SCH4U – UNIT 1STRUCTURE AND PROPERTIESCHAPTER 4 – CHEMICAL BONDING

Activity• With a partner discuss everything you remember about

chemical bonding• Eg. Types of bonds? why? What happens?

4.1 Types of Chemical Bonds• What are the two main types of chemical bonds?

• Ionic: chemical bond between oppositely charged ions• Electrostatic attraction

• Covalent: a chemical bond in which atoms share bonding electrons• Bonding Electron Pair: electron pair that is involved in

bonding

• Bond type depends on the attraction for electrons of the atoms involved• i.e. electronegativity

Ionic Compounds

How do these work?

Metal + Non-Metal Metal+ + Non-Metal-Low IE High IE Isoelectronic with noble gases

Low EA High EA

Opposites attract in no particular direction, considered non-directionalIons cling together in clusters known as crystals

• Get a lattice structure

• Lattice energy: energy change when one mole of an ionic substance is formed from its gaseous ions• Depends on:

• Charge on the ions• Size of the ions

Ionic Compounds and Bonding• Properties – WHY?

• Do not conduct electric current in the solid state

• Conduct electric current in the liquid state

• When soluble in water, form good electrolyte

• Relatively high MP and B

• Brittle, easily broken under stress

Covalent Bonds

Balance of attractive and repulsive forces

What are the forces acting here?

Octet Rule• Atoms share electrons so that they are surrounded by 8

electorns• # bonds = 8 - # valence electorn

• Example: Carbon, Oxygen, Nitrogen

• Two covalent bonds = double bond• Three covalent bonds = triple bond

Lewis Structures• Atoms and ions are stable if they have a full valence shell• Electrons are most stable when they are pair• Atoms form chemical bonds to achieve full valence shells

of electrons• Full valence shell may occur by an exchange or by

sharing electrons• Sharing – covalent; exchange - ionic

Polar Covalent Bonds• When electrons are shared unevenly in a covalent bond

• Example: HF, H2O

Coordinate Covalent Bonds• Both electrons are contributed by one atom

• Example:• NH4

+

• H3O+

• CO• N2O• NHO3

Resonance Structures• Single bonds are longer than double bonds, which are

longer than triple bonds

• Example: SO3

• Resonance Structure: Electron pair is shared over three bond evenly• Delocalized electrons

Less than 8• BeH2

• BCl3

More than 8• Octet rule only applies to the first two periods

• After that, can have expanded octets

• Example:• PF5

• BrF5

• SiF63-

Practice - Worksheet• H2

• F2

• OF2

• O2F2

Valence Bond Theory and Quantum Mechanics

• Covalent bonds occur when orbitals overlap and two electrons occupy the same region of space

• Example: H2

HF• What are the electron configurations for H and F?• How would the orbitals interact

H2O• What are the electron configurations for H and O?• How would the orbitals interact

Problem• We know from experiments in atomic structure that the

bond angle in H2O is 104.5°… not 90° as predicted by valence bond theory

• True for CH4 (109.5°) and NH3 (107.5°) – VBT always predicts bond angles of 90°

• So, we need a better theory…

Hybrization• Two problems still exist from Lewis Bonding Theory

1. Carbon atoms form 4 EQUAL C-H bonds in CH4 (or any other molecule)

• Not predicted due to electron configuration of C• Recall: s orbitals have lower energy than p orbitals,

therefore the bond length would be different

2. Existence of double and triple bonds

Hybridization of Carbon Orbitals• An s electron gets promoted to

the empty p-orbital

• This stabilizes the p- and s- orbitals and gives them all the same energy;• Half-filled subshells

• Called sp3 orbitals (HYBRID ORBITALS)

• Each sp3 orbital lies at 109.5°

Additional Hybrid Orbitals – sp - LINEAR

Additional Hybrid Orbitals – sp2 - PLANAR

Additional Hybrid Orbitals – sp3 - Tetrahedral

Double and Triple Bonds• Two types of orbital overlap exist

• What we have seen so far is one type

1. Sigma bonds: σ-bonds• End-on-end overlap of orbitals

2. Pi bonds:π-bonds• Sideways overlap of orbitals

Sigma Bonds• Occur in single bonds and account for the FIRST bond in

a double or triple bond

• Examples:

Pi Bonds• Occur when p-orbitals not on the bonding axis (py or pz)

overlap with each other

Making Double Bonds• Example: C2H4

• Draw a Lewis Structure

• What occurs with the C atoms hybridization?

• For double bonds, there must be one σ-bond from overlapping hybrid orbitals and one π-bond from overlapping py or pz orbitals

• Come from sp2 hybridized orbitals and result in trigonal

planar structures

Making Triple Bonds• Triple bonds have one σ-bond and two π-bonds; come

from sp-hybridized orbitals, and result in linear structures• Central atom has two un-hybridized p-orbitals

Practice• Explain the structure of the following molecules using

electron configurations, orbital hybridization and VBT.

• C2Cl4

• C2Cl2

• CO2

VSEPR Theory - Valence Shell Electron Pair Repulsion Theory

Work through VSEPR Chart• Fun times with molecular structure…

Practice Problems• Use Lewis Theory and VSEPR Theory to predict the

structure of the following molecules:

• Homework/Practice - Worksheet

Polar Molecules• Polar molecules are molecules where the electron charge

is not distributed evenly

Electronegativity and Polar Covalent Bonds

• Ionic Bond: ΔEN = >1.7• Electron transfer

• Polar Covalent Bond: ΔEN = 0.5-1.7• Electrons shared unevenly

• Pure Covalent Bond: ΔEN = 0.0-0.5• Electrons shared evenly

• Remember: Think of electrons as electron probabilities, electron cloud density is greater around one atom or another, therefore one gets a slight negative, the other slight positive charge

• Think of the scale as a continuum

Polar Molecules• Cannot exist if there are no polar bonds!

• Bond dipole: electronegativity difference of two atoms represented by an arrow pointing from the positive to the negative end (lower to higher EN)

• Non-polar molecule: either perfectly symmetrical so the bond dipoles cancel out, or when no polar bonds exist

• Polar molecule: occur when bond dipoles do not cancelout

• Example:• Determine the polarity of the following molecules

• H2O, CCl4, NH3, PCl5

• Practice:• CH3Cl, BeCl2, SiO2, BrF4

• CHF2Cl• CH3NH2

Intermolecular Forces• Forces that exist between molecules

• Three types:• Dipole-Dipole• Hydrogen Bonds• London Dispersion

• In order to determine the Intermolecular Forces (IMF), you need to first determine the polarity of the molecule

• Much weaker than covalent bonds

Dipole-Dipole Forces • Occur in polar molecules

• The slightly negative end on one molecule is attracted to the slightly positive end on another molecule

• Strength depends on the size of the dipole

London Dispersion Forces• Simultaneous attraction of the electrons in one molecule

to the nuclei in the surrounding molecules

• Increase as the number of electrons and protons in a molecule increase

• Exist in ALL molecules

• Weakest Force

Hydrogen Bonds• Attraction between H on one molecule

and O, N, or F on another molecule

• Strongest of the intermolecular forces

• Found in H2O, NH3, and HF, or whenever there is a –OH, -NH2 in a molecule

Predicting Strength of IMF• Use pol

Predicting Boiling Points • Boiling points increase as IMF strength increases

• Arrange the following molecules in order of increasing boiling points

1. SiH4, SnH4, GeH4, CH4,

2. C3H8, C2H4, C4H10

3. CH4, CCl3H, CBr3H

Practice• Determine the intermolecular forces that exist in each

molecule• CCl4, C5H12, CH3CH2OH

• Which molecules would have the strongest IMF• C2H5OH, C2H6, C2H5Cl• Explain you answer

Structure and Properties of Solids

• Different types of solids result depending on the type of bonding in the solid

• These solids have different properties

Ionic Crystals• Crystal Lattice

• Properties result from the lattice structure

• Brittle, high melting/boiling point, conduct electricity when dissolved in water or in liquid form, hard

Molecular Crystals• Arrangement of neutral molecules held together by weak

intermolecular forces• Properties vary depending on the strength of the IMF

Covalent Network Solids• Array of covalently bonded atoms, structure is held

together by covalent bonds• High MP/BP• Example: Silicate (SiO)

Carbon Network Solids• Diamond, Graphite, Carbon Nanotubes, Buckminster

Fullerenes• Explain the difference in properties between graphite and

diamond.

Metallic Crystals• Lots of electrons, but low ionization energy means they

are loosely held• Lots of empty valence orbitals with similar energy,

therefore electrons are free to move around• Strong, non-directional bonding

Properties of MetalsProperty Explanation

Shiny, Silvery

Flexible

Electrical Conductivity

Hard Solids

Crystalline

End of UNIT 1 – STRUCTURE AND PROPERTIES!!

• Your unit test is on: October 28

• Review package handed out