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Lewis Dot Structures of Covalent Compounds

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Title: Lewis Dot Structures of Covalent Compounds


1
Lewis Dot Structures of Covalent Compounds
  • Atoms are made up of protons, neutrons, and
    electrons. The
  • protons and neutrons are located at the center of
    the atom,
  • the nucleus. These electrons can be divided into
    core
  • electrons and valence electrons. The valence
    electrons are
  • the outermost electrons and are the ones involved
    in
  • chemical reactions

2
Electrons Electrons occupy most of the volume of
an atom They arrange themselves in shells at
varying distances from the nucleus
The Nucleus Protons and neutrons are located in
the nucleus (center) of the atom
n0
n0
n0
n0
n0
Valence electrons These are the outermost
electrons and the ones In chemical reactions
3
The number of valence electrons varies by
element. For the Main Group elements, the number
of valence electrons is equal to the Group
Number that the elements belong to. For
example, Sodium (Na) belongs to Group 1A and
therefore has 1 valence electron.
4
The Periodic Table
5
For example, Bromine (Br) belongs to Group VIIA
and therefore has 7 valence electrons. We can
represent the valence electrons of an atom using
a Lewis dot symbol, in which the element symbol
is surrounded by dots representing the valence
electrons. For example, Oxygen has six valence
electrons, so its Lewis dot symbol is


Note the six dots representing the six valence
electrons
6
For example, neon has eight valence electrons, so
its Lewis dot symbol is


For example, carbon has four valence electrons,
so its Lewis dot symbol is



7
How many valence electrons does Potassium (K)
have?

1
How many valence electrons does Antimony (Sb)
have?
5
How many valence electrons does Phosphorus (P)
have?
5
How many valence electrons does Magnesium (Mg)
have?
2
8
The Noble Gas elements in Group VIIIA have either
two valence electrons (He) or eight valence
electrons (Ne, Ar, Kr, Xe, and Rn). These
elements are extremely stable because they have
full valence shells- two electrons for He in the
first row and eight electrons in each of the
later rows. This is the basis for the Octet Rule
- elements tend to react in a way to attain the
electron configuration of Group VIIIA
9
Metallic elements at the left side of the
Periodic Table tend to lose one or more
electrons to form positive ions, such as Na and
Mg2, each of which has the electron
configurationof the Noble Gas that preceds it.
Nonmetals at the right side of the Periodic
Table tend to either gain electrons to form
negative ions such as F-, O2-, and N3- or to
share electrons in covalent bonds. This learning
objective describes how this is done

10
Covalent Bond When nonmetallic elements react
with other nonmetallic elements, they share
electrons in order to obtain eight valence
electrons.
Each fluorine atom has seven valence electrons.
They each require one more electron to satisfy
the Octet Rule.
11
The left fluorine atom now has a total of eight
electrons and the right fluorine atom now has a
total of eight electrons around it. When
nonmetallic elements react with other nonmetallic
elements, they share electrons in order to obtain
eight valence electrons.
12
The two electrons that form the covalent bond are
often Represented by a single line. The F2
molecule can be represented using a line and
dots to show the bonding pair and the six lone
pairs, respectively. This is called a Lewis dot
structure.
13
Multiple Covalent Bond
Some atoms have to share more than one electron
in order to satisfy the Octet Rule.
14
Each oxygen atom has six valence electrons. They
each require two more electrons to satisfy the
Octet Rule.
15
  • The left oxygen atom now has a total of eight
    electrons around it. The right oxygen atom now
    has a total of eight electrons around it.

16
The four electrons shared by the oxygen atoms
form a double bond.
The double bond is represented by two single
lines. Each line in the Lewis dot structure
represents two electrons
17
The element hydrogen is an exception to the Octet
Rule. It only needs two electrons, rather than
eight, to be stable.
The hydrogen atom has one valence electron. It
requires one more electron to be stable. The
fluorine atom has seven valence electrons. It
requires one more to satisfy the Octet rule.
18
The hydrogen atom now has a total of two
electrons around it and is stable. The fluorine
atom now has a total of eight electrons around it
and is stable.
19
  • The Lewis dot structure of the HF molecule shows
    a line and 6 dots to represent the bonding pair
    and the 3 lone pairs of electrons, respectively.

20
Rules for writing Lewis Dot structures
  • Rule 1
  • Add together the number of valence electrons for
    each atom
  • in the molecule. For example, CF4
  • Carbon has four valence electrons and each
    fluorine has
  • seven valence electrons 4 4(7)
  • 32

21
Rule 2 Write out the elements of the molecule so
that the least electronegative elements is in
the center surrounded by the other elements. For
example, CF4
22
Rule 3 Place a covalent bond between the central
atom and the outside atoms. Remember each
covalent bond contains two electrons.
23
The four covalent bonds use eight of the 32
valence electrons in CF4
  • This uses 24 electrons. There Are no electrons
    left, so this is The Lewis dot structure for CF4
  • Rule 4
  • There are 24 valence electrons remaining. Add
    electrons to
  • the outer atoms as lose pairs to satisfy the
    Octet Rule.

24
Rule 5 for example, NH3
  • First apply Rules 1-4 to the molecule
  • Rule 1 Count the valence electrons
  • Rule 2 Place the least electronegative element
    at the centre, except for H which is always an
    outer atom
  • Rule 3 Add covalent bonds between the centre
    atom and the outer atoms
  • Rule 4 Add lone pairs to the outer atoms
  • Rule 5 Add lone pairs to the centre atom

25
Rule 1 Nitrogen has 5 valence electrons and each
hydrogen has 1 valence electron The total number
of valence electrons 5 3 (1) 8 Rule
2 Hydrogen is always an outer atom and is never
at the centre of a molecule
26
Rule 3 Add the bonding electrons. This uses 6 of
the 8 valence electrons. Rule 4 The 2
remaining valence electrons are not added to the
outer atoms, because each H has its maximum of 2
valence electrons.
27
  • Rule 3
  • Add the bonding electrons. This uses 6 of the 8
  • valence electrons.
  • Rule 4
  • The 2 remaining valence electrons are not added
    to
  • the outer atoms, because each H has its maximum
  • of 2 valence electrons.

28
Rule 5
  • Place the remaining 2 Valence
  • electrons on the central
  • nitrogen atom

This is the Lewis structure For NH3
  • Rule 6
  • Check all atoms in the molecule to ensure that
    each has 8
  • electrons(2 for hydrogen). If an atom has fewer
    than 8
  • electrons, create double or triple bonds. (Note
    Double
  • bonds only exist between C,N,O and S atoms)

29
Apply rule 6 to the following CH4, CF4,
  • Hydrogen 1 bond 2 electrons (stable)
  • Carbon 4 bonds 8 electrons (stable)
  • Fluorine 1 bond 3 lone pairs 2 3 (2)
  • 8 electrons (stable)
  • Carbon 4 bonds 8 electrons (stable)

30
Example CH2O
  • Apply Rules 1-5 to the molecule
  • Rule 1 Count the valency electrons
  • Rule 2 Place the least electronegative element
    at the centre, except for H, which
    is always an outer atom
  • Rule 3 Add covalent bonds between the centre and
    the outer atoms
  • Rule 4 Add lone pairs to the outer atoms
  • Rule 5 Add lone pairs to the centre atom

31
Rule 1 Carbon has 4 valence electrons, each
hydrogen has 1 valence electron, and oxygen has 6
valence electrons. Total number of valence
electrons 4 2(1) 6 12 Rule 2 Carbon is at
the centre of the molecule because it is less
electronegative than oxygen. Hydrogen is always
an outer atom and is never at the centre of the
molecule.
32
  • Rule 3
  • Add the bonding electrons.
  • This uses 6 of the 12 valence
  • electrons
  • Rule 4
  • Add the remaining 6 lectrons to
  • the outer atom. Hydrogen does
  • not need any more electrons, but
  • Oxygen needs 6 to complete its
  • octet.

33
Rule 5 There are no valence electrons left to add
to the centre
  • Rule 6
  • Oxygen shares one of its lone pairs with C and O
    and give the desired 8 electron total

This is the Lewis dot Structure for CH2O
34
Exceptions to the Octet Rule
The Octet Rule applies to Groups IVA through VIIA
in the second row of the Periodic Table, but
there are exceptions to the rule among some other
elements. The following two cases are an
example Example BF3 Rule 1 Boron has 3 valence
electrons and each Fluorine has 7 valence
electrons Total number of electrons 3 3 (7)
24
35
  • Rule 2
  • Boron is at the centre of
  • the molecule because it is
  • less electronegative than
  • fluorine
  • Rule 3
  • Add the bonding electrons.
  • This uses 6 of the 24 valence
  • electrons

36
  • Rule 4
  • Add the remaining electrons
  • to the outer atoms. Each
  • Fluorine has the required 8
  • electrons
  • Rule 5
  • This uses the remaining
  • electrons leaving none to add
  • to the Boron central atom

37
  • Rule 6
  • Check the number of electrons around each atom.
    Each
  • Fluorine atom has 8 electrons, but the Boron Atom
    has only
  • 6. This is an exception to the Octet Rule. A BF
    bond is not
  • an option, because double bonds exist only
    between C,N,O,
  • and S atoms

This is the Lewis dot structure BF3
38
Example PF5
  • Rule 1
  • Phosphorus has 5 valence
  • electrons and each fluorine
  • has 7 valence electrons
  • Total number of electrons
  • 5 5(7) 40
  • Rule 2
  • Phosporus is at the centre
  • because it is less
  • electronegative than fluorine

39
  • Rule 3
  • Add the bonding electrons. This uses 6 of the 24
    valence
  • electrons.
  • Rule 4
  • Add the remaining electrons to the outer atoms.
    Each Fluorine requires 6 more electrons

40
  • Rule 5
  • This uses the remaining
  • electrons leaving none to
  • the central P atom
  • Rule 6
  • Check the number of electrons
  • around each atom. Each
  • Fluorine atom has 8 electrons,
  • but the phoshorus atom has 10.
  • This is an exception to the
  • Octet Rule.

41
Rule 6 Check the number of electrons around each
atom. Each fluorine atom has 8 electrons, but the
phoshorus atom has 10 . This is an exception to
the Octet Rule.
42
How Elements Form Compounds
Some atoms lose or gain electrons to become
stable charged particles called ions
When atoms loses electrons, they form positively
charged ions called cations
When atoms gain electrons, they form negatively
charged ions called anions.
43
Sodium chloride is a relatively harmless compound
because the sodium and chlorine atoms have stable
ions .
The compound formed is called an ionic compound
because it is made up of positive and negative
ions that have resulted from the transfer of from
a metal to a nonmetal.
The positive and negative ions are attracted to
each other because they have opposite charges.
44
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45
When ionic compounds are placed in water, the
ions separate and are surrounded by water
molecules. They are electrolytes. They are also
conductive
46
  • Ionic Charges and Chemical Families
  • Review
  • structure of the atom
  • How some atoms can form stable ions by gaining or
    losing electrons
  • You have also learned that the PT is a useful
    organizing tool for predicting behaviour of
    substances

47
The location of the alkalis metals (dark green),
the alkaline earth metals (light green), and the
halogens (red) in the PT
Activity
48
Ionic Compounds
There are over 100 elements in the PT Thousands
of different compounds are formed when these
elements combine. How can we name these
compounds? How can we write formulas to represent
them?
49
We have seen from past discussions that The PT
and a knowledge of the electronic structure could
be used to predict ionic charge of elements
Ionic charges (or valences) of some elements in
the PT
50
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51
  • Naming Ionic Compounds
  • The name of the metal first, followed by name of
    the of the nonmetal.
  • The ending of the name of the nonmetal changes
    and ends with ide

52
Names and Ionic Charges of some nonmetals
Name of element Symbol Ionic Charge Name in compound
Fluorine F 1- fluoride
Chlorine Cl 1- chloride
Bromine Br 1- bromide
Iodine I 1- iodide
Oxygen O 2- oxide
Sulfur S 2- sulfide
Nitrogen N 3- nitride
Phosphorus P 3- phosphide
53
Names and Formulas for Atoms with More Than One
Ionic Charge Some metals are able to form more
than one kind of ion.
For example, the element copper forms two
completely different compounds when it reacts
with chlorine One of the compound is white the
other is yellow
54
Ionic charge on the copper in the white compound
is 1 . Its chemical formula is CuCl The ionic
charge on the copper in the yellow compound is
2, its formula is CuCl2
Name of element Symbol Ionic charges Roman numeral
copper Cu 1,2 I,II
iron Fe 2,3 II,III
lead Pb 2,4 II,IV
tin Sn 2,4 II,IV
55
  • We have come across compounds such as
  • Calcium carbonate
  • Sodium bicarbonate
  • Calcium hydroxide, and copper sulfate
  • These names do not fit the naming so far

What are these compounds?
56
Polyatomic ions
  • They are pure substances
  • Involve combinations of metals with polyatomic
    ions
  • Groups of atoms that tend to stay together and
    carry an overall ionic charge

57
When a compound containing this ion is dissolved
in water, the positive metal ion and the nitrate
ion separate from each other but the nitrate ion
itself stays together as a unit surrounded by
water molecules
An example is The nitrate ion
58
Writing Formulas for Polyatomic Compounds The
ionic charges of polyatomic ions makes it
possible for them to form ionic compounds
Common Polyatomic ions and Their Ionic charges
Name of polyatomic ion Ion formula Ionic charge
nitrate NO3- 1-
hydroxide OH- 1-
bicarbonate HCO3- 1-
chlorate ClO3- 1-
carbonate CO32- 2-
sulfate SO42- 2-
phosphate PO43- 3-
59
When a polyatomic ion such as nitrate or sulfate
combines with other elements We follow the same
rules for writing formulas
60
What is the formula for the ionc compound formed
by sodium and a sulfate ion? Rule 1 write the
symbols of the metal and of the polyatomic
group Na SO4
Rule 2 write the ionic charges 1 2- Na SO4
61
Crisscross rule crisscross the ionic
charges 1 2- Na SO4
Note that polyatomic ions do not reduce .
Formula cannot be simplified Na1SO2 because SO4
is a group
The formula is Na2SO4
62
Try this what is the formula of lead(IV)
carbonate?
63
There are many types of polyatomic ions, but one
special group is known as the Oxyacids
Oxyacids are compounds formed when hydrogen
combines with polyatomic ions that contain
oxygen. Ionic charge for hydrogen in these
compounds is 1
Ion name Ion formula Ionic charge Oxyacid formula Oxyacid name
nitrate NO3- 1- HNO3 Nitric acid
nitrite NO2- 1- HNO2 Nitrous acid
chlorate ClO3- 1- HClO3 Chloric acid
carbonate CO32- 2- H2CO3 Carbonic acid
sulfate SO42- 2- H2SO4 Sulfuric acid
sulfite SO32- 2- H2SO3 Sulfurous acid
phosphate PO43- 3- H3PO4 Phosphoric acid
64
Molecular Compounds Imagine that you find an
unlabelled container of solid white crystals in
the kitchen. You are sure the crystals are either
salt or sugar A simple taste test will tell you
what the crystals are. But imagine you find the
same crystals in the lab. A taste is too
dangerous. What do you do? Dissolve the crystals
in water and test for conductivity. If it
conducts electricity, the compound must contain
ions Salt or sodium chloride is an ionic compound
65
In ionic compounds, metals with 1, 2, or 3
electrons in their outer shell lose electrons to
nonmetals, which often have 5, 6, or 7 electrons
in their outer shell.
If the solution does not conduct electricity, it
must be a different kind of compound
66
Most compounds you encounter every day do not
contain ions. Rather, they contain neutral groups
of atoms called molecules.
Sugar is a molecular compound. It is made up of
molecules in which nonmetal atoms, such as
hydrogen and oxygen share electrons to form
stable arrangements.
67
Water and carbon dioxide are also molecular
compounds, whether in in a gas, a liquid, or a
solid state, the particles in ionic and molecular
compounds are different as shown
Salt is an example of an ionic compound made up
of ions of opposite charge. Ice (H2O) is an
example of a molecular compound made up of
neutral molecules
68
Hydrogen gas is a molecule formed when two
hydrogen atoms combine. Each hydrogen atom has
one electron.
For the two hydrogen atoms to become stable, both
must gain an electron.
They do this by sharing a pair of electrons, one
from each atom
69
The result is a covalent bond--- a shared pair of
electrons held between two nonmetal atoms that
holds the atoms together in a molecule.
Many nonmetals form molecules in this way. For
example chlorine gas is a molecule that consists
of two chlorine atoms held together with a
covalent bond. Each chlorine atom has 7 electrons
in its outer orbit and needs to gain electron to
be stable
70
Many nonmetallic elements exist as covalently
bonded molecules. Table below lists elements that
form diatomic molecules.
Name of element Chemical symbol Formula and state at RT
Hydrogen H H2 (gas)
Oxygen O O2(gas)
Nitrogen N N2(gas)
Fluorine F F2(gas)
Chlorine Cl Cl2(gas)
Bromine Br Br2(liquid)
Iodine I I2(solid)
71
Molecular compounds are all around us a bottle of
soda contains water molecules, sucrose, glucose,
or fructose
72
Writing formulas for Molecular Compounds
Formulas can be written using a method similar to
the one used for ionic compounds. The number of
electrons that metals and nonmetals transfer to
become stable ions can be a clue to the formula
of an ionic compound. Similarly, the number of
electrons that a nonmetal needs to share to
become stable is a clue to the number of covalent
bonds it can form
The combining capacity of a nonmetal is a measure
of the number of covalent bonds that it will need
to form a stable molecule
73
Table 1 Combinig Capacities of Nonmetal Atoms
4 3 2 1
H
C N O F
Si P S Cl
As Se Br
I
74
Carbon has four electrons in its outer(valence)
orbit. If it lost 4 electrons, it would form a
positive ion. If it gained 4 electrons, it would
have the electron arrangement of neon and would
form a negative ion
It turns out that carbon cannot form either ion.
Instead it gains 4 electrons by sharing carbon
has a combining capacity of 4.
For example, when carbon shares one of its outer
orbit electrons with each of four different
hydrogen atoms, as shown in figure, the result is
methane CH4, the major component of natural gas
75
As a result of forming covalent bonds through
sharing electrons, the atoms end up with a stable
arrangement in their orbit similar to that of a
noble gas.
You can use the combining capacity to write the
formulas of molecular compound s without having
to consider the electronic structure
76
How would you write the formula for a compound
formed between Carbon and Sulfur?
Rule 1 Write the symbols, with the left hand
element from Table 1 with the combining capacities
  • 2
  • C S

Rule 2 Crisscross the combining capacities to
produce subscripts
  • 2
  • C S

The formula is C2S4
77
Rule 3 Reduce the subscripts if possible The
formula C2S4 is reduced to C1S2
Rule 4 Any 1 subscript is not needed.
The correct formula is CS2
78
Naming Molecular Compounds
Many molecular compounds have simple names. The
compound H2S is called hydrogen sulfide, much as
if it is ionic. Other molecular compounds have
names that are very familiar to us even though
they do not follow a system
Common names have been used for centuries for
water (H2O) ammonia (NH3), hydrogen peroxide
(H2O2) and methane (CH4)
79
The names of molecular compounds often contain
prefixes. These prefixes are used to count the
number of atoms when the same two elements form
different combinations.
For example , the gas that you exhale is carbon
dioxide (CO2) while the poisonous combination of
carbon and oxygen that can be formed in
automobiles is carbon monoxide
The prefixes di and mono differentiate
between the two molecules
80
Table 2 Prefixes in Molecular Compounds
Prefix Number Example (formula)
mon(o) 1 carbon monoxide(CO)
di 2 carbon disulfide (CS2)
tri 3 sulfur trioxide (SO3)
tetra 4 carbon tetrafluoride (CF4)
pent(a) 5 Phosphorus pentabromide (PBr5)
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