecular dipoles. These substances generally have higher melting and boiling points than non-polar
molecules with similar molecular weights.
Hydrogen Bonding is a specialized form of dipole-dipole attraction. It occurs as when O, N, and F are bonded to H, owing to the
large electronegativity difference between these elements and hydrogen: This is a stronger force than standard dipole-dipole
attraction. Molecules with hydrogen bonding will have boiling points and melting points quite a bit higher than molecules that
have only dipole-dipole or London dispersion forces. Hydrogen bonding is responsible for many of the unusual properties of
water.
O - H
N - H
F - H
3.5 - 2.1 = 1.4
3.1 - 2.1 = 1.0
4.1 - 2.1 = 2.0
Ionic Bonding generally occurs in compounds of metals and non-metals (salts). It is the result of the attraction of oppositely
charged ions. They come together in order to neutralize charge, and the attraction is relatively strong. The structures formed are
very orderly and are given the name crystal lattice. Ionic solids are called crystals. No sharing of electrons occurs between the
ions in the crystal lattice. As a result, ionic solids are brittle. Ionic solids conduct electricity only in the molten state, and not very
well. Ionic solids are characterized by very high melting and boiling points.
Metallic Bonding is the bonding which occurs between metals in the Periodic Table. It is characterized by close packing
of the atoms, with the electrons delocalized; that is, they are free to jump from atom to atom, filling unoccupied orbitals.
This free sharing of electrons allows metals to conduct electricity freely (copper conducts electricity 100 000 times better
than molten NaCl). The free electrons also act as a lubricant, allowing metal atoms to slide over one another without
affecting the integrity of the material. Thus metals are malleable and ductile. This bond is strong, giving most metals high
melting and boiling points.
Network Covalent Bonding This is the traditional covalent bond, expanded to 2 or 3 dimensions in a network which is
theoretically infinite (much like an ionic crystal lattice). Network solids include diamonds, graphite, quartz, and most
rocks. Because the covalent bond is stronger than any other bond, the network solid is very hard (diamonds are the hardest
substance known). Because electrons are held tightly in their bonds, network solids are brittle, and they do not conduct
electricity. Because the orientation of the atoms is very specific to the bonding orientation (tetrahedral, planar trigonal),
network solids form distinct crystals. Because of the strength of the bonds, network solids have very high melting and
boiling points
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