ch. 6. what is a chemical bond? mutual electrical attraction between the nuclei and valence...

CHEMICAL BONDINGCh. 6

What is a chemical bond? mutual electrical attraction between the nuclei and

valence electrons of different atoms that bind the atoms together

Why don’t noble gases do this?Already have filled s and p orbitals stable octet: 8 valence e- (or 2 if, you’re helium)

Atoms that don’t have a stable octet are more reactive

Key Point #1: By forming bonds with each other, most atoms reduce their potential energy, becoming more stable.

This is a chemical change! All chemical changes involve energy!

What types of bonds can be formed?

Metallic Bonding

• In a metal, the empty orbitals in the atoms’ outer energy levels overlap

Delocalized Electron: outer electron that does not belong to any one atom but can move freely through the metal’s network of empty atomic orbitals.

sea of electrons: mobile electrons around the metal atoms, which are packed together in a crystal lattice.

metallic bonding: chemical bonding that results from the attraction between metal atoms and surrounding sea of electrons

Key Point: In metallic bonding, valence electrons move freely throughout a network of metal atoms.

Unique Characteristics of Metals

Metals have many unique properties because of their sea of electrons

• Malleability: ability of a substance to be hammered or beaten into thin sheets

• Ductility: ability of a substance to be pulled into a thin wire

• Why? atoms can slide past one another along a plane without breaking bonds

Luster: shiny appearance• Why? Absorb a wide range of light frequencies,

many orbitals separated by small energy differences

Conductivity• Thermal: ability to conduct heat• Electrical: ability to conduct electricity

• Why? Electrons move easily through network of empty orbitals

Metallic Bonding Strength

The strength of metallic bonding is determined by the enthalpy of vaporization:

• the amount of energy required to vaporize (turn into a gas) 1 mol of a metal

In general, the strength of the metallic bond INCREASES moving left to right across the periodic table.• Soft metals (less dense) metals harder (more

dense) metals toward right

Properties of Metals: Malleability and Ductility

Properties of Metals: Surface Appearance

Properties of Metals: Electrical and Thermal Conductivity

Types of Bonds What type of bonds can be formed?

Ionic bond Covalent bond

○ Nonpolar covalent○ Polar covalent

Ionic bonding: bonds that result from electrical attractions between cations and anions

1 atom losses electrons 1 atom gains electrons

Covalent bonding: sharing of electrons between 2 or more atoms

Key Point 2: Rarely is bonding between atoms purely ionic or purely covalent. Instead, it usually falls somewhere between the two

extremes. Why?

Key Point 3: The extent of ionic or covalent bonding between two atoms can be estimated by calculating the difference in each elements’ electronegativity.

Covalent Bonding Large difference in E.N.: bond has more ionic

character Small difference in E.N: bond has more covalent

character

Types of Covalent Bonds Non-polar covalent bonding: both electrons equally

shared between atoms

Polar covalent bonding: unequal attraction for the shared electrons

6.1 Practice WorksheetPart 1

The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent.

The greater the electronegativity difference, the more ionic the bonding is.

If the calculated electronegative difference is…

> 1.7 : ionic bond is formed> 0.3 , < 1.7 : polar-covalent bond0 – 0.3 : non-polar covalent bond

Increasing difference in electronegativity

NonpolarCovalentshare e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Elements Electronegativity Electronegativity Difference

Bond Type

Element 1

Element 2

Mg to Cl

H to O

C to Cl

N to H

C to S

K to F

Na to Cl

H to H

1.2

2.5

3.0

2.5

.8

.9

2.1

2.1

3.0

3.5

3.0

2.1

2.5

4.0

3.0

2.1

1.8

1.4

.5

.9

0

3.2

2.1

0

Ionic

Ionic

Ionic

Polar covalent

Polar covalent

Polar covalent

Non-polar covalent

Non-polar covalent

Polyatomic Ions It is also possible if a compound contains polyatomic

ions, for both types of bonding to be present.

Monatomic Ions: Fe2+ , Na+, Cl-

Polyatomic Ions: PO43-, NH4

+ , NO-1

Groups of atoms are bonded covalent together, but because of few or more than expected valence electrons they have an overall charge (so they can also bond ionically with other ions)

Ex: Ca2+ and SO42- CaSO4 (metal & diff. nonmetals)

Classify the following as ionic, covalent, or both

1. CaCl2 = __________

(metal & nonmetal)

2. CO2 = __________

(nonmetal & nonmetal)

3. MgO = __________

(metal & nonmetal)

4. HCl = ___________

(nonmetal & nonmetal)

5. BaSO4 = ___________

(metal & diff. nonmetals)

6. H2O = ____________

(nonmetal & nonmetal)

7. SO3 = ___________

(nonmetal & nonmetal)

8. AlPO4 = ___________

(metal & diff. nonmetals)

Covalent

Covalent

Covalent

Ionic

Ionic

Both

Covalent

Both

Covalent BondingSection 6.2

What is a molecule? Neutral group of atoms that are held together by

covalent bonds.

Chemical formula: indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

Formation of Covalent Bonds The electrons of one atom and protons of the other

atom attract each another.

The two nuclei and two electrons repel each other.

These two forces cancel out to

form a covalent bond at a length

where the potential energy is

at a minimum.

Bond Length vs. Bond Energy Bond length (pm): distance between two bonded

atoms at their minimum potential energy

Bond energy (kJ/mol): energy required to break a chemical bond and form neutral isolated atoms.○ Breaking bonds: absorbs (requires) energy○ Forming bonds: releases energy

Key Point: As you increase the number of bonds between 2 atoms the bond energy increases, while the bond length decreases. This is an inverse relationship.

Bond Energies & Bond Lengths A. How many electrons are shared in a

single bond: double bond: triple bond:

B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break?

In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

Bond Energies & Bond Lengths A. How many electrons are shared in a

single bond: 2 e- double bond: 4 e- triple bond: 6e-

B. Which bond is shorter? C – C or C = C C. Which bond requires more energy to break? =

In addition to finding an ideal bond length, atoms also lower their potential energy by achieving a stable octet of 8 valence electrons

Octet Rule Octet Rule: chemical compounds tend to form so that

each atom has an octet of e-’s in its highest occupied energy level

Exceptions to the octet rule:Atoms that cannot fit eight electrons Atoms that can fit more than eight electrons

Hydrogen: 2e-Boron: 6e-Phosphorus, Sulfur, & Xenon: expanded

valence, more than 8e-

How can we represent molecules? Lewis Structures: formulas in which atomic symbols

represent nuclei and inner shells, which are surrounded by dot-pairs/dashes represent valence electrons

MOLECULAR GEOMETRY

Chapter 6.5

VSEPR THEORY

Lewis Structures are 2D but we live in a 3D world!molecular geometry: the three-dimensional

arrangement of a molecule’s atoms

What do those 3D structure/shapes look like??Follow the Valance Shell Electron Pair Repulsion

Theory or VSEPR○ Repulsion between the sets of valence electrons

surrounding an atom causes them to be oriented as far away from each other as possible

Why use VSEPR Theory? Key Point: VSERP Theory is used to predict the shape

of molecules based on the fact that electron pairs strongly repel each other.

Following VSEPR allows us to predict bond polarity:uneven distribution of electrons

AB2 – Linear

Atoms bonded to

central atom (B)

Number of Lone Pairs on central

atom (E)

Bond Angle

2 0 180˚

Central atom

Atoms/group of atoms attached to central atom

Other Linear Geometries The shape of two atoms bonded together is not

given in the chart.Ex: F2

What is the only possible shape a binary compound can have?

○ LINEAR!

AB2E1 – BentAtoms

bonded to central atom

(B)

Number of Lone Pairs

(E)Bond Angle

2 1 <120˚What happens to the bond angle between atoms as you increase the number of “lone pair electrons” on the central atom?

Bond angles decrease!

AB2E2 – Bent

Atoms bonded to

central atom (B)

Number of Lone Pairs

(E)Bond Angle

2 2 104.5˚

AB3 – Trigonal Planar

Atoms bonded to

central atom (B)

Number of Lone Pairs

(E)Bond Angle

3 0 120˚Shape is often associated with atoms that break octet rule, but doesn’t have to be

AB3E1 – Trigonal Pyramidal

Atoms bonded to

central atom (B)

Number of Lone Pairs

(E)Bond Angle

3 1 107˚

AB4 – TetrahedralAtoms

bonded to central atom

(B)

Number of Lone Pairs

(E)Bond Angle

4 0 109.5˚

Predicting Molecular Geometry

1. Draw Lewis structure for molecule.

VSEPR theory: if any lone pairs of electrons are found on the central atom, these electrons

decrease the bond angles of atoms attached to it.

2. Draw a revised Lewis structure to show more accurate geometry

SO O

AB2E

bent

S

F

F

F F

AB4E

tetrahedral

3. To indicate the polarity of the bonds, we use this symbol: __________________ , which always points toward the more electronegative element.

4. When multiple bonds are found in a molecule, we must identify polarity of each bond.

Predicting Molecular Polarity

H F

electron richregion

electron poorregion FH

e- riche- poor

d+ d-

5. Observe the overall polarity of the molecule. Think of it as “tug-of-war” for valence electrons between the various atoms.

Non-polar covalent molecules: If the atom is symmetrical and all atoms have an equal pull on electrons

Polar covalent molecules: If the atom is not symmetrical and/or the atoms do not all have an equal chance of winning the tug of war for electrons

Intermolecular Forces

Intermolecular forces: attractive forces between molecules.

Intramolecular forces: attractive forces within a molecule (the bonds)

intermolecular forces are much weaker than intramolecular forces

Intramolecular Forces (bond)

Intramolecular Forces

Intermolecular Forces

Strength of IMF

Hydrogen Bond Dipole – Dipole Induced Dipole London Dispersion

Forces

strongest

weakest

H F

Dipoles What is a dipole?

A polar moleculeUneven sharing of electrons so there is a

separation of charge

FH

electron richregion

electron poorregion

e- poor

d+ d-

Dipole-Dipole Forces

Attraction between two polar molecules

— + — +

Hydrogen Bonding

Special type of Dipole – Dipole Attraction between:

Hydrogen & Nitrogen/Oxygen/Fluorine

Induced Dipole Attraction between one polar and one nonpolar

molecule

— +

— + — +

Electrons shift toward positive end

of dipole

London Dispersion Forces

Attraction between two nonpolar molecules

— + — +

Electrons become

uneven and form a dipole

What does IMF effect? Viscosity Surface Tension Boiling Point

Boiling Point

Point at which liquid particles escape the surface of the liquid into the gas phase

Stronger IMF Higher Boiling Point

Stronger IMF Higher Surface Tension

Surface Tension

result of an imbalance of forces at the surface of a liquid.

Stronger IMF Higher Viscosity

Viscosity

Measures a fluid’s resistance to flow