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Study Guide for Exam #2
MSU LBS 171 – Profs. Robert LaDuca and Ryan Sweeder
-Molecule: chem. combo. of 2 or more atoms bonded together
-Element: matter that is made up of only one type of atom
-Compound: matter made up of 2 or more kinds of an atom
-Molecular formula: describes the composition of a specific molecular structure
Example: glucose is C6H1206
-Ionic compounds: compounds made of ions
-Ions: atom that has a charge (gained or lost e-)
-Cation: positively charged ion; anion: negatively charged ion
-Ionic compounds—naming (charge on cation must = charge on anion, since compounds are
neutral):
• CaO
calcium oxide
• MgS
magnesium sulfide
• Na2O
sodium oxide
• Cu2O
copper (I) oxide
• CuO
copper (II) oxide
-Compounds which are hydrated are those that crystallize with the addition of H2O
- 1 H2O monohydrate
2 H2O dihydrate
3 H2O trihydrate…….
-Periodic Trends:
• Atomic radius decreases L to R and increases down the periodic table.
o As you go across, the positive charge on the nucleus increases, pulling the
electrons closer to it.
o As you go down, you are increasing the number of energy levels on the atom,
increasing its size.
• Ionization energy (energy required to remove e- from atom in gas phase)…first ionization
energy increases L to R, and decreases down the periodic table.
o It gets harder to pull electrons from the atom as you move across the table, since
the positive nucleus has a stronger hold over the electrons in the atom.
o As you go down, the atom grows, since more energy levels are added, and it
becomes increasingly easier to pull an e- away, since it is not being held in by the
nucleus as tightly.
• Electron affinity increases L to R, and decreases down the periodic table.
o Effective nuclear charge increases L to R, and this makes it harder to ionize the
atom. However, this makes the atom more attractive to additional electrons.
Ionization energy and electron affinity are (for the most part) proportionately
related.
Electronegativity – The ability of an atom, when in a compound, to attract e- density to itself
•
Linus Pauling – 1930s – “The Nature of the Chemical Bond”
•
Traveling up and to the right on the periodic table increases the electronegativity (χ) with
Cs at the lowest value (0.8) and F at the highest (4.0).
Figure 1. Increasing Electronegativity in the Periodic Table
•
The difference in χ determines the type of chemical bond
For compounds containing NON-METALS ONLY
o χ ≤ 0.4 covalent bond (complete sharing)
o 0.5 ≤ χ polar covalent bond (unequal sharing)
For compounds containing METALS and NON-METALS
o χ ≥ 1.5
ionic bond (complete transfer of valence e-)
o 0.5 ≤ χ ≤ 1.5
Lewis Dot Structure – 1918
polar covalent bond (unequal sharing)
•
First level view of chemical bonding – each valence e- on an atom is a dot
•
Covalent or polar-covalent bond indicated by a dash (shared pair or e-)
•
Lewis structures work well for main group elements, but not transition metals
•
Drawing Lewis Structures
o Count valence electrons (VE) on individual atoms
o Put least electronegative atom (but not H) in center
o Connect all other atoms with single bonds
o Place lone pairs of e- on outer atoms to satisfy octet rule until no e- remain
o If remaining e-, place them on the central atom (only 8+ if it has a d-subshell)
o Calculate Formal Charge = (original # VE) – (# bonds) – 2 (# lone pairs)
o If central atom has + charge, form double bond. Recalculate. REMEMBER
ƒ
Atoms would prefer incomplete or expanded octet to a charge
ƒ
Cannot form expanded octet on first or second row elements
ƒ
If charge cannot be minimized, place negative FC on the more
electronegative atom
•
See lecture notes for examples of the formation of Lewis Dot Structures
Resonance
•
Localization of bonding e- throughout the molecule
•
Need more than one Lewis Structure to accurately represent the bonding
•
Combination of equivalent Lewis Structures
•
N ≡ N − O OR N = N = O
•
NOTE: bond orders are equal, actual orders found by
number of bonds
number of links
o Example: Ozone (O3) → O = O − O → three bonds, two links
o Bond lengths (also called bond order) =
3
= 1.5
2
Naming Molecular Compounds (non-metals ONLY)
•
Name of less electronegative atom stays the same
•
Name of more electronegative atom ends in –ide
•
Indicate the number of each atom by a Greek prefix
o CO2 – carbon dioxide (omit mono- from first element)
o N2O – dinitrogen monoxide
o P4O10 – tetraphosphorus decoxide
•
Acids
o Neutral molecular compounds that can liberate H+
o If parent anion ends in –ide, becomes hydro –ic acid
ƒ
Cl- – hydrochloric acid
ƒ
CN- (cyanide) – hydrocyanic acid
o If parent anion ends in –ate, becomes –ic acid
o If parent anion ends in –ite, becomes –ous acid
ƒ
HClO4 – perchloric acid
ƒ
HNO2 – nitrous acid
•
Acid – a compound that donates H+ in solution
•
Base – a compound that accepts H+ in solution
•
Amphoteric – water (H2O) can be either acidic or basic
Polyatomic Ions (LEARN THE SHEET! LEARN THE SHEET!)
Cation
Ion Name
Anion
Ion Name
NH4+
ammonium
Anion
CNCH3CO2CO3-2
HCO3-2
Ion Name
cyanide
acetate
carbonate
hydrogen carbonate
NO2NO3PO4-3
HPO4-2
H2PO4-
nitrite
nitrate
phosphate
hydrogen phosphate
dihydrogen phosphate
OHSO3-2
SO4-2
HSO4-
hydroxide
sulfite
sulfate
hydrogen sulfate
ClOClO2ClO3ClO4-
hypochlorite
chlorite
chlorate
perchlorate
CrO4-2
Cr2O7-2
MnO4-
chromate
dichromate
permanganate
Molecular Geometry – creates the 3D shapes molecules actually form
VSEPR – Valence Shell Electron Pair Repulsion
•
Number of e- groups determines basis shape of molecular compound
•
If lone pair or electron dot is present, molecular shape is different
-
basis shape
# e groups
2
3
4
5
6
free pair
shape
two free
pairs
linear
trigonal planar
bent <120
tetrahedral
trigonal pyramidal bent <109.5
trigonal bipyramidal
"see-saw"
"T-shaped"
octahedral
square pyramidal
square
three free
pairs
linear
Molecular Polarity – depends on two conditions
•
One or more polar bonds (χ ≥ 0.5)
•
Must be asymmetric
o Draw 2D structure / 3D structure
o Place polarity arrows (positive at atom with lower χ) on polar bonds
o Tug on arrows. If central atom moves, molecule is polar
Deeper View of Chemical Bonding
•
Linus Pauling – 1930s – bond forms due to valence orbital overlap
•
If covalent bond is symmetrical, it is a σ bond
o Single bonds are always σ bonds
•
Pauling stated that atomic orbitals can mix (or hybridize) to form new atomic orbitals that
overlap to create the correct geometries
# e- groups
2
3
4
5
6
•
hybrid
name
sp
sp2
sp3
sp3d
sp3d2
leftover
orbitals
two p's
one p
nothing
four d's
three d's
A double bond is formed by a hybrid bond and a p-orbital
o p-orbital bond is asymmetric
o Asymmetric bond called a π bond
References
Lecture notes from Prof. LaDuca
Figure 1. Increasing Electronegativity in the Periodic Table
<http://www.chem.ubc.ca/courseware/121/tutorials/exp7A/images/fig4.jpg >
Figure 2. 2-methyl propane
<http://library.tedankara.k12.tr/chemistry/vol3/Geometric%20isomers/z229.gif>
Figure 3. Isomers of Butene: cis- and trans<http://img.sparknotes.com/figures/8/8a5d5291809971d9fa98d928ddf79524/fig2_14.gif>
Study Guide Questions for Exam #2
1. Name each of the following ionic compounds
a) K2S
b) CoSO4
c) (NH4)3PO4
d) Ca(ClO)2
2. Draw the Lewis Dot Structure for each of the following molecules or ions
a) BrF3
b) BrF5
c) ClO3-
d) NH3
e) SiF6-2
f) O3
g) N2O
h) SO2
3. Give the basis shape and actual geometry of the molecules and ions from Problem 2.
4. Draw all of the resonance structures for NO35. Which molecules are polar?
a) H2O
b) CO2
c) NH3
e) CH3Cl
f) CH4
g) PCl3
d) NH2Cl
6. Compare the bond lengths of ClO2-, ClO3-, and ClO4-. Which ion has the longest bond
length? Which has the shortest bonds?
References
“Chemistry and Chemical Reactivity.” Kotz. Study Questions.