Octet rule / Chemical bonding
• The octet rule is a simple chemical rule of thumb
that states that atoms tend to combine in such a way
that they each have eight electrons in their valence
shells, giving them the same electronic configuration
as a noble gas. The rule is applicable to the maingroup elements, especially carbon, nitrogen, oxygen,
and the halogens, but also to metals such as sodium
or magnesium. In simple terms, molecules or ions
tend to be most stable when the outermost electron
shells of their atoms contain eight electrons
Octet rule
• The octet rule states that atoms are most stable when they
have a full shell of electrons in the outside electron shell. The
first shell has only two electrons in a single s subshell.
• Helium has a full shell, so it is stable, an inert element.
• All the other shells have an s and a p subshell, giving them at
least eight electrons on the outside. The s and p subshells
often are the only valence electrons, thus the octet rule is
named for the eight s and p electrons
Chemical reactions
• Chemical reactions take place because atoms
need to achieve a more stable electronic
configuration.
• Chemical reactions involve the re-distribution
of the electrons in the outter shell.
• Outter shell electrons are called valence
electrons
Octet rule / Periodic table
• All atoms ,except the group 8 atoms need to
achieve 8 electrons in the outter shell.
• In order to do these atoms have two choices,
ie to loose electrons or to gain electrons.
• Atoms in groups 1-3 will loose ,ie donate their,
electrons when they react.
• Atoms in groups 5-7 will accept, ie gain
electrons, when they react.
Valence shells
• The valence shell is the outermost shell of an
atom. It is usually said that the electrons in this
shell make up its valence electrons, that is, the
electrons that determine how the atom behaves
in chemical reactions.
• Atoms with complete valence shells,Group 8 or 0
= noble gases) are inert (unreactive).
• Those with only one electron in their valence
shells (alkalis) or just missing one electron from
having a complete shell (halogens) are the most
reactive.
Periodic table and valence
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On the Periodic Table with shell totals you can easily see the
octet rule.
A valence is a likely charge on an element ion.
All of the Group 1 elements have one electron in the outside
shell and they all have a valence of plus one.
Group 1 elements will lose one and only one electron, to
become a single positive ion with a full electron shell of eight
electrons (an octet) in the s and p subshells under it.
Example Na --- Na+ + e– 2,8,1
2,8
Group 2
• Group 2 elements all have two electrons in the
outer shell and all have a valence of plus two.
Beryllium can be a bit different about this, but
all other Group 2 elements can lose two
electrons to become +2 ions. They do not lose
only one electron, but two .
• Example Ca --- Ca2+ + 2e2,8 ,8,2
2,8,8
•
Group 3
• Group 3 elements have a valence of plus
three. Boron is a bit of an exception to this
because it is so small it tends to bond
covalently. Aluminum has a valence of +3.
• Example Al --- Al 3+ + 3e-
Group 4
• The smallest Group 4 elements, carbon and
silicon, are non-metals because it is difficult
to lose the entire four electrons in the outer
shell.
• Small Group 4 elements tend to make only
covalent bonds, sharing electrons. Larger
Group 4 elements have more than one
valence, usually including +4.
• We will explain this later
Group 5
• Small Group 5 elements, nitrogen and
phosphorus, are non-metals.
• They tend to either gain three electrons to
make an octet or bond covalently. The larger
Group 5 elements have more metallic
character.
• We will explain these later.
Group 6
• Small Group 6 elements, oxygen and sulfur,
tend to either gain two electrons or bond
covalently. The larger Group 6 elements have
more metallic character.
• Example O + 2e- --- O22 ,6,
2 , 8,
•
Group 7
• Group 7 elements all have seven electrons in
the outer shell and either gain one electron to
become a -1 ion or they make one covalent
bond.
• Example F + 1e- -- F–
2 ,7,
2 , 8,
• The Group 7 elements are diatomic gases due
to the strong tendency to bond to each other
with a covalent bond.
Group 0/8 = Inert
• All of the inert elements, the noble gases,
have a full octet in the outside shell (or two in
the first shell).
• They do not naturally combine chemically
with other elements.
• Inert = chemically un-reactive.
Chemical bonding: ionic
• Two basic types: ionic or covalent
• Ionic bonding: When atoms react they loose or gain
electrons. Atoms from groups 1-3 all react by
loosing/or donating their electrons.
• All metals behave this way, including Fe,Cu,Zn. On
the other hand atoms in groups 6 and 7 react by
accepting /or gaining electrons.
• These are the non-metals such as F, Cl, I, O, and S.
Ionic bonding/ ions
• All metals loose electrons to form positive
ions. These are called cations ,
• i.e Na+, Ca 2+, Fe2+, Al3+
• All non-metals gain electrons to form
negatively charged anions,
• i.e. Cl- ,S2- and O2-
Ionic bonds
Ionic bonding
• Ionic Bonding
– Some atoms gain electrons to become anions
– Others lose electrons to become cations
– Ions are attracted by their opposing charges
– Electrical Neutrality Maintained
– Most Important Bonding in Rocks and Minerals
Ionic bonds
• The Ionic Bond: Ionic bonds are formed when there
is a complete transfer of electrons from one atom to
another, resulting in two ions, one positively
charged and the other negatively charged. For
example, when a sodium atom (Na) donates the
one electron in its outer valence shell to a chlorine
(Cl) atom, which needs one electron to fill its outer
valence shell, NaCl (table salt) results. Ionic bonds
are often 4-7 kcal/mol in strength
Chemical bonding: covalent
• Covalent bonding is when atoms share their valence electrons
in order to have 8 in their outer shells.
• This kind of bonding takes place between non-metals groups
4, 5 and 6.
• This bonding does not form ions.
• Such bonds lead to stable molecules if they share electrons in
such a way as to create a noble gas configuration for each
atom.
• Each atom donates one electron to form a single covalent
bond
Lewis structure
• The Lewis structures are just an attempt to show these
valence electrons in a graphic manner as they are used to
combine with other elements.
• The element symbol is in the center and as many as four
groups of two electrons are shown as dots above, below, to
the right and left of the element symbol to show the valence
electrons.
• All of the inert gases (noble gases) have all eight of the
electrons around the element symbol, except for helium,
which has only two electrons even with a full shell. Below is a
demonstration of the noble gases written in Lewis structure.
Notice the electrons are in red just to emphasize them.
Lewis structure
Covalent bonding = sharing of
electrons
Dot/ cross diagram
Water: bond pair and lone pair
Diatomic molecules/elements
• Why are some elements diatomic?
• Hydrogen gas forms the simplest covalent
bond in the diatomic hydrogen molecule.
• The halogens such as chlorine also exist as
diatomic gases by forming covalent bonds.
• The nitrogen and oxygen which makes up the
bulk of the atmosphere also exhibits covalent
bonding in forming diatomic molecules.
Covalent bonds
Covalent bonding
• Covalent Bonding
– Electrons share electrons to fill incomplete shells
– Most Important Bonding in Organic Materials (and
Organisms)
Covalent bonds
• The Covalent Bond: Covalent Bonds are the strongest
chemical bonds, and are formed by the sharing of a pair of
electrons.
• The energy of a typical single covalent bond is ~80
kilocalories per mole (kcal/mol). However, this bond energy
can vary from ~50 kcal/mol to ~110 kcal/mol depending on
the elements involved.
• Once formed, covalent bonds rarely break spontaneously.
Octet rule
Ionisation energy
Ionization energy
• Definition: The first ionisation energy is the energy
required to remove the most loosely held electron
from one mole of gaseous atoms to produce 1 mole
of gaseous ions each with a charge of 1+.
• This is more easily seen in symbol terms.
• It is the energy needed to carry out this change per
mole of Na.
• Na (g) --- Na+
(g)
+ e-
Ionisation energy
• The state symbols - (g) - are essential. When
you are talking about ionisation energies,
everything must be present in the gas state.
• Ionisation energies are measured in kJ mol-1
(kilojoules per mole). They vary in size from
381 (which you would consider very low) up
to 2370 (which is very high).
• All elements have a first ionisation.
Ionization energy
• It is the amount of energy required to remove
one electron form the outter shell of an atom.
• An atom can have as many ionization energies
as it has electrons.
• Sodium has 11 electrons so it can have 11
ionization energies. First ,second third and so
on… till eleven.
Ionisation energies for Na
Patterns
• The first 20 elements
• First ionisation energy shows periodicity. That means that it
varies in a repetitive way as you move through the Periodic
Table. For example, look at the pattern from Li to Ne, and then
compare it with the identical pattern from Na to Ar.
• These variations in first ionisation energy can all be explained
in terms of the structures of the atoms involved.
Trends in ionisation energy
Ionization energy
• Notice that the highest ionization energies are
seen in the Group O elements, known more
commonly as the inert, or noble gases.
• This is because the atoms of these elements
have p-orbitals that are full of electrons. This
electronic configuration is particularly stable,
so that the atoms of these elements are
extremely reluctant to lose any electrons.
• as Na+ and K+.
Ionization energy
• At the other extreme, the lowest ionization
energies are seen in the Group 1 elements,
known as the alkali elements.
• These elements only have one electron in
their outermost orbital, and loss of this one
electron will give an ion with a stable, noble
gas-like, electronic configuration. Hence, they
readily give up this outer electron. This means
they readily form singly charged cations.
Explaining the pattern in the first few elements
• Hydrogen has an electronic structure of 1s1. It is a very small
atom, and the single electron is close to the nucleus and
therefore strongly attracted. There are no electrons screening
it from the nucleus and so the ionisation energy is high (1310
kJ mol-1).
• Helium has a structure 1s2. The electron is being removed
from the same orbital as in hydrogen's case. It is close to the
nucleus and unscreened. The value of the ionisation energy
(2370 kJ mol-1) is much higher than hydrogen, because the
nucleus now has 2 protons attracting the electrons instead of
1.
Ionization energy
• It is influenced:
– By distance of outer electron from nucleus
– The further away the less energy.
– The balance between remaining electrons and
protons in the nucleus .
Factors affecting the size of ionisation energy
• A high value of ionisation energy shows a high attraction
between the electron and the nucleus.The size of that
attraction depends:
• The charge on the nucleus.The more protons there are in the
nucleus, the more positively charged the nucleus is, and the
more strongly electrons are attracted to it.
• The distance of the electron from the nucleus.
• Attraction falls off very rapidly with distance. An electron
close to the nucleus will be much more strongly attracted
than one further away.
Trends in ionisation energy down a group
• As you go down a group in the Periodic Table
ionisation energies generally fall. Taking Group
1 as a typical example:
• Why is the sodium value less than that of
lithium?
Ionisation energy of Group 1
Trends in Group 1
• There are 11 protons in a sodium atom but
only 3 in a lithium atom, so the nuclear charge
is much greater. You might have expected a
much larger ionisation energy in sodium, but
offsetting the nuclear charge is a greater
distance from the nucleus and more
screening.
• Li 1s22s1 1st I.E. = 519 kJ mol-1
• Na 1s22s22p63s1 1st I.E. = 494 kJ mol-1
Explaining pattern of ionisation
energy
Explaining pattern of ionisation
energy
Ionisation energy and reactivity
Ionisation energies and reactivity
• The lower the ionisation energy, the more
reactive is the atom/element.
• You can explain the increase in reactivity of
the Group 1 metals (Li, Na, K, Rb, Cs) as you go
down the group in terms of the fall in
ionisation energy. Whatever these metals
react with, they have to form positive ions in
the process, and so the lower the ionisation
energy, the more easily those ions will form.
Sodium: very reactive
• Sodium ((Latin natrium), is a soft, silvery
white, highly reactive element and is a
member of the alkali metals within "group 1“
It has only one stable isotope, 23Na.
• Sodium quickly oxidizes in air and is violently
reactive with water, so it must be stored in an
inert medium, such as kerosene. Sodium is
present in great quantities in the earth's
oceans as sodium chloride (common salt).
Sodium: ionization energy
• Na (1s22s22p63s1) ----> Na+ (1s22s22p6) + e• First ionization energy for sodium
• Can have a maximum of 11 ionization
energies.
Sodium Na
Chemical bonding
Ionic bonding
C-H Covalent Bond
Carbon-Hydrogen Covalent
Bonding
Bond Number
Example
Energy
(kcal/mol)
single
H
|
H--C--H
|
H
~80
double
H H
| |
H--C==C--H
| |
H H
~150
triple
H
|
C
|||
C
|
H
~200
Covalent bonds