1

Ionic Bonding: Bonding that results from the electrical attraction between positive and negative ions.

Ionic Bonding will be assumed to occur when metals react with non-metals.

When a metal reacts with a non-metal, the metal will lose electrons to become a positive ion (cation) and the non-metal will gain electrons to become a negative ion (anion).

Formation of an Ionic Bond involves the transfer of electrons (from a metal atom to a non-metal atom).

Substances that contain ionic bonds usually form crystalline solids (like a grain of salt) and the chemical formula is called a formula unit.

2

Covalent Bonding: Bonding that results from the sharing of electron pairs between two atoms (the nucleus of each atom is attracting the same pair of electrons).

Covalent Bonding will occur when non-metals react with non-metals.

When a non-metal reacts with a non-metal, each atom will allow one or more of its electrons to be shared with the other atom (the electron will be in an orbital that is located between the two nuclei).

Formation of a Covalent Bond involves the sharing of electrons (between two non-metal atoms).

Substances that contain covalent bonds form molecular compounds which can be gases, liquids, or solids and these formulas are called molecules.

3

Metallic Bonding: Bonding that results from the nuclei of metal atoms being attracted to a “sea” of delocalized electrons (electrons that are temporarily released from a number of metal atoms)

Metallic Bonding will occur when metal atoms bond with each other or with atoms of other metals.

The delocalized electrons can move from one atom to another in the metal-this is why metals can conduct electricity and heat so well.

Formation of a Metallic Bond involves the sharing of electrons (between two or more metal atoms).

4

Technically, we use differences in electronegativity to determine whether a bond is covalent or ionic. If the difference in electronegativity between two atoms is greater than 1.7, the bond will be called ionic. If the difference is 1.7 or less, the bond will be called covalent.

Unless a question specifically mentions differences in electronegativity, we will assume that a bond between a metal and an non-metal will be ionic while a bond between two non-metals will be covalent.

5

Classify each of the following bonds as being ionic or covalent using electronegativity values from figure 3.11 on page 153.

CO KCl

NF CaBr

HC ClCl

LiO BaS

3.5 – 2.5 = 1.0

Therefore: Covalent

When you finish, use the metal-nonmetal rule to see if your answers agree.

6

We also use differences in electronegativity to determine whether a covalent bond is polar or non-polar.

A polar covalent bond has unequal sharing of the electrons between the atoms. This causes a small amount of charge to form on each atom in the bond.

A small amount of negative charge will form on the atom that is more electronegative, and a small amount of positive charge will form on the less electronegative atom.

7

Classify each of the following bonds as being polar covalent or non-polar covalent using electronegativity values from figure 3.11 on page 153.

CO SiH

NF ClBr

HC ClCl

CI ClS

3.5 – 2.5 = 1.0

Therefore: Polar Covalent

Use the electronegativity values to predict which atom in each bond would be a little bit negative and which would be a little bit positive.

8

Chemical formulas show the type and number of atoms present in one molecule (covalent) or formula unit (ionic) of a substance.

CH2Cl2 has 1 C atom, 2 H atoms and 2 Cl atoms in it

K2S has 2 K+1 ions and 1 S2 ion in it

For each formula below, determine how many atoms or ions of each element are present.

KNO3

P2O5

Cr(NO3)6

(NH4)2SO4

Fe2(SO3)3

Co(S2O3)3

P3,4,5 Start

9

(a) The Interaction of Two Hydrogen Atoms

(b) Energy Profile as a Function of the Distance Between the Nuclei of the Hydrogen Atoms

Energy of a Covalent Bond

The Energy change from zero to the lowest part of the curve is the energy released when the two hydrogen atoms form a bond.

10

Energy of a Covalent Bond

Bond Energy (or Bond Disassociation Energy) is the energy required to break a bond (a positive number)

However, when to atoms come together to make a bond, energy will be released (a negative number) and the amount of energy released will be equal to the amount of energy that would be required to break that bond.

In general, the stronger the bond (higher bond energy) the shorter the bond length is. Bond length is the distance between two nuclei in a bond.

Endothermic: a process where energy flows into the system (the system has more energy at the end than at the beginning)-positive energy change

Exothermic: a process where energy flows out of a system (the system has less energy at the end than at the beginning)-negative energy change

11

Bond Strength depends upon distance between atoms and the number of electrons being shared between the two atoms.

Smaller atoms can form stronger bonds than larger atoms. So, shorter bonds are generally stronger bonds.

The more bonds between two atoms, the stronger the overall bond is. So, a double bond is stronger than a single bond and a triple bonds is stronger than a double bond as long as the atoms involved are about the same size.

One measure of bond strength is the amount of energy required to break the bonds-this is called the bond dissociation energy.

P2 start

12

Bond breaking is always endothermic

Endothermic means that energy is taken in as the process happens. In other words, we must add energy to make the process happen.

Bond making is always exothermic

Exothermic means that energy is given off (exits) as the process happens.

An endothermic process would feel cold as it happened while an exothermic process would feel hot as it happened.

13

Using the bond lengths given, predict which bond is stronger in each pair.

CO CN

143 pm 147 pm

BrBr ClCl

229 pm 199 pm

CC CH

154 pm 109 pm

HF HCl

92 pm 127 pm

14

Review of Lewis DOT Diagrams: A diagram where the valence electrons of an atom are shown as “dots” located around the element symbol

· O ·..

..

Remember that the number of “dots” is equal to the total of the “s” and “p” electrons in the valence shell of the atom.

For example: Oxygen has an electron configuration of: 1s2, 2s2, 2p4

Since the second period is the valence shell, oxygen would haveSix dots (2s2, 2p4 = 2 + 4 = 6).

Practice: write dot diagrams for each of the following elements:

H, B, C, N, S, Cl, Xe

15

Octet Rule: when atoms form covalent bonds, atoms will share enough electrons to make each atom “feel” that it has an octet of electrons (like a noble gas configuration).

Most atoms wants to have 8 valance electrons!

Predicting the number of bonds for period 2 non-metals.

8 – (# of Valence e1) = (# of e1 needed)

The (# of e1 needed) is also the # of bonds the atom will usually form

C N O F

8 – 4 = 4 8 – 5 = 3 8 – 6 = 2 8 – 7 = 1

Based on this analysis, we would predict that C will form 4 bonds when it bonds to other atoms and that nitrogen will form 3 bonds and so on for the other atoms of period 2.

16

Since Hydrogen is in the 1st period, it will only be able to “feel” like it has 2 electrons (like the noble gas He). We can call this the Duet Rule: hydrogen will share enough electrons to make each hydrogen atom “feel” that it has a duet of electrons (like the noble gas configuration of He).

Bonding in Water (H2O)H · · O · · H

..

..

If oxygen shares one electron with each hydrogen, two bonds will be formed-one each between the oxygen and each hydrogen

H O H..

..

Counting electrons that each atoms “feels” that it has around itself:

each H the O

2 8

17

H O H..

..

The solid lines represents single bonds (sigma, bonds) and

each bond has 2 electrons in it.

The pairs of dots are called “lone pairs” and represent 2 electrons held in an orbital.

This type of drawing of water is called a Lewis Structure.

18

Lewis Structure Practice:

NH3 CBr4 SF2

PCl3 SiH4 NF3

19

Multiple Bonds between two atoms: double and triple bonds

Many times, it becomes necessary for two atoms to share more than one pair of electrons between the two atoms. This results in multiple bonds.

Two bonds makes a double bond.

Three bonds makes a triple bond.

In all cases, the first bond between two atoms is called a sigma bond () and all additional bonds are called pi bonds ().

· O ·..

..· O ·

..

..· O

..

..O ·..

..1st

bond O

..

..O..

..2nd

bond

neitherO atomhas “8”

BothO atomshave “8”

Double bond

20

A bond is “end-to-end” overlap of hybrid orbitals (the single bond)

A bond is “edge-to-edge” overlap of “p” orbitals (a double bond)

=C CThis would be a carbon-carbon double bond

21

If atoms in a compound can not satisfy the octet rule by forming only sigma (single) bonds, it will form pi bonds (double or triple bonds) if possible.

Example: CO2 Total valance electrons needed: 4 from C and 12 from O = 16

· C ·.

.· O ·

..

..· O ·

..

..

C .

.O ·

..

..· O

..

..

Arranging atoms with the atom that needs the most bonds in the center:

Form single bonds to connect all the atoms together:

Notice that C “thinks” it has 6 and each O “thinks” it has 7. Neither atom has met the octet rule yet.

C O..

..O ..

..

Use double bonds to get each atom to satisfy the octet rule:

Notice that C “thinks” it has 8 and each O “thinks” it has 8. Both atoms have met the octet rule.

16 Valance electrons present

Still only 16 Valance electrons present

22

C O O

C O..

..O ..

..

We can come up with the same Lewis structure for CO2 using a different approach.

C 1 * 4 = 4O 2 * 6 = 12

Total number of electrons needed in the Lewis structure: 16

Place carbon in the center (only 1 carbon atom, but 2 oxygen atoms) and attach both oxygen atoms to the carbon atom.

Try to satisfy the octet rule for all atoms by placing the 12 remaining electrons as lone pairs around the atoms.

::..

..

..

..C O O

The two bonds use 4 electrons so there are 12 remaining.

Both oxygen atoms have 8, but carbon only has 4. This is an indication that multiple bonds will be needed.

Remove 1 lone pair from each oxygen and replace them with a bond between the C and the O. This will give all atoms 8 electrons and we will use all 16 of the valence electrons we started with.

23

Lewis structure practice (compounds that require a double or triple bond)

C2H4 CS2

N2H2 SO2

24

Making Lewis Structures:1. Determine the total number of electrons that must appear in the structure.

All valence electrons from each atom must appear.

2. Place the atoms in a reasonable pattern-usually the element that has only 1 atom present will go in the center (except H-it can never go in the center).

In general, the atom that is present in the least number will be in the “middle” and other atoms will be attached to it.

3. Connect all atoms with single bonds.

4. Use the remaining electrons to try to make each atom achieve an octet by placing lone pairs of electrons around the outer atoms (duet for Hydrogen). If you run out of valence electrons before every atom has an octet, this is an indication that double or triple bonds are needed.

5. Remove one lone pair (at a time) from an outer atom and draw a double bond between the outer atom and the central atom. Repeat this as needed until all atoms have 8.

25

Most of the non-metal atoms in period 2 of the periodic table will form enough bonds to allow the atom to “feel” that it has 8 valance electrons when the bonding is complete.

There are Exceptions to the Octet Rule:

Boron is an exception and can not usually obtain an octet of electrons (8 valance electrons). Boron will often end up with 6.

Hydrogen (a period 1 element) can never have more than 2 electrons so it will always form only 1 bond when it bonds to another atom.This is the duet rule!

Non-metal atoms in the 3rd period or higher (periods 3 through 6) can actually obtain more than 8 valance electrons by using their “d” orbitals. They will expand their octet to reduce “formal charge”.

26

Formal Charge

“formal charge” can be used to determine which atoms in a polyatomic ion have charge

Formal charge = valance electrons of the atom originally minus the electrons it “owns” in the structure

Note: in a Lewis structure, an atom “owns” lone pair electrons and ½ of the bonding electrons to the atom

O

O

OC..

:

..

.. ..

..::

1

1

Example:

The oxygen with a double bond has two lone pairs (4 e1)and two bonds (4 e1). It “owns” the lone pair electrons and ½ of the bonding electrons ½*4 = 2. Therefore, the oxygen with the double bond has a formal charge equal to 6 – 6 = 0.

Oxygen always has 6 valance electrons

Each oxygen with single bonds has three lone pairs (6 e1) and one bond (2 e1). It “owns” the lone pair electrons and ½ of the bonding electrons ½*2 = 1. Therefore, the oxygen with the single bond has a formal charge equal to 6 – 7 = 1.

Note: any atom in its “normal” bonding pattern will always have a formal charge of zero!

27

An atom will expand its octet in an attempt to avoid a “formal charge”.

Formal Charge is a bookkeeping method that compares the number of electrons the atom is supposed to “own” to the number of electrons it actually “owns” in the diagram.

The formula for formal charge is: (alternate explanation)

(valence e) – ½(bonding e) – (# of e in lone pairs)

O

O

O

OP ..

:..

..

:..

..: :

..

: :

For P: 5 – ½(8) – 0 = +1

For each O: 6 – ½(2) – 6 = 1

The sum of all formal charges must equal the charge of the ion or compound!

3

(+1) + (1) + (1) + (1) + (1) = 3

28

O

O

O

OP ..

:..

..

:..

..: :

: :

For P: 5 – ½(10) – 0 = 0

For 3 of the O: 6 – ½(2) – 6 = 1

3

(0) + (1) + (1) + (1) + (0) = 3

Now if P expands its octet:

For the double bonded O: 6 – ½(4) – 4 = 0

This structure is “better” than the previous onebecause fewer atoms have a formal charge that is not zero.

29

3

O

O

O

OP ..

:..

..

:..

..: :

: :O

O

O

OP ..

:..

..

:..

..: :

..

: :

3

Resonance: resonance structures are equivalent structures that only differ in the placement of some electrons.

3

O

O

O

OP..

:..

..

:..

..: :

:

O

O

O

O..

:..

..

:..

..: :

: :P

3

O

O

O

O P..

:..

..

:..

..::

:

3

The more resonance structures there are, the lower the energy is, the more stable the ion is.

All of these structures have 32 electrons.

30

Resonance practice. Draw resonance structures for the following ions:

C2H3O2

SO3

SO4 NO3

Remember that atoms in resonance structures must still obey the ideas of Lewis structures.In addition, the resonance structure that has negative charge on the more electronegative atom is better than one where the negative charge is on the less electronegative atom.

31

VSEPR Model: Valence Shell Electron Pair Repulsion Model

The purpose of this model is to predict the shape of molecules.

Shape includes a name for the shape and the bond angles between bonded atoms.

Be

Cl

Cl

Bond Angle: the angle in degrees between any two terminal atoms that are bonded to the same “central atom”-see “Be” and “B” below.

B

F

F F

180o120o

32

VSEPR Model: Valence Shell Electron Pair Repulsion Model

The shape of a molecule depends upon the number of electronic groups that surround the nuclei of the atoms in the molecule.

Like charges repel each other.

Pairs of electrons in bonds and lone pairs of electrons are electronic groups.

Imagine that we have two negative 1 charges attached to a nail with stings of equal length.

When we “let go” of the charges, what will they do?

They will move away from each other until they are as far apart as the “strings” allow.

What will this look like?

33

Two Charges held loosely.

let go

Farthest apart is a line.Imagine the “nail” is a nucleus.

Three Charges held loosely.

let go

Farthest apart is a triangle.Imagine the “nail” is a nucleus.

34

If we tried this charges on a string experiment with four, the shape we get when the charges are as far apart as possible will be three dimensional-not two dimensional like the first two examples.

Imagine, the “C” atom as the nail, and the bonds to the “H” atoms as the strings. When the charges are as far apart as possible they will take the shape of a tetrahedron. The blue lines represent the edges of the tetrahedron.

35

With five charges, the shape becomes two triangular pyramids that are stuck together. This shape is called a trigonal bipyramide.

36

With six charges, the shape becomes two square pyramids that are stuck together. This shape is called an octahedron.

37

This gives us a total of five shapes. We will rename the shapes slightly to get the “Electronic Shapes” that will be used to describe molecules.

A line becomes:

A triangle becomes:

A tetrahedron becomes:

A triangular bipyramid becomes:

A octahedron becomes:

Linear

Trigonal Planar

Tetrahedral

Trigonal Bipyramidal

Octahedral

38

Summary of possible shapes:

electronicshape

# ofgroups

examplemodel

Hybridization

sp

sp2

sp3

sp3d

sp3d2

BondAngle

180o

120o

109.5o

90o, 120o

90o

39

Hybridization: process of combining atomic orbitals on an atom to make new “hybrid” orbitals to be used for bonding.

Hybridization always starts with the lowest energy orbitals available (s then p then d). If two orbitals are combined, then two new hybrid orbitals are produced.

Examples:

Combining one “s” and one “p” orbital gives two “sp” hybrid orbitals

Combining one “s” and two “p” orbitals gives three “sp2” hybrid orbitals

Combining one “s” and three “p” orbitals gives four “sp3” hybrid orbitals

Combining one “s” three “p” and one “d” orbitals gives five “sp3d” hybrid orbitals

Combining one “s” three “p” and two “d” orbitals gives six “sp3d2” hybrid orbitals

40

If we focus on the number of electronic groups attached to an atom, we can predict hybridization just by counting the groups.

As stated on the previous slide: 2 groups would need “sp” hybridization; 3 groups would need “sp2” hybridization; 4 groups would need “sp3” hybridization; 5 groups would need “sp3d” hybridization; 6 groups would need “sp3d2” hybridization;

If we know hybridization, we can predict the electronic and molecular shapes.

If we know the electronic shape, we can predict bond angles.

An electronic group is a bonded atom or a lone pair.

CO2 CO3 CCl4 PCl5 SF6

If all of the electronic groups are bonded atoms, then the molecular shape is the same as the electronic shape.

In the examples below, there are no lone pairs on any “central atom”.

41

Subgroups within the electronic shapes: When one or more of the electronic groups around an atom is a lone pair, the molecular shape is a subset of the electronic shape. Subgroups of the “Tetrahedral” electronic shape are shown below.

Since these “new” shapes are all based on the tetrahedral shape because of the four attached electronic groups, the bond angels will be close to 109.5o.

Since lone pairs “push” harder than an atom, the bond angels will be less than the normal 109.5o when lone pairs are present.

42

CH4: no lone pairs

Tetrahedral electronic and molecular shape.

NH3: one lone pair

Tetrahedral electronic shape, but pyramidal molecular shape.

H2O: two lone pairs

Tetrahedral electronic shape, but bent molecular shape.

43

Effect of “lone pairs” on bond angle

44

Subgroups of the “Trigonal Bipyramidal” electronic shape are shown below.

“Trigonal bipyramidal”5 groups-no lone pairs

“See-Saw”5 groups-one lone pair

“T-shape”5 groups-two lone pairs

“Linear”5 groups-three lone pairs

Molecular Shape

45

Why is the shape linear when there are three lone pairs with the trigonal bipyramidal electronic shape?

The shape that allows the lone pairs to be as far apart as possible is the shape that the molecule will take.

The linear arrangement of I atoms allows the three lone pairs to be 120o from each other instead of 90o.

The T-shape arrangement of I atoms allows two lone pairs to be 120o from each other instead of 90o.

46

Subgroups of the “octahedral” electronic shape are:

Octahedral (no lone pairs)

Square Pyramidal (one lone pair)

Square Planar (two lone pairs)

Square planar allows the two lone pairs to be 180o apart instead of 90o.

47

Molecular Shape has an impact on polarity.

So far we have talked about polar bonds (like the CO bond) where the greater electronegativity of oxygen causes the oxygen to become partially negative and the lesser electronegativity of carbon causes it to become partially positive.

In molecules, all of the polar bonds interact with each other. Depending upon shape, these interactions can cause some molecules with polar bonds to be non-polar molecules.

Let’s look at water and carbon dioxide to discuss these issues:

48

In general, if a molecule has one of the five “perfect” shapes and all of the outer atoms are identical, the molecule will be non-polar because all of the individual polar bonds will cancel out.

The perfect shapes are: lineartrigonal planartetrahedraltrigonal pyramidaloctahedral

In general, if a molecule does not have polar bonds, it can not be polar.

In general, if a molecule has one of the “non-perfect” shapes or all of the outer atoms are not identical, the molecule will be polar because all of the individual polar bonds will not cancel out.

The imperfect shapes are: benttrigonal pyramidalsee-sawT-shapesquare pyramidal

49

The polarity of molecules plays a large role in many physical properties like:

boiling point, freezing point, density, and solubility

Remember, there is really only one force in chemistry (opposite charges attract each other). Polarity and geometry determine how much opposite charge is present in molecules.

If polar molecules are present, the force of attraction is called dipole-dipole interactions. These interactions are relatively strong.

If non-polar molecules are present, the force of attraction is called London dispersion forces (or just dispersion forces). These interactions are relatively weak.

If “H” is bonded to “F”, “O”, or “N”, the dipole-dipole force of attraction is particularly strong and is called hydrogen bonding (not as strong as covalent, ionic, or metallic bonding).

50

Na

+11

Na+1

+11

11 protons 11 electrons

0 charge

11 protons10 electrons+1 charge

Review of Positive ion formation:

+ Energy + 1 e1

Na Na+1+ Energy + 1 e1

Ionization Summary (called ionization energy)

1s2, 2s2, 2p6, 3s1 1s2, 2s2, 2p6

51

F

+9 +9

F 1

9 protons9 electrons0 charge

9 protons10 electrons1 charge

Review of Negative ion formation:

F + Energy+ 1 e1

Ionization Summary (called electron affinity)

1s2, 2s2, 2p5 1s2, 2s2, 2p6

+ 1 e1 + Energy

F 1

52

Electron configurations of ions from the representative elements.

Most Representative elements obtain a Nobel Gas electron configuration when they become ions.

Transition elements usually do not obtain a Nobel Gas electron configuration when they become ions.

53

Diagrams of LiF salt

Structures represent theCrystal Lattice

Top is a wire-frame picture.Ions at corners and “wires” indicating ionic bonds.

Bottom is a space fillingmodel where ions are shown“touching”. The pointswhere the ions touch is wherethe ionic bonds are located.

LiF is called the formula unit (smallest ratio of ions that describes the ionic compound)

54

Therefore, Lattice Energy depends upon:

Distance between the nucleiAnd

Charges on the ions

Where q = charge on an ion d = distance between nuclei

Shorter distance-higher energy

Greater Charge-higher energy

Force ≈(q1)*(q2)

d2

Attractive force between ions is related to Lattice Energy

55

The Energy Changes Involved in the Formation of Lithium Fluoride from Its Elements

1) making Li a gas

2) making Li an ion

3) breaking F2 into 2 F

4) making F an ion

5) Reacting Li+ and F

Ene

rgy

Giv

en O

ff

Lat

tice

Ene

rgy

56

The relatively high lattice energy of ionic compounds (compared to covalent molecules) gives ionic compounds many of their properties.

Ionic Compounds: All are hard, brittle solids All have high melting points All have very high boiling points If they dissolve in water, they all produce ions in solution.

Compounds that produce ions in solution are called electrolytes.

Water solutions of electrolytes conduct electricity, water solutions of non-electrolytes do not conduct electricity