Honors Chemistry Chapters 7 and 8(and a small portion of Chapter 9) Notes
(Student’s edition)
Chapter 7 problem set:53, 60, 67, 73, 83
Chapter 8 problem set:40, 43, 45, 47, 54, 56, 58, 61, 63, 65-68, 74, 79, 80
Useful diagrams: This book has excellent figures and tables – most of them have something to offer you intellectually.
Start of Ch 9 Manipulating Polyatomic Ions and Naming Monatomic Ions (part of Ch9)
-Know the basic polyatomic ions:
Name / Formula / Name / FormulaHydroxide / OH-1 / Nitrate / NO3-1
Ammonium / NH4+1 / Phosphate / PO4-3
Acetate / C2H3O2-1 / Chromate / CrO4-2
Carbonate / CO3-2 / Dichromate / Cr2O7-2
Sulfate / SO4-2 / Chlorate / ClO3-1
Permanganate / MnO4-1 / Oxalate / C2O4-2
Tartrate / C4H4O6-2 / Cyanide / CN-1
- Rule #1Change the number of oxygens:
remove one oxygen = change ending of name to ____
remove two oxygens = change ending of name to ____ and
beginning of name to ______
add one oxygen = change beginning of name to _____
Examples:
Name / Formula / Name / Formula / Name / Formula / Name / FormulaChlorate / ClO3-1 / Sulfate / SO4-2 / Nitrate / NO3-1 / Phosphate / PO4-3
Chlorite / Sulfite / NO2-1 / PO3-3
Hypochlorite / Hyposulfite / NO-1 / PO2-3
Perchlorate / Persulfate / NO4-1 / PO5-3
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- Rule #2Other family members:
Elements near each other in the same column tend to form similar
polyatomic ions.
Examples:
Name / Formula / Name / Formula / Name / FormulaChlorate / ClO3-1 / Sulfate / SO4-2 / Phosphate / PO4-3
Fluorate / Selenate / Arsenate
Iodate
Bromate
- Rule #3Add Hydrogen
Add only one H = change the beginning of the name to ____ and
make the charge one ____ negative (due to hydrogen’s
positive one charge)
Examples:
Name / Formula / Name / FormulaCarbonate / CO3-2 / Sulfate / SO4-2
Bicarbonate / HSO4-1
-Combine rules 1, 2, and 3
Examples:
Ex1: What is the formula for hypoiodite?
Find on the periodic table. It is near __. Chlorine forms
chlorate (______). Therefore, Iodine forms iodate (___).
The –ite and hypo- in hypoiodite mean that iodate has lost
______.
Hypoiodite =
Ex2: What is the formula for Biperselenate?
Find Se on the periodic table. It is near __. Sulfur forms
sulfate (____). Therefore, selenium forms selenate (____).
The per- in biperselenate means that selenate has gained
______. Also, the bi- means that it has gained a
______(don’t forget to change the ______!).
Biperselenate =
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-Monatomic Ions
For nonmetals, almost all single names that end with ____ indicates a single
charged atom.
Simply write the ______and the ______. The periodic table column indirectly
indicates the element’s ______. Remember, elements want to have __
electrons in their outer shell (______). For example, column #1 elements have a __ charge, column #2 elements have a __ charge, column
#3 = ___, column #5 = ___, column #6 = ___, and column #7 = ___.
Examples:
Ex1: What is the formula for chloride?
Ex2: What is the formula for an aluminum ion?
Ex3: What is the name of the S-2 anion?
Ex4: What is the name of the Mg+2 cation?
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NIB – an introduction: Attachment Between Atoms
-Most elements are not found in nature. They are
Chemical Bond -
.
-Types of chemical bonds:
Ionic -
Covalent -
Metallic -
8.4 Polar Bonds and Polar Molecules
-Polar -
-Nonpolar -
-There are two ways to predict polar vs nonpolar ( and covalent vs ionic)
#1 Use electronegativity difference
nonpolar covalent
0.4 - 1.7 - polar covalent
greater than 1.7 - ionic
Examples:
NaClHClCl2
Cl = Cl = Cl =
Na = H = Cl =
#2 - There is an easier way to predict
Ionic = or
Polar Covalent =
Nonpolar Covalent =
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7.1 and 7.2 Ions and Ionic bonds and Ionic Compounds
-Ionic compound - a substance composed of positive and neg. ions so that the charges
are . It involves a of electrons.
Ca+2 with Cl–1 will form the compound . It takes
chlorine ions to cancel out the the +2 charge on the calcium ion.
-Formula unit - lowest whole # ratio of ions
- Metals - electrons - why? They form cations
- Nonmetals - electrons - why? . They form anions
- Metals lose electrons until they become like a .
(usually 8 valence electrons)
-Nonmetals gain electrons until they do the same.
-Both go to s2p6 - valence electrons - called a
-The tendency to arrange electrons so each atom has 8 is called the .
-The formation of an ionic bond:
Na wants to ______
1s 2s 2p 3s
Cl wants to ______
1s 2s 2p 3s 3p
an easier way.....
-The ionic bonding picture looks like this....
Ex1: NatoCl
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-Other examples of ionic bonding pictures:
Ex2: BatoCl
Ex3: AltoN
Ex4: NatoS
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Ex5: Al toS
-Ions make up substances - not atoms
-Crystal Lattice - regular repeating pattern of in an ionic substance
-Ions are held in fixed position by .
-A amount of energy is needed to break this structure. Because of this, ionic compounds have a high .
-Ionics can’t conduct electricity as , but they can when dissolved in .
-Property summary of ionic compounds - hard, shatter (not ), conduct when in , high melting point, no smell (low )
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NIB(although this is covered somewhat in section 8.2)
-Energy is involved in all .
Na + Cl yields NaCl + 769 kJ
-Lattice energy - energy released when an forms.
NaCl = - 769 kJ/moleNaF = - 922 kJ/moleKCl = -718 kJ/mole
smaller ions have
7.3 Bonding in Metals
-“Sea of electrons theory”
-Properties of metals -
- Alloy - solution
- Amalgam - solution
-Alloys are important because they have properties that are better than the properties of
the .
8.1 and 8.2 Molecular Compounds and The Nature of Covalent Bonding
-Molecule - smallest quantity of matter that can exist by itself and still retain the
properties of that substance. Usually used when describing a covalently bonded substance.
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-monatomic molecules -
-diatomic molecules -
-polyatomic molecules -
-chemical formulas show the relative #’s of atoms in a chemical compound
ex.C6H12O6C = __, H = __, O =__
Pb(NO3)2Pb = __, N = __, O =__
(NH4)2Cr2O7 N = __, H = __, Cr = __, O =__
-The formation of a covalent bond: bond length
attraction
repulsion=
-Bond Length vs. Bond Energy
Bond length = Bond Energy = Strength
-Diatomic Molecules and Orbital Notation (Orbital overlap or notation diagrams):
H2
1sThis is a bond.
1s
O2
1s 2s 2pThis is a bond.
1s 2s2p
N2
1s 2s 2pThis is a bond.
bond length
bond energy
1s 2s2p
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- Octet Rule-
______.
-Electron cloud representations (overlapping orbitals):
F2 HCl
______
- HF - orbital overlap diagram
H
1s
F
1s 2s2p
Dot Diagrams of molecules and polyatomic ions
-Lewis Structures for covalent compounds
-Basic rules (try to use as many as the rules as possible)
1. Each atom wants 8 electrons (except H wants 2).
2. Each atom goes for close to the right # of bonds.
3. The least electronegative atoms goes in the middle.
- The atom that makes the most bonds goes in the middle. H always on the outside.
- The single atom (the atom that does not have a subscript after it) goes in the middle.
- Attempt to make the structures as symmetrical as possible. Place the
atoms in order (left, right, bottom, and top) around a central atom.
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-Examples: Draw the following Lewis structures
Ex1:H2O
Ex2:PCl3
Ex3:SiH2F2
Ex4:CS2
Ex5:C2H6
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Ex6: C2H4
Ex7:C2H2
Ex8:CH2O
Ex9:HCN
Ex10:FON
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-Drawing polyatomic ions:
less bonds than atoms want = charge more bonds than atoms want = charge
Ex11: PO4-3
-coordinate covalent bond -
Ex12:NH4+
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Ex13:OH-1
Ex14:sulfate
Ex15: nitrate
Ex16:nitrite
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Ex17:carbonate
Ex18:bicarbonate
Ex19H2SO4
Ex20:H3PO4
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Exceptions to the Rule of 8
-Resonance - refers to bonding in molecules that be represented by Lewis
structure.
Ex1: SO2
Ex2: SO3
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-Originally it was thought that electrons in the second bond of a double bond “ ” back and forth between on both sides of the sulfur atom. Experiments show that each bond is equal so the electrons must be (another example is O3).
-Exceptions to the rule of 8
Group 13 - Boron is small - it can only accommodate 3 pairs of electrons.
For example, BH3 violates the octet rule:
Groups 15, 16, 17 can find themselves in “expanded” valence shells.
Expanded means that the shells hold more than 8 electrons.
examples:
SF6
PCl5
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XeF4
Some molecules have odd #’s of electrons
For example, NO has 15 total electrons and 11 valence electrons.
Also, NO2 has 17 valence electrons.
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-Related idea - Paramagnetism - atoms that interact weakly with a .
This results from electrons - magnetic fields cancel out from ______
spin of the electrons.
D block elements tend to be paramagnetic. These unpaired electrons also cause the coloring of these compounds.
Hybridization
-Carbon’s full electron configuration is:
Also, carbon’s orbital notation:
1s 2s2p
It looks like it wants to make 2 bonds. However, experiments show it makes 4 identical bonds.
-Hybridization - of atoms in a
chemical reaction.
Carbon will hybridize and have an electron configuration of:
These orbitals have different character than just “s” or just “p”.
They form a new type of orbital called the . Carbon has orbitals.
Carbon’s hybridized orbital notation:
Carbon combines with hydrogen to make methane CH4:
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Methane’s Lewis Structure:
also happens in Be:
1s 2s2p
after hybridization:
example - BeH2 - violates the octet rule - sp hybrid
BeH2 Lewis Structure:
also happens in B:
1s 2s2p
after hybridization:
example - BF3 - violates the octet rule - sp2 hybrid – draw the Lewis structure:
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8.3 and 8.4 Bonding Theories and Polar Bonds and Molecules
Shapes of Molecules - VSEPR Model
-Valence shell electron pair repulsion theory (vesper) -
Note: single bonds, double bonds, triple bonds, and non-bonding electron pairs
all equal one RHEDs (region of high electron density).
Example / RHEDs / Shape / Angle(s) and drawingAB / 1 / Linear / A B
AB2 / 2 / Linear / B A B
AB3 / 3 / Triganal Planar / B
A
B B
AB4 / 4 / Tetrahedron
AB5 / 5 / Triganal Bipyramid
AB6 / 6 / Octahedron
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-Examples: Predict the shapes of the following (show all work):
Ex1:CCl4
Ex2:HBr
Ex3:SO3
Ex4:SO2
Ex5:H2S
Ex6: NH3
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Ex7:ClO4-1
Ex8:PF5
-dipole - when electrons are distributed
-Intermolecular forces(IMF)- forces that molecules together
- happens between compounds
- intermolecular forces - can be
-Intramolecular forces– bonds (ionic, covalent, metallic)
- happens within a molecule or compound
- always
……
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- dipole-dipole (the first type of IMF)
Ex1: predict the IMF that occurs with HCl
Ex2: predict the IMF that occurs with H2S
Hydrogen Bonding(the second type of IMF) H-bonding is a “super-duper” dipole-dipole
-Ex3:predict the IMF that occurs with HF
-Are we having FON? - H-bonding happens any time H is bonded to .
-Why? A large difference in between F, O, or N and H results in one
end of the molecule being very , while the other end is very .
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-Why N and not Cl? They have identical electronegativities!
Well, N is so much than Cl so the negative charge is spread over a smaller
area which exerts more force.
-Effect of H-bonds on physical properties:
H-bonding tends to cause the following in substances:
Boiling Point
Heat of Vaporization
Vapor Pressure
Melting Point
Also,
H-bonds causes water to when it freezes.
H-bonding is also responsible for the of proteins.
Molecular Substances
-We have learned that polar molecules have IMF holding them
together.
-We have learned about H-bonding or “super-duper” IMF.
-These two types of IMF usually result in substances being at room temp.
-A polar molecule has dipole-dipole or hydrogen bonding IMFs.
-Most nonpolar covalent substances are at room temp. as the forces holding them together are not enough to keep the molecules - hence they are .
-O2, H2, N2 - straight substances
-CO2 - have dipoles, but due to its molecular geometry
-now, a third type of IMF:
Van der Waals Forces (London Forces) -
Ex4: - CO2
more electrons = attraction
Thus, bigger atoms have Van der Waals forces
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Network Solids
-They are covalent or covalent .
-Examples:
-They form do not form , molecules. It is one continuous .
-Properties:
-They make good
NIB Summary - see, chemistry makes sense after all!
IMF / Molecule Type / Molecule / Boiling Pt. (Co)London / Noble Gas / He / -269
London / Noble Gas / Ar / -186
London / Nonpolar / H2 / -253
London / Nonpolar / O2 / -183
London / Nonpolar / Cl2 / -34
H-bonding / Polar / HF / 19.5
Dipole-dipole / Polar / ICl / 97
Ionic / NaCl / 1413
Ionic / MgF2 / 2237
Metallic / Cu / 2567
Metallic / Fe / 2750
Bond Energy
- basic idea - what is the strength of chemical bonds?
-bond energy - energy needed to a bond - measured in kJ/mole
-bond strength and stability:
stronger bond - stable -needs more to break the bond
weaker bond - takes energy to break the bond so the chemical is
-chemical changes favor energy states - reactions are favored
NIB Bond Strength
- which is stronger? - single, double, or triple bond?
-which is shortest bond length? s, d, or t?
-which is stronger, short or long bonds?
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