Unit 1 Notes - Structure and Properties of Matter

• 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
• 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-MetalLow IE
Low EA
High IE
High EA
Isoelectronic with noble gases
Opposites attract in no particular direction, considered nondirectional
Ions 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
• # 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
• Covalent bonds occur when orbitals overlap and two
electrons occupy the same region of space
• Example: H2
• What are the electron configurations for H and F?
• How would the orbitals interact
• What are the electron configurations for H and O?
• How would the orbitals interact
• 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…
• 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
• 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
• 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
• 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
• 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
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
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
• Determine the intermolecular forces that exist in each
• 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
• 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
• Explain the difference in properties between graphite and
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 Metals
Shiny, Silvery
Hard Solids
• Your unit test is on: October 28
• Review package handed out

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