unit 2: chemical bonding -...

Unit 2: Chemical Bonding

Chemistry2202

Outline

Bohr diagrams

Lewis Diagrams

Types of Bonding

Ionic bonding

Covalent bonding (Molecular)

Metallic bonding

Network covalent bonding

Types of Bonding (cont’d)

London Dispersion forces

Dipole-Dipole forces

Hydrogen Bonding

VSEPR Theory (Shapes)

Physical Properties

Bohr Diagrams (Review)

How do we draw a Bohr Diagram for

- The F atom?

- The F ion?

Draw Bohr diagrams for the atom and

the ion for the following:

Al S C l Be

Lewis Diagrams

LD provide a method for keeping

track of electrons in atoms, ions, or

molecules

Also called Electron Dot diagrams

the nucleus (P& N) and filled energy

levels are represented by the element

symbol

Lewis Diagrams

dots are placed around the element

symbol to represent valence electrons

Lewis Diagrams

eg. Lewis Diagram for F

F•

••

•••

•

lone pair

lone pair

lone pair

bonding

electron

Lewis Diagrams

lone pair – a pair of electrons not

available for bonding

bonding electron – a single electron

that may be shared with another atom

Lewis Diagrams

eg. Lewis Diagram for C

C•

••

•

Lewis Diagrams

eg. Lewis Diagram for P

P•

••

••

Lewis Diagrams

eg. Lewis Diagram for Na

Na•

Lewis DiagramsFor each atom draw the Lewis diagram

and state the number of lone pairs and number of bonding electrons

Li Be Al Si

Mg N B O

Complete bonding worksheet #1

Lewis Diagrams for Compounds

draw the LD for each atom in the

compound

The atom with the most bonding

electrons is the central atom

Connect the other atoms using single

bonds (1 pair of shared electrons)

In some cases there may be double

bonds or triple bonds


eg. Draw the LD for:

PH3

CF4

Cl2O

C2H6

C2H4

C2H2


eg. Draw the LD for:

NH3 SiCl4 N2H4 HCN

SI2 CO2 N2H2 CH2O

POI CH3OH

N2 H2 O2


A structural formula shows how the

atoms are connected in a molecule.

To draw a structural formula:

replace the bonded pairs of electrons

with short lines

omit the lone pairs of electrons


Why is propane (C3H8) a gas at STP while

kerosene (C10H22) a liquid?

Why is graphite soft enough to write with while

diamond is the hardest substance known even

though both substances are made of pure

carbon?

Why can you tell if it is ‘real gold’ or just ‘fool’s

gold’ (pyrite) by hitting it with a rock?

‘As Slow As Cold Molasses’

‘All Because of Bonding’

‘liquids’ @ -30 ºC

Bonding

Bonding between atoms, ions and

molecules determines the physical and

chemical properties of substances.

Bonding can be divided into two

categories:

- Intramolecular forces

- Intermolecular forces

Bonding

Intramolecular forces are forces of

attraction between atoms or ions.

Intramolecular forces include:

1. ionic bonding

2. covalent bonding

3. metallic bonding

4. network covalent bonding

Bonding

Intermolecular forces are forces of

attraction between molecules.

Intermolecular forces include:

5. London Dispersion Forces

6. Dipole-Dipole forces

7. Hydrogen Bonding

Ionic and Covalent Bonding

ThoughtLab p. 161

Identify #’s 1 - 6

Ionic Bonding

Occurs between cations and anions –

usually metals and non-metals.

An ionic bond is the force of attraction

between positive and negative ions.

Properties:

conduct electricity as liquids and in solution

hard crystalline solids

high melting points and boiling points

brittle

In an ionic crystal

the ions pack tightly

together.

The repeating 3-D

distribution of

cations and anions

is called an ionic

crystal lattice.

Ionic Bonding

Each anion can be

attracted to six or

more cations at

once.

The same is true

for the individual

cations.

Ionic Bonding

Ionic Bonding

Covalent Bonding

Occurs between non-metals in molecular

compounds.

Atoms share bonding electrons to become

more stable (noble gas structure).

A covalent bond is a simultaneous

attraction by two atoms for a common pair

of valence electrons.

Covalent Bonding

Molecular compounds

have low melting and

boiling points.

Exist as distinct

molecules.

Covalent Bonding

Molecular

compounds do

not conduct

electric current

in any form

Property Ionic Molecular

Type of

elements

Metals and

nonmetals

Non-Metals

Force of

Attraction

Positive ions attract

negative ions

Atoms attract a

shared electron

pair

Electron

movement

Electrons move

from the metal to

the nonmetal

Electrons are

shared

between atoms

State at room

temperature

Always solids Solids, liquids,

or gas

Property Ionic Molecular

Solubility Soluble or low

solubility

Soluble or

insoluble

Conductivity in

solid state

None None

Conductivity in

liquid state

Conducts None

Conductivity in

solution

Conducts None


Metallic Bonding (p. 171)

metals tend to lose valence electrons.

valence electrons are loosely held and

frequently lost from metal atoms.

This results in metal ions surrounded by

freely moving valence electrons.

metallic bonding is the force of attraction

between the positive metal ions and the

mobile or delocalised valence electrons

Metallic Bonding

Metallic Bonding

This theory of metallic bonding is called

the ‘Sea of Electrons’ Model or ‘Free

Electron’ Model

Metallic Bonding

This theory accounts for properties of metals

1. electrical conductivity

- electric current is the flow of charge

- metals are the only solids in which charged

particles are free to move

2. solids

- Attractive forces between positive cations and

negative electrons are very strong

Metallic Bonding

3. malleability and ductility

- metals can be hammered into thin

sheets(malleable) or drawn into thin

wires(ductile).

- metallic bonding is non-directional such that

layers of metal atoms slide past each other

under pressure.

Network Covalent Bonding (p. 199)

occurs in 3 compounds (memorize these)

diamond – Cn

carborundum – SiC

quartz – SiO2

large molecules with covalent bonding in 3-d

each atom is held in place in 3-d by a network of other atoms

Network Covalent bonding

Properties:

the highest melting and boiling points

the hardest substances

brittle

do not conduct electric current in any

form

Strongest

1. Network Covalent (Cn ,SiO2 , SiC)

2. Ionic bonding(metal & nonmetal)

3. Metallic bonding (metals)

4. Molecular (nonmetals)

Weakest

MP

& B

P d

ecre

ase

s

Valence Shell Electron Pair

Repulsion theory (VSEPR)

The shape of molecules is determined by

the arrangement of valence electron pairs

around the atoms in a compound.

There are 5 shapes that can be

determined by the # of bonds and # of

lone pairs on the central atom.

1. Tetrahedral (4 bonds; 0 lone pairs)

2. Pyramidal (3 bonds; 1 lone pair)

3. V-shaped (2 bonds; 2 lone pairs)

4. Trigonal Planar (3 bonds; 0 lone pairs)

5. Linear (2 bonds; 0 lone pairs)

For each molecule below draw the

Lewis diagram, structural diagram

shape name.

HOCl H2Se H2O2

NBr3 C2F4 C2H6

CHCl3 CH3OH PBr3

I2 SiH4 HCN

HSiHO C2H2

Complete molecular models lab

Electronegativity (EN - p. 174)

EN is a measure of the attraction that

an atom has for shared electrons.

A higher EN means a stronger

attraction or electrostatic pull on

valence electrons

EN values increase as you move:

- from left to right in a period

- up in a group or family

Increases

The Bonding Continuum

Electronegativity & Covalent Bonds

1. polar covalent bond

- a bond between atoms with different EN

- the shared electron pair is attracted

more strongly to the atom with the

higher EN

ClH

δ−δ+

Electronegativity (p. 174)

polar covalent bond

a covalent bond between atoms with different EN

the shared electron pair is attracted more strongly to the atom with the higher EN

nonpolar covalent bond

a covalent bond between atoms with the same EN

bond dipole

An arrow drawn in the direction of the atom with the greatest EN

Complete: #’s 7 – 9 on p.178

Weakest

Covalent (nonmetals)

→ London Dispersion

(all molecules)

→ Dipole-Dipole

(polar molecules)

→ H bonding

(H-N, H-O, H-F)

p. 226 #13

Omit parts g), j) – o), q), u), & v)

- Answers on p. 815 for #13

- Incorrect answers

c), d), & s)

Electronegativity and Ionic Bonds

Because the EN of metals is so

low, metals lose electrons to form

cations

Nonmetals gain electrons to form

anions because the EN of

nonmetals is relatively high

Electronegativity and Ionic Bonds

When ions form, the resulting

electrostatic force is an ionic bond

Electronegativity and Covalent Bonds

Atoms in covalent compounds can

either have:

the same EN

eg. Cl2 , PH3, NCl3 different EN

eg. HCl


Atoms that have the same EN attract

the shared valence electrons to the

same extent.

Covalent bonds resulting from equal

sharing of the bonding electron pairs

are called Nonpolar Covalent

Bonds


Atoms that have different EN attract

the shared pair of valence electrons

at different strengths

The atom with the higher EN exerts

a stronger attraction on the shared

electron pair

eg. H2O


Since the oxygen atom has a higher EN

the bonding electrons will be pulled

closer to the oxygen atom

This results in slight positive and

negative charges within the bond.

These charges are referred to as

“partial charges” and are denoted

with the Greek letter delta (δ).


The region around the oxygen atom

will be slightly negative, and around

the hydrogens will be slightly positive


The symbol, δ+ represents a partial

positive charge (less than +1) and

δ− represents a partial negative

charge (less than −1).

Since the bond is polarized into a

positive area and a negative area

the bond has a “bond dipole”.


The arrow points to

the atom with the

higher EN.

p.178


Covalent bonds resulting from

unequal (electronegativities)

sharing of bonding electron

pairs are called Polar

Covalent Bonds

Electronegativity Homework

#’s 7, 8, & 9 - p. 178

#’s 1, 2, & 3 - p. 180

Bond Energy (pp. 179-180)

1. Describe the forces of attraction and repulsion present in all bonds.

2. What is bond length?

3. Define bond energy.

4. Which type of bond has the most energy?

5. How can bond energy be used to predict whether a reaction is endothermic or exothermic?

Test Outline

Bohr Diagrams (atoms & ions)

Lewis Diagrams (Electron Dot)

Ion Formation

Ionic Bonding, Structures & Properties

Covalent Bonding, Structures & Properties

Test Outline

Metallic Bonding Theory& Properties

Network Covalent Bonding & Properties

Electronegativity

Bond Dipoles & Polar Molecules

VSEPR Theory

LD, DD, & H-bonding

Predicting properties (bp, mp, etc.)

Molecular Dipoles

The vector sum of all the bond

dipoles in a molecule is a Molecular

Dipole

A Polar Molecule has a molecular

dipole that points toward the more

electronegative end of the molecule.

eg. H2O

Molecular Dipoles

NonPolar Molecules DO NOT have

molecular dipoles. This occurs when:

- the bond dipoles cancel

- there are no bond dipoles

eg. CO2 PH3

Molecular Dipoles

To determine whether a molecule is

polar:

- draw the Lewis diagram and the

structural diagram

- draw the bond dipoles and determine

whether they cancel

Intermolecular Forces

Strongest bonds; Highest mp and bp

1. Network Covalent (Cn SiO2 SiC)

2. Ionic bonding(metal & nonmetal)

3. Metallic bonding (metals)

4. Molecular (nonmetals)

Weakest bonds; Lowest mp and bp

- Intermolecular forces present

To compare mp and bp in covalent

compounds you must use:

- London Dispersion forces (p. 204)

(all molecules)

- Dipole-Dipole forces (pp. 202, 203)

(polar molecules)

- Hydrogen Bonding (pp. 205, 206)

(H bonded to N, O, or F)

Intermolecular Forces (p. 202)


Covalent compounds have low mp and

bp because forces between molecules

in covalent compounds are very weak.

Intermolecular forces were studied

extensively by the Dutch physicist

Johannes van der Waals

In his honor, two types of intermolecular

force are called Van der Waals forces.


Intermolecular forces can be used to

account for the physical properties of

covalent compounds.


Van der

Waals


• LD forces exist in ALL molecular elements

& compounds.

•The positive charges in one molecule

attract the negative charges in a second

molecule.

• The temporary dipoles caused by electron

movement in one molecule attract the

temporary dipoles of another molecule.


The strength of these forces depends on:

a) the number of electrons

more electrons produce stronger LD

forces that result in higher mp and bp

eg. CH4 is a gas at room temperature.

C8H18 is a liquid at room temperature.

C25H52 is a solid at room temperature.

Account for the difference.


Two molecules that have the same number

of electrons are isoelectronic

eg. C2H6 and CH3F


b) shape of the molecule

molecules that “fit together” better will

experience stronger LD forces

eg. Cl2 vaporizes at -35 ºC while C4H10

vaporizes at -1 ºC. Use bonding to

account for the difference.

2. Dipole-dipole Forces

- occur between polar molecules

- the δ+ end of one polar molecule is

attracted to the δ- end of another polar

molecule (& vice-versa)

eg. Which has the higher boiling point

CH3F or C2H6 ?

p. 202

3. Hydrogen Bonds

- a special type of dipole-dipole force

(about 10 times stronger)

- only occurs between molecules that

contain a H atom which is directly

bonded to F, O, or N

ie. the molecule contains at least one

H-F, H-O, or H-N covalent bond.

3. Hydrogen Bonds

-the hydrogen bond occurs between the

H atom of one molecule and the N, O,

or F of a second molecule.

eg. Arrange these from highest to

lowest boiling point

C3H8 C2H5OH C2H5F

p. 206

NOTE: To compare covalent compounds

you must use:

- London Dispersion forces

(all molecules)

- Dipole-Dipole forces

(polar molecules)

- Hydrogen Bonding



p. 210

Intermolecular Forces1. Use intermolecular forces to explain the following:

a) Ar boils at -186 °C and F2 boils at -188 °C .

b) Kr boils at -152 °C and HBr boils at -67 °C.

c) Cl2 boils at -35 °C and C2H5Cl boils at 13 °C .

2. Examine the graph on p. 210:

a) Account for the increase in boiling point for the

hydrogen compounds of the Group IV elements.

b) Why is the trend different for the hydrogen

compounds of the Group V, VI, and VII elements?

c) Why are the boiling points of the Group IVA

compounds consistently lower than the others.

3.Which substance in each pair has the higher

boiling point. Justify your answers.

(a) SiC or KCl

(b) RbBr or C6H12O6

(c) C3H8 or C2H5OH

(d) C4H10 or C2H5Cl

2. Examine the graph on p. 210:

a) Account for the increase in boiling point for

the hydrogen compounds of the Group IV

elements.

b) Why is the trend different for the hydrogen

compounds of the Group V, VI, and VII

elements?

c) Why are the boiling points of the Group IVA

compounds consistently lower than the other

compounds.

Dipole-Dipole Forces

In the liquidstate, polar molecules (dipoles) orient themselves so that oppositely charged ends of the molecules are near to one another.

Summary

The types of bonding/forces ranked from

strongest to weakest are:

Strongest - Network Covalent

- Ionic

- Metallic

Weakest - Covalent

NOTE: To compare covalent compounds

you must use:

- London Dispersion forces

(all molecules)

- Dipole-Dipole forces

(polar molecules)

- Hydrogen Bonding


p. 226 #’s 13 & 14


The electrostatic attractions

between these oppositely charged

ends of the polar molecules are

called dipole-dipole forces.


Results of dipole-dipole attractions:

polar molecules will tend to attract one another more at room temperature than similarly sized non-polar molecules

energy needed to separate polar molecules is therefore higher than for non-polar molecules of similar molar mass


Results of dipole-dipole attractions:

The melting points and boiling

points of substances made of polar

molecules are higher than for

substances made of non-polar

molecules.

Ion-Dipole Forces

An ion-dipole force is the force of

attraction between an ion and a polar

molecule (a dipole).

Ion-Dipole Forces

NaCl dissolves in water because the attractions between the Na+ and Cl- ions and the partial charges on the H2O molecules are strong enough to overcome the forces that bind the ions together.

Induced Intermolecular Forces


Induction of electric charge occurs

when a charge on one object

causes a change in the distribution

of charge on a nearby object. (for

example, the balloon)


There are two types of charge -

induced dipole forces:

1. An ion-induced dipole force

results when an ion in close

proximity to a non-polar molecule

distorts the electron density of the

non-polar molecule


The molecule then becomes momentarily polarized, and the two species are attracted to each other. (ie. hemoglobin)

2. In a dipole-induced dipole forcethe charge on a polar molecule is responsible for inducing the charge on the non-polar molecule.

Dispersion (London) Forces

Bond vibrations, which are part of the normal condition of a non-polar molecule, cause momentary, uneven distribution of charge;

a non-polar becomes slightly polar for an instant, and continues to do so in a random but constant basis.


At the instant that one non-polar

molecule is in a slightly polar

condition, it is capable of inducing

a dipole in a nearby molecule

This force of attraction is called a

dispersion force.


Two factors affecting the magnitude of dispersion forces are:

1. The number of electrons in the molecule:

Vibrations within larger molecules that have more electrons than smaller molecules can easily cause an uneven distribution of charge.


The dispersion forces between these

larger molecules are thus stronger,

which has the effect of raising the

boiling point for larger molecules.

2. The shape of the molecule:

A molecule with a spherical shape has a

smaller surface area than a straight

chain molecule that has the same

number of electrons


Therefore, the substance with molecules that have a more spherical shape will have weaker dispersion forces and a lower boiling point.

London dispersion forces are responsible for the formation and stabilization of the biological membranes surrounding every living cell.

Hydrogen Bonding

In order to form a hydrogen bond, a

hydrogen atom must be bonded to

a highly electronegative atom such

as oxygen, nitrogen, or fluorine.

Hydrogen Bonding

These bonds are very polar, and

since hydrogen has no other

electrons, the positive proton, H+,

is exposed and can become

strongly attracted to the negative

end of another dipole nearby

Hydrogen Bonding

A hydrogen bond is an

electrostatic attraction between the

nucleus of a hydrogen atom,

bonded to fluorine, oxygen, or

nitrogen and the negative end of a

dipole nearby.

δ+δ+

δ+δ−δ−

…

Hydrogen Bonding

In biological systems, these polar

bonds are often parts of much

larger molecules (ie. N H bonds

and H O bonds found in biological

molecules)

Hydrogen Bonding in Water

Hydrogen bonds between the

hydrogen atoms in one water

molecule and the oxygen atom in

another account for many unique

properties of water.

δ+δ+

δ+δ−δ−

…


In liquid water, each water molecule is hydrogen bonded to at least four other water molecules.

The large number of bonds between water molecules makes the net attractive force quite strong


the strong attractive forces are

responsible for the relatively high

boiling point of water.

The water molecules are farther

apart in ice then they are in liquid

water making ice less dense than

liquid water.


Hydrogen bonds

force water

molecules into the

special hexagonal,

crystalline structure

of ice when the

temperature is

below 4 degrees

celcius.

unit 2: chemical bonding -...

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