Unit 4: Chemical Bonds Chapter 7-9

Objectives

26 Identify the number of valence electrons for elements and their Lewis dot structure

27 Define the terms cation and anion including radius size and charge

28 Determine the isoelectronic electron configurations for atoms and their ions including the ionic charges

29 Identify the properties of ionic bonds

30 Predict the shape of molecules using the VSEPR theory

31 Identify the bonds between certain elements within a compound as non-polar, polar, or ionic

32 State and identify the three intermolecular forces including London dispersion forces and how they affect melting points, dipole forces, and hydrogen bond forces

33 Convert between formula and chemical name for covalently bonded molecules, binary ionic compounds, polyatomic ionic compounds, and hydrates

34 Identify the dissociation factor of compounds

26 Valence Shells

Before we discuss bonds, we need to

determine the number of electrons in the

valence shell of an ion.

This is accomplished by looking at the

electron configuration.

Valence Shells

O: 1s2 2s2 2p4

Consider oxygen

It’s outer energy level is the second energy

level.

This tells us there are 6 valance electrons.

Lewis Dot Structures

Once we know the number of valence

electrons, it is possible to give a visual

representation.

This is called a Lewis Dot structure.

It requires a dot for each valence electron

surrounding the element symbol.

Paired electrons should still be depicted.

Lewis Dot Structures

Oxygen has 6 valence electrons.

If we look at the orbital notation for just its valence

shell, we get the following:

2p ___ ___ ___

2s ___

Its Lewis Dot structure would look as follows.

O

Lewis Dot Structures

O

Notice how oxygen has two electrons that are not paired up.

This indicates that oxygen would like to gain two more electrons so it has 8 total electrons.

All elements are most stable with 8 valence electrons.

This is known as the Octet Rule.

◦ There are 5 exceptions: hydrogen, helium, lithium, beryllium, and boron.

◦ These prefer to have 2 electrons in their valence shell.

27 Cations and Anions

Atoms can gain or lose electrons.

When atoms lose electrons, they become

positive and are called cations.

When atoms gain electrons, they become

negative and are called anions.

Ionic Radii

Nuclear charge holds electrons a certain

distance from the nucleus.

As a cation is formed, there are less

electrons for the nucleus to hold.

This allows the nucleus to pull the outer

energy levels slightly closer.

Ionic Radii

As an anion is formed, there are more

electrons for the nucleus to hold.

The nucleus does not have enough charge to

hold the extra electrons as close, and as a

result, the radius increases slightly.

28 Isoelectronic Configurations

As discussed in Unit 3, each atom has an

electron configuration to show where each

electron belongs.

◦ For example: Al: 1s2 2s2 2p6 3s2 3p1

When ions are formed, the electrons are

either added to the last energy levels or

taken from the last energy levels.

◦ For example: assume we take three electrons

from aluminum.

Al+3: 1s2 2s2 2p6 Aluminum’s 3s and 3p orbitals

are now empty.

Isoelectronic Configuration

The term isoelectronic can be broken down

into:

◦ Iso: same

◦ Electronic: electrons

Therefore, the term means the same electron

configuration as another element.

In chemistry, this refers to a noble gas.

Isoelectronic Configuration

If we look at an element, the number of

electrons it holds are close to a noble gas.

This means it will tend to gain or lose

electrons until it matches that noble gas.

◦ For example: Oxygen has 8 electrons and is thus

close to neon’s 10.

◦ Neon has a configuration of 1s2 2s2 2p6 .

◦ Oxygen has a configuration of 1s2 2s2 2p4 .

◦ For oxygen to be isoelectronic with neon, it

needs two more electrons, thus oxygen tends to

gain electrons to get: O-2: 1s2 2s2 2p6

Charges

The elements marked below will always carry the charge indicated.

The elements in white can have charges that vary.

◦ These will be determined with a Roman numeral.

Transition

Metals

29 Ionic Properties

Every ionic compound will follow certain properties.

They are:

◦ Ionic compounds form crystalline structures.

◦ Ionic compounds are brittle.

◦ Ionic compounds have high melting and boiling points.

◦ Ionic compounds as solids will not conduct electricity.

◦ Ionic compounds in solution will conduct electricity.

30 Molecular Shapes

Covalently bonded molecules will create

different shapes.

These shapes are controlled by the bonds

formed and the paired electrons.

To determine the shapes, the valence shell

electron pair repulsion (VSEPR) theory is

used.

VSEPR Theory

VSEPR Theory states that the valence

electrons in a molecule will position

themselves so they are as far away from the

other electrons as possible.

To determine the shapes of molecules using

this theory, the number of bonds and pairs

of electrons must be determined.

VSEPR Theory

Recall Lewis Dot structures from the

previous unit.

◦ Each dot represents a valence electron.

Oxygen has six valence electrons so its

Lewis Dot structure is as follows:

O These paired electrons are

called lone pairs because they

belong to only one atom.

VSEPR Theory

The Lewis Dot structures can be used to

determine the bonds created.

Take water which is H2O.

O H O H

H

H When a bond is formed, the

electrons are shown with a

single line between the

atoms.

VSEPR Theory

Notice that the water molecule still contains two

sets of lone pair electrons.

These electrons will force the hydrogens to create

a bent shape.

O H

H

VSEPR Theory

There are five shapes the basic covalent

molecules will create.

Each can be determined by looking at the

central atom of the molecule.

The two components to look at are the

number of atoms bound and the number of

lone pairs on the central atom.

VSEPR Theory

Shape Lone Pairs of

electrons on

Central Atom

Atoms bonded

to the Central

Atom

Example

Linear 0 2 CO2

Bent 1 or 2 2 H2O or HNO

Trigonal Planar 0 3 BF3

Trigonal Pyramidal 1 3 NH3

Tetrahedral 0 4 CH4

In addition, any molecule that has only two atoms will be linear.

i.e.: oxygen gas, O2

31 Intramolecular Bonds

When a bond is created, it is often stated

that the atoms share their electrons.

While this is true to some degree, the

sharing is not always equal.

Each atom has its own electronegativity.

◦ The tendency of an atom to attract a bonded

electron to itself.

◦ The greater an atom’s electronegativity, the more

time the electrons will spend near that atom.

Bond Strength

This pulling of the electrons towards one

atom can create partial charges.

Each bond can be classified as either

nonpolar covalent, polar covalent, or ionic.

To determine how atoms share, compare

their electronegativity values.

◦ The greater the difference, the more ionic

character will be present in the molecule.

Intramolecular Bonds

Nonpolar Polar Covalent Ionic

Covalent

0 0.5 2.1

Difference in Electronegativity

As we look at the difference in electronegativity, we can use the chart

above to determine the type of bond.

Take water for example: H2O

◦ The bonds formed are between hydrogen and oxygen.

◦ They have the following electronegativities: 2.2 for H and 3.4 for O

◦ The difference between the two is 1.2 and thus the bond is polar covalent.

◦ Because oxygen has the larger electronegativity, the electrons spend more time near

oxygen. This creates a partial negative charge on the oxygen atom and a partial

positive charge around the hydrogen atom.

32 Intermolecular Forces

While intramolecular forces occur inside

a molecule, intermolecular forces occur

between molecules.

Three intermolecular forces exist:

◦ London Forces

◦ Dipole Forces

◦ Hydrogen bonding

Intermolecular Force Strength

Each of the intermolecular forces hold

molecules together.

However, certain forces are stronger.

◦ London is the weakest.

◦ Dipoles use the partial charges as an attractive

force making them stronger than London.

◦ H-Bonding is the strongest because of the

partial charges and the use of hydrogen.

Determining Intermolecular Forces

To determine the intermolecular force on a

molecule, it is necessary to know whether it

is polar or not.

◦ If polar, the molecule will have a partial positive

and partial negative side.

◦ This can be determined using the

electronegativities and the Lewis Dot structure.

Determining Intermolecular Forces

Water has a Lewis Dot structure as shown.

Oxygen has an electronegativity of 3.4 while hydrogen has an electronegativity of 2.2.

◦ This would mean that oxygen is partial negative and the hydrogens are each partial positive.

Because the molecule has a postive and negative side, it is considered polar.

If the molecule would have had the same charge on the entire outside, it would be considered non polar.

Intermolecular Forces

Force Strength Polar/Nonpolar Unique

Characterestics

H-Bonding Strongest Polar Must contain H

and either O, N, F

Dipole Forces

Medium Polar

London Forces

Weak Nonpolar

Dipole forces and H-Bonding are the only forces that are polar, but H-bonding

has element requirements. If a molecule is polar but does not contain one of

the elements listed above, it must be a dipole.

33 Writing Binary Formulas and

Names

Binary compounds refer to compounds that

contain 2 elements.

When writing the name of a binary

compound, list the first element exactly as it

appears on the Periodic Table.

For the second element, drop the ending of

the element’s name and add –ide.

Binary Names-Examples

NaCl

sodium chloride

(chlorine drops the –ine)

CaBr2

calcium bromide

(bromine drops the –ine)

Writing Binary Formulas and

Names

Writing the formulas from the names

requires the use of the charges.

It is important to balance the positive charge

with the negative charge.

This is done by adding subscripts to the

elements.

Binary Formulas - Examples

calcium phosphide

Calcium has a +2 charge.

Phosphorus has a -3 charge

To balance their charges, we need to have multiple atoms.

If we add another calcium, we will have an overall charge of +4 to -3.

Let’s add another phosphorus to give us a overall charge of +4 to -6.

Since we are only off by 2, adding another calcium will balance the charges at +6 to -6.

Ca3P2

Writing Formulas

We got to the answer on the last slide by using logical method of adding one atom at a time.

This can also be done by looking for the least common multiple.

Since we had charges of 2 and 3, the least common multiple is 6.

◦ (2 x 3 = 6)

◦ (3 x 2 = 6)

Therefore, the atom with the charge of 2 requires 3 atoms and the atom with the charge of 3 requires 2.

Ca3P2

Writing Formulas

magnesium bromide

MgBr2

Magnesium has a +2 charge.

Bromine has a -1 charge.

The least common multiple is two.

◦ 1 x 2 = 2

◦ 2 x 1 = 2

Transition Metals

If you recall from the slide on charges,

transition metals did not have a defined

charge.

Their charges vary and thus a Roman

numeral is used to determine their charge.

This Roman numeral is always listed directly

after the metal.

Transition Metals

Fe2O3

For the formula above, we know Fe is iron and O is oxide (oxygen as an ion)

Oxygen has a -2 charge and since there are 3, this compound has a overall -6 charge.

Since we have to have a +6 charge as well, we have to consider a number times 2 to give 6. (? x 2 =6)

In this case, the answer would be three.

Therefore, the name of the compound is iron (iii) oxide.

Transition Metals

Manganese (ii) chloride

In this compound, Manganese has a +2

charge and chlorine has a -1 charge.

Therefore, the least common multiple is 2.

MnCl2

Polyatomic Ions

Some elements will combine covalently

(Unit 6) and still have a charge.

As long as they have a charge, they can

create ionic compounds.

They are treated as though they are single

entities.

Naming the Polyatomics

From a formula, naming the polyatomic ionic

compounds requires element to be named

and the polyatomic ion.

For instance:

CaSO4

◦ Ca represents calcium

◦ SO4 represents sulfate

The name of this compound is calcium sulfate.

Writing formulas

Writing formulas from the name works the

same as binary compounds.

◦ Determine the charge on each.

◦ Find the least common multiple.

◦ Add the proper subscripts.

If more than one polyatomic ion is required,

add parenthesis around the ion before

adding the subscript.

Writing formulas-Example

Iron (iii) nitrate

Iron (iii) refers to Fe+3

Nitrate refers to NO3-1

Therefore, the least common multiple is 3

and three nitrates are required.

Fe(NO3)3

The parenthesis tells us that there are 3 N and 9 O.

Hydrates

Hydrates are unique ionic compounds that

attract water.

Each hydrate is surrounded by a certain

number of water molecules.

These water molecules need to be identified

in both the formula and the name.

Hydrates

To indicate a hydrate, a dot is used to indicate a weak bond.

The number of water molecules are indicated with a numeric prefix.

The ionic part of the compound is named as previously described.

For example: copper (ii) sulfate pentahydrate

CuSO4 • 5H2O

Molecular Nomenclature

When naming covalent molecules, first

identify each element.

◦ If there is more than one of the first element,

add the appropriate prefix to the front of its

name.

The second element should always includes

its prefix and its ending should change

to –ide.

Molecular Nomenclature

CO2

To name this compound, first identify each element.

Since there is only one of the first element, no prefix is needed.

There are two of the second element so the prefix di- will be added. Notice the ending of oxygen has already be changed.

carbon oxide di

Molecular Nomenclature

Writing the formulas will simply work in the

opposite direction.

Identify and write the symbol for each

element.

Use the prefixes to determine the subscript

of each element.

Molecular Nomenclature

Tetraphosphorus decoxide

First, record the symbols for each element.

Tetra- indicates four so the subscript on phosphorus will be four.

Deca- means ten so the subscript on oxide will be ten

P O P4O10

34 Dissociation Factors

Dissociation factors describe how many

pieces an ionic compound can divide into.

This is calculated by adding the subscripts of

each ion.

Be careful because the one’s are omitted

when writing formulas.

Dissociation Factors

Assume we have calcium chloride:

CaCl2

If this molecule breaks apart, we will have 1

calcium ion and 2 chloride ions.

This means the dissociation factor is 3.

Dissociation Factors

The same idea applies to polyatomics:

Ca(NO3)2

If this molecule breaks apart, we will have 1

calcium ion and 2 nitrate ions.

This means the dissociation factor is 3.

The polyatomic ions do not break apart.

This concludes the tutorial on

measurements.

To try some practice problems, click here.

To return to the objective page, click

here.

To exit the tutorial, hit escape.

Prefixes

mono 1

di 2

tri 3

tetra 4

penta 5

hexa 6

hepta 7

octa 8

nona 9

deca 10

Return

Definitions-Select the word to return to the tutorial

Valence Shell: Outer most energy level of an

atom