IONIC BONDING NOTES
Valence electron – electrons in the outermost (highest) energy level
Lewis dot structure – representation of the valence electrons in an atom
Octet rule - atoms tend to gain or lose e- so that they have 8 valence e- (noble gases have 8 & are stable)
Ionic compounds – form crystal structure
They are held together by opposite charges to form stable lattice structure
Opposite charges attract – energy released as they attract, this
overcomes any + H that may result from a rxn.
Lattice Energy – energy required to completely separate 1 mol of an ionic cmpd into its constituent gaseous
ions.
All are +, thus they require E to be added to the system (bond) to break apart
Depends on Q & d (Hey! That’s Coulomb’s Law!!)
Attractive energy primarily depends on charges because difference between radii is negligible
Eg. 1 Which has the largest Lattice Energy – Al2O3, KCl or MgBr2?
Review ions
Lattice energy encourages ion formation, but is not strong enough to strip e- a from noble gas core
Oftentimes it is not strong enough to remove all e- from transition metal to get to noble gas core
Loose highest energy level e- 1st, then removal of d subshell e- depends on stability of the ion
Eg. 2 Write the electron configuration of Hg+.
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Review ion size
Eg.3 How many valence electrons do P3- & Cl- have?
Why do they differ in size?
What metal ion has same # valence e?
Ch. 8 You should know how to do these but they are not hmwk 27-30,37,43,45a,49,52-54
Lewis Structures
Atoms form bonds because they are more stable, ie. at a state of less energy
We are working w/ covalent molecules only from here on out in this unit.
Review Lewis Dot Diagram
Eg. 1 Write the Lewis dot diagram for the following elements.
a) C
b) P
c) Cl
Review of covalent bonding & valence electrons
How many valence electrons does N have?
How many does it need to attain a noble gas configuration?
How many electrons are available for sharing?
How many covalent bonds will form?
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Lewis Structures cont’d
Eg. 2 Draw the Lewis structure for PCl3
1) Add all valence electrons of constituent atoms & divide by 2 to get the number of e- pairs.
2) Write the atom symbols in the correct arrangement
Formulas written so that they are in bonding order (molecular structure)
Central is atom written 1st
Exception: H & acids. For acids with O, the H is bonded to an O that’s connected to the central
atom.
3)
4)
5)
6)
Form single bonds btw bonded atoms.
Complete the octets of atoms attached to central w/ lone pairs
Remaining pairs go on central atom. THIS IS ALWAYS TRUE!!!
If the octet isn’t met, use lone pairs to form double or triple bonds
Eg. 3 Draw Lewis structures for the following molecules
a) F2
b) O2
c) HClO3
d) C2H4
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For ions add or subtract electrons where appropriate
Eg. 4 Draw the Lewis structure for NO3-
Eg. 5 Draw the Lewis structure for SO2
Based on experimentation all SO bonds have same length & strength (half way between single &
double bond)
Resonance structures – different Lewis structures for same molecule
Eg. 7 Draw the Lewis structure for the following
a) CH3CO2-
b) HNO3
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Exceptions to the Octet Rule
Odd #e, not all e- are paired few exceptions
PO2
Can’t satisfy octet, few exceptions & the most common are B molecules.
BCl3 Resonance would force Cl to have double bond this is not likely w/ halogen, stability of no octet is
lower than other electron arrangements. Experimental evidence supports this.
More than octet, most of the exceptions
SF4, XeF4
Practice: Z Ch8 79, 81, 87
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Bond Strength Notes
Atoms form bonds because they are more stable, i.e.at a state of lower energy
Bond energy – amount of E needed to break all bonds in 1 mole of a molecular substance to produce the
gaseous constituant atoms.
In terms of H, all bonds require E to break them so the process is + or endothermic.
The amount of energy required to break the bonds in a molecule helps explain the stability of the
molecule. If a lot of E is required to break the bonds, then the molecule has a more stable bond. Molecules
commonly found in nature are those with a bond arrangement that is the most stable for that molecular
formula.
Bond energy values are listed in a table in your book or other resources. All values are averages because
surrounding atoms affect E required to break the bond.
For example the energy to break the O-H bond in water is (HO-H) 492 kJ/mol, but for methanol it is
(CH3O-H) 435 kJ/mol.
This is NOT considered a large difference
Calculating average bond energies
Eg 1
What is the bond energy for C=O if H for CO2  C + 2O is 1598 kJ?
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Eg. 2 What is the bond energy for the C-H bond if H for the equation C2H4  2C + 4H is 1652 kJ?
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Using Bond Energies to Calculate H for a rxn
Bond E can be used to calculate H if Hfo is not known – Oh NO! We’re not done w/ that?!
rctnt bonds broken -  E prdct bonds formed
NOTICE THIS IS DIFFERENT, WHY?
Eg. 3
Use bond energies to calculate H for the combustion of etOH (ethanol) vapor.
C2H5OH(g) + 3O2(g)  2CO2 + 3H2O
Eg. 4
BLB #52
Bond E & length are related. How?
Practice: Z Ch8 64 66 67 71 76
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LEWIS STRUCTURES & MOLECULAR GEOMETRY NOTES
Atoms form bonds because they are more stable and therefore at a state of less energy
Keep in mind we are working w/ covalent molecules.
Molecular shape depends on bond angles & length
Shape plays an important role in interactions/functioning of molecule (like what it dissolves in).
Molecules w/a central atom use VSEPR theory to describe (valence shell electron pair repulsion)
Pairs of e- repel, so they position themselves as far away as possible, thus repulsion is lowest
Shape depends on # & position of bonds and lone pairs
Electron Geometries
Shape names & angles
Tetrahedral 109.5
Trigonal planar 120
We use atom positions to define molecular shape
Molecular shape
Repulsion order high to low
lone-lone
lone-bond
bond-bond
H2O prime example – bent, lone pairs affect angle size
# of bonds also affect angle size
How to determine the shape of molecule
1) Draw the Lewis structure, include lone pairs on central atom
2) Arrange the e-pairs so it minimizes repulsion (pick the correct electron geometry)
3) Describe shape using the molecular shape (look at number of atoms for the electron geometry &
use the appropriate molecular shape)
Eg.
CH4, BF3, NI3, CS2, NO2-, NO3-
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Eg.
Effect of bond angles C2H4
Exceptions to the octet rule
If there is > octet
5 e-prs trigonal bipyramidal
lone prs take equatorial positions
6 e-prs octahedron; all equidistant, does not matter where lone prs are
Eg. SCl4 (5, see saw), ICl3 (5, t-shaped), IF5 (6, square pyramidal)
No Central atom
Look at the geometry of each “central” atom
Eg. CH3-CH3, C3H5 (trigonal planar mid C) HC=C-CH=CH2 (linear C) H2N-CH2-COOH
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Eg. Predict bond angles of ammonia (107), water (104) acetic acid (109.5,180, >120 c=o, <109.5 c-o-h)
Review Polarity
Degree of polarity or dipole moment
Polarity of the molecule depends on polarity of bonds & the shape of the molecule. A molecule can have a
polar bond but due to shape can be nonpolar.
Eg. CH4, CS2, H2S, C2H2Cl2 shape determines polarity
Practice: Z Ch 8 107-114
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BONDING HYBRIDIZATION NOTES
Atoms form bonds because they are more stable and therefore at a state of less energy
Keep in mind we are working w/ covalent molecules.
Ionic substances form crystal lattices due to attraction of opposite ions.
VSEPR is incomplete & doesn’t explain how bonding occurs between atoms of a mlcl, i.e. how orbitals
overlap to share
Valence bond theory –
e- is shared in the space btw the nuclei where the orbitals of the 2 atoms overlap
e- attracted to the + nucleus of both atoms so e- are [ ] in the region between the 2 nuclie thereby
holding the 2 atoms together.
Fig 9.11
Hybrid orbitals
The easy way.
Hybrid orbitals take on a blended shape.
Eg. HCl
Eg. CO2
Eg. BCl3
Eg. CH4
Eg. N2
Eg. C2H2
Eg. PF5
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Z Ch 9 23 (for 87), 33, 59, 61