ObjectivesDefine chemical bond.

Explain why most atoms form chemical bonds.

Describe ionic and covalent bonding.

Explain why most chemical bonding is neither purely ionic nor purely covalent.

Classify bonding type according to electronegativity differences.

Introduction to Chemical BondingChapter 6

What is bonding?

• Bonding is the “glue” that hold two or more elements together

• This “glue” is most likely formed as a result of a chemical reaction

• Bonding and molecular structure play a central role in determining the course of chemical reactions

What is a bond?

• A bond can be thought of as a force that holds groups of two or more atoms together and makes them function as a unit

• Example : water

OH H

Bonds require energy to break and release energy when made

Why do atoms bond?

Atoms bond because bonding lowers the atom’s potential energy.

Bond Energy is the energy required to break chemical bonds-is an endothermic process-requires energy-endothermic process-bond energy values are postive

Bond Formation-is a energy releasing process-exothermic process-the values for bond formation are negative

Bond LengthIs the distance between bonded atoms

Bond length and bond energy are indirectly related.

IONIC BONDING

Chemical bonding resulting from the electrical attraction between cations and anions.

Three Types of Bonds

• Ionic • Covalent• Metallic

Visual Concepts

Ionic Bonding

Chapter 6

Ionic Bonds

• Na and Cl– Na is a metal and likes to lose one electron– Cl is a nonmetal and likes to gain one electron– the final ionic compounds is NaCl

Na+ Cl-+ NaClThe electrostatic interaction keeps them together!

Ionic Bonds

• Na looses an electron and chlorine gains it!

• They do this to achieve an octet!

Na Cl

Covalent Bonds

• Covalent Bonds–exist between nonmetals bonded

together–form when atoms of nonmetals

share electrons–electrons can be shared equally or

unequally

Types of Covalent

• Polar• Nonpolar• Network Covalent

Visual Concepts

Covalent Bonding

Chapter 6

Metallic Bonds

• Metallic bonds exist between metals• Occur when two metals, usually the same

metal, are bonded together

IONIC OR COVALENT?

How can you tell?

The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

Visual Concepts

Using Electronegativity Difference to Classify Bonding

Chapter 6

Electronegativity Difference

• If the difference in electronegativities is between:– 1.7 to 4.0: Ionic– 0.3 to 1.7: Polar Covalent– 0.0 to 0.3: Non-Polar Covalent

Example: NaClNa = 0.8, Cl = 3.0Difference is 2.2, sothis is an ionic bond!

Visual Concepts

Comparing Polar and Nonpolar Covalent Bonds

Chapter 6

Ionic Vs. Covalent Bonding

Bonding and Ionic CompoundsChapter 6

Bonds Between Atoms

Covalent

Ionic

Polyatomic Ions

Metallic

Molecular Substance Polar

Nonpolar

Network Solids

What are we going to learn about???

Marriage

Divorce

Forming of a bond is like marriage

• More stable• exothermic

The breaking of a bond relates to a divorce.

• Less stable• Endothermic

Compare and Contrast

• Ionic and Covalent Compounds

PROPERTIES OF IONIC COMPOUNDS

Ionic compound – is composed of positive and negative ions that are equal in charge.

Formula Unit – is the simplest collection of atoms from which an ionic compound’s formula can be established.

Properties of Ionic Bonds

• What is an Ionic Bond?

- An Ionic Bond is a chemical bond resulting from the TRANSFER of electrons from one bonding atom to another

• When is an ionic bond formed?

- An ionic bond is formed when a cation (positive ion) transfers electrons to an anion (negative ion).

What are some characteristics of an ionic bond?

1. Exist as crystalline units at room temperatures

2. Brittle3. Have higher melting

points and boiling points compared to covalent compounds . Melting points are 1000 ˚C

4. Conduct electrical current in molten or solution state but not in the solid state

5. The smallest piece of an ionic compound is a formula unit.

What are some characteristics of an ionic bond?

6. A formula unit is the smallest collection of atoms from which a formula can be established.

7. Usually dissolves in polar solvents such as water. This is called dissociation.

8. The best way to test for an ionic compound is electrical conductivity in solution.

Formula Unitsof salt

Dissolving of salt in water - Dissociation

Lattice energy- is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.

Properties of Ionic Compounds

Lattice energy values are negative.

The more negative the value the strongerthe ionic bond.

Properties of Ionic Compounds

Compound Calculated Lattice Energy

NaCl −756 kJ/molLiF −1007 kJ/molCaCl2 −2170 kJ/mol

Covalent Bonds• What is an Covalent Bond?

- A covalent bond is a chemical bond resulting from SHARING of electrons between 2 bonding atoms.

• What forms a covalent bond?

- A covalent bond is formed between two nonmetals.

There are five different categories associated with covalent bonds. What are the 5 different categories?

Covalent

Molecular Substance Polar

NonpolarCoordinate Covalent

Network Solids

What are some characteristics of a covalent bond?

1. Very strong2. Low melting and boiling points Below 500 ˚C4. Exist as solids, liquids, or

gases at room temperature5. Do not conduct electricity

unless it is a molecular electrolyte (acid or base)

6. Molecular electrolytes undergo ionization in polar solvents acids and bases.

Ionization of Molecular Acids in Water.

Melting and Boiling Points of Compounds

Section 3 Ionic Bonding and Ionic CompoundsChapter 6

Covalent Bonds can have multiple bonds, so you should be familiar with the following…

Single Covalent Bond- chemical bond resulting from sharing of an electron pair between two atoms.

Double Covalent Bond- chemical bond resulting from sharing of two electron pairs between two atoms.

Triple Covalent Bond-chemical bond resulting from sharing of three electron pairs between two atoms.

Triple bonds are the strongest and shortest

c

First, we are going to look at Polar Covalent…

What is polar covalent?

-Polar covalent is a description of a bond that has an uneven distribution of charge due to an unequal sharing of bonding electrons.

The boy is not equally sharing with anyone else but rather taking all the food for himself.

Next, we are going to look at Non-Polar Covalent…

What is non-polar covalent?

-Non polar covalent is a covalent bond that has an even distribution of charge due to an equal sharing of bonding electrons.

This couple is non- polar because they are sharing the drink equally between them.

Now, we are going to look at Network Solids…

What is a Network Solid?

-A network solid is a solid that has covalently bonded atoms linked in one big network or one big macromolecule.

Name 3 Characteristics of a Network Solid.

1. Poor conductors of heat and electricity

2. Hard / Strong

3. High melting and boiling points

Diamond and graphite are examples of network covalent compounds.

Just as a summary to what each bond looks like…

The Metallic-Bond Model

The chemical bonding that results from the attraction

between metal atoms and the surrounding sea of electrons is called metallic bonding.

Properties of Metallic Bond

Metallic BondingWhat is a Metallic Bond?

- A metallic bond occurs in metals. A metal consists of positive ions surrounded by a “sea” of mobile electrons.

Name 4 Characteristics of a Metallic Bond.

1. Good conductors of heat and electricity

2. Great strength

3. Malleable and Ductile

4. LusterThis shows what a metallic bond might look like.

Properties of Metals: Malleability and Ductility

Malleability is the ability of a substance to be hammered or beaten into thin sheets.

Ductility is the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire.

Metallic Bonding

Chemical bonding is different in metals than it is in ionic, molecular, or covalent-network compounds.

The unique characteristics of metallic bonding gives metals their characteristic properties, listed below.

electrical conductivity

thermal conductivity

malleability

ductility

shiny appearance

Metallic Bonding

Why:a. are metals malleable?

b. are metals ductile?

c. are metals shiny?

Intermolecular Forces

• An attractive force that operates between molecules

• There are 3 kinds of intermolecular forces:– London dispersion force– Dipole-dipole force– Hydrogen-bonding force

London Dispersion Forces• Instantaneous dipole

– A temporary dipole formed when the electrons in an atom or nonpolar molecule happen to be more on one side in an instant in time, causing it to be more negative than normal and the opposite side positive

• Induced dipole– Positive end of the dipole exerts an attractive

force on nearby electrons, causing an adjacent atom to develop into another temporary dipole

London Dispersion Forces

– The attraction between temporary dipoles– Occurs between atoms and molecules– Only intermolecular force in nonpolar

substances– Tend to be stronger the larger the atom or

molecule (the more electrons are in the atom or molecule)

– Relatively weak forces

Dipole-Dipole Forces

• Attraction between polar molecules

• Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule

• Generally stronger than London dispersion forces

Fig. 5-8, p. 113

Hydrogen Bonding

• Special type of dipole-dipole force

• Only occurs in molecules that contain hydrogen bonded to a small, highly electronegative element (N, O, F)

• Stronger than a regular dipole-dipole force

Fig. 5-10, p. 114

Type of Force

Type of Interaction Occurrence

London dispersion

force

A temporary dipole in one molecule induces

the formation of a temporary dipole in a

nearby molecule and is attracted to it.

All atoms and molecules

Dipole-Dipole Force

Polar molecules (permanent dipoles) attract one another

Polar molecules

Hydrogen-Bonding Force

Two dipoles, one containing hydrogen to

an electronegative element and the other

containing an electronegative element,

attract one another.

Polar molecules containing unpaired

molecules and a hydrogen bonded to nitrogen, oxygen, or

fluorine

Electron Dot

How to Draw Lewis Structures

linear 180o

BeCl2

valence e- = 2 + (2 x 7) = 16e-

Cl....

..BeCl....

..

fewer than 8e-

valence pairs on Be bonding e-

linear molecule

two

linear 180o

CO2

valence e- = 4 + (2 x 6) = 16e-

CO....

.. O....

.. CO..

O..

.. ..

valence pairs on C ignore double bondstwo

single and double bonds same

linearmolecular shape

molecular geometry linear

120o

SO2

valence e- = 6+ (2 x 6) = 18e-

valence pairs on Sthree

one lone pair

molecular shape bent

S O....

..O....

:

SO....

.. O....

..:

SO...... O

..

..

:

two bonding pairs

< 120o

109.5o

NH3

valence e- = 5+ (3 x 1) = 8e-

valence pairs on Nfour N HH

H

:

< 109.5o

molecular shape trigonal pyramid

one lone pair

three bonding pairs

tetrahedral 109.5o

CH4

valence e- = 4+ (4 x 1) = 8e-

valence pairs on CfourC HH

H

H109.5o

molecular geometry

molecular shape tetrahedral

tetrahedral

We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds.

Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

What Proof Exists for Hybridization?

Lets look at a molecule of methane, CH4.

What is the expected orbital notation of carbon in its ground state?

(Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Can you see a problem with this?

Carbon ground state configuration

You should conclude that carbon only has TWO electrons available for bonding. That is not not enough!

How does carbon overcome this problem so that it may form four bonds?

Carbon’s Bonding Problem

The first thought that chemists had was that carbon promotes one of its 2s electrons…

…to the empty 2p orbital.

Carbon’s Empty Orbital

However, they quickly recognized a problem with such an arrangement…

Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

A Problem Arises…

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy.

But what about the fourth bond…?

Unequal bond energy

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron.

Such a bond would have slightly less energy than the other bonds in a methane molecule.

Unequal bond energy

This bond would be slightly different in character than the other three bonds in methane.

This difference would be measurable to a chemist by determining the bond length and bond energy.

But is this what they observe?

Unequal bond energy

The simple answer is, “No”.

Chemists have proposed an explanation – they call it Hybridization.

Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.

Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.

Enter Hybridization:

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals.

These new orbitals have slightly MORE energy than the 2s orbital…

… and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals