Download Review for Exam 1

Review for Exam 1
CH 1-2 Concepts to know
 Classification of matter: pure substances & mixtures
 Homogeneous vs Heterogeneous
 Distinguish the difference between chemical and physical
properties & changes
 We represent uncertainty with significant figures
 You do not need to memorize Sig Fig rules
 Scientific Notation
 Conversions within the metric system and non metric units
 Temperature conversions
 Density & Specific gravity
 Familiarity with how compounds will be drawn
 Molecular formulas
 Structure of an atom: protons, neutrons, electrons
 Atomic number, isotope mass number, atomic weight
 Navigate the periodic table: properties shared within a group,
trends, metals/metalloids/nonmetals
 Determine valance electrons & draw electron dot
representations
 Ionization Energy & Atomic Size
Conversions & Equations To Memorize
Unit Conversions
Equations
For metric units (m, kg, s, K, mole):
mega (M) 106
kilo (k) 103
centi (c) 10-2
milli (m) 10-3
micro (μ) 10-6
nano (n) 10-9
Pico (p) 10-12
Density = mass / Volume
d = m/V
dH2O = 1 g/mL = 1 g/cm3
Time conversions: dhrms
1 mL = 1
cm3
T(kelvin) = T(°C) + 273.15
Specific Gravity =
density substance / density of water
y x 10x
Coefficient:
A number
between
1 and 10
Exponent:
Any positive
or negative
whole number
Elements & Molecules
A
Z
X
Elements on
the Periodic
Table
X = Element symbol (ie O = oxygen)
A = Isotope Mass Number = # protons + # neutrons
Z = Atomic Number = # protons
6
atomic number
C
element symbol
12.01
atomic weight (amu) = weighted
average of atomic weight of isotopes
Molecular Formula: AxBy
Drawing Molecules:
Methane
CH4
H
H
C
H
H
Ex: CH3O
Properties of Metals, Nonmetals, Metalloids
Metals
• Metallic luster,
malleable, ductile,
hardness variable
• Conduct heat and
electricity
• Solids at room
temperature with
the exception of Hg
• Chemical reactivity
varies greatly: Au,
Pt unreactive while
Na, K very reactive
Nonmetals
• Brittle, dull
• Insulators, nonconductors of
electricity and heat
Metalloids
• Properties
intermediate
between metals and
nonmetals
• Chemical reactivity
varies
• Metallic shine but
brittle
• Exist mostly as
compounds rather
then pure elements
• Semiconductors:
conduct electricity
but not as well as
metals: examples
are silicon and
germanium
• Many are gases,
some are solids at
room temp, only Br2
is a liquid.
Valence Electrons
Example: Determine the valence electrons
of Selenium (Se):
1. Find Se on the periodic table
2. Focus on just the column Se is in
3. Column number indicates number of eElectron Dot Symbols:
Represent the valence electrons by
drawing them around the element
symbol for Selenium.
Se
Periodic Trends
Size
INCREASING
Ionization Energy
INCREASING
CHAPTER 3-4: Concepts to Know
 The difference between ionic and covalent bonds
 Define cations and anions
 Predict cation/anion charge using the octet rule or group number
 Familiar with metals with multiple potential charges (do not
need to memorize)
 Determine ionic compound formulas from the name of a compound
or from the elements that compose it.
 Criss-cross rule
 Naming ionic compounds and covalent molecules
 Familiar with polyatomic ions (do not need to memorize but must
be able to recognize)
 Draw lewis dot structures
 Determine molecular geometry
 Identify polar bonds
 Determine dipole moment of molecules
Need to Memorize
Ionic vs Covalent Bonding
Ionic Bonds result from electrostatic attraction
between a cation and anion: metal-nonmetal (with the
exception of NH4+ and H3O+ cations).
Covalent bonds result from the sharing of electrons
between two atoms: nonmetal-nonmetal.
Li
F
Ionic Bonds
Covalent Bonds
Naming
HOW TO Name an Ionic Compound
Step [1]
Determine the charge on
the cation.
Step [2]
Name the cation and the
anion
 If the cation could be
multiple charges indicate
the charge with roman
numerals or with a –ous /
-ic suffix.
Step [3]
Write the name of the
cation first then the name
of the anion
HOW TO Name a Covalent Molecule
Step [1]
Step [2]
Name the first nonmetal by
its element name and the
second using the suffix
“-ide.”
Add prefixes to show the
number of atoms of each
element.
Predicting Cations & Anions
the cation charge = the group number
the anion charge = 8 – group number
Octet Rule
The octet rule: a main group element is especially stable
when it possesses an octet of e− in its outer shell.
octet = 8 valence e−
Exceptions (need to memorize):
F
F
B F
HO
P OH
OH
only 6 e− on B
O
O
10 e− on P
HO
S OH
O
12 e− on S
Ionic Compound Formulas
HOW TO Write a Formula for an Ionic Compound
Step [1]
Identify which element is the cation
and which is the anion.
Step [2]
Determine how many of each ion type is
needed for an overall charge of zero.
 When the cation and anion have different
charges, use the ion charges to determine
the number of ions of each needed.
Step [3]
To write the formula, place the cation
first and then the anion, and omit charges.
Lewis Dot Structures
NH3
Step [1] Arrange the atoms next to each
other that you think are bonded
together. Place H and halogens on
the periphery, since they can only
form one bond.
Step [2] Count the valence electrons. The
−
sum gives the total number of e
that must be used in the Lewis
structure. For each atom the
number of bonds = 8 – valence
electrons.
N
H
H
H
Nitrogen has 5 valence electrons, so
it will have 8 – 5 = 3 bonds.
Hydrogen will have 2-1 = 1 bond.
There are 8 total valance electrons
H
N
H
Step [3] Arrange the electrons around the
atoms. Place one bond (two e−)
between every two atoms. Use all
remaining electrons to fill octets
with lone pairs, beginning with
atoms on the periphery.
H
1 lone pair: 2
3 bonds:
6
Total e8
= total valence e-
Resonance Structures
Resonance structures exist when there are multiple lewis dot structures
with different electron arrangements with the same connectivity between
atoms. Resonance structures help us understand delocalization
(spreading) of charge within a molecule that stabilizes the anion or cation.
Other Examples: CO32- and O3
Molecular Shape
Periodic Trend: Electronegativity
Electronegativity
INCREASING
Polarity
1. Assess the relative electronegativity of atoms bonded
together, if there is a difference it is a polar bond.
2. Indicate polar bonds with δ+ / δ - or
3. If polarity of bonds does not cancel draw the overall
dipole moment of the molecule using
Electron density is disproportionately
distributed over the molecule. Above red
indicates partial negative charge, or greater
electron density, and blue indicates partial
positive charge.
Effectively oxygen is hogging the electrons
CH 5 Concepts to know
 Define a chemical reaction
 Correctly write a chemical reaction
 Balance reactions by inspection
 Calculate molecular mass for any compound or molecule
 Apply mole ratios within molecules and between molecules.
 Solve stoichiometry problems
 Convert between mass and moles
 Identify limiting reagent
 Calculate percent yield
 Identify reduction and oxidation equations and pick out the
compound being reduced or oxidized
Smith. General Organic & Biolocial Chemistry 2nd
20
Need to Memorize
6.02 x 1023 is Avogadro’s number.
æ actual yield ö
÷÷ ´ 100%
percentage yield = çç
è theoretical yield ø
Oxidation is the loss of electrons from an atom.
Reducing agents are oxidized
Reduction is the gain of electrons by an atom.
Oxidizing agents are reduced.
Smith. General Organic & Biolocial Chemistry 2nd
21
Writing and Balancing Equations
aA (physical state) + bB (state)  cC (state) + dD (state)
HOW TO Balance a Chemical Equation
Step [1] Write the equation with the correct formulas.
•The subscripts in a formula can never be changed
to balance an equation, because changing a
subscript changes the identity of a compound.
Step [2]
Balance the equation with coefficients one
element at a time.
Step [3]
Check to make sure that the smallest set
of whole numbers is used.
Smith. General Organic & Biolocial Chemistry 2nd Ed.
22
Solve Stoichiometry Problems
aA +
mass A
FW/MM
bB
mass B
FW/MM
moles A

a:b
moles B
cC
+
mass C
mass D
FW/MM
b:c
moles C
dD
FW/MM
c:d
moles D
a:c
a:d
mol
g
mol
g
Note:
FW/MM means
Formula wt. or
Molar mass
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Limiting Reactant


Compare the actual amount of each reactant
to the amount required in the balanced
equation to determine how many times the
“reaction can be run”
Use the amount of the limiting reactant to
calculate how much product can be produced
æ actual yield ö
÷÷ ´ 100%
percentage yield = çç
è theoretical yield ø
Redox Half Reactions
Cu2+ gains 2 e−
Zn2+ + Cu
Zn + Cu2+
Zn loses 2 e–
Each of these processes can be written as an
individual half reaction:
Zn2+ + 2 e−
loss of e−
Oxidation half reaction:
Zn
Reduction half reaction:
Cu2+ + 2e−
gain of e−
Smith. General Organic & Biolocial Chemistry 2nd Ed.
Cu
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