# Chemistry Midterm Study Guide

```Chemistry 530 Midterm Review
Exam on Monday Dec 16 at 8:30
Section I = 75 pts
 Consists of 75 multiple choice questions worth 1 pt. each
 Questions cover all material covered during the second semester
 Questions will be similar to multiple choice problems in the book and review guides on-line
 Practicing the review guides is the best approach; answers will be posted on-line
Section II = 95 pts
 Consists of 19 short answer problems totaling 95 points
 Questions cover the whole semester
 Examples of types of questions
o Metric Conversions
o Volume Calculations
o Percent Composition
o Atomic Mass
o Electron configuration
o Nomenclature
o Balancing Equations
o Predicting the products of a reactions
o Identifying types of reactions
o Mole Conversions
o Molar Mass of a compound
o Molar Mass of a gas at STP
o Density of a gas at STP
o Empirical &amp; Molecular Formula
Study approach
 First, read through your notes. Once thoroughly and once more quickly.
 Go through the review packet and make sure you are comfortable with all of the terms and how to do
problems
 Do the review packets for help with multiple choice
 Do questions on old tests for practicing written problems
 Make sure you have practiced AT LEAST one of the types of example problems
Chapter 1 – Introduction to Chemistry
 Chemistry: the branch of science that studies the composition and properties of matter.
 Scientific Method:
o Observation
o Hypothesis
o Controlled Experiment
 Theory versus Natural Law
Prerequisite Science Skills
 Rules for Significant Figures:
1) All non-zero digits are significant.
2) Zeros are significant only when they are between two non-zero digits
or if they occur after a decimal point and after a non-zero digit.
3) In multiplication or division, you cannot end up with more significant figures than you started with.
4) In addition or subtraction, you cannot have more digits to the right of the decimal than you started with.
 Exponential numbers
 Scientific Notation = D.DD x 10n where 1.00 &lt; D.DD &lt; 9.99
Prerequisite Science Skills Practice Problems
1. Circle which of these measurements is NOT uncertain?
67 g
4.5 mL
457 K
22 people
6.7 x 10
11
moles
2. Convert the following into standard notation:
-5
4.89 x 10 nm=
2
7.9 x 10 mm =
3. Convert the following into scientific notation:
0.00056 mg =
78,341 kg =
4. Give the number of significant figures in each of these measurements:
46 crucibles =
607 g =
0.00900 mL =
4.90 x 10 -5 s =
5. Subtract 45.6 from 67.89 and round off the answer. Answer = =
6. Multiply 240 by 2.10 and round off the answer. Answer =
Chapter 2 – The Metric System
 Units:
Unit
Length
Mass
Volume
Time
Temperature
Amount
Energy

Metric Unit Conversions
Prefix
Symbol
kilok
centic
millim
microμ
nano-
n
Metric
m
g
L
s
C
mol
cal
SI
m
kg
L
s
K
mol
J
Conversion Factor
1 km = 1000 m
1 cm = 0.01 m
1 mm = 0.001 m
1 μm = 0.000001 m
1 nm = 0.000000001 m
Conversions
1000 g = 1 kg
K = C = 273.15
cal = 4.184 J





Percent = % = Part / Total x 100
Volume = length x width x height
Volume by displacement
Density = Mass / Volume = M / V
Temperature
o F = 1.8 C + 32
o C = ( F – 32 ) / 1.8
o K = C = 273
Chapter 2 Practice Problems
7. Make the following unit conversions:
5 g into kg
40 mL into L
319 nm into cm
316 K into Celsius
8. A student is trying to determine the identity of an unknown metal.
He finds that its mass is 20.0 g and its volume is 1.0 mL. The density of the metal is:
Known Densities of Metals
Gold
Mercury
Silver
Copper
Aluminum
3
19.3 g/cm
3
13.6 g/cm
3
11.4 g/cm
3
10.5 g/cm
3
8.9 g/cm
3
2.7 g/cm
The metal that the student has is:
9. An alloy contains 12 parts platinum and 4 parts gold. Find the percent of platinum and gold in this alloy.
Chapter 3 – Matter and Energy
 Element: a pure substance that cannot be broken down by a chemical reaction.
 Compound: a pure substance that can be broken down into two or more substances by a chemical reaction.
 Mixtures:
o Homogeneous (the same properties throughout)
o Heterogeneous (different properties throughout).
 Physical States: Solid (s), Liquid (l), Gas (g), Aqueous (aq).
 Chemical vs. Physical properties and changes.
 Metals:
 Non-Metals:
 Semi-metals or Metalloids:
 Law of Conservation of Mass: Matter can neither be created nor destroyed.
 Law of Conservation of Energy: Energy can neither be created nor destroyed.
 Potential vs. Kinetic Energy
Chapter 3 Practice Problems
10. Give examples of the following:
Element =
Compound =
Homogeneous mixture =
Heterogeneous mixture =
11.
How many atoms are in a glucose molecule, Al (C2H3O2 ) 3 ?
# Al = __________
# C = ___________
#H = ___________
# O = ___________
12.
Name the following phase changes:
Solid  Liquid =
Liquid  Gas =
Gas Solid =
Solid  Gas =
Gas  Liquid =
Liquid  Solid =
13. Consider a glass of water. What are three of its physical properties? What are three of its chemical properties?
14. What are three properties of a gas? Three properties of a liquid? Three properties of a solid?
Chapter 4 – Models of the Atom
 Atom
+
o proton (p , 1 amu)
o
o neutron (n , 1 amu)
o electron (e , 0 amu)
 Basic idea of theories of the atom
 Bohr models of the atom
 Atomic notation
 Atomic number
 Mass number
 Isotopes
o Atomic mass calculations based on isotope abundance
 Electromagnetic spectrum
 Quantum mechanical model of atom
o orbitals (n = 1, 2, 3, 4 …)
o sub-orbitals (s, p, d, f)
 know how many electrons each holds
 know shapes of s and p orbitals
 electron configuration for atoms and ions
o know order of filling
 Heisenberg’s Uncertainty Principle
Chapter #4 -Practice Problems:
15. Complete the following table:
Symbol
Atomic Number
Mass Number
Protons
Al
Neutrons
14
33
75
19
9
16. Write the atomic symbol of an element which has 6 protons, 6 electrons and 7 neutrons.
17. Copper has two natural isotopes: Cu-63 (62.930 amu) and Cu-65 (64.928 amu). Calculate
the atomic mass of copper given the abundance of Cu-63 is 69.09%.
18. Give the maximum number of electrons for the following orbitals:
 s = ____________; Shape of s orbital = _______________
 p = ____________; Shape of p orbital = _______________
 d = ____________
 f = ____________
19. Give the maximum number of electrons for the following energy levels:
 1st= ____________
 2nd= ____________
 3rd= ____________
 4th= ____________
20. Write the electron configurations for:
 Br =
3 N = __________________________________
 Pt = __________________________________
+
 Cu = __________________________________
Chapter 5 – The Periodic Table
 Periodic table
o groups (alkali, alkaline earth metals, halogens, noble gases)
Electrons



o periods
o representative elements v. transition elements
Periodicity (periodic trends) - know how each increases/decreases, and what each is
o metallic character (same as atomic radius)
o ionization energy
o electronegativity
Valence electrons
Lewis dot structures
Chapter #5 - Practice Problems:
21. How many electrons are in the valence shells of these atoms and ions? Write the electron configuration using
Noble Gas configuration.
 Cl
= ____ valance electrons,
o

Mg
= ____ valance electrons,
o

Cr
Se
2-
Electron configuration =
= ____ valance electrons,
o

Electron configuration =
= ____ valance electrons,
o

Electron configuration =
Zn 2+
o
Electron configuration =
= ____ valance electrons,
Electron configuration =
22. Draw Lewis electron-dot structures of the atoms and ions given above.
Cl
Mg
Se 2-
Zn 2+
Cr
23. Put these atoms and ions in order from lowest to highest atomic radius.
22+
Ca
Ag
Se
Ba
O 
24. Circle the atom or ion with the highest ionization energy in each pair.
2+
–
Zn or Zn
Si or S
F or F
25. Put these atoms and ions in order from highest to lowest electronegativity.
+
3I
Fr
Cu
N
P 
26. Name the group to which each of these atoms belongs.
 F =

Na =

Mg =

Fe =

Pb =
Chapter 6: Language of Chemistry
 Ions:
o Cation = a positively charged ion.
o Anion = a negatively charged ion.
 Ions are formed by gaining or losing electrons.
 Know monoatomic ion’s name, formula and charge (Use location on Periodic Table )
 Memorize polyatomic ion’s name, formula and charge
 Forming Ionic Compounds - Cross over the charge, not the + or Example: Magnesium is a 2+ cation = Mg 2+. Nitrate is a 1- anion = NO3 -.
The chemical formula is Mg(NO3)2.
Nomenclature:
I.
Binary Ionic Compounds (metal + non-metal)
Name the metal cation first.
Name the non-metal anion second, change the suffix to –ide.
Example: CaCl2 is calcium chloride.
II.
Binary Ionic Compounds with a metal of variable charge (metal + non-metal)
Name the metal cation first.
Give the charge on the metal using Roman Numerals (I), (II), (III), (IV), (V).
Name the non-metal anion second, change the suffix to –ide.
Example: Fe2O3 is iron (III) oxide.
III.
Ternary Ionic Compounds (containing polyatomic ions.)
Name the cation first, using Roman Numerals if necessary.
Name the anion second, do not change the suffix if it is a polyatomic ion.
Examples: Ni3(PO4)2 is nickel (II) phosphate; NH4I is ammonium iodide.
IV.
Binary Molecular Compounds (non-metal + non-metal)
Name the most metallic element first (the one furthest to the left table).
Name the least metallic element second, change the suffix to –ide.
Use prefixes to indicate the number and type of each atom
(mono-, di-, tri-, tetra- penta-, hexa-, hepta-, octa-, nona-).
Example: SF6 is sulfur hexafluoride (or monosulfur hexafluoride).
V.
Acids- Can often be recognized by hydrogen beginning a formula.
Binary acid – To name
1. changing the word hydrogen to hydro2. change the ending: -ide to –ic acid
Ternary oxyacids - To name
1. drop the word hydrogen,
2. change the ending:
–ate changes to –ic acid.
–ite, changes to –ous acid.
You should be able to name these five common acids:
HCl
HNO3
H2SO4
H3PO4
HC2H3O2
VI.
Bases - Can often be recognized by hydrogen ending a formula. Bases are named just like other ternary
ionic compounds (see III); the exception is NH3, which is ammonia.
Chapter 6 Practice Problems
27. Name the following compounds:
CCl4
MgCr2O7
NiSO4
NaOH
H2SO4
28. Write formula for the following compounds:
Zinc sulfite
Silicon dioxide
Iron (II) phosphide
Potassium dichromate
Ammonium nitrate
Chapter 7: Chemical Reactions
 You should be able to predict and balance equations:
 Make sure there are the same number of atoms of each type on each side by using coefficients.
Types of Reactions:
A.
Combination Reactions
Two or more reactants  One product
2Mg(s) + N2(g)  2MgN(s)
B.
Decomposition Reactions
One reactant  Two or more products
2NaHCO3(s)  Na2CO3(s) + H2O(l) + CO2(g)
Na2CO3(s)  Na2O(s) + CO2(g)
2NaO(s)  2Na(s) + O2(g)
C.
Single-replacement Reactions
One element replaces another element in a compound.
2Na(s) + ZnCl2(aq)  2NaCl(aq) + Zn(s)
Predictable: Consult the activity series More active metals can replace less active ones.
D.
Double-replacement Reactions
Two ionic compounds exchange partner ions, forming two new compounds.
2AgNO3(aq) + CuCl2(aq)  2AgCl(s) + Cu(NO3)2(aq)
Predictable: Swap the ions to make two neutral compounds.
Consult the solubility rules to determine if the products are solid (s) or aqueous (aq).
E.
Neutralization Reaction
Acid and Base produce water and a salt (ionic compound)
Chapter 7 Practice Problems
29. For the following reactions:
1. Predict the products of the following reactions
2. Indicate physical states for double- replacement reactions (consult the Solubility Rules).
3. Balance the chemical equations.
4. Indicate the type of chemical reaction for each
a.
O2(g) 
Na(s) +
Type:
b.
H2O(l) 
Ag(s) +
Type:
c.
MgBr2(aq) +
K2SO4(aq) 
Type:
d.
Al(s) +
Pb(NO3)2(aq) 
Type:
e.
MgCO3(s) 
Type:
f.
H2O(aq) + KMnO4(aq) 
Type:
Chapter 8 – The Mole Concept






23
Avogadro’s Number = 6.02 x 10
particles (atoms, ions, formula units or molecules)
Mole
Molar Mass = MM = grams
(Get this from the Periodic Table)
Molar Volume of a gas at STP = 22.4 Liters
Density of a gas @ STP = Molar Mass / Molar Volume = MM / 22.4 L = g/L
Percent Composition:
% Element = Molar Mass of Element / Molar mass of Compound x 100


Empirical Formula: Gives the lowest whole number ratio of atoms in a compound.
Molecular Formula: Either the same as the empirical formula, or a simple whole-number multiple of it
Chapter 8 Practice Problems
30. How many grams are in 1.204 x 10
24
molecules of calcium hydroxide, Ca(OH)2 (MW = 74g/mol)?
31. How many moles are in 11.2 L of methane (CH4) gas?
32. Calculate the percent composition of heptane, C7H16.
33. Calculate the molar mass for Freon-12 if its density is 5.40 g/L at STP.
33. Nicotine has a molar mass of 160 g / mole. If the percent composition is 74.0 % C, 8.7 % H, and 17.3 % N, what
is the molecular formula of nicotine? (HINT: Determine empirical formula first)
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