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Chapter 8:   Covalent Bonding Chapter 8:   Covalent Bonding

Chapter 8: Covalent Bonding - PowerPoint Presentation

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Chapter 8: Covalent Bonding - PPT Presentation

Matter takes many forms in nature In this chapter we are going to learn to distinguish the type of compound that we have already studied the ionic compound which contains oppositelycharged particles metal ID: 700413

atoms electrons bond atom electrons atoms atom bond bonds molecular molecule polar covalent valence compound hydrogen molecules ionic single

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Slide1

Chapter 8: Covalent Bonding

Matter takes many forms in nature

: In this chapter, we are going to learn to distinguish the type of compound that we have already studied, the “

ionic

compound

” (which contains oppositely-charged particles: metal

cations

and non-metal anions), from a different type of compound – a “

molecular

compound

”. Additionally, we are going to focus on a type of molecular compound known as a

binary

molecular compoundSlide2

II. Binary compounds

:

A “binary” compound contains

atoms

from

two

different elements.

A.

NaCl

”, “CaF

2

”, and “Al

2

O

3

” (3 ionic compounds) are binary

ionic

compounds. “NH

4

Cl” is an ionic compound, but because it contains more than 2 elements, it is

not

a binary ionic compound.Slide3

B. “N

2

O

5

”, “SF

6

”, and H

2

O” (3 molecular compounds) are binary molecular compounds.

“C

6

H

12

O

6

” is a molecular compound, but because it contains more than 2 elements it is

not

a binary molecular compound.Slide4

Comparison/Contrast between an ionic compound and a molecular substance.

A. Molecular substances are made of

molecules

.

1. There is no “molecule” in an ionic compound.Slide5

B. A

“molecule” contains a specific

number

of atoms, connected in a specific manner, to give a specific shape.

If even one atom is “missing” or “different”, the molecule would be an entirely different substance. Slide6

1. Not so with an ionic compound: In an ionic compound there is a specific

ratio

of atoms.. In salt (

NaCl

) for example, there is a ratio of 1 Na for every 1 Cl.

If a clump of salt lost 1 Na and 1

Cl

, it would still be the same original substance:

NaClSlide7

C. The formula of a “molecule” should never be simplified. C

2

H

8

is not the same substance as CH

4

.

1. The formula of an ionic compound should always be simplified. Ba

2

O

2

is the same substance as BaO.

D. A molecule will

not

crack apart.

1. Ionic compounds can crack apart if hammered….. If the cations come close too close together and the anions come too close together, the structure cracks apart.Slide8

E. Molecular substances may have

low

melting points.

1. All ionic compounds have a very high melting point.

F. Molecular substances may, at room temperature, be found as

solids, liquids,

or

gases.

1. All ionic compounds are solids (at room temperature).Slide9

G. Molecular substances contain atoms which are held together by

covalent

bonds.

1. Ionic substances are held together by ionic bonds. Slide10

IV. Covalent bonding:

The type of bonding that occurs within a molecular substance, in which atoms

share

their valence electrons in order to become more stable.

A. Occurs between

atoms

of

nonmetallic

elements.Slide11

B. Not all “molecules” or “molecular substances” are compounds!

In addition to the binary molecular compounds that we will study, there are 7 nonmetallic elements found in nature (in their elemental form) as pairs of atoms. These are the 7 “

diatomic

” elements:

N

2

, O

2

, F

2

, Cl

2

, Br

2

, I

2

, H

2

.Slide12

V. An important review:

A. Metallic elements: Found to the

left

side of the staircase boundary on the periodic table.

Non-metallic elements

:

Elements found to the

right

side of the staircase boundary on the periodic table.

2.

Hydrogen

is a nonmetallic element also.Slide13

VI. The octet rule:

When a molecule is formed: “

Nonmetal atoms

share

electrons in covalent bonds in order to obtain a full

octet

of electrons.” An octet = 8 valence electrons.

A. Exception: A hydrogen atom will end up with a total of

two

electrons by sharing with 1 other atom. Slide14

B. There are a few other notable

exceptions

to the octet rule:

1. A few molecular compounds which contain an odd number of valence electrons are known to exist.

2. A few molecular compounds have either a boron or an aluminum atom with 6 valence electrons.Slide15

2. A

few molecular compounds have a central atom with 10 or 12 valence electrons.

(1) One common example is “

sulfur hexafluoride

”.

In this compound, the central sulfur atom contains 6 x 2 =

12

valence electrons. Be sure to remember that this compound is an example in which the

central

atom does

not

follow the octet rule

.Slide16

VII. Types of covalent bonds.

A.

Single

covalent bond

1 shared pair of valence electrons:

2 dots, or a single dash

, represent 2 electrons that are simultaneously being attracted by, or “shared” by, the nuclei of two neighboring atoms. Slide17

1. The

formula in the center is a type of structural formula called a

Lewis dot structural formula

”.

The formula on the right is the

molecular

formula.

H

H

H

H

H

2

Slide18

B.

Double

covalent bond

two pairs of shared valence electrons: 4 dots or 2 parallel dashes

.

C C C C C

2

H

4

H H

H H

H H

H HSlide19

C.

Triple

covalent

bond

three pairs of shared valence electrons: 6 dots or 3 parallel dashes

.

N N or N N N

2

Notice the

two

unshared

pairs

” of electrons (one pair is to the far left and one pair to the far

righ

)t of the nitrogen structure.

You may never use a long dash to represent an unshared pair of electrons

.Slide20

Unshared pairs of electrons don’t bond the atoms together….but, the repulsive forces of unshared pairs of electrons do dramatically influence the

shape

of a molecule! Slide21

D. Notice how an ion can react with a molecule to generate a polyatomic ion. In the example below, a hydrogen ion bonds to a molecule of ammonia(NH

3

) to make the

ammonium ion (NH

4

)

+

:

H

+

+ N H

H

N H

H

H

H

H

+

]

[Slide22

VIII. Drawing a Lewis Dot Structure:

A. Certain elements are known as “central” atoms…. They will be found in the center of a structure. The first element given in a formula is usually the central atom (exception: hydrogen and the halogens).

1. Position the central atom in the center of your work space.Slide23

B. Hydrogen

and the halogens are known as “peripheral” atoms. They will be found only connected to one other atom.

Position

hydrogen and halogen atoms so that they “touch”, or “go around” only 1 other atom.Slide24

C. Add up all the valence electrons. Position the valence electrons as dots around the atom they belong to - the valence electrons may never leave the original atom.

Position the dots to form a “doorway” with 4 sides, in which the symbol of the element appears centered in the doorway.

Start with no more than 2 dots on each side of the 4 sided doorway.Slide25

D. If you can’t easily achieve a Lewis dot structure which has each atom (other than hydrogen) surrounded by 8 dots by doing what is described above, then you either need a double bond (2 pairs of shared electrons) or a triple bond (3 pairs of shared electrons).

For CO

2

, you will need two double bonds.Slide26

1. To make a double bond, move

one

un

-shared electron”

simultaneously from each of two neighboring atoms

, and place those 2 electrons in between the two neighboring atoms.

2. To make a triple bond, start with a double bonded pair of atoms, and simultaneously move

one more

unshared electron

from each of the two

atoms. Reposition those two electrons in between the atoms.Slide27

Important points regarding nonmetal atoms and their bonding charcteristics:Slide28

Atoms of the Following Nonmetallic Elements

:

Have This Number of Electrons (dots)

when

in a Stable Structure:

Type of Bonds Permitted by these atoms:

Things to Remember about atoms of these elements; or to remember about a specified molecule.

Hydrogen

2

Single covalent

Diatomic element;

H

2

is a “linear” diatomic

molecule

Boron, Aluminum

8 or 6

Single covalent

When only 6 electrons surround a boron or an aluminum atom, the molecule’s shape will be “

trigonal

planar” (a flat pancake).

 Slide29

Atoms of the Following Nonmetallic Elements

:

Have This Number of Electrons (dots)

when

in a Stable Structure:

Type of Bonds Permitted by these atoms:

Things to Remember about atoms of these elements; or to remember about a specified molecule.

 

Sulfur

8 EXCEPT with

“SF

6

” when there are 12

Single, double, and/or triple covalent

In sulfur hexafluoride sulfur does NOT follow the octet rule. This is one “

exception”

to the octet rule.

The Halogen family: F, Cl, Br, I, At

8

Single covalent

All are diatomic elements; and,

F

2

, Cl

2

, Br

2

, I

2

, At

2

are all linear molecules, with only single bonds.Slide30

Atoms of the Following Nonmetallic Elements

:

Have This Number of Electrons (dots)

when

in a Stable Structure:

Type of Bonds Permitted by these atoms:

Things to Remember about atoms of these elements; or to remember about a specified molecule.

 

The Noble Gas family:

8….Except for helium (2). Noble gas atoms don’t form compounds.

Do NOT form compounds easily (no bonds).

Always found as single atoms in the gaseous state.

Nitrogen and oxygen

8

Single, double, and/or triple covalent

Diatomic elements; linear molecules. N

2

has one triple bond, while O

2

has 1 double bond.

All other nonmetal atoms

8

Single, double, and/or triple covalent

 Slide31

IX. Lewis Dot structural formulas for polyatomic ions:

A. Covalent bonds occur

within

a polyatomic ion (

not

between

polyatomic ions).

B. When

drawing polyatomic ions, place the

first

element in the

center

of the structure, and place the second element

around

the first element (placing 1 atom of the second element along each different

side

of the first element).Slide32

C. When the charge of a polyatomic ion is +, you need to

subtract

the indicated number of electrons from the total of the valence electrons in the molecule. So, for +1 ions: take away

1

electron from the molecular ion’s number of valence electrons.

D. When the charge of a polyatomic ion is –, you need to

add

the indicated number of electrons to the molecular ion’s number of valence electrons. Slide33

1

. If the charge is 1

-

, then add

1

more electron to the molecule’s total number of valence electrons.

2. If the charge is 2

-

, then add

2

more electrons; if the charge is 3

-

, then add

3

more electrons.

E. Last, for a polyatomic ion: Draw a large

bracket

around the ion; and, place its charge at upper right.Slide34

# of N valence electrons: 5 5

# of H valence electrons: 4 x 1 = 4 4

Charge of ion = +1, therefore less 1 -1

Therefore, total = 8

Ammonium ion

(NH

4

)

+

H N H

H

H

+

]

[Slide35

Steps for Dot Structures:

Step 1: total # valence electrons.

Step 2. Position central atoms:

carbon atoms form a straight line;

assume only single bonding.Slide36

Step 3. Position other atoms; remember “special” molecules.

A. Peripheral atoms:

Hydrogen and the halogens-

connect to only 1 other atom

,

use only 1 single bond

.

B. binary polyatomic ions: first element is central, second element is peripheral.

assume all single bonds.

C. hydrocarbons” – molecular formula gives list of atoms (from left to right) connecting to each central atom (usually carbon)

CH

3

CH

2

OH means “first carbon touches 3 H atoms, second carbon touches 2 H atoms, then there is an O touching an H.

Assume all

single bonds

.

D. Memorize

: CO

2

C in the middle;

use two double bondsSlide37

Step 4: Make each atom stable.

Work from left to right:

Assume all single bonds.

Position unshared pairs to provide octets.

Exception: Hydrogen atoms = only 2 dots.

Step 5: Count the dots you’ve used.

Make sure the # you used = the # you were supposed to. Erase extras.Slide38

Step 6: Make corrections -

If your structure “needs” 2 extra dots, it really needs a double bond….

Erase 2 unshared dots, and share them

(as part of a double bond).

If your structure “needs” 4 extra dots, it really needs a triple bond.....

Erase 4 unshared dots, and share them (as part of a triple bond).

Every atom should now be stable.Slide39

X. VSEPR Theory –

V

alence

S

hell

E

lectron

P

air

R

epulsion theory

. [Remember: Like charges repel!]

A. A theory to

predict

the 3-dimensional geometry,

ie

.

the“

shape

” of a moleculeSlide40

1. The theory is based on “electrostatic repulsion”: Molecules will adjust their

shape

to keep the negatively-charged pairs of

valence

electrons as

far apart

as possible from each other.

B. When NOT to use

VSEPR

theory:

When there are only 2 atoms in a molecule. These molecule’s shapes are called

linear

– it doesn’t matter if there are single bonds, double bonds, triple bonds, or unshared electron pairs.Slide41

C. Using VSEPR theory:

1. Draw the Lewis

dot structure

for the molecule.

2. Identify its

central

atom.

3. Identify the sets of valence electrons as one of two possibilities

:

A. Those

connecting

two atoms.

B. Those that do not connect two atoms

. These are

called “

unshared

pairs”.Slide42

4. The unshared pairs found on a central atom strongly

repel

each other; and molecules that would otherwise be linear, will be forced into a

bent

(or

angular

) shape

.

5. Unshared pairs also cause a molecule that would be shaped like a flat triangle

(

trigonal

planar

), to be forced into a

not flat

(

trigonal

pyramidal

) shape.Slide43

6.

Count

the number of connections

separately

from the number of unshared pairs

.

1 single bond counts as

1

connection.

1 double bond counts as

1

connection.

1 triple bond counts as

1

connection.

Each unshared

set of 2 dots

counts as

1

un-shared pair.Slide44

D Predicting Shapes Using VSEPR Table

Read horizontally across the table.

Connections To

the Central Atom

Unshared Pairs of Electrons Around Central Atom

Molecular Shape

Around Central Atom

2

0

Linear

3

0

Trigonal

Planar

4

0

Tetrahedral, 109.5

o

2

1 or 2

Bent

3

1

PyramidalSlide45

Shapes:

,

E.

Shap

e

s:

Linear diatomic

Linear triatomic

Trigonal Planar

Bent

Pyramidal

Linear diatomic

Tetrahedral

Linear triatomic

Trigonal planar

bent

Trigonal pyramidal

tetrahedralSlide46

“Molecular Polarity”

– A term that is used to distinguish two types of molecules…. Based on the presence or absence of a

separation

of

charge

.

Some

molecules

show

characteristics indicating that they have

oppositely

-charged ends (a positive end and a negative end). This is called a separation of the charges (or “separation of charge”).Slide47

Other molecules

show characteristics indicating that their structure doesn’t have a separation of charge, or their

structure

hides

the presence of their oppositely-charged ends.Slide48

How to determine a molecule’s polarity.

The first part of determining a molecule’s polarity is to calculate each individual bond’s polarity.

Be careful with the vocabulary being used –

An

individual

bond’s polarity is called the “

bond polarity”

The polarity of the entire molecule is called the “

molecular polarity”Slide49

To calculate a bond polarity, first identify the “

electronegativity

value” of each of the 2 atoms in the bond you are working on.

The

electronegativity value

is number (from 0 to 4) which informs us

of an atom’s ability to

attract

electrons when in a compound

. Slide50

The electronegativity value is given on your periodic table, side 2, within each element’s square…..upper right corner of the square, in black print.

The closer an element’s electronegativity is to “

4

, the better that an atom of that element will attract electrons when that atom is found in a compound

.

The closer an element’s electronegativity is to “

0

”, the

less

likely it is for that atom to be able to attract electrons when in a compound.Slide51

After identifying the two electronegativity values, subtract the two electronegativity values. Take the absolute value of the answer (make the answer positive). Slide52

After subtracting and taking the absolute value of the two electronegativity values,

then

think about your answer,

and determine whether your answer indicates that there is, or that there is not, a situation in which one of the two atoms in the bond “overpowers” the other atom in terms of electron attracting ability

. If an atom is able to overpower the other atom, it will “hog” the electrons, as opposed to sharing the electrons equally with the other atom. You will see (on the next page) that I’ve placed a “δ–”sign next to an oxygen atom, and a “δ+” sign next to 2 hydrogen atoms, in a sketch of a water molecule. By doing this, I am indicating that the oxygen atom is hogging the negatively-charged valence electrons belonging to the 2 hydrogen atoms; and, the 2 hydrogen atoms are both overpowered and “partially lose” their own valence electrons to the oxygen atom.Slide53

δ

-

indicates the “partial negative” atom; and,

δ

+ indicates the “partial positive” atom

. Slide54

Finally, you are now able to conclude that a bond is either “nonpolar covalent” or “polar covalent”.

A bond is to be called “

non-polar

covalent

” when

the two atoms share the electrons more-or-less

equally

.

Labeling a bond as nonpolar covalent means that

the difference in electronegativity values fell between 0 and 0.4

Labeling a bond as nonpolar covalent means that

the electrons are distributed practically equally along the bond

.

Labeling a bond as nonpolar covalent means

that there is no separation of charge

in that particular bond.Slide55

A bond is to be called

polar

covalent

” when the electrons are NOT distributed equally along the bond, rather,

the electrons are found much or most of the time toward the atom that has the higher

electronegsativity

value

.

This will occur when difference in electronegativity values is greater than 0.4, but less than 2.

We place a δ – next to an atom to identify it as the “partial negative” atom; and, we place a δ + next to an atom to identify it as the “partial positive” atom in a polar covalent bond.Slide56

A bond is designated as

ionic

when one atom has stripped the other of some of its valence electrons, yielding a

cation

and an anion. Then, the

cation

and anion are attracted, and more

cations

and anions are attracted and an ionic compound forms.

This should occur when the difference in electronegativity values is greater than 2.

We studied this type of compound previously. Ionic compounds have very different characteristics than compounds in which polar covalent and nonpolar covalent bonds exist.Slide57

The second part of determining the “molecular polarity” (a

molecule’s

polarity

)

is visualizing the effect of the individual bond polarities in conjunction with the

shape

of the molecule.

Some common occurrences are listed below:

If the molecule’s shape is either

tetrahedral

,

linear triatomic

or

trigonal

planar

and if

all

the bonds in the structure are

identical

, then the overall molecule’s polarity (the molecular polarity) is

nonpolar

Slide58

This is because, if every bond is nonpolar, there isn’t a substantial separation of charge to begin with.

And, if every bond is polar (as long as each bond

has the same 2 atoms

), the polarity of each bond will be

cancelled

due to the symmetry of the polar bonds around a central atom. When this symmetry exists, the

δ

– charge(s)

cancel

the

δ

+ charge(s) in each bond.Slide59

If the molecule’s shape is either

tetrahedral ,

trigonal

planar,

or

linear

, and the structure contains

non

-identical bonds in which 1 is polar, then the overall (molecular) polarity is polar (because 1 polar bond won’t have another bond to be cancelled with).

If the molecules’ shapes are pyramidal or bent, and one or more of the bonds is polar, the overall (molecular) polarity is “

polar

”. This is because in molecules having these shapes, the bond polarity is not situated symmetrically relative to the molecule’s center, so the polarity doesn’t cancel.Slide60

Polar bonds

 NonPolar molecule

Polar bonds  Polar molecule.

Polar bonds combine to cancel out; Polar bonds do

NOT

cancel out;

Producing

Non-polar molecules

: Producing

Polar Molecules

:Slide61

“Like dissolves like”:

An expression stating that

ionic

compounds and molecules whose molecular polarity is

polar

will be able to dissolve only in solvents containing

polar

molecules; and, conversely molecules whose molecular polarity is

nonpolar

will be able to dissolve only in solvents containing

nonpolar

molecules

.Slide62

Some characteristics of a water molecule

:

A

H

2

O

molecule has 2 identical polar bonds, a bent structure, and its molecular polarity is “polar”.

The partial negative (or slightly negative) region of a water molecule is the area closest to the

oxygen

atom.

The partial positive (or slightly positive) region of a water molecule lies within the area closest to the two

hydrogens

atoms.Slide63

Because of having the above characteristics, water behaves in an unusual manner:

When placed between a positively charged metal plate and a negatively charged metal plate, the water molecules all line up, with their

positively

-charged region attracted toward the negative plate and their

negatively-

charged region attracted toward the positive plate.Slide64

Inter

- molecular Bonds (or forces of attraction)

are DIFFERENT from the covalent bonds you have been studying so far!

The covalent bonds you have been studying so far (single, double, triple bonds) are bonds

within

a molecule; these would be called______

intra-molecular

bonds.

Inter

-molecular bonds (forces) are the attractions

between

2 molecules

.

They hold

separate

molecules together.Slide65

They are

weaker

than polar and non-polar covalent bonds (single/double/triple).

They are

weaker

than the ionic bonds which connect

cations

to anions.

Breaking an inter-molecular bond is a

physical

change; whereas, breaking an intra-molecular bond is a chemical change.Slide66

When you boil water (to generate water vapor), or melt ice (to generate liquid water), you are breaking inter-molecular bonds. When you treat a water molecule with electricity, you destroy the water molecule and generate from it oxygen gas and hydrogen gas….this is the breaking of intra-molecular bonds….this is a chemical change.Slide67

Types of intermolecular bonds:

Hydrogen

bond: Is the

strongest

intermolecular force; and, it is perhaps the most important intermolecular bond, as it is necessary for life as we know it.

Hydrogen bonds take effect when you have a combination of a few certain atoms in a

polar molecule

.

Slide68

The

polar

molecule must contain a

hydrogen

atom, and that hydrogen atom must be connected to one of the highly electronegative atoms listed below:

Fluorine, oxygen

,

OR

nitrogen

.

Ex: H

2

O, HF, NH

3

Slide69

Hydrogen bonds determine the properties of water and biological molecules (such as proteins).

Cause water to predominate as a

liquid

(rather than as a gas) on earth.

Cause ice to

expand

upon freezing (rather than contract as the kinetic molecular theory would predict).

Holds the

DNA

double helix structure together.Slide70

Time-permitting, we will also learn about one of the weaker types of intermolecular forces of attraction by doing a lab in which we use evaporation rates to identify between hydrogen bonding and a weaker force, present between nonpolar molecules.