Chemical Bonding Chapter 8 & 9. I. Overview/Types of Compounds

Chemical Bonding

Chapter 8 & 9

I. Overview/Types of Compounds

Overview of Bonds• Chemical Bonds are attractions

between atoms or ions in a compound• A bond forms when valence electrons

are either shared or transferred from one atom to another

• Most elements are more stable when bonded to other atoms, which is why bonds are formed.

Compounds

Before we start learning details about bonding, let’s review

compounds!

A. Definition of a Compound

• Remember, a compound consists of 2 or more elements chemically bonded together. It is always neutral.

EX: H20, NaCl, Sb3(PO4)5

B. Types of Compounds

2 types of chemical compounds

1. Molecular or Covalent Compound:

• consists of only non-metal atoms that share electrons which form a covalent bond.

– Examples: H20, C8H18

– smallest unit is called a molecule – molecules also include diatomics

(ex. H2)

2. Ionic Compound:• a compound that is formed from

the attraction between ions of opposite charges. It is held together by an ionic bond and usually contains a metal and nonmetal.

Ionic Compounds• Examples: NaCl, KBr• smallest unit is called a formula

unit• Cations are + ions. This happens

when an atom loses electrons. (ca+ion)

• Anions are – ions. This happens when an atom gains electrons. (a negative ion)

Practice1. Determine whether the following are

ionic or molecular covalent compounds

a. N2O5 b. PbNO5

c. KF d. AgCl

e. PCl3

ionic

molecular

molecular

ionic

ionic

always1+

2+ 3+ 3- 2- 1-

Practice2. Write the ion formed by these

elements

a. iodine

b. potassium

c. sulfur

d. gallium

e. strontium

S-2

I-

Sr2+

Ga+3

K+

C. Properties of Compounds

• ionic and molecular covalent compounds typically have very different properties

• this stems from the fact that the ionic attraction between formula units is far stronger than the attraction between covalently-bonded molecules

Comparison of PropertiesCharacteristic Molecular

CompoundsIonic Compounds

representative unit

Molecule Formula Unit

type of elements Nonmetals only Metals and Nonmetals

type of bonding Covalent (sharing)

Ionic (transfer)

physical state at RT

S, L, or G Solid only

melting point <300 °C ** > 300 °C

hardness soft hard hard

conduct electricity?

no yes (when dissolved in water)

Examples Candle wax Rock salt (NaCl)

**would melt in a bunsen burner

Ionic Compounds = transfer of electrons

Na Cl+ -

NaClIs formed

Molecular Compounds = share electrons

H O

H2OH

II. Types of BondsWe will learn about 4 different types of bonds: ionic, non-polar covalent,

polar covalent, and metallicBonds between two elements can be categorized based on the difference between the electronegativities of

those two elements.

A. Ionic Bonds

Ionic Bonds• Ionic bonds are the result of the

transfer of electron. It holds an ionic compound together.

• Remember, an ionic compound is a compound formed from the attraction between a cation and anion, which usually includes a metal and non-metal.

Ionic Bonds

• After a cation and anion are bonded to form an ionic compound, the net charge is zero

• The Electronegativity difference between the two atoms bonded is >1.7

[ ]-

Ionic Bonds

Na Na+

Cl Cl

Ionic Compounds = transfer of electrons

Na Cl+ -

NaClIs formed

Drawing ionic compounds

[ ]Cl + -

NaRemember when drawing ionic

compounds, only the anion is in brackets & has its valence electrons showing

Another Example

2+[

]O

2-

Ca

Ionic Bonds Strongest type of bond between

atoms. This results in high melting points and high boiling points.

Formula units form a crystalline structure

a120o

c

a

B. Covalent Bonds

• A covalent bond is formed by the sharing of electron-pairs within molecules

• Covalent bonds are formed between non-metals

Ex. H OH

Covalent Bonds

• Two atoms can share more than one covalent bond between them– Single bonds form between two atoms that

SHARE 1 PAIR of electrons. – Double bonds form between two atoms that

SHARE 2 PAIRS of electrons.

– Triple bonds form between two atoms that

SHARE 3 PAIRS of electrons.

Triple Bonds have the shortest bond length

Bond Length Bond Energy

H H Longest Weakest(163 kJ/mol)

O O Medium (418 kJ/mol)

N ≡ N Shortest Strongest (945 kJ/mol)

• There are two types of covalent bonds: non-polar covalent and polar covalent

Non-Polar Covalent Bonds

• electronegativity difference between atoms: 0-0.3

• even distribution of electrons • symmetrical

observed in all diatomic atoms (BrINClHOF elements)

Br I ClN HOF

Polar vs Non-PolarCovalent Bonds

Polar Covalent Bonds

• electronegativity difference between atoms: 0.3-1.7

• polar means “having opposite ends”

• unequal sharing of charges

Remember Electronegativity?

© 2003 Prentice Hall


• A dipole is formed

• the partial charges can be shown with the following symbols: δ - and δ+

Unequal sharing of electrons



• electrons are more attracted to the more electronegative element

• Opposites attract concept is observed

Polar or Non-Polar?

Has an even distribution of charge

Non-Polar!!

Polar or Non-Polar?

Has a ΔEN = 0

Non-Polar!!

Polar or Non-Polar?

Has a δ+ and δ- end

Polar!!

Polar or Non-Polar?

I2

Non-Polar!!

Polar Means:

“Having Opposite Ends”

(like N and S Poles)

C. Metallic Bonds

C. Metallic Bonds

• A metallic bond is the bond that holds two metals together. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

Metallic Bonds• The bonding

electrons are shared over the whole metal sample. This is known as electron delocalization.

• Because of the nature of metallic bonding, metals are malleable and ductile.

• -malleable = ability of a substance to be hammered or beaten into thin sheets

• -ductile = ability of a substance to be drawn or pulled through a small opening to produce a wire

Summary ChartDon’t forget your general rules of Ionic

compounds = Metal + Non-Metal

Type of Bond (between 2 atoms)

ΔEN(Paulings)

Metallic bondnot quantified

using EN

Ionic(transfer of electrons)

>1.7

Polar Covalent(unequal sharing of

electrons)0.4 – 1.7

Non-Polar Covalent(equal sharing of electrons)

0 – 0.4diatomicsUse a table/chart with electronegativity values to determine the difference

Electronegativity Values

0.3 1.7


H – F ΔEN1.9

H – Cl ΔEN0.9

H – BrΔEN0.7

H – IΔEN0.4

Most Polar Least Polar

• Although a compound is an ionic compound, it can still possess covalent bond characteristics and vice versa

Ex. H – F

Ionic Character Increases

Covalent Character Increases

2.1 4.0

ΔEN = 1.9

While this is a molecular compound, it possess ionic bond characteristics

III. Lewis Structures

III. Lewis Structures

• Only for molecular compounds!

• Recall drawing electron dot diagrams using the number of valence electrons (Roman Numerals on PT)

Lewis Structures are drawings!!

• Symbol = Nuclei

• Dots = Valence Electrons

• Dashes = Shared Electron Pair

• Electron Pair = unshared (aka “lone pair”)

= electron pair

1 bond = 2 shared electrons

• Electron pairs that are shared can be replaced with a line to represent a covalent bond. Up to three bonds can be shared between two elements

• Single bonds are formed between two elements that have 2 shared electrons (Ex. )

• Double bonds are formed between two elements that share 4 valance electrons (Ex. )

• Triple bonds are formed between two elements that share 6 valence electrons (Ex: )

Triple Bonds have the shortest bond length

Bond Length Bond Energy

N N Longest Weakest(163 kJ/mol)

N N Medium (418 kJ/mol)

N ≡ N Shortest Strongest (945 kJ/mol)

Steps to Drawing Lewis Structures

• Choose the center atom (Tips: choose the one you have less of, and carbon loves to be the center of attention)

• Draw an electron dot structure for the central atom• Draw the remaining electron dot structures near the

central atom’s unpaired electrons• Create bonds between the unpaired electrons• Rearrange your structure so it looks “neat” (spread the

atoms out)• Add lone pairs so elements 1-5 have “duets” and all

other elements have “octets”

Disclaimer: These rules actually only work for the examples we’re doing in our class. There are far more extensive rules that we will ignore for the sake of simplification.

Octet Rule• Octet Rule: chemical compounds tend to

form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

-Elements 1-5 need to be surrounded by 2 electrons-All other elements need to be surrounded by 8 electrons

Same as


Examples

CCl4

C

Cl

Cl

Cl

Cl

Examples

C2H2

C HH C

C HH C

Are the duets & octets complete?

ExamplesCO2

OO C

OO C


C OO

ExamplesCO2

OO C


Exceptions to the octet rule (incomplete octets)

Be HH

B F

F

FBF3

BeH2

Exceptions to the Octet Rule

FF

F

BBe HH

BF3

3 valence &3 bonded atoms

BeH2

2 valence &

2 bonded atoms

What does polar and non-polar mean again?

Review on how to draw structures (Ex. NH3)

N H

H

H

Polar

N H

H

H

S SNon-Polar

EN = 2.1

EN = 3.0EN = 2.1EN = 2.1

Polar & Non-Polar Molecules

• Like bonds, molecules can be polar (uneven distribution of electrons) or non-polar (even distribution)

• Non-polar molecules are symmetrical and have no lone pairs around the center atom

• Polar molecules have lone pairs around the central atom or an uneven distribution of electrons

NP OO C

PracticeIdentify each molecule as polar or

non-polar

FF

F

BNP

B F

F

FB

F

F

F

P


non-polar

NP

C2HCl


non-polar

P

C ClH C

Polar & Non-Polar Molecules


Dipoles

• Some molecules have dipoles that cancel out and have a zero net dipole

• The poles can be combined to form a net dipole

H HO

Positive end

Negative end

Polar or Non-polar Molecules?

CCl4

C

Cl

Cl

Cl

Cl EN = 2.5

EN = 3.0

EN = 3.0

EN = 3.0

EN = 3.0

C2H2

C HH C


CO2

OO C


N H

H

H

A better way to draw it

NH

HH

Br BrBr Br

B F

F

F

Draw BF3

B

F

F

F

IV. VSEPR Theory

VSEPR Theory

• VSEPR stands for Valence Shell Electron Pair Repulsion

• VSEPR is used to describe the 3D orientation of the electron regions.

• Shared electrons that make up covalent bonds and lone pairs will repel each other as much as possible

B

F

F

F NH

HH

BF3 vs NH3

Shape: Trigonal Planar

Shape: Trigonal pyramidal

Lone pairs repel more than bonds do

BF3 vs NH3

Shape: Trigonal Planar Shape: trigonal pyramidal

Look at the chart on page 7 of your notes

Ahhh! Memorization!?

• You will be responsible for predicting shapes with a * next to the name of shape without any notes

• You should be able to determine the shape of non - * compounds given this chart

1 lone pair

Trigonal planar*

/bent*

Linear*

1 lone pair 2 lone pairs

tetrahedral*

/bent*

Documents

Chemical Bonding Chapter 8 & 9. I. Overview/Types of Compounds