Chapter 6

Chemical Bonding

Section 6-1

Introduction to Chemical Bonding

Types of Chemical Bonding

Very rarely do atoms exist by themselves.

They usually exist in combinations with each other.

These combinations are held together by chemical bonds.

What is a Chemical Bond?

It is an electrical attraction between the nuclei of one atom and the valence electrons of another that binds the atoms together.

It occurs to increase stability.

There are Two Types of Chemical Bonds.

Ionic Bonding – The bonding that occurs between large numbers of metal cations and nonmetal anions. It is the attraction of opposite charges.

Covalent Bonding – The bonding that occurs between two elements that share electron pairs.

Usually, a Bond is neither purely ionic or covalent.

The ionic or covalent character of a bond is determined by electronegativity.

What is electronegativity?

Electronegativity

The ability of an atom to attract electrons to itself.

If the electronegativity difference is less than 0.5, it is covalent.

If it is between 0.6 and 1.9, it is polar.

If it is greater than 2, it is ionic.

What is a Polar Covalent Bond?

It is a bond where the electrons are pulled more closely toward the electronegative element.

Polar Covalent Bonds, cont.

This gives the electronegative atom a slightly negative charge, and the electropositive atom a slightly positive charge.

It is a bond where electrons are shared unequally.

Section 6-2

Covalent Bonding and Molecular Compounds

Molecular Compounds

Many compounds are made up of molecules.

A molecule is a neutral group of atoms held together by covalent bonds.

A chemical compound whose simplest units are molecules is called a molecular compound.

Chemical Formulas

A chemical compound is represented by a chemical formula.

Chemical formulas indicate the number of each type of atom in a chemical compound by using atomic symbols and numerical subscripts.

A molecular formula represents a molecular compound.

A Sample Molecular Formula

The molecular formula H2O is the formula for water.

It indicates that water is composed of two atoms of hydrogen, and one atom of oxygen.

Diatomic Molecules

There is a special class of molecule called a diatomic molecule.

These are formed by atoms that cannot exist alone, so if they are pure, they must be in pairs.

Diatomic Molecules

There are seven diatomic molecules.These are: H2, N2, O2, F2, Cl2, Br2, and I2

These molecules are held together by purely covalent bonds.

Every Covalent Bond has Two Properties

Bond LengthBond Energy

Bond Length

The average distance between two bonded atoms.

Measured in picometers (pm)

Bond Energy

The energy required to break a chemical bond.

Measured in kilojoules per mole. (kJ/mol)

These two properties are characteristic of every bond.

Table 6.1 on page 168 shows some common values.

Noble Gases

Noble Gases do exist by themselves.

This is because they have 8 valence electrons.

Other atoms must achieve this through covalent or ionic bonding.

This helps them satisfy the octet rule.

The Octet Rule

Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons has 8 valence electrons in its outer energy level.

The exceptions are H, Li, Be, and B, which can have less than 8.

Extended Valence

Some elements can have more than 8 valence electrons when they combine with elements like F, O, or Cl.

These cases of extended valence are rare, and are not important for us.

Electron Dot Notation

Covalent bonding usually only involves an atom’s valence electrons.

We use electron dot notation to keep track of these electrons.

Electron dot notation is an electron configuration in which valence electrons are shown as dots around the element’s symbol.

The number of dots is equal to the group number for all group A elements.

This can also be used for molecules.

Electron Dot Notation

In a molecule, a pair of dots between two atoms represents a shared pair of electrons, or what we call bonding electrons.

A pair of electrons around an element that are not shared with another element are non – bonding electrons or unshared electron pairs.

Lewis Structures

The bonding pair of electrons can also be drawn as a dash.

This then becomes a Lewis Structure.Lewis structures are formulas in which

bonding electrons are drawn as a dash and nonbonding electrons are drawn as dots.

Structural Formulas

Lewis structures can also be drawn without dots around them.

These are structural formulas.

Multiple Bonds

Up until now, every molecule we have looked at has a single pair of shared electrons between atoms.

It is possible to have more than one pair that are shared between the same two atoms. These result in multiple bonds.

If two pairs of electrons are shared, that results in a double bond.

Multiple Bonds, cont.

If three pairs of electrons are shared, it is a triple bond.

The octet rule must still be followed in multiple bonded molecules.

Anytime C, N, or O are in a molecule, multiple bonds are possible.

Resonance

Some molecules cannot be accurately shown with one structure.

These are molecules with alternating single and double bonds.

Neither form really exists. The true form is a combination of the two.

This is called resonance.

Resonance

Each structure is called a resonance structure.

The true form of the molecule is called a resonance hybrid.

A double headed arrow between Lewis structures indicates that they are resonance forms.

Section 6-3

Ionic Bonding and Ionic Compounds

Ionic Compounds

To this point, we have only discussed molecular compounds. A second type of compound is an ionic compound.

This is a compound made up of positive and negative ions that combine into a neutral substance.

These exists as crystals of many positive and negative ions.

Ionic Crystals

A three – dimensional arrangement of ions in a regular, repeating pattern.

Instead of writing an entire crystal, we usually represent an ionic compound with a formula unit.

Formula Unit

A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be determined.

This ratio is determined by ionic charge.

Crystal Lattice

An ionic crystal is known as a crystal lattice.

For example, in NaCl, each sodium attracts six negative chlorines, and every chlorine attracts six positive sodiums.

All of these additional bonds makes a crystal very strong and stable.

Lattice Energy

It takes a great deal of energy to break down an ionic compound.

This energy is known as lattice energy.Lattice energy is the energy released

when one mole of a crystal forms.

What are the differences between Ionic and Molecular Compounds?

Ionic Compounds: Have High Melting

Points. Have High Boiling

Points. Are Very Hard and

Brittle. Conduct Electricity in

water.

Molecular Compounds:

Have Low Melting Points.

Have Low Boiling Points.

Are Soft. Do not Dissolve in

Water.

Why do Ionic Compounds Conduct Electricity?

Scientists have known for a long time that ionic substances conduct electricity when they are dissolved in water. In fact, another term for an ionic substance is an electrolyte.

The problem was, no one knew how or why this occurred.

Svante August Arrhenius (1859 – 1927)

Arrhenius was one of the scientists studying electrolytes.

He concluded that electrolytes split, or dissociated in water into positive and negative ions.

The degree to which this happened depended on the ionic character of the substance, and its concentration.

Svante August Arrhenius (1859 – 1927)

The work of Arrhenius not only described how ionic compounds are put together, but also greatly affected both chemistry and biology, allowing us to describe simple reactions like why adding ionic compounds increases freezing and boiling points, to a greater understanding of digestion and absorption.

Polyatomic Ions

A charged group of covalently bonded atoms are known as a polyatomic ion.

There are many polyatomic ions, and you will be responsible for memorizing a few of them.

Polyatomic Ions you Should Know

NH4+1 – Ammonium

H3O+1 – Hydronium

CO3-2 – Carbonate

ClO3-1 – Chlorate

ClO2-1 – Chlorite

CrO4-2 – Chromate

HCO3-1 – Hydrogen

Carbonate

HSO4-1 – Hydrogen

Sulfate OH-1 – Hydroxide NO3

-1 – Nitrate

NO2-1 – Nitrite

PO4-3 – Phosphate

SO4-2 – Sulfate

SO3-2 - Sulfite

Section 6-4

Metallic Bonding

Metallic Bonding

There is a third type of bonding that is different from ionic or covalent bonding. This is metallic bonding.

Metallic bonding accounts for the different properties metals have.

Metallic Bonding

Most metals have few valence electrons.The s-block metals have empty p

subshells, and the transition metals have empty p and d subshells. These empty orbitals tend to overlap.

This allows valence electrons to freely move across the metal from atom to atom.

The Sea of Electrons

These electrons are said to be delocalized.

The bonding that results from the attraction of metal atoms to the sea of electrons is metallic bonding.

Metallic Bonding Accounts for Several Unusual Properties of Metals

Conductivity Malleability Ductility The electrons surround

each atom, and keep opposite charges from contacting each other.

Allows the atoms to slide past each other, without repelling.

Section 6-5

Molecular Geometry

The properties of molecules depends not only on their bonding, but also on their molecular geometry.

Molecular geometry is determined mostly by molecular polarity – the uneven distribution of molecular charge.

VSEPR Theory

The theory used to determine molecular geometry is VSEPR Theory.

VSEPR stands for Valence Shell Electron Pair Repulsion.

It states that the valence level electron repulsion around an atom causes the sets of electrons to be oriented as far apart as possible.

There are six geometric shapes molecules form that we will deal with in this class.

Shape 1

Any molecule made of only two atoms is called linear diatomic.

Shape 2

Any molecule made of three atoms with no pair of free electrons on the center atom is linear triatomic.

An example would be Calcium Fluoride, CaF2.

Shape 3

Any molecule with three atoms and a pair of free electrons on the center atom is Bent Triatomic.

An example of this would be water, H2O.

Shape 4

Any molecule with four atoms and no free electron pairs on the center atom is trigonal planar.

An example of this is Aluminum Chloride, AlCl3.

Shape 5

Any molecule with four atoms and a pair of free electrons on the central atom is called pyramidal.

An example of this is ammonia, NH3.

Shape 6

Any molecule made of five atoms is tetrahedral.

An example of this is methane, CH4.

There are other types of VSEPR shapes, but they only apply to molecules with extended valence, and will not be studied in this class.

Hybridization

The VSEPR shapes of molecules are possible because the orbitals in the molecules become rearranged.

This is a process called hybridization. Hybridization is the blending of two or

more different atomic orbitals to create orbitals of identical shape and energy.

There are three types of hybridization

Sp3 – one s and three p orbitals in an energy level combine to form four equal orbitals.

Sp2 – one s and two p orbitals in an energy level combine to form three equal orbitals.

Sp – one s and one p orbital in an energy level combine to form two equal orbitals.

Sp3 Hybridization

The sp3 hybridization results in all single bonds.

It forms the tetrahedral, pyramidal, or bent VSEPR shapes.

Sp2 Hybridization

The sp2 hybrid results in either single bonds, or a double bond.

It forms the trigonal planar VSEPR shape.

Sp Hybridization

The sp hybrid allows for single, double, or triple bonding.

It results in molecules with a linear VSEPR shape.

There are two types of covalent bonds.

Sigma Bonds – form from the head on overlap of two atomic orbitals.

Pi Bonds – form from the sideways overlap of two unhybridized p orbitals.

A single bond is one sigma bond. A double bond is one sigma and one pi bond. A triple bond is one sigma and two pi bonds.

Intermolecular Forces

These are forces between two or more molecules.

These are relatively weak forces, but in large numbers, can affect things like boiling and melting points.

There are three types of Intermolecular Forces.

Dipole – Dipole ForcesHydrogen BondingLondon Dispersion Forces

Dipole – Dipole Forces

Forces of attraction between polar molecules.

Polar molecules have equal but opposite charges separated by a short distance.

The negative region of one polar molecule attracts the positive region of another.

We use arrows to point to the negative pole. Not all polar molecules are dipoles.

Hydrogen Bonding

A very strong dipole – dipole force.A hydrogen ion in one polar molecule is

attracted to a strongly electronegative element (O, Cl, F, N) in another.

Hydrogen bonding is what causes ice to expand when it is frozen. It also causes high boiling and melting points in water.

London Dispersion Forces

Intermolecular attraction resulting from the constant motion of electrons.

Creates instantaneous dipoles.

This acts only on noble gases and nonpolar molecules.