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CHEMISTRY REVIEW
Atoms, Stoichiometry, Reactions and Bonding
MODERN VIEW
• The atom is mostly
empty space.
• Two regions
• Nucleus- protons and
neutrons.
• Electron cloud- region
where you might find an
electron.
SUB-ATOMIC PARTICLES
• Z - atomic number = number of protons determines
type of atom.
• A - mass number = number of protons + neutrons.
• Number of protons = number of electrons if neutral.
A
SYMBOLS
X
Z
23
Na
11
CHEMICAL BONDS
• The forces that hold atoms together.
• Covalent bonding - sharing electrons.
• Makes molecules.
• Chemical formula- the number and type of atoms in a
molecule.
• C2H6 - 2 carbon atoms, 6 hydrogen atoms,
• Structural formula shows the connections, but not
necessarily the shape.
Quantum Numbers
Each electron in an atom has a unique
set of 4 quantum numbers which describe
it.
 Principal quantum number
(n)
 Angular momentum quantum number (l)
 Magnetic quantum number (m)
 Spin quantum number
(s)
Assigning the Numbers
 The three quantum numbers (n, l, and m) are
integers.
 The principal quantum number (n) cannot be
zero.
 n must be 1, 2, 3, etc.
 The angular momentum quantum number (l )
can be any integer between 0 and n - 1.
 For n = 3, l can be either 0, 1, or 2.
 The magnetic quantum number (ml) can be any
integer between -l and +l.
 For l = 2, m can be either -2, -1, 0, +1, +2.
Principle, angular momentum, and magnetic
quantum numbers: n, l, and ml
E
Periodicity
TRENDS IN ATOMIC SIZE
• Influenced by three factors.
• Energy Level
• Higher energy level is further away.
• Charge on nucleus
• More charge pulls electrons in closer.
• Shielding
• Layers of electrons shield from nuclear
pull.
SHIELDING
• The electron on the outside
energy level has to look
through all the other energy
levels to see the nucleus.
• A second electron has the
same shielding.
GROUP
TRENDS
• As we go down a
group
• Each atom has
another energy
level,
• So the atoms get
bigger.
H
Li
Na
K
Rb
PERIODIC TRENDS
•
•
•
•
As you go across a period the radius gets smaller.
Same energy level.
More nuclear charge.
Outermost electrons are closer.
Na
Mg
Al
Si
P
S Cl Ar
IONIC SIZE
• Cations form by losing electrons.
• Cations are smaller that the atom they
come from.
• Metals form cations.
• Cations of representative elements have
noble gas configuration.
IONIC SIZE
• Anions form by gaining electrons.
• Anions are bigger that the atom they come from.
• Nonmetals form anions.
• Anions of representative elements have noble gas
configuration.
IONIZATION
ENERGY
• The amount of energy required to completely remove an
electron from a gaseous atom.
• Removing one electron makes a +1 ion.
• The energy required is called the first ionization energy.
IONIZATION ENERGY
• The second ionization energy is the energy required to
remove the second electron.
• Always greater than first IE.
• The third IE is the energy required to remove a third
electron.
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11810
14840
3569
4619
4577
5301
6045
6276
WHAT DETERMINES IE
• The greater the nuclear charge the greater IE.
• Distance from nucleus increases IE
• Filled and half filled orbitals have lower energy, so
achieving them is easier, lower IE.
• Shielding
GROUP TRENDS
• As you go down a group first IE
decreases because
• The electron is further away.
• More shielding.
PERIODIC TRENDS
• All the atoms in the same period have the same energy
level.
• Same shielding.
• Increasing nuclear charge
• So IE generally increases from left to right.
• Exceptions at full and 1/2 fill orbitals.
Ne
He
First Ionization
energy
N F
 Na
has a lower
IE than Li
are s1
 Na has more
shielding
 Greater distance
 Both
H
C O
Be
Li
B
Na
Atomic
number
Electron Affinity - the energy change
associated with the addition of an electron
 Affinity tends to increase across a period
 Affinity tends to decrease as you go down
in a group
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
ELECTRONEGATIVITY
• The tendency for an atom to attract electrons to itself
when it is chemically combined with another element.
• How fair it shares.
• Big electronegativity means it pulls the electron toward
it.
• Atoms with large negative electron affinity have larger
electronegativity.
GROUP TREND
• The further down a group the farther the electron is
away and the more electrons an atom has.
• More willing to share.
• Low electronegativity.
PERIODIC TREND
• Metals are at the left end.
• They let their electrons go easily
• Low electronegativity
• At the right end are the nonmetals.
• They want more electrons.
• Try to take them away.
• High electronegativity.
Ionization energy,
electronegativity
Electron affinity INCREASE
Atomic size increases,
shielding constant
Ionic size increases
Prepare yourself to
^
C
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light.
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Types of electromagnetic radiation:
PROBLEM 37
The energy (E ) of electromagnetic
radiation is directly proportional to the
frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
 = frequency, in units of hertz (hz, sec-1)
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table
Relating Frequency, Wavelength
and Energy
c 
E  h
Common re-arrangements:
E
hc

hc

E
PROBLEM 43
PES
• method that provides information on all the
occupied energy levels of an atom (that is,
the ionization energies of all electrons in the
atom) is known as photoelectron
spectroscopy; this method uses a photon (a
packet of light energy) to knock an electron
out of an atom.
PHOTOELECTRON SPECTRUM
The photoelectron spectrum is a plot of the number of electrons
emitted versus their kinetic energy. In the diagram below, the
“X” axis is labeled high to low energies so that you think about
the XY intersect as being the nucleus.


http://www.chem.arizona.edu/chemt/Flash/ph
otoelectron.html
2p
61s
21-
2s
3p
3s
Orbital names s,
p, d, and f stand
for names given
to groups of lines
in the spectra of
4s the alkali metals.
Early chemists
called the line
groups sharp,
principal,
diffuse, and
fundamental.
Interpretations from the data:
1. There are no values on the y axis in the tables above. Using the Periodic Table
and Table 1, put numbers on the y axis.
2. Label each peak on the graphs above with s, p, d, or f to indicate the suborbital
they represent..
3. What is the total number of electrons in a neutral potassium atom?
IONS
• Atoms or groups of atoms with a charge.
• Cations- positive ions - get by losing electrons(s).
• Anions- negative ions - get by gaining electron(s).
• Ionic bonding- held together by the opposite charges.
• Ionic solids are called salts.
POLYATOMIC IONS
• Groups of atoms that have a charge.
• Yes, you have to know common ones.
Periodic Table
• Conductors
• Lose electrons
• Malleable and ductile
METALS
• Brittle
• Gain electrons
• Covalent bonds
NONMETALS
SEMI-METALS OR
METALLOIDS
Alkali Metals
Alkaline Earth Metals
Halogens
Transition metals
Noble Gases
Inner Transition Metals
+1+2
-3 -2 -1
NAMING COMPOUNDS
• Two types
• Ionic - metal and non metal or polyatomics.
• Covalent- we will just learn the rules for 2 non-metals.
IONIC COMPOUNDS
• If the cation is monoatomic- Name the metal (cation)
just write the name.
• If the cation is polyatomic- name it.
• If the anion is monoatomic- name it but change the
ending to –ide.
• If the anion is poly atomic- just name it
• Practice.
• Two words, with prefixes.
• Prefixes tell you how many.
COVALENT
COMPOUNDS
• mono, di, tri, tetra, penta, hexa, septa, octa, nona,
deca
• First element whole name with the appropriate prefix,
except mono.
• Second element, -ide ending with appropriate prefix.
• Practice
ACIDS
• Substances that produce H
water.
• All acids begin with H.
• Two types of acids:
• Oxyacids
• Non-oxyacids
+ ions when dissolved in
NAMING ACIDS
• If the formula has oxygen in it
• write the name of the anion, but change
• ate to -ic acid
• ite to -ous acid
• Watch out for sulfuric and sulfurous
• H2CrO4
• HMnO4
• HNO2
NAMING ACIDS
• If the acid doesn’t have oxygen
• add the prefix hydro• change the suffix -ide to -ic acid
• HCl
• H2 S
• HCN
HYDRATES
• Some salts trap water when they form crystals.
• These are hydrates.
• Both the name and the formula needs to indicate how
many water molecules are trapped.
• In the name we add the word hydrate with a prefix
that tells us how many water molecules.
HYDRATES
• In the formula you put a dot and then write the number
of molecules.
• Calcium chloride dihydrate = CaCl22O
• Chromium (III) nitrate hexahydrate = Cr(NO3)3 6H2O
THE MOLE
• The mole is a number.
• A very large number, but still, just a number.
• 6.022 x 10
23 of anything is a mole
• A large dozen.
• The number of atoms in exactly 12 grams of carbon-12.
MOLAR MASS
• Mass of 1 mole of a substance.
• Often called molecular weight.
• To determine the molar mass of an element, look on
the table.
• To determine the molar mass of a compound, add up
the molar masses of the elements that make it up.
FIND THE MOLAR MASS
OF
• CH4
• Mg3P2
• Ca(NO3)3
• Al2(Cr2O7)3
• CaSO4 · 2H2O
D
PERCENT COMPOSITION
• Percent of each element a compound is composed
of.
• Find the mass of each element, divide by the total
mass, multiply by a 100.
• Easiest if you use a mole of the compound.
• Find the percent composition of CH4
• Al2(Cr2O7)3
• CaSO4 · 2H2O
WORKING BACKWARDS
• From percent composition, you can determine the
empirical formula.
• Empirical Formula the lowest ratio of atoms in a
molecule.
• Based on mole ratios.
• A sample is 59.53% C, 5.38%H, 10.68%N, and 24.40%O
what is its empirical formula.
• A 0.2000 gram sample of a compound (vitamin C)
composed of only C, H, and O is burned completely
with excess O2 . 0.2998 g of CO2 and 0.0819 g of H2O
are produced. What is the empirical formula?
EMPIRICAL TO MOLECULAR
FORMULAS
• Empirical is lowest ratio.
• Molecular is actual molecule.
• Need Molar mass.
• Ratio of empirical to molar mass will tell you the
molecular formula.
• Must be a whole number because...
EXAMPLE
• A compound is made of only sulfur and oxygen. It is
69.6% S by mass. Its molar mass is 184 g/mol. What is its
formula?
CHEMICAL EQUATIONS
• Are sentences.
• Describe what happens in a chemical reaction.
• Reactants  Products
• Equations should be balanced.
• Have the same number of each kind of atoms on both
sides because ...
ABBREVIATIONS
• (s) 
• (g)
• (aq)
• heat
•
D
• catalyst
MEANING
• A balanced equation can be used to describe a
reaction in molecules and atoms.
• Not grams.
• Chemical reactions happen molecules at a time
• or dozens of molecules at a time
• or moles of molecules.
STOICHIOMETRY
• Given an amount of either starting material or product,
determining the other quantities.
• use conversion factors from
• molar mass (g - mole)
• balanced equation (mole - mole)
• keep track.
EXAMPLES
• One way of producing O2(g) involves the
decomposition of potassium chlorate into
potassium chloride and oxygen gas. A 25.5 g
sample of Potassium chlorate is decomposed. How
many moles of O2(g) are produced?
• How many grams of potassium chloride?
• How many grams of oxygen?
EXAMPLES
• A piece of aluminum foil 5.11 in x 3.23 in x
0.0381 in is dissolved in excess HCl(aq). How
many grams of H2(g) are produced?
• How many grams of each reactant are
needed to produce 15 grams of iron form
the following reaction?
Fe2O3(s) + Al(s)  Fe(s) + Al2O3(s)
YIELD
How much you get from an chemical reaction
LIMITING REAGENT
• Reactant that determines the amount of product
formed.
• The one you run out of first.
• Makes the least product.
• Book shows you a ratio method.
• It works.
• So does mine
LIMITING REAGENT
• To determine the limiting reagent requires that you do
two stoichiometry problems.
• Figure out how much product each reactant makes.
• The one that makes the least is the limiting reagent.
EXAMPLE
• Ammonia is produced by the following
reaction
N2 + H2  NH3
What
mass of ammonia can be produced from a
mixture of 100. g N2 and 500. g H2 ?
• How much unreacted material remains?
EXCESS REAGENT
• The reactant you don’t run out of.
• The amount of stuff you make is the yield.
• The theoretical yield is the amount you
would make if reaction went to
completion.
• The actual yield is what you make in the
lab.
PERCENT YIELD
• % yield = Actual
Theoretical
• % yield =
x 100%
what you got
x 100%
what you could have got
TYPES OF REACTIONS
 Precipitation reactions
• When aqueous solutions of ionic compounds are
poured together a solid forms.
• A solid that forms from mixed solutions is a precipitate
• If you’re not a part of the solution, your part of the
precipitate
PRECIPITATION
REACTIONS
• NaOH(aq) + FeCl3(aq) NaCl(aq) + Fe(OH)3(s)
• is really
+
-
+3
-
• Na (aq)+OH (aq) + Fe
+ Cl (aq) 
+
Na (aq) + Cl (aq) + Fe(OH)3(s)
• So all that really happens is
+3
• OH (aq) + Fe
 Fe(OH)3(s)
• Double replacement reaction
SOLUBILITY RULES
All nitrates are soluble
+
Alkali metals ions and NH4 ions are soluble
+
+2
Halides are soluble except Ag , Pb , and
+2
Hg2
Most sulfates are soluble, except
+2
+
+2
+2
+2
+2
Ag ,Pb , Hg
,Sr , Ca
and Ba
SOLUBILITY RULES
 Most hydroxides are slightly soluble
(insoluble) except NaOH and KOH
 Sulfides, carbonates, chromates, and
phosphates are insoluble
 Lower number rules supersede so Na2S is
soluble
THREE TYPES OF
EQUATIONS
• Molecular Equation- written as whole formulas, not the
ions.
• K2CrO4(aq) + Ba(NO3)2(aq) 
• Complete Ionic equation show dissolved electrolytes as
the ions.
+ + CrO -2 + Ba+2 + 2 NO - 
4
3
+
BaCrO4(s) + 2K + 2 NO3
• 2K
• Spectator ions are those that don’t react.
THREE TYPE OF
EQUATIONS
• Net Ionic equations show only those ions that react, not
the spectator ions
• Ba
+2 + CrO -2  BaCrO (s)
4
4
• Write the three types of equations for the reactions
when these solutions are mixed.
• iron (III) sulfate and potassium sulfide Lead (II) nitrate
and sulfuric acid.
TYPES OF REACTIONS
 Acid-Base
• For our purposes an acid is a proton donor.
• a base is a proton acceptor usually OH
-
• What is the net ionic equation for the reaction of
HCl(aq) and KOH(aq)?
• Acid + Base  salt + water
• H
+ + OH-  H O
2
ACID - BASE REACTIONS
• Often called a neutralization reaction Because the
acid neutralizes the base.
• Often titrate to determine concentrations.
• Solution of known concentration (titrant),
• is added to the unknown (analyte),
• until the equivalence point is reached where enough
titrant has been added to neutralize it.
TITRATION
• Where the indicator changes color is the endpoint.
• Not always at the equivalence point.
• A 50.00 mL sample of aqueous Ca(OH)2 requires 34.66
mL of 0.0980 M Nitric acid for neutralization. What is
[Ca(OH)2 ]?
• # of H
+ x M x V = # of OH- x M x V
A A
B B
ACID-BASE REACTION
• 75 mL of 0.25M HCl is mixed with 225 mL of 0.055 M
+
Ba(OH)2 . What is the concentration of the excess H
or OH ?
TYPES OF REACTION
 Oxidation-Reduction called Redox
• Ionic compounds are formed through the transfer of
electrons.
• An Oxidation-reduction reaction involves the transfer of
electrons.
• We need a way of keeping track.
OXIDATION STATES
• A way of keeping track of the electrons.
• Not necessarily true of what is in nature, but it works.
• need the rules for assigning (memorize).
 The oxidation state of elements in their standard states
is zero.
 Oxidation state for monoatomic ions are the same as
their charge.
OXIDATION STATES
 Oxygen is assigned an oxidation state of -2 in its covalent
compounds except as a peroxide.
 In compounds with nonmetals hydrogen is assigned the
oxidation state +1.
 In its compounds fluorine is always –1.
 The sum of the oxidation states must be zero in
compounds or equal the charge of the ion.
OXIDATION STATES
• Assign the oxidation states to each element in the
following.
• CO2
-
• NO3
• H2SO4
• Fe2O3
• Fe3O4
E
OXIDATION-REDUCTION
• Transfer electrons, so the oxidation states change.
• Na + 2Cl2  2NaCl
• CH4 + 2O2  CO2 + 2H2O
• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.
• OIL RIG
• LEO GER
HALF-REACTIONS
• All redox reactions can be thought of as happening in
two halves.
• One produces electrons - Oxidation half.
• The other requires electrons - Reduction half.
• Write the half reactions for the following.
• Na + Cl2  Na
-
+ + Cl-
-
• SO3 + H+ + MnO4-  SO4 + H2O + Mn
+2
BALANCING REDOX
EQUATIONS
• In aqueous solutions the key is the number of electrons
produced must be the same as those required.
• For reactions in acidic solution an 8 step procedure.
 Write separate half reactions
 For each half reaction balance all reactants except H
and O
 Balance O using H2O
ACIDIC SOLUTION
 Balance H using H+
 Balance charge using e Multiply equations to make electrons equal
 Add equations and cancel identical species
 Check that charges and elements are balanced.
PRACTICE
• The following reactions occur in aqueous solution.
Balance them
-
-
• Cr(OH)3 + OCl + OH 
-2
CrO4 + Cl + H2O
-
+2 Mn+2 + Fe+3
+2 + NO(g)
• Cu + NO3  Cu
-2  PbSO
• Pb + PbO2 + SO4
4
+2 + NaBiO  Bi+3 + MnO • Mn
3
4
• MnO4 + Fe
REDOX TITRATIONS
• Same as any other titration.
• the permanganate ion is used often because it is its
+2 is colorless.
own indicator. MnO4 is purple, Mn
When reaction solution remains clear, MnO4 is gone.
• Chromate ion is also useful, but color change,
orangish yellow to green, is harder to detect.
EXAMPLE
• The iron content of iron ore can be determined by
titration with standard KMnO4 solution. The iron ore is
+2
dissolved in excess HCl, and the iron reduced to Fe
ions. This solution is then titrated with KMnO4 solution,
+3 and Mn+2 ions in acidic solution. If it
producing Fe
requires 41.95 mL of 0.205 M KMnO4 to titrate a solution
made with 0.6128 g of iron ore, what percent of the ore
was iron?
B
CHAPTER 8
Bonding
WHAT IS A BOND?
• A force that holds atoms together.
• We will look at it in terms of energy.
• Bond energy - the energy required to break a bond.
• Why are compounds formed?
• Because it gives the system the lowest energy.
IONIC BONDING
• An atom with a low ionization energy reacts with an
atom with high electron affinity.
• The electron moves.
• Opposite charges hold the atoms together.
COULOMB'S LAW
• E= 2.31 x 10-19 J · nm(Q1Q2)/r
• Q is the charge.
• r is the distance between the centers.
• If charges are opposite, E is negative
• exothermic
• Same charge, positive E, requires energy to bring them
together.
• endothermic
WHAT ABOUT COVALENT
COMPOUNDS?
• The electrons in each atom are attracted to the nucleus of
the other.
• The electrons repel each other,
• The nuclei repel each other.
• They reach a distance with the lowest possible energy.
• The distance between is the bond length.
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Bond Length
Internuclear Distance
Energy
Bond Energy
0
Internuclear Distance
Bond Length Diagram
COVALENT BONDING
• Electrons are shared by atoms.
• There are two extremes.
• In between are polar covalent bonds.
• The electrons are not shared evenly.
• One end is slightly positive, the other negative.
• Indicated using small delta d.
Electronegativity:
D
Ionic Bonds
 Electrons are transferred
 Electronegativity differences are
generally greater than 1.7
 The formation of ionic bonds is
always exothermic!
DIPOLE MOMENTS
• A molecule with a center of negative charge and a
center of positive charge is dipolar (two poles),
• or has a dipole moment.
• Center of charge doesn’t have to be on an atom.
• Will line up in the presence of an electric field.
2.1
2.5
H
S
S
H
H
d-
S
H
d+
H
d+
d-
S
H
d+
H
d+
H
d+
H
d+
d-
S
d-
S
H
H
d-
S
H
d+
d+
d+
H
d+
H
H
d+
d+
d-
S
d-
S
H
d+
H
d+
HOW IT IS DRAWN
H
d+
d-
S
H
d+
WHICH MOLECULES HAVE
DIPOLE MOMENTS?
• Any two atom molecule with a polar bond.
• With three or more atoms there are two considerations.
• There must be a polar bond.
• Geometry can’t cancel it out.
GEOMETRY AND POLARITY
• Three shapes will cancel them out.
• Linear
GEOMETRY AND POLARITY
• Three shapes will cancel them out.
• Planar triangles
120º
GEOMETRY AND POLARITY
• Three shapes will cancel them out.
• Tetrahedral
GEOMETRY AND POLARITY
• Others don’t cancel
• Bent
GEOMETRY AND POLARITY
• Others don’t cancel
• Trigonal Pyramidal
IONS
• Atoms tend to react to form noble gas configuration.
• Metals lose electrons to form cations
• Nonmetals can share electrons in covalent bonds.
• When two non metals react.(more later)
• Or they can gain electrons to form anions.
IONIC COMPOUNDS
• We mean the solid crystal.
• Ions align themselves to maximize attractions between
opposite charges,
• and to minimize repulsion between like ions.
• Can stabilize ions that would be unstable as a gas.
• React to achieve noble gas configuration
PERIODIC TRENDS
• Across the period nuclear charge increases so they get
smaller.
• Energy level changes between anions and cations.
Li+1
Be+2
B+3
C+4
N-3
O-2
F-1
SIZE OF ISOELECTRONIC IONS
• Iso - same
• Iso electronic ions have the same # of electrons
• Al
+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
• All have 10 electrons.
2 2
• All have the configuration 1s 2s 2p
6
SIZE OF ISOELECTRONIC IONS
• Positive ions have more protons so they are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
USING BOND ENERGIES
• We can find DH for a reaction.
• It takes energy to break bonds, and end up with atoms (+).
• We get energy when we use atoms to form bonds (-).
• If we add up the energy it took to break the bonds, and
subtract the energy we get from forming the bonds we get
the DH.
• Energy and Enthalpy are state functions.
FIND THE ENERGY FOR THIS
2 CH2 = CHCH3 + 2NH3 + O2
 2 CH2 = CHC  N + 6 H2O
C-H 413 kJ/mol
C=C 614kJ/mol
N-H 391 kJ/mol
C-C 347 kJ/mol
O-H 467 kJ/mol
O=O 495 kJ/mol
CN 891 kJ/mol
LEWIS STRUCTURE
• Shows how the valence electrons are arranged.
• One dot for each valence electron.
• A stable compound has all its atoms with a noble gas
configuration.
• Hydrogen follows the duet rule.
• The rest follow the octet rule.
• Bonding pair is the one between the symbols.
RULES
• Sum the valence electrons.
• Use a pair to form a bond between each pair of
atoms.
• Arrange the rest to fulfill the octet rule (except for H
and the duet).
• H 2O
• A line can be used instead of a pair.
• POCl3 P is central atom
-2 S is central atom
• SO4
-2
• SO3
S is central atom
-2 P is central atom
• PO4
• SCl2
S is central atom
EXCEPTIONS TO THE OCTET
• BH3
• Be and B often do not achieve octet
• Have less than and octet, for electron deficient
molecules.
•
•
•
•
SF6
Third row and larger elements can exceed the octet
Use 3d orbitals?
I3
EXCEPTIONS TO THE OCTET
• When we must exceed the octet, extra electrons go
on central atom.
• ClF3
• XeO3
-
• ICl4
• BeCl2
RESONANCE
• Sometimes there is more than one valid
structure for an molecule or ion.
• NO3
• Use double arrows to indicate it is the
“average” of the structures.
• It doesn’t switch between them.
• NO2• Localized electron model is based on
pairs of electrons, doesn’t deal with odd
numbers.
•
FORMAL
CHARGE
The difference between the number of valence
electrons on the free atom and that assigned in the
molecule.
• We count half the electrons in each bond as
“belonging” to the atom.
-2
• SO4
• Molecules try to achieve as low a formal charge as
possible.
• Negative formal charges should be on electronegative
elements.
EXAMPLES
• XeO3
• NO4
-3
• SO2Cl2
VSEPR
• Lewis structures tell us how the atoms are connected to
each other.
• They don’t tell us anything about shape.
• The shape of a molecule can greatly affect its
properties.
• Valence Shell Electron Pair Repulsion Theory allows us
to predict geometry
VSEPR
• Molecules take a shape that puts electron pairs as far
away from each other as possible.
• Have to draw the Lewis structure to determine electron
pairs.
• bonding
• nonbonding lone pair
• Lone pair take more space.
• Multiple bonds count as one pair.
VSEPR
• The number of pairs determines
• bond angles
• underlying structure
• The number of atoms determines
• actual shape
VSEPR
Electron Bond
pairs
Angles
2
180°
Underlying
Shape
Linear
3
120°
4
109.5°
Tetrahedral
5
90° &
120°
6
90°
Trigonal
Bipyramidal
Octagonal
Trigonal Planar
A
C
C
B
SP HYBRIDIZATION
• End up with two lobes 180º apart.
• p orbitals are at right angles
• Makes room for two p bonds and two
sigma bonds.
• A triple bond or two double bonds.
SP2 HYBRIDIZATION
• When three things come off atom.
• trigonal planar
• 120º
• on s one p bond
SP2 HYBRIDIZATION
• C2H4
• Double bond acts as one region of electrons.
• trigonal planar
• Have to end up with three blended orbitals.
• Use one s and two p orbitals to make sp
• Leaves one p orbital perpendicular.
2 orbitals.
SP3 HYBRIDIZATION
• We blend the s and p orbitals of the valence electrons
and end up with the tetrahedral geometry.
• We combine one s orbital and 3 p orbitals.
•
3
sp
hybridization has tetrahedral geometry.
HOW WE GET TO HYBRIDIZATION
• We know the geometry from
experimentation
• We know the orbitals of the atom
• hybridizing atomic orbitals can explain
the geometry.
• So if the geometry requires a
tetrahedral shape, it is sp3 hybridized
• This includes bent and trigonal
pyramidal molecules because one of
the sp3 lobes holds the lone pair.
3
SP D HYBRIDIZATION
• The model predicts that we must use the d orbitals.
3
• sp d hybridization, ex: PCl5
• Trigonal bipyrimidal
• can only s bond.
• can’t p bond.
• basic shape for five bonds.
• gets us to six bonds around
• octahedral
SP3D2
TWO TYPES OF BONDS
• Sigma bonds from overlap of orbitals.
• Between the atoms.
• Pi bond (p bond) above and below atoms
• Between adjacent p orbitals.
• The two bonds of a
double bond.
• C can make two s and two p
• O can make one s and one p
O
C O
CO2
Let’s Try Some Problems
B
B
C
D
A
A
B
D
B
B
B
D
D
E
B
D
E
D