chapter 6 chemical bonding. section 1: introduction to chemical bonding

Chapter 6

Chemical Bonding

Section 1:

Introduction to chemical bonding


What is a chemical bond???

A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together


Why do atoms bond?

They are working to achieve more stable arrangements where the bonded atoms will have lower potential energy than they do when existing as individual atoms- increase stability.


Types of Chemical Bonding:

1. Ionic – an electrical attraction that forms between cations (+) and anions (-)

2. Covalent – are formed when electrons are shared between atoms

3. Metallic – formed by many atoms sharing many electrons


However…. Bonds are never purely

covalent or purely ionic. The degree of ionic-ness or

covalent-ness depends on property of electronegativity.

Degree of Ionic/Covalent Character in Chemical Bonds

Ionic

Polar-Covalent

Nonpolar-Covalent

100%

50%

5%

0%


Recall what electronegativity is:

The ability or degree of attraction that an atom has to electrons that are within a bonded compound.

(see page 161)


To determine the degree of ionic-ness or covalent-ness you must take each of the electronegativities for the elements in the compound and subtract them.


If difference is 0-0.3 = nonpolar covalent

If difference is 0.4 – 1.7 = polar covalent

Above 1.8 = Ionic

Ionic/Covalent Character Due to Electronegativity Differences

Ionic

Polar-Covalent

Nonpolar-Covalent

100%

50%

5%

0%

3.3

1.7

0.3

0


Sulfur + Hydrogen

Sulfur + Cesium

Sulfur + Chlorine

2.5 - 2.1 = 0.4 Polar Covalent

2.5 - 0.7 = 1.8 Ionic

2.5 – 3.0 = 0.5

Polar Covalent

Introduction to chemical bondingIn general however…

If bonding elements are on opposite sides of the periodic table (metal with a nonmetal) then they tend to be ionic.

If elements are close together (nonmetal to nonmetal), then they tend to be covalent.

Section 2:

Covalent Bonding & Molecular Compounds

Covalent Bonding

What is a molecule?

A neutral group of atoms that are held together by covalent bonds.

May be different atoms such as H2O or C6H12O6

May be the same atoms such as O2

Covalent Bonding

Molecular compounds are made of molecules ….. Not ions!

We represent covalent or molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH4 - methane

Covalent Bonding

Diatomic molecules are those elements that exist in pairs of like atoms that are bonded together.

There are 7 diatomic molecules:

H2 N2 O2 F2 Cl2 I2 Br2

Big 7

Covalent Bonding

Formation of a covalent bond:When atoms are far apart they do

not attract – potential energy is zero.

As they come closer the electrons are attracted to protons but electrons and electrons repel – but e- to p attraction is stronger!

Covalent Bonding

The electron clouds of the bonded atoms are overlapped and form a “bond length.”

Covalent Bonding

Energy is released when these atoms join together with a bond.

Energy must be added to separate these atoms into neutral isolated atoms – called bond energies.

Bond energy is expressed in kilojoules per mole.

Covalent Bonding

Octet Rule – Atoms will either gain, lose, or share electrons so that their outer energy levels will contain eight electrons (H is an exception since it can only have 2 in the outer level).

These electrons that are being gained, lost, or shared are represented by using the electron dot diagrams.

Examples of electron dot notations

1 valence electron

3 valence electrons

5 valence electrons

7 valance electrons

X

X

X

X

Covalent Bonding

Shared electron pairs and unshared pairs:

Cl:ClShared pair

Unshared pairs

Covalent Bonding

These electron dot representations are called Lewis structures.

Dots represent the valence electrons

Covalent Bonding

Lewis structures can also be represented using structural formulas.

Dashes indicate bonds of shared electrons (unshared e- are not shown

Cl - Cl One pair (2 e-) is shared here.

Steps To Drawing Lewis Structures

Calculate the number of valance electrons.

Arrange atoms. Compare number of electrons used with

number of electrons available. Check octet rule. Change dots to dashes where

appropriate.

Covalent Bonding

Lewis structure for ammonia (NH3)

Covalent Bonding

Practice: Draw Lewis structure for

methane CH4

Ammonia NH3

Hydrogen Sulfide H2S

Phosphorus trifluoride PF3

More Guidelines

H and halogen atoms usually bond to only one other atom in a molecule and are usually on the outside or end of a molecule (each only need 1 electron to form stable octet and electronegativity)

More Guidelines

The atom with the smallest electro-negativity is often the central atom

When a molecule contains more atoms of 1 element than the other, these atoms often surround the central atom

Covalent Bonding

Some atoms can form multiple bonds – especially C, O, & N.

Double bonds are bonds that share 2 pair of electrons

C=C means C::CTriple bonds share 3 pair

C≡C means C:::C

Covalent Bonding

Resonance:Some substances cannot be

drawn correctly with Lewis structure diagrams

Some electrons share time with other atoms – ex. Ozone – O3

Covalent Bonding

Electrons in ozone may be represented as: O = O–O

Other times it may be represented as O–O=O

Actually these structures are shared – electrons “resonate” (go back & forth) between them

Section 3:

Ionic Bonding and Ionic Compounds

Section 3: Ionic Bonding & Compounds

Ionic compounds are formed of positive and negative ions

When combined these charges equal zero

Ex: Na = 1+

Cl = 1-0 charge


Ionic substances are usually solids

Ionic solids are generally crystalline in shape

An ionic compound is a 3-D network of + and – ions that are attracted to each other


Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.


Ionic substances are not referred to as “molecules”

Ionic substances are referred to as “formula units”

A formula unit is the simplest ratio of the ions that are bonded together.


The ratio of ions depends on the charges.

What would result when F-

combines with Ca2+?

CaF2


When ions are written using electron dot structures the dots are written and symbols for their charges.

Na. Na+

Cl -

Compared to molecular compounds, ionic compounds:

Have very strong attractions Are hard, but brittle Have higher melting points and

boiling points When dissolved or in the molten

state they will conduct electricity

Polyatomic Ions:

A group of atoms covalently bonded together but with a charge.

Sulfate SO42-

Carbonate CO32-

Nitrate NO3-

Ammonium NH4+

Section 4:

Metallic Bonding

Metallic Bonding

Metals are excellent electrical conductors in the solid state.

This is due to highly mobile valence electrons that travel from atom to atom.

e-

Metallic Bonding

Generally metals have either 1 or 2 s electrons

p orbitals are vacantMany are filling in the d levelElectrons become delocalized

and move between atoms (sea of electrons)

Metallic Bonding

A metallic bond is the mutual sharing of many electrons among many atoms.

Metallic Properties

High electrical conductivityHigh thermal conductivityHigh lusterMalleable (can be hammered or

pressed into shape)Ductile (capable of being drawn or

extruded through small openings to produce a wire)

Metallic Bond Strength

Varies with nuclear charge and number of electrons shared.

High bond strengths result in high heats of vaporization (when metals are changed into gaseous phase)

Section 5:

Molecular Geometry

Molecular geometry…

A molecule’s properties depend on bonding of atoms, but also the molecular geometry.

Molecular geometry…

Is the three dimensional arrangement of a molecule’s atoms in space.

VSEPR Theory

Valence Shell Electron Pair Repulsion

Electrons around a nucleus repel each other to be as far away from each other as possible.

VSEPR Theory

Types of e- Pairs Bonding pairs - form bonds Lone pairs - nonbonding e-

Lone pairs repel

more strongly than

bonding pairs!!!

Draw the Lewis Diagram. Tally up e- pairs on central atom.

double/triple bonds = ONE pair

Shape is determined by the # of bonding pairs and lone pairs.

Know the common shapes & their bond angles!

Common Molecular Shapes

2 total

2 bond

0 lone

LINEAR180°BeH2

3 total

3 bond

0 lone

TRIGONAL PLANAR

120°

BF3


4 total

4 bond

0 lone

TETRAHEDRAL

109.5°

CH4


4 total

3 bond

1 lone

TRIGONAL PYRAMIDAL

107°

NH3


4 total

2 bond

2 lone

BENT

104.5°

H2O


PF3

4 total

3 bond

1 lone

TRIGONAL PYRAMIDAL

107°

F P FF

Examples

CO2

O C O2 total

2 bond

0 lone LINEAR

180°

Examples

Hybridization

Explains how atom’s orbitals become rearranged to form covalent bonds.

Hybridization is the mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energies.

Hybridization

Methane (CH4) is an example of hybridization: Carbon’s normal configuration is

2s22p2

In methane all the electrons in the 2nd energy level become equal in energy and is referred to as sp3

Intermolecular Forces

11.2

Intermolecular forces are attractive forces between molecules.

Intramolecular forces hold atoms together in a molecule.

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

Generally, intermolecular forces are much weaker than intramolecular forces.

“Measure” of intermolecular force

boiling point

melting point

Intermolecular Forces:

Strong IM forces exist in polar molecules.

Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)

Types of Intermolecular Forces

Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid

11.2

Types of IMF

Dipole-Dipole Forces

+ -

View animation online.

Types of Intermolecular Forces

Dipole ForcesAttractive forces between an ion and a polar molecule

11.2

Ion-Dipole Interaction


Another IM force is Hydrogen bonding.

Is the strongest type of dipole-dipole force

Explains high boiling points of H-containing substances such as water and ammonia


In hydrogen bonding, a hydrogen atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.

Types of Intermolecular ForcesHydrogen Bond (strongest)

11.2

The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. IT IS NOT A BOND.

A H…B A H…Aor

A & B are N, O, or F

Types of IMF

Hydrogen Bonding

Intermolecular Forces:London/ Dispersion/VanDerWaals

forces: Are very weak bonds Occur due to the fact that since

electrons are in constant motion that briefly there are moments where electrons are unevenly distributed and thus the molecule briefly has a charged area.

Types of IMF

London Dispersion Forces

View animation online.

chapter 6 chemical bonding. section 1: introduction to chemical bonding

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