Chapter 4

Structure and Properties of Ionic and Covalent Compounds

Denniston Topping Caret

4th Edition

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

4.1 Chemical Bonding

• Chemical Bond - the force of attraction between any two atoms in a compound.

• Interactions involving valence electrons are responsible for the chemical bond.

4.1

Che

mic

al B

ondi

ngLewis Symbols

• Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and valence electrons as dots.

4.1

Che

mic

al B

ondi

ng Principal Types of Chemical Bonds: Ionic and Covalent

• Ionic bond - a transfer of one or more electrons from one atom to another.

• forms attractions due to the opposite charges of the atoms.

• Covalent bond - attractive force due to the sharing of electrons between atoms.

1

4.1

Che

mic

al B

ondi

ng Let’s examine the formation of NaClNa + Cl NaCl

IONIC BONDING

Sodium has alow ionization energyit readily loses this electron .

When Sodium loses the electron, it gains the Ne configuration.

Na Na+ + e-

Chlorine has a high electron affinity.

When chlorine gains an electron, it gains the Ar configuration.

:

..

..Cl: e

..

..Cl:

4.1

Che

mic

al B

ondi

ngEssential Features of Ionic Bonding

• Atoms with low I.E. and low E.A. tend to form positive ions.

• Atoms with high I.E. and high E.A. tend to form negative ions.

• Ion formation takes place by electron transfer.

• The ions are held together by the electrostatic force of the opposite charges.

• Reactions between metals and nonmetals (representative) tend to be ionic.

4.1

Che

mic

al B

ondi

ngCOVALENT BONDING

Let’s look at the formation of H2

H + H H2

• Each hydrogen has one electron in it’s valance shell.

• If it were an ionic bond it would look like this:

:H H H H

• However, both hydrogen atoms have the same tendency to gain or lose electrons.

• Both gain and loss will not occur.

4.1

Che

mic

al B

ondi

ng • Instead, each atom gets a noble gas configuration by sharing electrons.

H:H H H

The shared electrons pair is a

Covalent Bond

Each Hydrogen atom now has two electrons around it and has a He configuration

4.1

Che

mic

al B

ondi

ngFeatures of Covalent Bonds

• Covalent bonds tend to form between atoms with similar tendency to gain or lose electrons.

• The diatomic elements have totally covalent bonds (totally equal sharing.)

H2, N2, O2, F2, Cl2, Br2, I2

:..

..F:

..

..F: :

..

..F

..

..F:

Each fluorine has eight electrons around it. Ne’s configuration.

HOFBrINCl

4.1

Che

mic

al B

ondi

ngOther examples:

H:..

..O:H

..

..O 2H

H H:

..C:H

.C 4H

H

4.1

Che

mic

al B

ondi

ngPolar Covalent Bonding and

Electronegativity

The Polar Covalent Bond

• Ionic bonding involves the transfer of electrons.

• Covalent bonding involves the sharing of electrons.

• Polar covalent bonding - bonds made up of unequally shared electron pairs.

1

4.1

Che

mic

al B

ondi

ng :..F:H :

..F H

These two electrons

are not sharedequally.

• The electrons spend more time with fluorine.

• This sets up a polar bond • A truly covalent bond can only occur

when both atoms are identical.• Electronegativity is used to determine if

a bond is polar and who gets the electrons the most.

somewhat negatively chargedsomewhat positively charged

4.1

Che

mic

al B

ondi

ng • Electronegativity - a measure of the ability of an atom to attract electrons in a chemical bond.

electronegativity increases elec

tron

egat

ivit

y in

crea

ses

4.1

Che

mic

al B

ondi

ng• The greater the difference in

electronegativity between two atoms, the greater the polarity of a bond.

• Which would be more polar, a H-F bond or a H-Cl bond?

• H-F …4.0 - 2.1 = 1.9

• H-Cl… 3.0 - 2.1 = 0.9

• Therefore, the HF bond is more polar than the HCl bond.

4.2 Naming Compounds and Writing Formulas of Compounds

• Nomenclature - the assignment of a correct and unambiguous name to each and every chemical compound.

• We will learn two systems– one for naming ionic compounds and – one for naming covalent compounds.

Sec

tion

4.2

I. Ionic Compounds

A. Writing Formulas of Ionic Compounds from the Identities of the Component Ions.

B. Writing Names of Ionic Compounds from the Formula of the Compound

C. Writing Formulas of Ionic Compounds from the Name of the Compound

IV. Covalent Compounds

A. Naming Covalent Compounds

B. Writing Formulas of Covalent Compounds

Outline of Nomenclature Topics

Sec

tion

4.2

I. Ionic Compounds

• Metals and nonmetals usually react to form ionic compounds.

• The metals are the cations and the nonmetals are the anions.

• The cations and anions arrange themselves in a regular three-dimensional repeating array called a crystal lattice.

Sec

tion

4.2

• Here is the crystal lattice of sodium chloride (table salt.)

• Formula - the smallest whole number ratio of ions in the crystal.

• The formula for sodium chloride is NaC.

Sec

tion

4.2

A. Writing Formulas of Ionic Compounds

• Determine the charge of the ions (usually can be obtained from the group number.)

• Cations and anions must combine to give a formula with a net charge of zero

• It must have the same number of positive charges as negative charges.

Sec

tion

4.2 Predict the formula of the ionic

compounds formed from the following combination of ions:

1. sodium and sulfur

2. magnesium and oxygen

3. aluminum and oxygen

4. barium and fluorine

Sec

tion

4.2

B. Writing Names of Ionic Compounds from the Formula

Stock System:

1. Name cation followed by the name of anion.

2. Give anion the suffix -ide.

• Examples:

• NaCl is sodium chloride.

• AlBr3 is aluminum bromide.

2

Sec

tion

4.2

3. If the cation of an element has several ions of different charges (as with Transition metals) use a Roman numeral after the metal name.

• Roman numeral gives the charge of the metal.

• Examples:

• FeCl3 is iron(III) chloride

• FeCl2 is iron(II) chloride

• CuO is copper(II) oxide

Sec

tion

4.2

Common Nomenclature System

• Use -ic to indicate the higher of the charges.

• Use -ous to indicate the lower of the charges.

• Examples:

• FeCl2 is ferrous chloride

• FeCl3 is ferric chloride

• Cu2O is cuprous oxide

• CuO is cupric oxide

Sec

tion

4.2

• Note that the common system requires you to know the common charges of the metals and use the Latin names of the metals.

• Monatomic ions - ions consisting of a single atom.

• Examples• K+ potassium ion

• Ba2+ barium ion

• O2- oxide ion

• Table 4.2 gives common monatomic cations and anions.

Sec

tion

4.2

• Polyatomic ions - ions composed of 2 or more atoms bonded together.

• Within the ion, the atoms are bonded using covalent bonds.

• The ions will be bonded to other ions with ionic bonds.

• Examples:

• NH4+ ammonium ion

• SO42- sulfate ion

Sec

tion

4.2

Table 4.3 Common Polyatomic Cations and Anions

Ion Name

NH4+ ammonium

NO3- nitrate

SO42- sulfate

OH- hydroxideCN- cyanidePO4

3- phosphateCO3

2- carbonateHCO3

- bicarbonateCH3COO- (or C2H3O2

-) acetate

Table 4.3 in your textbook gives a more complete list of polyatomic ion.

Sec

tion

4.2

Name the following compounds using the Stock System

1. NH4Cl

2. BaSO4

3. Fe(NO3)3

4. CuHCO3

5. Ca(OH)2

Sec

tion

4.2

C. Writing Formulas of Ionic Compounds from the Name

• Determine the charge on the ions.

• Write the formula so the compounds are neutral.

• Example:Barium chloride:

Barium is +2, Chloride is -1

Formula is BaCl2

2

Sec

tion

4.2

Give the formula for the following ionic compounds:

1. sodium sulfide

2. copper(I) carbonate

3. magnesium phosphate

4. chromium(II) sulfate

Sec

tion

4.2

II. Covalent Compounds

• Covalent compounds are usually formed from nonmetals.

• Molecules - compounds characterized by covalent bonding.

• not a part of a massive three dimensional crystal structure.

Sec

tion

4.2

A. Naming Covalent Compounds

1. The names of the elements are written in the order in which they appear in the formula.

2. A prefix indicates the number of each kind of atom• mono- 1 hexa- 6

• di- 2 hepta-7

• tri- 3 octa-8

• tetra- 4 nona-9

• penta- 5 deca- 10

3

Sec

tion

4.2

3. If only one atom of a particular kind is present in the molecule, the prefix mono- is usually omitted from the first element.

Example: CO is carbon monoxide

4. The stem of the name of the last element is used with the suffix -ide

5. The final vowel in a prefix is often dropped before a vowel in the stem name.

Sec

tion

4.2

Name the following covalent compounds

1. SiO2

2. N2O5

3. CCl4

4. IF7

Sec

tion

4.2

B. Writing Formulas of Covalent Compounds

• Use the prefixes in the names to determine the subscripts for the elements.

• Examples:• nitrogen trichloride:

• NCl3

• diphosphorus pentoxide

• P2O5

• Some common names that are used:• H2O water

• NH3 ammonia

• C2H5OH ethanol

• C6H12O6 glucose

2

4.3 Properties of Ionic and Covalent Compounds

• Physical State– Ionic compounds are solids at room

temperature– Covalent compounds are solids, liquids and

gases

• Melting and Boiling Points– melting point - the temperature at which a

solid is converted to a liquid– boiling point - the temperature at which a

liquid is converted to a gas

4

4.3

Pro

per

ties

• Melting and Boiling Points– Ionic compounds have much higher melting

points and boiling points than covalent compounds.

– Due to the large amount of energy required to break the electrostatic attractions between ions.

– Ionic compounds typically melt at several hundred degrees Celsius.

• Structure of Compounds in the Solid State– Ionic compounds are crystalline– Covalent compounds are crystalline or

amorphous - have no regular structure.

4.3

Pro

per

ties

• Solutions of Ionic and Covalent Compounds– Ionic compounds often dissolve in water,

when they do they dissociate - form positive and negative ions in solution.

– Electrolytes -ions present in solution allow the solution to conduct electricity.

– Covalent solids usually do not dissociate and do not conduct electricity - nonelectrolytes

4.4 Drawing Lewis Structures on Molecules and Polyatomic Ions

Lewis Structure Guidelines

1.Use chemical symbols for the various elements to write the skeletal structure of the compound.– the least electronegative atom will be placed in

the central position,– hydrogen and fluorine occupy terminal positions,– carbon often forms chains of carbon-carbon

covalent bonds.

5

4.4

Dra

win

g L

ewis

S

tru

ctu

res

2.Determine the total number of valence electrons associated with each atom in the compound.– for polyatomic cations, subtract one electron for

every positive charge;– for polyatomic anions, add one electron for

every negative charge.

3.Connect the central atom to each of the surrounding atoms using electron pairs. Then give each atom an octet.– Remember, hydrogen needs only two electrons

4. Count the number of electrons you have and compare to the number you used.

• If they are the same, you are finished.• If you used more electrons than you

have add a bond for every two too many you used. Then give every atom an octet.

• If you used less electrons than you have….(see later when discuss exceptions to the octet rule)

5. Check that all atoms have the octet rule satisfied and that the total number of valance electrons are used.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Using the guidelines presented, write Lewis structures for the following:

1. H2O

2. NH3

3. CO2

4. NH4+

5. CO32-

6. N2

4.4

Dra

win

g L

ewis

S

tru

ctu

res

• Single bond - one pair of electrons are shared between two atoms

• Double bond - two pairs of electrons are shared between two atoms

• Triple bond - three pairs of electrons are shared between two atoms

6

Lewis Structure, Stability, Multiple Bonds, and Bond Energies

NNor ..N

..N

OOor ..O::

..O H - Hor H:H

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Bond energy - the amount of energy required to break a bond holding two atoms together.

triple bond > double bond > single bond

Bond length - the distance separating the nuclei of two adjacent atoms.

single bond > double bond > triple bond

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Lewis Structures and Resonance

• Write the Lewis structure at CO32- on

your paper.

• If you look at the people around you they probably put the double bond in different places.

• Who is right?

• You all are.

• Experimental evidence shows all bonds are the same length. Shorter than single bond but longer than a double bond.

• None of the three exist but the actual structure is an average or hybrid of these three Lewis structures.

• Resonance - two or more Lewis structures that contribute to the real structure.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

:..O:C::

..O: :

..O:C:

..O: O::C:

..O:

: :: : :

..O: :O: :

..O:

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Lewis Structures and Exceptions to the Octet Rule

1. Incomplete Octet - less then eight electrons around an atom other than H.

• Let’s look at BF3

2. Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight electrons.• Let’s look at NO

3. Expanded Octet - elements in 3rd period and below may have 10 and 12 electrons around it.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

• Expanded octet is the most common exception.

• Write the Lewis structures of SF6 and SF4

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Lewis Structures and Molecular Geometry: VSEPR Theory

• VSEPR theory - valance shell electron pair repulsion theory.

• This is used to predict the shape of the molecules.

• Electrons around the central atom (both bonding and nonbonding pairs) arrange themselves so they can be as far away from each other as possible.

7

4.4

Dra

win

g L

ewis

S

tru

ctu

res

• We will do several examples in order to see how the geometry is determined.

• Let’s do BeH2 on the overhead.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

• Let’s do BF3 on the overhead.

• Let’s do CH4 on the overhead

4.4

Dra

win

g L

ewis

S

tru

ctu

res

• Let’s do NH3 on the overhead

• Let’s do H2O on the overhead

Trigonal Pyramidal

Bent

4.4

Dra

win

g L

ewis

S

tru

ctu

res

The basic procedure to follow to determine the shape:

1. Write the Lewis structure.

2. Count the number of bonded atoms and non-bonded electrons around the central atom.• 2 legs - linear

• 3 legs - trigonal planer

• 4 legs - tetrahedron

3. Look at the atoms and name the shape.• These would include: linear, trigonal planer,

bent, trigonal pyramid, tetrahedon.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

PCl3

SO2

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Lewis Structures and Polarity

• Polar molecules when placed in an electric field will align themselves in the field.

• Molecules that are polar behave as a dipole (two “poles”)

• One end is positively charged the other is negatively charged.

• Nonpolar molecules will not align themselves in an electric field.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

To determine if a molecule is polar or has a dipole:

• Write the Lewis structure.

• Draw the Geometry.

• Use the following symbol to denote the polarity of each bond.

This end of the arrow points to the negative end of the bond.

The more electronegative atom attracts the electrons more strongly towards it.

This end of the arrow points to the positive end of the bond.

The less electronegative atom.

4.4

Dra

win

g L

ewis

S

tru

ctu

res

Determine the polarity of the following molecules:

1. O2

2. HF

3. CH4

4. H2O

4.5 Properties Based on Molecular Geometry

• Intramolecular forces - attractive forces within molecules. (Chemical bonds)

• Intermolecular forces - attractive forces between molecules.

• We will look at how intermolecular forces affect:– I. Solubility.– II. Boiling points and melting points.

8

9

I. Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature.

• “Like dissolves like” – polar molecules are most soluble in polar

solvents– nonpolar molecules are most soluble in

nonpolar solvents.

• Will ammonia (NH3) dissolve in water?

• Yes, both molecules are polar.

4.5

Pro

per

ties

Bas

ed o

n

Mol

ecu

lar

Geo

met

ry

4.5

Pro

per

ties

Bas

ed o

n

Mol

ecu

lar

Geo

met

ry

4.5

Pro

per

ties

Bas

ed o

n

Mol

ecu

lar

Geo

met

ry

• What do you know about oil and water?

• “They don’t mix”

• Why?

• Because water is polar and oil in nonpolar.

4.5

Pro

per

ties

Bas

ed o

n

Mol

ecu

lar

Geo

met

ry II. Boiling Points and Melting Points

• The greater the intermolecular force the higher the melting point and boiling point.

• Two factors to consider:

• Larger molecules have higher m.p. and b.p. than smaller molecules.

• Polar molecules have higher m.p. and b.p. than nonpolar molecules (of similar molecular mass.)

The End Chapter 4