1

Chemistry 2202

Unit II: Chemical Bonding Name:________________

Questions to Consider :

- What holds matter together?

- What type of bonding is found in a given chemical compound?

- How does the type of bonding affect the physical and chemical properties of

substances?

- Why does a particular compound have the properties it does?

Chemical bonds are attractive forces that hold all substances together.

There are two types of forces that hold substances together to form compounds. They are;

1. Intramolecular forces:

- Bonding between atoms or ions.

- Forces within molecules, or formula units, of a compound.

- forces that hole atoms to each other

- atoms within H2O molecule are held together by intramolecular bonding

2. Intermolecular Forces

- Bonding between molecules

- Evidence of this type of bonding is available via observations of surface

tension in water, changes in states of materials under specific conditions,

heat vapourization, etc.

- H2O molecule are held to each other by intermolecular bonding

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1. Intramolecular Forces

We consider four types of intramolecular forces that exist in our environment. These

forces are vital in holding molecules together, and in holding formula units that contain

atoms and ions together. The relative strengths of these intramolecular forces are listed in

the following table

Bond Strength

The strength of Intramolecular bonds, within molecules, depends on the type of bonding that

exists.

Type of Bonding and Relative Strength

Bond Type Example

Strongest Bonds

Network Covalent - diamond and graphite

Ionic

Metallic

- brass, for example, is a

combination of CuZn

Weakest Bonds

Covalent

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Ionic Bonding: - Second strongest type of Intramolecular force.

- It is bonding that usually occurs between positively charged

metallic ions (cations) and negatively charged non-metallic ions,

or complex ions (anions).

- ex.

Compound Formula Compound Name

NaCl,

MgCO3,

AlF3

Recall: - An ION is an atom that lost (+) or gained (-) electrons

- An ionic bond is formed as a result of a force of attraction between

oppositely charged ions.

Chloride Atom (to form an ion will gain one electron) Chloride

Recall: You know that the electrons in atoms exist in specific energy levels. Also, each

energy level has a limit to the number of electrons that can occupy it. Only two

electrons are allowed to occupy the lowest energy level while eight electrons can

exist in the second, and third, energy levels (energy levels are also referred to as

shells). When the maximum number of electrons is in an energy level, it is said to

be a filled energy level or “closed shell.”

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Electron Transfer:

Note 1: It is only the electrons in the outermost energy level, or shell, that participate in

bonding. These outer electrons are called valence electrons.

Note 2: An Ionic Bond is the attraction between cations and anions. An Ionic Bond is the

force of attraction between cations and anions that have transferred electrons.

Note 3: There are no distinct pairs of sodium and chloride ions that you could identify as

a molecule. A crystal of solid sodium chloride with a mass of about 1 mg contain

about 1 x 1021

sodium ions and an exactly equal number of chloride ions.

In a crystal, the ions pack together in a way that allows the positive and negative

charges to be positioned as close together as possible. The 3-D array of

alternating positive and negative ions is called a crystal lattice.

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Crystal Lattice: Ionic compounds exist as a crystal lattice structure. The attraction between

(+) and (-) forces results in a 3-D, definite, geometric, repeating pattern of

alternating cations and anions.

Note: Each sodium ion is equally attracted to six different chloride ions and each chloride ion

is equally attracted to six different sodium ions. There are no distinct pairs of sodium

and chloride ions that you could identify as a molecule. When you write NaCl, you are

simply stating the ratio of sodium to chloride ions is one to one.

Example: CaF2 Ca2+

and F- ions will pack together in a way that allows the

positive and negative ions to be as close together as possible.

Empirical Formulas: The chemical formulas that you write for ionic compounds are

called empirical formulas. These formulas represent the whole

number ratio between the positive and negative ions in the smallest

neutral unit of an ionic crystal lattice—the formula unit.

Formula Unit: Iionic compounds exist simply as ratios of one atom to another

atom. There aren’t any distinct ionic entities that exist in nature.

The formula unit of an ionic compound simply represents the

atoms present in a compound, and the lowest relative quantities of

each.

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Properties of Ionic Compounds:

State: Ionic compounds are always SOLID at STP because of strong attraction

between ions. High melting and boiling points.

Ionic solids are Brittle. When bent, ‘like’ charges can become aligned and

repel each other. The crystal then breaks along the line where the charges

are repelling giving these substances a brittle character.

Conductors: Ionic Compounds conduct electricity when dissolved in water (aq) or in

molten state (l). When dissolved in water, the ions are surrounded by

water molecules but they still carry a net charge. The ions easily move

through water toward oppositely charged electrodes.

They do not conduct electricity in solid form because the valence electrons

are held within the individual ions in the lattice and these charged ions

are NOT free to move about.

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Metallic Bonding: - Third strongest type of bonding.

- It is bonding between metals such as Fe, Na, Mg, etc.

- Metallic cations attract valence electrons from other metallic

cations.

- Because the metallic atoms tend to have low attractions for

Valence electrons they are free to move from one valence orbital to

another in a “sea of delocalized electrons.” The resulting force of

attraction between these delocalized electrons and the atom nuclei

keeps these metallic substances together.

Examples: Metallic alloys (ex. Steel) and CuZn (brass)

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Properties of Metals:

Ductile and Malleable: The positive metal exist in fixed arrays, or layers (LIKE

SOLDIERS LINED UP FOR INSPECTION). When stress

is applied to a metal, such as hammering, one layer of

positive nuclei can slide across another layer. The layers

move without breaking the array because the delocalized,

freely moving valence electrons, continue to exert a

uniform attraction on the positive ions. For this reason,

metals do NOT shatter immediately along a clearly defined

point of stress but bend to a new shape.

Conduct Electricity The ability to conduct electrical current requires that

negative charges move freely and independently of each

other. Valence electrons can move freely throughout the

metallic structure carrying electric charge from one place

to another.

Conduct Heat Free moving electrons will collide with particles of

adjacent hot objects and receive kinetic energy in

collisions. The electrons then move freely throughout the

metal and pass the kinetic energy on to other particles by

colliding with them.

Shiny Metals are shiny because of the very strong absorption of

light by the delocalised bonding electrons. When light falls

on a metal it is almost totally absorbed since the bonding

electrons can jump up to a broad band of energy levels

allowing energy changes corresponding to the full range of

frequencies in the visible region of the spectrum.

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Covalent (Molecular) Bonding: - Fourth strongest type of bonding.

- The sharing of electrons between two or more non-

metals produces covalent bonds. All covalent

compounds share electrons to fill their valence orbit,

leading a stable electron configuration. If two non-

metals bond, both atoms share electrons.

Properties of Molecular Compounds:

- Solid, Liquid or Gas at Room Temperature. When molecular substances melt or

boil, the intramolecular covalent bonds between the atoms do not break, but

instead, intermolecular bonds between the particles must break.

- Low melting/boiling points

- Non-Electrolytes because the electrons are all localized within the molecules.

- Soluble or Insoluble in Water

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Lewis Diagrams: Converting Energy Level(Bohr) Diagrams into Lewis Dot Diagrams

A convenient method of tracking electrons in covalent bonds is to use Lewis Diagrams.

According to the theory of quantum mechanics, an orbital is defined as a region in space

in which an electron with a given energy is likely to be found. Electrons are thought to

occupy a region in space in somewhat the same way that clouds occupy regions of the

atmosphere. Since chemical reactions are thought to involve only the outer or valence

electrons, the discussion of orbitals is usually restricted to only the valence orbitals or an

atom.

Orbital Number of Electrons

in the Orbital

Description of

Electrons

Type of Electrons

Empty 0

Half-filled 1 Unpaired Bonding

Filled 2 Lone pair Non-bonding

According to the theory of quantum mechanics, the number and occupancy of valence

orbitals in the representative elements are determined by the following theoretical rules;

1 There are four valence orbitals in the valence level of atoms

representative elements. Hydrogen, which has a very simple structure, is

an exception. It only has one valence orbital.

2 An orbital may contain 0, 1, or 2 electrons.

3 Electrons occupy any empty valence orbitals before forming electron

pairs.

4 A maximum of eight (8) electrons can occupy orbitals in the valence level

of an atom. This is known as the octet rule:

Exception: hydrogen can gain, lose, or share 1 e-.

Valence Electrons: For representative elements, number of valence e- is equal to the

last digit of the group #

Group Number Number of Valence Electrons

Group 1 1 valence e-

Group 2 2 e-

Group 13 3 e-

Group 14 4 e-

Group 15 5 e-

Group 16 6 e-

Group 17 7 e-

Group 18 8 e- stable structure – valence shell filled

Ca (2), Li (1), F (7), P (5), Ar (18), C (4), Al (3), Se (6)

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Lewis Diagrams for Atoms

- Created by American chemist G.N. Lewis in 1916.

- A convenient way of tracking electrons in covalent bonds

Structure: Element symbol: to represent nucleus and inner filled orbits.

Valence electrons: electrons represented by dots above, below, left,

and right of element symbol. Put one electron in

each orbital at a time, starting at top.

Drawing Lewis(Electron Dot) Diagrams for Atoms:

- Write the element symbol to represent the nucleus and any filled energy levels of

the atom.

- Use a dot to represent each valence electron.

- Start by placing a single valence electron into each of the four valence orbitals

(represented by the four sides of the element symbol).

- If additional locations are required for electrons, start filling the four orbitals

with a second electron until up to eight positions for valence electrons have been

occupied.

Lone Pair: a pair of electrons that exist in a valence orbital that do not

take part in bonding with other electrons.

Bonding Electron: a single electron that does bond with other electrons.

Bonding Capacity: the capacity for an atom to bond with other atoms. Equal

to the number of bonding electrons. Ex. O has 2 bonding

electrons, therefore has a bonding capacity of two

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Example Lewis Dot Diagrams

Questions

1, For atoms of each of the following elements, determine the number of lone pairs and

bonding electrons.

Element Lone Pairs Bonding Electrons

Sn

Ca

Cl

Al

I

Ar

H

Ba

Br

C

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Bonding Electrons and Lone Pairs

Element # Bonding Electrons # Lone Pairs

Carbon 4 bonding electrons 0 lone pairs

Nitrogen 3 bonging electrons 1 lone pair

Sulfur 2 bonding electrons 2 lone pairs

Iodine 1 bonding electron 3 lone pairs

Complete the following chart by filling in the blank spaces.

Element Group # valence

electrons

Lewis Diagrams # Bonding

Electrons

# Lone Pairs

C 14

4

4 0

O 16

6

2 2

Si 14

4

4 0

Ne 18

8

0 4

Your Turn:

Element Group # valence

electrons

Lewis Diagrams # Bonding

Electrons

# Lone Pairs

S 16 6

2 2

Cl 17 7

1 3

P 15 5

3 1

I 17 7

1 3

Note: It is only atoms of the group VIIIA, or “noble gases, that are never found in nature

bonded to any other atoms. The outer energy level of all the noble gases is filled. Ex. He

2 e-. All others outer energy level has eight electrons. A filled outer energy level

makes atoms very stable and creates NO tendency to form bonds with other atoms.

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Chemistry 2202

Unit 2: Lewis Diagrams and Atoms Name:__________________ (lewis and atoms)

Lewis dot diagrams provide a convenient way for keeping track of the origin, and

distribution, of valence electrons that are a part of bonding in molecular compounds.

These diagrams were first introduced by G.N. Lewis in 1916 as a means to better

understand the processes taking place in covalently bonded compounds.

Complete the chart below by filling in the blank spaces.

Atom

Group

# Valence

Electrons

# Valence

Orbitals

# of

Lone

Pairs

# Bonding

Electrons

Lewis diagram

of Atom

H

1 A

1

1

0

1

H

C

N

O

VI A

6

4

2

2

O

F

S

Si

P

Cl

Br

Se

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Lewis Diagrams for Molecules:

Octet Rule: Atoms of all elements other than the noble gases form bonds in a way that

will create a noble gas configuration—a filled outer energy level. This

tendency is the basis of the Octet Rule

When bonds form, atoms share electrons in a way that creates an octet, or

filled outer energy level for atoms involved in bonding (exceptions are H

and He; filled with 2e-).

Hydrogen is unique because it can gain or lose only one electron. If hydrogen loses an

electron, it has no electrons. Positively ionized hydrogen is simply a proton. If hydrogen

gains an electron, it has two electrons in its filled outer energy level. The negatively

charged hydrogen ion is called a hydride ion.

Lewis Diagrams for Molecules:

Lewis diagrams allow for the visual representation of lone pairs and bonding electrons

between bonded atoms. To draw Lewis Diagrams for atoms use the following criteria.

Consistent with the octet rule, atoms share a number of electrons to attain eight valence

electrons around each atom, except for hydrogen.

Drawing Lewis Diagrams for Molecules:

- Draw Lewis diagram for each atom.

- Central atom is one with most bonding electrons

- Connect all other atoms using Single bonds

- Create double or triple bonds with any extra bonding electrons

- Check diagram by counting to see if all valence shells are FULL (8e-)

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Chemistry 2202

Unit 2: Lewis Diagrams and Atoms Name:__________________

Practice problems:

For each of the following molecule compounds, draw the Lewis (electron dot) diagram

for it.

Compound Name Lewis Diagram

a. hydrogen bromide

··

H : Br :

··

b. sulfur dibromide

c. carbon tetrachloride

d. phosphorus trichloride

·· ·· ··

: Cl : P : Cl :

·· ·· ··

: Cl :

··

e. dihydrogen sulfide

··

H : S : H

··

f. carbon dioxide

g. carbon tetraiodide

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Multiple Bonds:

When nonmetal atoms have less than 7 electrons in their valence level, they have more

than one unpaired electron available to form covalent bonds. The result is the sharing of

more than one pair of electrons with another atom in order to achieve an octet.

Example 1: Oxygen (O2) - six valence electrons

Bonding electrons _______

Lone Pairs _______

Note: To achieve a stable octet, with another oxygen atom, multiple sharing and

multiple bonds will be created between the two

atoms.

Example 2: Nitrogen (N2) - five valence electrons

Bonding electrons _______

Lone Pairs _______

Note: Atoms of nitrogen have one lone pair and three bonding electrons. To

achieve a stable octet, with another nitrogen atom, multiple sharing and

multiple bonds will be created between the two atoms.

Example 3: Carbon Dioxide CO2

Note: Atoms of unlike elements can also form double and triple bonds

with each other.

Single Covalent Bond: One pair of shared electrons is considered to be one bond

and it therefore called a “Single Covalent Bond.”

Double Covalent Bond: When atoms share four electrons, they are sharing two

pairs of electrons and this combination is called a “Double

Covalent Bond.”

Triple Covalent Bond: Atoms sharing three pairs of electrons

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Structural Formulas:

Structural formulas are used to demonstrate how atoms are connected to each

other to form a molecule. To draw a structural diagram simply ignore the lone

pairs and bonding electrons an use a single dash(line) to represented the bond

between each atom.

For simplification, use one single line to represent a single bond ( — ), two lines

for a double bond ( = ), and three lines for a triple bond ( ). Lone pairs are

usually omitted from the structures. However, they are still important to

remember when you consider the overall shape of the molecules.

Lewis Diagram Structural Diagram

Example #1 - Methane CH4

Example #2 - Ammonia NH3

Example #3 - Water H2O

Example #4 - Carbon tetrachloride CCl4

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Examples:

SiBr4

PI3 SiS2 HPO

Lewis Diagram

Lewis Diagram

Lewis Diagram

Lewis Diagram

Structural

Diagram

Structural Diagram

Structural

Diagram

Structural

Diagram

CBr4

PCl3 C2H4 HPS

Lewis Diagram

Lewis Diagram

Lewis Diagram

Lewis Diagram

Structural

Diagram

Structural Diagram

Structural

Diagram

Structural

Diagram

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Chemistry 2202

Bonding: Lewis and Structural Diagrams (lewis and structural_2) Name:__________________

Name

Molecular

Formula

Lewis Diagram

of Atoms

Lewis Diagram of

Molecule

Structural Diagram

H2(g)

Chlorine

HCl(g)

Methane

NI3(g)

Oxygen

difluoride

N2H4(l)

Hydrogen

peroxide

Ethane

C2H6(g)

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Chemistry 2202

Unit 2: Lewis and Structural Diagrams Name:__________________ (lewis and structural_1)

Molecular

Compound

Lewis Diagram of

Atoms

Lewis Diagram of

Molecule

Structural Diagram

of Molecule

N2

CH2O

C2H4

SCl2

PH3

CI4

FCl

CO2

C2H2

SiBr4

PI3

HPO

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Valence-Shell Electron-Pair Repulsion (VSEPR) Theory:

VSEPR - is the study of the three-dimensional shapes of molecules.

- can be used to predict molecular geometry,

- can be determined from Lewis-dot diagrams of molecules

- incorporates the arrangement of electron pairs, and bonded pairs, around

each central atom.

Shape Diagrams:

One of the more obviously important properties of any molecule is its shape. Clearly it is

very important to know the shape of a molecule if one is to understand its reactions. It is

also desirable to have a simple method to predict the geometries of compounds. For

main group compounds, the VSEPR method is such a predictive tool and unsurpassed as

a handy predictive method.

It is a remarkably simple device that utilizes a simple set of electron accounting rules in

order to predict the shape of, in particular, main group compounds. Organic molecules

are treated just as successfully as inorganic molecules.

The underlying assumptions made by the VSEPR method are the following.

- Atoms in a molecule are bound together by shared electrons, called bonding

pairs. More than one set of bonding pairs may bind any two atoms together

(multiple bonding).

- Some atoms in a molecule may also possess pairs of electrons not involved in

bonding. These are called lone pairs or non-bonded pairs.

- The bonding pairs and lone pairs around any central atom in a molecule take up

positions in which their mutual interactions are minimized. The logic here is

simple. Electron pairs are negatively charged and will get as far apart from each

other as possible.

- Lone pairs occupy more space than bonding electron pairs.

- Double bonds occupy more space than single bonds.

The bonding pairs and lone pairs of electrons in the valence level of an atom repel one

another due to their negative charges. The pairs of electrons take up positions as far from

each other as possible about the spherical central atom while remaining in the molecule.

A lone pair (LP) will spread out more than a bond pair (therefore, repulsion is greatest

between lone pairs (LP-LP).

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Shape Diagrams Shape Representation Properties

Example:CCl4

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

Example: NH3

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

Example: H2O

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

Example: HPO

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

Example: CH2O

Example: C2H4

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

Example: CO2

Example: HF

shape will be _____________________

__ lone pairs around

central atom

__ atoms bonded to

central atom(bonding

groups)

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Molecule Shape Geometry: The order of decreasing repulsion can be expressed as follows:

Compound Lewis

Diagram for Molecule

# of

Lone

Pairs

# of

Bonding

Electron

Groups

Name of

Shape

Shape Diagram

CH4

C2H4

H2O

NH3

C2H2

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Chemistry 2202

Lewis and Shape Diagrams Name:_____________________ (lewis and shapes) Formula

Lewis Diagram

for Molecule

For each central atom

Name to

represent

shape

Shape diagram

# of lone

pairs

# bonded

groups

NI3

CF4

OCl2

C2H2

FCN

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Formula

Lewis Diagram

for Molecule

For each central atom

Name to

represent

shape

Shape diagram

# of lone

pairs

# bonded

groups

N2H4

C2H4

CS2

SiH2O

SiCl4

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Chemistry 2202

Lewis and Shape Diagrams Name:_____________________ (lewis and shapes) Formula

Lewis Diagram

for Molecule

For each central atom

Name to

represent

shape

Shape diagram

# of lone

pairs

# bonded

groups

SiCl2S

CH3Cl

N2

C2F4

CS2

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Formula Lewis Diagram

for Molecule

For each central atom Name to

represent

shape

Shape diagram

# of lone

pairs

# bonded

groups

C2H3Cl

C2F2

C3H8

C3H7OH

PF3

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Chemistry 2202

Stereochemistry Lab Name:___________________ (Stereochemistry Lab) Purpose - to predict the shapes around central atoms using VSEPR theory

- to construct molecular model and test predictions of the shapes of molecules

Prelab - complete the pre lab columns below by constructing models of the molecules listed, and

draw a shape diagram for each model constructed. Name the actual shape and compare with

the predicted name.

Pre Lab Exercise

Lab Observations

Formula Lewis Diagram

For each central atom

Name to

represent

predicted

shape

Shape diagram

Name of

Actual

Shape

# of lone

pairs

# bonded

groups

NBr3

SiF2S

CF4

OCl2

C2F2

HOF

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NHF2

C2IBr

C2HF3

CHClBr2

H2O2

CO2

N2H3 F

CH3 OH

C3H8

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Polarity of Molecules

Atoms have different abilities to attract electrons. For example, the farther away from the

nucleus electrons are, the weaker is their attraction to the nucleus. Also, inner electrons

(those closer to the nucleus) shield the valence electrons form the attraction of the

positive nucleus.

Electronegativity - refers to the relative attraction that an atom has for shared

electrons in a covalent compound. It refers to the relative ability of

an atom to attract a pair of bonding electrons in its valence level.

Electronegativity Scale- developed by Linus Pauling.

- fluorine has the highest electronegativity, 4.0

- cesium has the lowest electronegativity, 0.7

- metals tend to have low electronegativities

- nonmetals tend to have high electronegativities.

The more electronegative an atom is the more strongly it will be able to pull

covalently shared electrons toward it. This will result in the formation of a ‘polar’

molecule.

Polar molecule- a molecule with a displacement of charge

Non-Polar Covalent Bonds

- covalent bonds between two non-metallic atoms with the same electronegativity.

- equal sharing of electrons between the bonded pair of atoms - examples include H2, N2, O2, SCS

Polar Covalent Bond

- bonding between two non-metallic atoms with different electronegativities.

- there is un-equal sharing of the bonding electron pair

- the bonded electron pair is displaced toward the more electronegative atom creating a

dipole.

– high electronegativity – partially negative

– low electronegativity – partially positive

- examples include HBr, H2O, HCl, NH3

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Attraction for Bonded Electrons : Tug of War

H ---- F

2.1 4.0

On one side of this tug of war there are two students tugging on the bonded

electrons, and on the other side there are four. In which direction will the bonded

electrons be displaced? What will be the result of this displacement?

In the above example the fluoride atom will attract the shared electrons more strongly

than the hydrogen atom. Thus, the shared pair are pulled closer to the more

electronegative atom. This results in the fluorine end of the molecule becoming partially

negative relative to the hydrogen end. The resulting dipole, along with the shape, causes

the molecule to be polar.

Polar Molecule A molecule in which the bond dipoles present do not cancel each

other out and thus result in a molecular dipole and a polar

molecule.

Bond Dipole - indicates the direction of displacement of charge in a

covalently bonded substance. Criteria regarding dipoles include;

- they are represented by an arrow in a Lewis diagram

- each polar bond has a bond dipole

- all polar molecules have bond dipoles

- many non-polar molecules will bond dipoles

Molecular Dipole - the result of bond dipoles in a molecule

- bond dipoles may or may not cancel out

- dipoles cancelled out - Non Polar

- dipoles not cancelled out - Polar

Example 1: H ----- H There will be no displacement of bonded electrons,

2.1 2.1 therefore the molecule is non-polar

--------->

Example 2: H ----- F In this case there will be a displacement of charge,

2.1 4.0 and there will be a bond dipole and the molecule

will be polar. The arrow indicates the shared pair is attracted

more strongly to the fluorine also. This also means that the

fluorine will have a net negative charge and the hydrogen a net

positive charge.

33

Examples

CO2 - is a linear molecule with two bond dipoles

- dipoles do cancel each other out

- there is no displacement of charge over the length of the molecule

- this is a Non Polar molecule.

HCN - is a linear molecule with two bond dipoles

- the dipoles do not cancel each other out

- this is a Polar molecule.

H2O - the water molecule is bent

- the individual bond particles do not offset, or cancel, each other.

- the geometry of the molecule does not allow the bond polarities to cancel

- as a result, water is a polar molecule.

CH4 - methane has a tetrahedral shape with four bond dipoles

- the dipoles all point toward the inner region of the molecule

- the whole inner region has a net negative charge.

- the whole outer surface, however, has a net positive charge.

- as a result, this molecule is Non Polar.

CH3Cl - is also a tetrahedral molecule with four bond dipoles.

- three of the dipoles point toward the inner region of the molecule

- one dipole points toward the chlorine atom.

- one end of this molecule will have a net negative charge

- the other end a net positive charge

- this molecule is Polar.

34

Chemistry 2202

Lewis and Shape Diagrams Name:_____________________ (electronegativity and polarity_2)

It should be remembered that in a covalent bond there is neither a gain or loss of electrons, but

instead, there is a sharing of electrons. The shared electrons can be displaced, however, away

from the least electronegative atom and toward the more electronegative atom. This, essentially,

creates a dipole within the individual covalent bond and may, in conjunction with the shape of

the molecule, result in the formation of a polar molecule.

1. Complete the chart below by filling in the appropriate information in the spaces provided.

Compound

Structural Formula w. arrows to represent Bond

Dipoles

Polarity

HF

NH3

CH4

CO2

N2

OF2

C2H6

CH3Cl

35

Chemistry 2202

Lewis and Shape Diagrams Name:_____________________ (electronegativity and polarity_3)

Complete the chart below by filling in the appropriate information in the spaces provided.

Compound

Lewis Diagram

of Molecule

Shape around

Central Atom

Shape Diagram

w. Bond Dipoles

Polarity

NH3

N2

CBr

OCl2

SiCl4

CHI3

C2H3Cl

C2H4

CH3OH

36

Dipole Strength

The degree to which a molecule is polar can be determined by comparing the difference

in electronegativity between the two atoms that are sharing the bonded pair of electrons.

This results in the formation of a Bond Continuum that is based on the idea of charge

separation due to differences in the electronegativity of individual atoms within a

molecule.

Non Polar

Covalent Bond

Polar

Covalent Bond

Ionic

Bond

no charge separation

bonding e- equally

shared

some charge separation

bonding e- not equally

shared

complete charge separation

bonding e- not equally shared

Ex. H2

Ex. HF

Ex. NaF

The amount of separation, and the location of a molecule within the bond continuum, can

be determined using the following formula;

EN = ENA - ENB - where EN refers to the electronegativity of the molecule. This

difference in electronegativity value will determine the degree of

polarity.

Electronegativity Value

Type of Compound

EN greater than 1.7 the molecule is mostly ionic. Shared electrons transferred

to the more electronegative atom.

EN is 1.7 to 0.5 the molecule is polar

EN less than 0.5 the molecule is slightly polar

EN is 0 the molecule is non-polar

Example1: KF

K ------ F

.82 3.98

EN = ENA - ENB

= 3.98 - 0.82

= 3.16 (Ionic compound)

Example 2: O2

O ------- O

3.44 3.44

EN = ENA - ENB

= 3.44 - 3.44

= 0 (Non Polar Covalent)

Example 3: NH

N-------H

Example 4: PH

P-------H

37

Chemistry 2202

Lewis and Shape Diagrams Name:_____________________ (electronegativity and polarity_1)

1. The shared electrons in a covalent bond will be displaced toward the atoms with the

higher electronegativities. For each of the following, identify the individual polarities of

the atoms, the polarity of the bond, the EN value and the location on the Bond

Continuum. Shared Electrons

w. dipole

Polarity of

Covalent

Bond

Bond

Electronegativity

Value (EN)

Location on Bond

Continuum

H - F

2.1 - 4.0

--------->

polar

EN = 4.0 - 2.1

= 1.9

Slightly or mostly

ionic (Hydrogen bond)

N - H

B - F

S - O

P - H

Si - C

Cu - Br

N - I

38

Br - Cl

N - I

Br - Cl

C - H

O - H

C - Cl

C - O

2. List the elements located in period 2 and their respective Pauling electronegativities.

3. List he elements located in Group VIIA and their respective Pauling electronegativities.

4. Explain why cesium and francium are the most reactive metals.

5. Explain why fluorine is the most reactive non-metal.

39

Network Covalent Bonding:

Strongest type of Intramolecular force.

Network covalent creates molecules where all atoms are linked together by individual

covalent bonds. This results in very large molecules with very high bonding strengths.

Because all atoms are bound in a three-dimensional structure these substances are the

hardest naturally occurring substances with the highest boiling points.

Substances with network covalent bonds do, however, tend to be brittle because the

shared electrons are rigidly held in place.

This type of bonding is found in four allotropes of Carbon. Allotropes are molecular forms of

the same element that have different physical and chemical properties.

diamond (Cn): - each carbon is covalently bonded to four other carbon atoms

- each molecule is constructed with a repeating 3-D rigid structure

- does not melt and is the hardest naturally occurring substance

- this material is so dense it slows the speed of light

- it is capable of breaking light into its component parts and re-reflecting it

- used in jewellery, drill-bits, etc.

- 90% of all industrial diamond is synthetic(man-made)

graphite (Cn): - each C atom is covalently bonded to three other C atoms

- C atoms are covalently bonded to form a 2-D hexagonal shape structure

- it contains three very strong covalent bonds and one weak one

- soft - this material is in your pencil

- slippery – this substance is used in many lubricants

Fullerines, (C60)- also called ‘buckminsterfullerine’ or ‘bucky balls’

- an extensive class of large C60, C70, C74, C80 type compounds

- only recently discovered – in the 1980’s

- currently little commercial use though potential is significant

- potential use in fractional distillation to produce unique hydrocarbons

- potential use as lubricants

- made up of C atoms in the shape of 12 pentagons and 20 hexagons –

like a soccer ball

- NOVA – Diamond Deception.

Nanotubes - very tiny structures

- C400 and greater

- like fullerine network that has been stretched out into a cylinder shape

- very high strength structures unmatched by current materials

- 500 times stronger than steel

- 100 timel lighter than steel

- great potential in clothing, combat gear, various fibres, drug delivery,

vehicle panels, cosmetics, clothing, etc

- http://nanotechweb.org/

40

Network Covalent Bonding is also found in;

silicon carbide (SiC) - also called carborundum

- like carbon but alternating Si and C atoms

- used in abrasive tools such as grindstones

silicon dioxide, (SiO2) - also known as quartz

- each Si atom is surrounded by four O atoms

- used extensively in jewellery and electronic equipment

Macromolecules - a macromolecule is a molecule with a large molecular mass,

- the term is generally restricted to polymers

- these include proteins, starches, lipids and nucleic acids (such as DNA)

- some of these are called "biomacromolecules"

- synthetic examples include plastics.

Plastic is the general term for a wide range of synthetic or semisynthetic polymerization

products.They are composed of organic condensation or addition polymers and may

contain other substances to improve performance or economics. There are many natural

polymers generally considered to be "plastics". Some household plastics include:

The term macromolecule is also sometimes used to refer to aggregates of two or more

macromolecules held together by intermolecular forces rather than by chemical "bonds".

Substances that are composed of macromolecules often have unusual physical properties.

The properties of liquid crystals and such elastomers as rubber are examples.

Liquid crystals are substances that exhibit a phase of matter that has properties

between those of a conventional liquid, and those of a solid crystal. For instance,

a liquid crystal (LC) may flow like a liquid, but have the molecules in the liquid

arranged and oriented in a crystal-like way.

The term elastomer is often used interchangeably with the term rubber. They are

amorphous polymers existing above their glass transition temperature, so that

considerable segmental motion is possible. Their primary uses are for seals,

adhesives and molded flexible parts.

41

Chemistry 2202

Unit 2: STSE: Common Bonds Name:_________________

Directions: Answer the following questions in the spaces provided.

Part 1: Bonding in the Molecules of Life

1. What does DNA stand for? _____________________________________

2. What are the 3 parts that make up a nucleotide?

i. __________________________________

ii. __________________________________

iii. __________________________________

3. What are the four types of DNA nitrogenous bases?

i. ________________________

ii. ________________________

iii. ________________________

iv. ________________________

4. How many hydrogen bonds hold together?

a. The A-T base pair? _____________________

b. The C-G base pair? _____________________

5. In messenger RNA, the nitrogeneous base uracil (U) takes the place of thymine. Thus,

when mRNA reads the DNA code, the A links to the U. Given the structure of uracil,

draw the structures of A and U showing how the two form a base pair, including all

hydrogen bonds. Ensure the hydrogen bonds are clearly labelled.

42

6. Why is uracil complementary only to adenine, and not guanine or cytosine?

________________________________________________________________________

________________________________________________________________________

Part 2: Bonding in the Noble Gases: The Canadian Connection

7. Define Ainert.@ ________________________________________________________________________

________________________________________________________________________

8. Is it true or false to describe noble gases as inert? Explain.

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

9. Name the noble gas compound that Professor Neil Bartlett synthesized in 1962.

________________________________

Part 3: Why does Ice Float?

10. What are 2 physical properties of ice that make it different that water?

i. ________________________________________

ii. ________________________________________

11. Write the equation for density.

12. What is happening with hydrogen bonds in water in the liquid state. Explain and draw a

diagram.

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

13. What is happening with hydrogen bonds in water in the liquid state. Explain and draw a

diagram.

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

43

14. What do ice and snowflakes have in common?

________________________________________________________________________

________________________________________________________________________

Part 4: Buckminsterfullerines

15. Write the formula for buckminsterfullerene: _________________________

16. Describe the shape of a fullerene molecule.

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

17. What are some possible applications of fullerene molecules?

i. _____________________________________________

ii. _____________________________________________

18. In producing buckminsterfullerene for the first time, an inert gas was used in the

container of vapourized carbon atoms. Why would using a simple air filled container not

work?

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

44

2. Intermolecular Forces – bonding between molecules

Water and hydrogen sulfide have the same molecular shape. They are both v-

shaped, or bent. However, H2O with a molar mass of 18g, is a liquid at room

temperature, while H2S , with a molar mass of 34g, is a gas. As well, water has a

boiling point of 100oC, while hydrogen sulfide has a boiling point at -61

oC. These

physical properties are difficult to explain considering what we have learned

about intramolecular forces.

Note: It is the intramolecular force that creates the molecule, it is not the

force that holds molecules to each other.

It is important to remember that pure covalent bonds are not held together by

ionic bonds in lattice type structure. They do, however, form solids, liquids and

gases at room temperatures. Something must hold these molecules in close

proximity to each other when molecular compounds are in the liquid and solid

state.

Note: The force that does hold these molecules together is called an

intermolecular force – meaning between molecules.

The intermolecular forces that hold molecules to each other were studied

extensively by a Dutch physicist named Johannes van der Waals. The

intermolecular forces he studied are often called ‘Van der Waals forces’. Johannes

van der Waals showed that there were three different types of intermolecular

forces, namely;

1. London Dispersion forces

2. Dipole – Dipole forces

3. Hydrogen bonding

45

1. London Dispersion Forces

- this type of intermolecular forces is found in all molecular compounds

- Weak, simultaneous(at the same time) attractions will form between the

electrons in molecules to the protons in neighbouring molecules

- These IMFS exist between all molecules, regardless of polarity

1

2

3

4

Strength of London Dispersion Forces and Number of Electrons

London Dispersion forces tend to be relatively weak – especially when the

molecular compounds are in the liquid or gaseous state. The strength of these

forces is proportional to the numbers of electrons in the compound. Molecular

compounds with larger numbers of electrons tend to have stronger LD

intermolecular forces and subsequently higher boiling points.

Compound Number of electrons Boiling point

Fluorine (F2) 18 electrons -188oC

Iodine (I2) 106 electrons 184oC

46

Noble Gases and Strength of London Dispersion Forces

Noble Gas Number of Electrons Boiling Point

helium -269°C

neon -246°C

argon -186°C

krypton -152°C

xenon -108°C

radon -62°C

All of these elements have, or are, monatomic molecules.

The reason that the boiling points increase as you go down the group is that the

number of electrons increases, and so also does the radius of the atom. The more

electrons you have, and the more distance over which they can move, the greater

the number of attractions exist and therefore the greater the dispersion forces.

Isoelectronic Molecules: Molecules with the same number of electrons,

isoelectronic, would be expected to have same or

similar boiling points as they should have the same

strength of LD intermolecular forces.

Compound Number of electrons Boiling point

Argon 18 electrons -186oC

Fluorine, F2 18 electrons -188oC

These forces are critical in the stabilization of biological membranes of living

cells, or lipids.

47

Strength of London Dispersion Forces and Shapes of Molecules

Molecule complexity is a contributing factor, as well, to the strength of London

Dispersion Forces.

Less complex molecules, or uniquely shaped molecules, require greater amounts of

energy to achieve separation. Lesser complexity, or unique shape, results in higher

boiling and melting points.

Long thin molecules can develop stronger temporary dipoles due to electron cluster

movement than short fat ones containing the same numbers of electrons. Long thin

molecules can also lie closer together - these attractions are most effective if the

molecules are really close. Long thin molecules will have stronger London Dispersion

Forces. (Think of attractions between sumo wrestlers compared to WWE wrestlers )

Example: the hydrocarbon molecules butane and 2-methylpropane both have

a molecular formula C4H10, but the atoms are arranged differently.

In butane the carbon atoms are arranged in a single chain, but 2-

methylpropane is a shorter chain with a branch.

Butane has a higher boiling point because the London dispersion

forces are greater. The molecules are longer and smaller and can

lie closer together than the shorter, fatter 2-methylpropane

molecules.

48

2. Dipole – Dipole Intermolecular Forces

- these types of forces are found only in molecules which are polar

- the force of attraction involves the permanent net negative end of one

molecule being attracted to the net positive end of a neighbouring

molecule, and vice-versa. The molecules will orient themselves so that

oppositely charged ends are close to each other.

There is no reason why this has to be restricted to two molecules. As long as the molecules are

close together this synchronised movement of the electrons can occur over huge numbers of

molecules.

- dipole – dipole forces add to the already present LD intermolecular bond

strength resulting in molecules having higher boiling and melting points.

This is because the energy requirement needed to separate the molecules

is greater.

Compound Number of

electrons

Boiling point

Kr isoelectronic -152oC

HBr isoelectronic -67oC

The HBr, in the above example, has a higher boiling point because of the

added strength of a dipole – dipole intermolecular bond.

49

3. Hydrogen Bonding

Compare the two graphs above and explain why there is a different trend in the boiling

point for each of the following ‘groups’ of compounds.

Compound CH4 SiH4 GeH4 SnH4

# of

Electrons

Boiling

Point

Compound HF HCl HBr HI

# of

Electrons

Boiling

Point

The hydrogen bond is a special case of a dipole – dipole force, but with a

significantly stronger bond (about 10X stronger). It involves a direct bond

between hydrogen atoms and a highly electronegative atom such as oxygen,

nitrogen or fluorine. Because of the nature of hydrogen, and because it has only

one electron that has now taken part in the bonding process, the hydrogen’s

positive nucleus becomes attracted to the neighboring molecules net negative end.

This results in a highly polar bond. It is the strongest of the intermolecular forces.

50

Substances with hydrogen bonds

- will have stronger intermolecular forces

- will require greater amounts of energy to break

bonds between molecules

- will therefore have higher boiling and melting points.

Water, for example, has a very high boiling point

relative to the number of electrons. It has a high

boiling point because it has the added intermolecular

bond strength of a hydrogen bond.

This strong hydrogen bond gives water, for example, some unusual properties, one being the

expansion of water molecules as they freeze. This allows the ice to float, rather than sink,

allowing the frozen water to provide a blanket of insulation against extensive freezing of our

water bodies. The consequences to aquatic life, as we currently know it, would be catastrophic.

Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being

constantly broken and reformed in liquid water. If you liken the covalent bond between the

oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On

the same scale, van der Waals attractions represent mere passing acquaintances!

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four hydrogen bonds with surrounding

water molecules. There are exactly the right numbers of + hydrogens and lone pairs so that

every one of them can be involved in hydrogen bonding.

This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride. In the

case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only

has one lone pair. In a group of ammonia molecules, there aren't enough lone pairs to go around

to satisfy all the hydrogens.

In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right

number of each. Water could be considered as the "perfect" hydrogen bonded system.

51

Chemistry 2202

Unit 2: Bond Type Name:__________________ (bondtype)

Identify the types of bonding or forces found in each of the following:

Chemical

Formula

Intramolecular Bonding

Intermolecular Bonding

Network

Covalent

Ionic

Bonding

Metallic

Bonding

Covalent

Bonding

London

Dispersion

Diploe-

dipole

Hydrogen

Na

KCl

H2O

SiO2

CH4

NH3

SiC

CH3OH

HCN

CO2

OF2

Ca(OH)2

CH3F

52

Strength of Bonds: and Type of Intramolecular Force.

We can use the boiling and melting points of compounds to compare the strength of

intramolecular forces between atoms and the strength of intermolecular forces between

molecules. When comparing th b.p. amd m.p. of a compound we are really measuring the bond

energy within these substances.

Bond Energy - the energy required to break a bond apart

- reflects, exactly, the amount of energy required to form a bond

- energy released when a bond is formed (exothermic rxn)

- energy absorbed when a bond is broken (endothermic rxn)

Bond Strength

Bond Type Bond Strength Examples

Network Covalent Bonding Strongest Bonds Cn, SiC, SiO2

Ionic Bonding NaCl, MgSO4

Metallic Bonding Na, Al,

Covalent Bonding Weakest Bonding H2O, CO

- Any network covalently bonded substance, in a list of substances, will have the stronger

bonds and will have the higher boiling point. Comparing the boiling points of NaCl and

SiC would reveal that SiC has the higher boiling point because it has network covalent

bonding.

- Any metallically bonded substance that is present with a network covalently bonded

substance and an ionically bonded substance will have the third strongest bonds and

therefore the third highest boiling point.

The compounds Na, SiO2 and KCl, when ranked from highest to lowest boiling

point, would be SiO2, KCl and Na.

53

Strength of Bonds: and Type of Intermolecular Force.

Any molecular compounds that are not network covalent solids will have the lowest

boiling points. To compare the boiling points of these molecular compounds we need to

look at the types of intermolecular forces that are present.

Bond Type Relative Strength

London Dispersion - present in all molecules

- bond strength is based on # of electrons and shape

- if iso-electronic check for other Intermolecular bonding

Dipole-Dipole - present in polar substances

- adds slightly to intermolecular bond strength

- increases b.p. of substances slightly

Hydrogen - present in molecules containing H-F, H-O and H-N

- adds significantly to intermolecular bond strength

- increases b.p. of substances significantly

Note: Substances that are iso-electronic and polar, but do not contain a hydrogen

bond, the substance with the higher b.p. will be the one with the greatest

difference in electronegativity, or the one that has a peculiar shape that results in

more intermolecular bonds between specific molecules.

54

Example Questions

1. Compare the bond strengths of each of the following by ranking from lowest

boiling point to highest boiling point.

Ag, PH3, LiCl, SiO2

What type of bonding does each have?

2. The boiling point of C2H6 is -87oC while the boiling point of CH3F is -78

oC.

Account for this difference. What type of bonding does each molecule have - Intermolecular

Are there London Dispersion Forces present - yes

Compare strengths of the LD forces - isoelectronic ?

Is there dipole - dipole intermolecular force - Compare strengths of the DD forces

Is there hydrogen bond intermolecular force

Compare strengths of the hydrogen bond forces

Account for this difference in boiling point.

3. The compounds C2H5F and C2H5OH are both molecular substances with

different boiling points. Identify the types of intermolecular forces that exist

between individual molecules and identify which has the higher boiling point.

Be able to explain your answer.

What type of bonding does each molecule have - Intermolecular

Are there London Dispersion Forces present - yes

Compare strengths of the LD forces - isoelectronic ?

Is there dipole - dipole intermolecular force -

Compare strengths of the DD forces

Is there hydrogen bond intermolecular force Compare strengths of the hydrogen bond forces

Predict which substance will have the higher boiling point.

55

Chemistry 2202

Unit 2: Bond Strength Name:__________________ (bondstrength_1)

1. The boiling point of argon is -186oC and the boiling point of fluorine is -188

oC. Are these

two molecular compounds expected to have similar boiling points? Explain.

2. Krypton has a boiling point of -152oC and hydrogen bromide has a boiling point of -

67oC. Explain why these two compounds have different boiling points.

3. The boiling point of chlorine is -35oC and the boiling point of C2H5Cl is 13

oC. Account

for their difference in boiling points.

4. Given the compounds C2H3Cl and C2H3I, which would have the higher boiling point?

Explain.

5. The boiling point of BrF is -20oC and the boiling point of C3H8 is -45

oC. Account for

their difference in boiling points.

6. The boiling point of C2H6 is -87oC and the boiling point of CH3F is -78

oC. Account for

their difference in boiling points.

56

Chemistry 2202

Unit 2: Bond Strength Name:__________________ (bondstrength_2) 1. CH4 has a boiling point of is -162

oC and the boiling point of C5H12 is 36

oC. Account for

their difference in boiling points.

2. The boiling point of C3H8 is -45oC and the boiling point of hutane,C5H12F is -38

oC.

Account for their difference in boiling points.

3. Identify the types of intramolecular and/or intermolecular forces found in each of the

following compounds and rank from highest to lowest boiling points.

Compound

Type of

Intramolecular Force

Type of

Intermolecular Force

Ranking from

Highest to Lowest NaCl

C2H5OH

C2H3Cl

C3H8

SiO2

4. Given the boiling points of the compounds in the chart below, predict from the list the

boiling point of C6H5Br.

Compound

Boiling Point (

oC)

C6H5F

85

oC

C6H5Cl

132

oC

C6H5I

188

oC

A. 98 oC b. 122

oC c. 156

oC d. 249

oC e. 337

oC

57

Unit 2: Bond Strength Chemistry 2202 Name:__________________ Molecular

Compound

Number

of

Electrons

Boiling

Point

(oC)

London

Dispersion

Forces

Dipole-dipole

Intermolecular

Forces

Hydrogen

Bonding

F2(g)

- 188

Cl2(g)

- 35

Br2(g)

59

I2(g)

184

ClF(g)

- 101

BrF(g)

- 20

BrCl(g)

5.0

ICl(g)

97

IBr(g)

116

CH4(g)

- 162

C2H6(g)

- 87

C3H8(g)

- 45

C4H10(g)

- 0.5

C5H12(l)

36

CF4(g)

- 129

CCl4(l)

77

CBr4(s)

189

CH3F(g)

- 78

CH3Cl(g)

- 24

CH3Br(g)

3.6

CH3I(l)

43

CH3OH(l)

65

58

C2H5F(g)

- 38

C2H5Cl(g)

13

C2H5Br(l)

38

C2H5I(l)

72

C2H5OH(l)

78

Using the information in the preceding table, complete the following questions in the space

provided. (bondstrength_3)

1. Compare the boiling points of BrF(g) and C3H8(g). Account for the difference in their

intermolecular bond strength.

2. Dimethyl ether, (CH3)2O(g), has a boiling point of - 24.9 oC. Compare with the boiling

point of ethanol and account for the difference.

3. Explain the trend, that exists within molecular compounds, that compares the number of

electrons and the strength of intermolecular forces.

4. Methanol,CH3OH(l), and ethanol,C2H5OH(l), each have the least number of electrons, in

their respective series, but the highest boiling points. Account for this difference.

5. Explain the difference in boiling point between C2H6(g) and CH3F(g).

6. Explain the difference in boiling point between Cl2(g) and C4H10(g).

7. Explain the difference in boiling point between BrCl(g) and C2H5Br(l).

59

Chemistry 2202 Unit 2: Bond Strength Name:__________________ (bondstrength_4)

1. Draw lewis, structural and shape diagrams for the following molecules.

Compound

Lewis diagram

Structural diagram

Shape Diagram

Polarity

CH3Cl

NBr3

H2S

C2H6

2. With reference to electronegativity, what is the difference between a covalent and an

ionic bond?

3. Ionic compounds have high melting and boiling points. Why is this so?

4. Distinguish between intramolecular and intermolecular forces.

60

5. Methane has a boiling point of about -190oC and ammonia has a boiling point of about -

30oC. Account for this difference.

6. A. What is an intermolecular dipole-dipole force?

B. Show a dipole-dipole intermolecular force with the aid of a diagram.

C. What conditions are necessary for a dipole-dipole intermolecular force to exist?

7. The boiling point of water, a molecular compound, is 100oC and the boiling point of

NaBr, an ionic compound, is 1390oC. Chemists, however, believe that the bonds within

the water molecules are stronger than those in the NaBr crystal. Explain how this

difference in boiling points does not contradict the belief of chemists.

61

Chemistry 2202 Bonding Review

(bonding review)

Intramolecular Bonding

Network Covalent - Strongest type of bonding

- Bonding within the compounds Cn, SiC and SiO2

Ionic - Ions are formed when electron(s) are transferred from one atom to another

- Metal atoms lose electrons to form cations (positive) and nonmetal atoms

gain electrons to form anions (negative)

- Ionic bonds are formed by the attraction of oppositely-charged ions to

each other

- The second strongest type of bonding

Metallic - Positive ions attract valence electrons which are free to move from

one empty valence orbital to another

- The third strongest type of bonding

Covalent - Sharing of electrons between 2 nonmetallic atoms

- Occurs in molecular substances

Intermolecular (Covalent)

London Dispersion Forces

- occurs among all molecular substances

- the attraction of positive nuclei of one molecule to the electrons of another

molecule (& vice-versa)

- strength of these forces depends on the number of electrons a substance contains,

such that the greater the number of electrons, the stronger the London forces

among the molecules of that substance

Dipole-dipole Forces

- only occurs among polar molecules

- the partial-positive end of one polar molecule is attracted to the partial-negative

end of another polar molecule (& vice-versa)

Hydrogen Bonds

- a special type of dipole-dipole force (about 10 times stronger)

- only occurs among molecules that contain a H atom which is directly bonded to a

highly electronegative atom ( F, O, N) ie. the molecule contains at least one H-F,

H-O, or H-N bond.

62

Chemistry 2202 Bonding Review

(bonding review)

Criteria for Determining Strength of Bonds - determining Boiling Point (b.p.) or Melting Point (m.p.)

Any network covalent solid (eg. Cn , SiC or SiO2) will have the highest b.p.

Any ionic substance that is present with a network covalent solid will have the second highest

b.p.

(eg. NaCl with SiC). If a network covalent solid is not present then the ionic substance

will have the higher b.p.

Any metallic substance that is present with a network covalent solid and an ionic compound will

have the third highest b.p. (eg. Na with SiC and NaCl). If it is present with only one of these it

will have the second highest b.p., while if neither a network covalent solid nor an ionic substance

is present the metallic substance will have the highest b.p.

Any molecular substances that are not network covalent solids will have the lowest b.p.Of these,

to determine the substance with the highest b.p., identify the types and relative strengths of the

intermolecular forces (IMF) present:

London forces

- present in all molecules

- count the number of electrons; if no other IMF are present the sustance

with the greatest number of electrons will have the highest b.p.

- if substances have the same # of electrons (isoelectronic) then determine if

other IMF are present.

Dipole-dipole forces

- present in polar substances, in addition to London forces, thus this

substance has the higher b.p.

- if both substances are polar, then determine if Hydrogen bonds are present

Hydrogen bonds

- present if molecules contain a H-F, H-O or H-N bond

- this substance will have the higher b.p., since it contains all 3 types of IMF

ie. London forces, Dipole-dipole forces and Hydrogen bonds

Note: For substances that are isoelectronic and polar but do not contain Hydrogen bonds, the

substance with the higher b.p. will be the one that is most polar ie. has the greatest

difference in electronegativities between its' atoms.