Chapter 13 Bonding: General Concepts - UCSBChapter 13: Bonding: General Concepts Take Away From Chapter 13 –Types of Bonding Types of Bonding Ionic Bonds (metal/non metal) Be able

Chapter 13

Bonding:

General Concepts

o Types of Bonding

o Deviations from ideal

bonding

o Lewis Structures

o Enthalpy of Reaction

o Shapes of Molecules (VSEPR)

o Polar Molecules

1

Big Idea: Bonds are formed from

the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence

electrons participate

in bonding. The shape

of the molecules

maximize the distance

between areas of high

electron density.

Chapter 13: Bonding: General Concepts

Types of Bonding

Ionic Bonds: Formed when a lower

energy can be achieved by the

complete transfer of one or more

electrons from the atoms of one element

to those of another; the compound is

then held together by electrostatic

attraction between the ions.

Covalent Bonds: Formed when the lowest

energy structure can be achieved by

sharing electrons.

2


Types of Bonding

Ionic Bonds: Tend to be between a metal

and a non metal.

Metals: Usually lose their electrons.

Nonmetals: Usually accept additional

electrons.

3

Note: In general, atoms gain or lose electrons until they have the same number of

electrons as the nearest noble gas


Types of Bonding

Element

Electron

Configuration

(Atom)

Gain/Lose

Electrons

Ion

Formed

Electron

Configuration

(Ion)

S

K

I

4

What ions do atoms form?


Student Question

Types of Bonding

Cations are ___________ than their parent atom.

a) Larger

b) Smaller

5


Types of Bonding

Isoelectronic: Ions containing the same

number of electrons.

6


Types of Bonding

Ionic Solids: Assembly of cations and anions

stacked together in a regular array.

Ionic Compounds are represented with formula

units (lowest ratio of types of atoms in the

compound).

Ionic Compound

Covalent Compound

7

AB2 2AB2 3AB2

CD 2CD 3CD+ -+ -

+-

+ -+-

+ -

Note: Ionic solids are example of crystalline arrays in which the overall charge on

an ionic solid is neutral.


Types of Bonding

What is the formula unit of a compound

formed from P and Na?

8


Student Question

Types of Bonding

Which of the following is a correct ionic

formula?

a) Al2O3b) Mg2O2c) Mg2F

d) BeO2e) None of the above

9


Types of Bonding

Coulomb Potential Energy

𝑬 =𝒒𝟏𝒒𝟐𝟒𝝅𝜺𝒐𝒓

r = separation

q =charge

10


Types of Bonding

Steps to calculate the energy needed to form an ionic bond:

Step 1: Standard states to gaseous single atom state

Na(s) Na(g) 97 𝑘𝐽𝑚𝑜𝑙

½F2(g) F(g) 80.𝑘𝐽

𝑚𝑜𝑙

Step 2: Both atoms have to form ions

Na(g) Na+(g) + e-(g) 494 𝑘𝐽𝑚𝑜𝑙

(ionization energy)

F(g) + e-(g) F-(g) -323 𝑘𝐽𝑚𝑜𝑙

(electron affinity)

Step 3: The ions need to come together to form a

crystal (Lattice Energy)

Na+(g) + F-(g) NaF(s) -923 𝑘𝐽𝑚𝑜𝑙

Total Reaction

Na(s) + ½F2(g) NaF(s)

97 𝑘𝐽𝑚𝑜𝑙

+ 80. 𝑘𝐽𝑚𝑜𝑙

+ 494 𝑘𝐽𝑚𝑜𝑙

+ −323 𝑘𝐽𝑚𝑜𝑙

+−923 𝑘𝐽𝑚𝑜𝑙

= −575 𝑘𝐽𝑚𝑜𝑙

11

Note: When energy is

released, the sign is negative

because no work is needed to

make the reaction happen.


Types of Bonding

12

Sublimation (step 1)

X-X bond strength (step 1)

Ionization energy (step 2)

Electron affinity (step 2)

Lattice Energy (step 3)

Born-Haber Cycle

Heat of Formation


Types of Bonding

What is holding ionic solids together?

Coulombic Potential Energy

𝐸𝑃,12 =𝑍1𝑒 𝑍2𝑒

4𝜋𝜀∘𝑟12= +𝑍 −𝑍 𝑒

2

4𝜋𝜀∘𝑑=−𝑍2𝑒2

4𝜋𝜀∘𝑑

Z1 & Z2 = charge of ions

The total potential energy is the sum of all the

potential energies

𝐸𝑃 =1

4𝜋𝜀∘−𝑍2𝑒2

𝑑+𝑍2𝑒2

2𝑑−𝑍2𝑒2

3𝑑+𝑍2𝑒2

4𝑑⋯ = − 𝑍

2𝑒2

4𝜋𝜀∘𝑑1−

1

2+1

3−1

4+⋯

𝐸𝑃 = −𝑙𝑛 2𝑍2𝑒2

4𝜋𝜀∘𝑑

Need to multiply by 2 to account for the other half of the

line.

𝐸𝑃 = −2𝑙𝑛 2𝑍2𝑒2

4𝜋𝜀∘𝑑

13

Note: 𝑙𝑛 2 = 1 − 12+ 1

3− 1

4+⋯


Types of Bonding

Once neighboring ions come into contact they

start to repel each other.

𝐸𝑃∗ ∝ 𝑒− Τ𝑑 𝑑

∗

d=0 𝑒− Τ𝑑 𝑑∗= 1 max repulsion

d>1 𝑒− Τ𝑑 𝑑∗= 1 repulsion decreases

The potential energy of an

ionic solid is a combination of

the favorable Coulombic

interaction of the ions and the

unfavorable exponential

increase which results when

the atoms touch. The ideal

bond length occurs at the

minimum potential energy.

Energy Minimum Occurs: 𝐸𝑝,𝑚𝑖𝑛 = −𝑁𝐴 𝑍1𝑍24𝜋𝜀∘𝑑

1 − 𝑑𝑑∗𝐴

14

Note: d* is a constant

that is commonly

taken to be 34.5 pm


Types of Bonding

Covalent Bond: A pair of electrons shared

between two atoms (occurs between

two non metals)

15

Note: In covalent bond formation, atoms go as far as possible toward completing

their octets by sharing electron pairs.


Take Away From Chapter 13 – Types of Bonding

Types of Bonding

Ionic Bonds (metal/non metal)

Be able to write electron configuration of ions. (30,31,32,

33,34)

Be able to predict size of ions.(27, 28, 29)

Be able to predict formula unit ionic compound.(37)

Covalent Bonds (non metal/non metal)

16

Numbers correspond to end of chapter questions.

Chapter 13

Bonding:

General Concepts

o Types of Bonding

o Deviations from ideal

bonding

o Lewis Structures



o Polar Molecules

17


the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence



of the molecules



electron density.


Deviations From Ideal Bonding

Electronegativity (χ): The ability of an

atom to attract electrons to itself when it

is part of a compound

18

Note: The atom with

higher electronegativity

has a stronger

attractive power on

electrons and pulls the

electrons away from the

atom with the lower

electronegativity.

Difference in

ElectronegativityType of Bond

> 1.8 Mostly Ionic

0.4-1.8 Polar Covalent

< 0.4 Mostly

Covalent

0 Non-polar

Covalent


Deviations From Ideal Bonding

What makes covalent bonds partly ionic?

Electric Dipole: A positive charge next to an

equal but opposite negative charge.

Electric Dipole Moment (𝝁): The magnitude of the electric dipole [units debye (𝐷)].

Polar Covalent Bond: A covalent bond between

atoms that have partial electric charges.

19

Note: The dipole moment associated with H-Cl is about 1.1 𝐷.


Deviations from Ideal Bonding

What makes ionic bonds partly covalent?

Polarizability (𝜶): The ease with which the electron cloud of a molecule can be distorted.

As the cation’s positive

charge pulls on the

anion’s negative

electrons, the spherical

electron cloud of the

anion becomes distorted

in the direction of the

cation. This causes the

bond to have covalent

bond properties.

20

Note: The larger the anion the easier it is to distort the electron cloud.


Deviations from Ideal Bonding

Consider the following bonds:

C-O, Mg-O, N-O, O-O, and F-O

Which are

non polar bonds

ionic bonds

polar covalent bonds

21


Take Away From Chapter 13 – Deviations from Ideal Bonding

Deviations from Ideal Bonding Electronegativity (17)

Know that covalent bonds can have ionic because of

dipole moments Be able to identify the most polar bond. (18,22,23)

Polarizability

Know that ionic bonds can have covalent character

because of polarizability

22


Chapter 13

Bonding:

General Concepts

o Types of Bonding

o Deviations from Ideal

Bonding

o Lewis Structures



o Polar Molecules

23


the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence



of the molecules



electron density.


Lewis Structures

Lewis Symbols: The chemical symbol of an

element, with a dot for each valence electron.

Step 1: Determine the number of valence e- from

electron configuration.

Step 2: Place 1 dot around the element for each

valence electron.

24


Lewis Structures

Examples

P

I

25


Lewis Structures

Drawing Ionic Lewis Structures

Step 1: Determine the electron configuration of

the elements in the compound.

Step 2: Determine the electron configuration of

the ions that the elements forms.

Step 3: Draw the Lewis symbols for the ions. Do

not forget to include charge. The cations should

have no electrons around them and the anions

should have 8 electrons around them.

Step 4: Organize the Lewis symbols such that

cations are next to anions.

26


Lewis Structures

Lewis Structures Ionic Compound Examples

CaCl2

MgO

27


Lewis Structures

General Rules (Covalent Lewis Structures)

All valence electrons of the atoms in the Lewis

structures must be shown.

Generally, electrons are paired. Except for

odd electron molecules such as NO and NO2.

Generally, each atom has 8 electrons in its

valence shell with the exception of H which

only needs 2 valance electrons.

Multiple bonds (double and triple bonds) can

be formed.

Show atoms by their chemical symbols (ex. H)

Show covalent bonds by lines (ex. F–F)

Show lone pairs of electrons by pairs of dots (ex. :)

28


Lewis Structures

Drawing Covalent Lewis Structures (for structures that obey octet rule)

Step 1: Count the number of valence electrons

on each atom; for ions adjust the number of

electrons to account for the charge.

Step 2: Calculate the number of electrons that

are needed to fill each atom’s octet (or duplet,

in the case of H).

Step 3: Calculate the number of bonds:

# 𝐵𝑜𝑛𝑑𝑠 =𝑊𝑎𝑛𝑡𝑒𝑑 𝑒− 𝑆𝑡𝑒𝑝 2 −𝑉𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− 𝑆𝑡𝑒𝑝 1

2.

Step 4: Calculate the number of electrons left

over: # 𝑒− = 𝑉𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝐵𝑜𝑛𝑑𝑠 .

Step 5: Place bonds/electrons around elements

so that octets/duplet are satisfied.29


Lewis Structures

HF

30

#(Type) Total

Step 1: Valence e-

Step 2: Wanted e-


Lewis Structures

N2

31

#(Type) Total

Valence e-

Wanted e-


Lewis Structures

Tips When Drawing Lewis Structures

How to pick central atom:

Choose the central atom to be the atom with the lowest

ionization energy (atom closest to the lowest left hand

corner of the periodic table)

Arrange the atoms symmetrically around the central

atom

32

Examples:

SO2 would be arranged OSO not SOO

Note: Acids are an exception to the rule because H is written first in acids.

Note: In simple formulas the central atom is often written first, followed by the

atoms attached to it.


Lewis Structures

H2O

33

#(Type) Total

Valence e-

Wanted e-


Lewis Structures

NH4+

34

#(Type) Total

Valence e-

Wanted e-


Lewis Structures

HCN

35

#(Type) Total

Valence e-

Wanted e-


Lewis Structures

What is the Lewis Structure for SCN-?

# 𝑏𝑜𝑛𝑑𝑠 =𝑤𝑎𝑛𝑡𝑒𝑑 𝑒− − 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒−

2=

24−16

2= 4

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝑏𝑜𝑛𝑑𝑠 = 16 − 2 4 = 8

Which structure is most likely?

36

1(S) 1(C) 1(N) 1(e-) Total

Valence e- 1(6) 1(4) 1(5) 1 16

Wanted e- 1(8) 1(8) 1(8) 24

___


Lewis Structures

Formal Charge: The electric charge of an atom

in a molecule assigned on the assumption that

the bonding is nonpolar covalent.

Formal Charge = Valence e- – e- Surrounding Atom

Generally, compounds with the lowest formal

charges possible (charges closest to 0) are

favored.

37

Note: The formal charge on neutral molecules must add up to zero.

Note: The formal charge on ions must add up to the charge on the ion.


Lewis Structures

38

___


Lewis Structures

Formal charge and oxidation numbers both give

us information about the number of electrons

around an atom in a compound.

Formal Charge Exaggerates Covalent Character

Assumes that the electrons are shared equally by all atoms.

Oxidation Number Exaggerates Ionic Character

Assumes that octets are complete filled and all

electrons must only belong to one atom.

39

0 0 0

-2 4 -2

O●●●●

●●

●● C

2-4+

O●●●●

●●

●●

2-


Lewis Structures

Double bonds

are not twice as

strong as a

single bonds.

Triple bonds are

not three times

as strong as a

single bond.

40


Lewis Structures

NO3-

41

#(Type) Total

Valence e-

Wanted e-


Lewis Structures

Resonance: A blend of Lewis structures

into a single composite hybrid structure.

Resonance Hybrid: The composite

structure that results from a resonance.

Delocalized Electrons: Electrons that are

spread over several atoms in a molecule.

42


Lewis Structures

What is the Lewis Structure for SCN-?


2=

24−16

2= 4

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝑏𝑜𝑛𝑑𝑠 = 16 − 2 4 = 8

Which structure is most likely?

43

1(S) 1(C) 1(N) 1(e-) Total

Valence e- 1(6) 1(4) 1(5) 1 16

Wanted e- 1(8) 1(8) 1(8) 24

___

(ADDITION)

0 0 -1 -1 0 0 1 0 -2


Student Question

Lewis Structures

Which of the following has the longest

carbon-oxygen bond?

Hint: You must draw the Lewis structures.

a) CO

b) CO32-

c) CO2d) CH3OH

44


Lewis Structures

What is the Lewis Structure for PO43-?


2=

40−32

2= 4

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝑏𝑜𝑛𝑑𝑠 = 40 − 2 4 = 32

45

1(P) 4(O) 1(e-) Total

Valence e- 1(5) 4(6) 1(3) 32

Wanted e- 1(8) 4(8) 40

3-


Lewis Structures

When the central atom in a molecule has empty

d-orbitals, it may be able to accommodate 10,

12, or even more electrons, this is referred to as

an expanded valence shell.

Size also plays a role in how many atoms can fit

around a given molecule.

46

Note: This only applies to nonmetal atoms in Period 3 and later

Examples:

PCl5 Known to exist

NCl5 Not known to exist (N is too small for the to fit 5 Cl atoms around it)

Note: On homework problems only expand octets if there is no other way to

accommodate electrons or if the problem tells you to minimize the formal

charge.


Lewis Structures

What is the Lewis structure for PO43- that

minimizes formal charges?

47


Take Away From Chapter 13 – Lewis Structures

Lewis Structures

Be able to draw Lewis symbols (atoms).

Be able to draw Lewis structures of ionic compounds.

Be able to draw Lewis structures of covalent

compounds. (63,64)

Know how to calculate formal charges.(85)

Identification of most likely Lewis structure.

Know when multiple resonance structures are possible for a

compound. (66,67,71)

Know when atoms can expand their octets (group 3 and

greater). (86)

Know how to estimate the length of bonds.

Triple < Double < Single (79,80,84)

48


Chapter 13

Bonding:

General Concepts

o Types of Bonding


Bonding

o Lewis Structures



o Polar Molecules

49


the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence



of the molecules



electron density.


Enthalpy of Reaction

Enthalpy Change (∆𝐻): The amount of heat evolved or absorbed in a reaction carried out at

constant pressure

Dissociation Energy (𝐷): The energy required to separate bonded atoms

50

Avg. Bond Energies 𝒌𝑱𝒎𝒐𝒍

H-H 432 O-H 467

H-F 565 O-O 146

H-Cl 427 F-F 154

C-H 413 Cl-Cl 239

C-C 347 O=O 495

C-N 305 C=O* 745

C-O 358 N≡N 941N-H 391 C≡C 839

N-N 160 C≡N 891

*C=O (CO2) = 799

∆𝐻𝑟𝑥𝑛 =𝐷 𝑏𝑟𝑜𝑘𝑒𝑛 −𝐷 𝑓𝑜𝑟𝑚𝑒𝑑



What is the energy change associated with the

following reaction?

C2H2(g) + 5

2O2(g) 2CO2(g) + H2O(g)

51


Take Away From Chapter 13 – Enthalpy of Reaction


Know how to calculate Δ𝐻 from bond dissociation energies .

△𝐻 = σ𝐷 𝑏𝑟𝑜𝑘𝑒𝑛 − σ𝐷 𝑓𝑜𝑟𝑚𝑒𝑑 (47,49,52)

52


Chapter 13

Bonding:

General Concepts

o Types of Bonding


Bonding

o Lewis Structures



o Polar Molecules

53


the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence



of the molecules



electron density.


Shapes of Molecules (VSEPR)

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

Extends Lewis’s theory of bonding to account for

molecular shapes by adding rules that account for

bond angle.

Rule 1: Regions of high electron concentration

(bonds and lone pairs on the central atom) repel

one another and to minimize their repulsion,

these regions move as far apart as possible while

maintaining the same distance from the central

atom.

Rule 2: There is no distinction between single and

multiple bonds: a multiple bond is treated as a

single region of high electron concentration.

54



Rule 3: All regions of high electron density, lone

pairs and bonds, are included in a description of

the electronic arrangement, but only the

positions of atoms are considered when

reporting the shape of a molecule (molecular

shape).

Rule 4: The strengths of repulsion are in the order

lone pair – lone pair > lone pair – atom > atom –

atom.

55



Assigning Shape and Bond Angles of Molecules

Step 1: Draw the Lewis structure.

Step 2: Assign the electronic arrangement

around the central atom (linear, trigonal planer,

tetrahedral)

Step 3: Identify the molecular shape (linear,

bent, trigonal planer, trigonal pyramidal,

tetrahedral)

Step 4: Figure out the bond angle (allow for

distortion)

56

Note: Electronic arrangement includes all areas of electron density

(lone pairs and bonds).

Note: Molecular shape includes only bonds.



Possible Electronic Arrangements

57



Possible Molecular Shapes

The names of the shapes of simple molecules and their bond angles. Lone pairs of electrons are not shown.

58



59

No Lone PairsTetrahedral

1 Lone PairTrigonal Pyramidal

2 Lone PairsBent

No Lone PairsTrigonal Planar

1 Lone PairBent

No Lone PairsLinear

3 Areas of

Electron

Density

2 Areas of

Electron

Density

4 Areas

of

Electron

Density



60

No Lone Pairs

Trigonal

Bipyramidal

2 Lone

Pairs

T-Shaped

3 Lone

Pair

Linear

1 Lone

Pair

Seesaw

5 Areas of

Electron Density



61

No Lone

Pairs

Octahedral

1 Lone Pair

Square

Pyramidal

3 Lone

Pairs

T-Shaped

4 Lone Pair

Linear

2 Lone Pairs

Square Planer

6 Areas of Electron

Density



Step 1: Draw Lewis Structure

Determine number of bonds

# 𝑏𝑜𝑛𝑑𝑠 =𝑤𝑎𝑛𝑡𝑒𝑑−𝑣𝑎𝑙𝑒𝑛𝑐𝑒

2

Determine number of electrons

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 − 2 𝑏𝑜𝑛𝑑𝑠

Step 2: Determine Electronic Arrangement

Step 3: Determine Molecular Shape

Step 4: Determine Angle

62

CH4 1(C) 4(H) Total

Valence e-

Wanted e-



Step 1: Draw Lewis Structure (Obeys Octet Rule)

Determine number of bonds

# 𝑏𝑜𝑛𝑑𝑠 =𝑤𝑎𝑛𝑡𝑒𝑑−𝑣𝑎𝑙𝑒𝑛𝑐𝑒

2

Determine number of electrons

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 − 2 𝑏𝑜𝑛𝑑𝑠

Step 2: Determine Electronic Arrangement

Step 3: Determine Molecular Shape

Step 4: Determine Angle

63

SO32- 1(S) 3(O) 2(e-) Total

Valence e-

Wanted e-



What is the shape of SO32- if you minimize the formal

charges

Lewis Structure

Electronic Arrangement

Molecular Shape

Angle

64


Student Question


What is the most likely shape of ICl4-?

Helpful Hint: Make sure that your formal

charges are minimized.

a) Octahedral

b) Trigonal Planar

c) Seesaw

d) Tetrahedral

e) None of the Above

65


Take Away From Chapter 13 – Shape of Molecules (VSEPR)

Shape of Molecules (VSEPR) (96,99,100)

Know how to determine electronic arrangement.

Linear, trigonal planar, tetrahedral, trigonal bipyramidal, or

octahedral.

Know how to determine molecular shape.

Linear, angular, trigonal planar, trigonal pyramidal, T-

shaped, tetrahedral, seesaw, square planar, trigonal

bipyramidal, square pyramidal, or octahedral.

Know how to determine bond angles.

66


Chapter 13

Bonding:

General Concepts

o Types of Bonding


Bonding

o Lewis Structures



o Polar Molecules

67


the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence



of the molecules



electron density.


Polar Molecules

Polar Bond: A bond that has a dipole moment

Nonpolar Bond: A bond that does not have a dipole moment

Polar Molecule: A molecule with a nonzero electric dipole moment.

68

Examples:

→H-O

The arrow points in the direction of electron movement

Note: Nonpolar molecules can have polar bonds.

For structure with only 2 types of atoms

• If molecular shape is one of the electronic arrangements

(nonpolar)

• Other most likely polar


Polar Molecules

69

Is CO2 polar or nonpolar? Is H2O polar or nonpolar?


Polar Molecules

70

For the structure XA3 where A is more electronegative

that X is the structure polar or non polar?

X

A

AA


Student Question

Polar Molecules

Is PCl4- polar or nonpolar?

Helpful Hint: Make sure that your formal

charges are minimized.

a) Polar

b) Nonpolar

71


Take Away From Chapter 13 – Polar Molecules

Polar Molecules

Be able to determine if a molecule is polar or non polar.

(101,102)

72


Documents

Chapter 13 Bonding: General Concepts - UCSBChapter 13: Bonding: General Concepts Take Away From Chapter 13 –Types of Bonding Types of Bonding Ionic Bonds (metal/non metal) Be able