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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

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lewis dot structures of covalent compounds
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

slide2
p+

p+

p+

p+

p+

p+

p+

e-

e-

e-

e-

e-

e-

e-

e-

e-

e-

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

slide3
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.

slide5
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

slide6
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 :

slide7
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

slide8
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
slide9
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

slide10
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.

slide11
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.

slide12
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.

multiple covalent bond
Multiple Covalent Bond

Some atoms have to share more than one electron in order

to satisfy the Octet Rule.

slide14
Each oxygen atom has six valence electrons. They each

require two more electrons to satisfy the Octet Rule.

slide15
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.
slide16
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

slide17
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.

slide18
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.

slide19
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.
rules for writing lewis dot structures
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

slide21
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

slide22
Rule 3

Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two

electrons.

the four covalent bonds use eight of the 32 valence electrons in cf 4
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.

rule 5 for example nh 3
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
slide25
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

slide26
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.

slide27
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.

rule 5
Rule 5

Placethe 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)

apply rule 6 to the following ch 4 cf 4
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)
example ch 2 o
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

slide31
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.

slide32
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.

slide33
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

slide34
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

slide35
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

slide36
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

slide37
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 B=F bond is not

an option, because double bonds exist only between C,N,O,

and S atoms

This is the Lewis

dot structure BF3

example pf 5
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

slide39
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

slide40
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.

slide41
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.

slide42
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.

slide43
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.

slide45
When ionic compounds are placed in water, the ions separate and are surrounded by water molecules. They are electrolytes.

They are also conductive

slide46
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
slide47
The location of the alkalis metals (dark green), the alkaline earth metals (light green), and the halogens (red) in the PT

Activity

slide48
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?

slide49
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

slide51
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”

slide53
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

slide54
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

slide55
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?

slide56
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
slide57
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

slide58
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

slide59
When a polyatomic ion such as nitrate or sulfate combines with other elements

We follow the same rules for writing formulas

slide60
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

slide61
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

slide63
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+

slide64
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

slide65
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

slide66
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.

slide67
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

slide68
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

slide69
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

slide70
Many nonmetallic elements exist as covalently bonded molecules. Table below lists elements that form diatomic molecules.
slide71
Molecular compounds are all around us a bottle of soda contains water molecules, sucrose, glucose, or fructose
slide72
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

slide74
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

slide75
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

slide76
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

slide77
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

slide78
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)

slide79
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

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