Chemical vs. Physical Properties
Physical vs. Chemical
EC Version
A physical property is a property that can be observed and
determined without destroying the object. For example; color, shape,
mass, length and odor are all physical properties.
Changes
A chemical property indicated how a substance reacts with something
else. The original substance is fundamentally changed in observing a
chemical property.
In a physical change, the original substance still exists and has
only changed in form. In a chemical change, bonds have been
broken or formed and a new substance is produced. Chemical
changes are also accompanied by an energy exchange.
Classify the following properties as either chemical or physical
properties.
Classify the following as chemical or physical changes
Property
Blue color
Physical
Chemical
1. Salt dissolves in water
2. Hydrochloric acid reacts with potassium hydroxide to produce
salt and water
Density
3. A pellet of sodium is sliced in two
Flammability
4. Water is heated and changed to steam
Solubility
5. Potassium chlorate is heated and decomposed to potassium
Reacts with acid to form H2
chloride and oxygen
Supports combustion
6. Iron rusts.
Sour taste
7. When placed in water a sodium pellet catches fire as H2 gas is
liberated and sodium hydroxide forms.
8. Evaporation
Melting point
Reacts with water to form gas
Reacts with base to form water
Hardness
Boiling point
Can neutralize a base
Luster
Review Part 1: Chapters 1 – 3, 13
9. Ice melting
10. Milk sours
11. Sugar dissolves in water
12. Wood rotting
13. Grass growing via photosynthesis
14. A tire is inflated with air
1
EC Version
15. Food is digested in the stomach
16. Water is absorbed by a paper towel
Mixtures or Pure Substances
All matter can be classified as either a mixture or a pure substance. If
it is a mixture it can further be classified as homogenous or
heterogenous. If it is a pure substance it can be either a compound or
an element.
Matter
Pure Substance
Mixture
(separated by physical means)
(made in any proportion)
Elements
On PT
Symbols
One type
Of atom
Compounds
Homogenous
Heterogenous
2 or more atoms
bonded together
Formulas H 2O
well blended
solution/alloys
identifiable parts
easily separated
ice cream
Air
Iron
Brass
Solids/Liquids/Gases
The Kinetic Molecular Theory helps to describe the behavior
and characteristics of solids, liquids and gases. Basically it
states that all the properties unique to a particular state of
matter are due to:
1. The spacing of the particles
2. The motion of the particles
3. The energy of the particles
In the table below first identify the state of matter the property is
referring to. It may refer to more than one state. Then identify the
part of the kinetic molecular theory responsible for this property.
Property
Classify the following as pure substances or mixtures.
Type of
Matter
Pure Substance
Mixture
Element Compound Heterogeneous Homogenous
Chlorine
Water
Soil
Sugar water
Oxygen
Carbon
dioxide
Rocky road
Review Part 1: Chapters 1 – 3, 13
State of Matter
Due to spacing, motion
and/or energy
Compressible
Diffuse
High Density
Organized
Take shape of
container
Expansion
Definite volume
Very low density
2
EC Version
Fluid
Have mass
Undergo boiling
Review Part 1: Chapters 1 – 3, 13
3
Atomic Structure
An atom is made up of protons and neutrons (both found in the
nucleus) and electrons (found in the orbitals surrounding the
nucleus). The atomic number of an element is equal to the number of
protons. If the atom is neutral then the atomic number also equals the
number of electrons. A charge written in the upper right corner
indicates that the number of electrons has been altered. The mass
number (different than the average atomic mass) is the sum of the
protons and neutrons.
Nitrogen- 15 (+3) cation
Mass Number  15
N +3 ion charge
Atomic # 
7
7 protons
8 neutrons (15-7)
4 electrons (normally 7 but +3 means loses 3 electrons)
Complete the following table
Atomic
Mass
#
#
#
Element/Ion
Number Number Protons Neutrons Electrons
1
H
1 +1
H
12
C
7
Li +
35
Cl -1
39
K
24
Mg +2
74
As -3
108
Ag
108
Ag +1
33 -2
S
238
U
Review Part 2: Chapters 4 -7, 25
Isotopes and Average Atomic
EC Version
Mass
Elements come in a variety of isotopes, meaning they are made up of
atoms with the same atomic number but different mass numbers due
to varying numbers of neutrons. A weighted average can be taken of
the mass numbers of each isotope. The average is called the average
atomic mass and appears in the box on the periodic table.
Cesium has three naturally occurring isotopes, cesium-133, cesium132 and cesium- 134. The percent abundance of each is 75 %, 20 %
and 5 % respectively. Determine the average atomic mass of cesium.
133 x 0.75 = 99.75
132 x 0.20 = 26.4
134 x 0.05 = 6.7
132.85 amu
Determine the average atomic mass of the following elements.
1. 80 % 127 I, 17 % 126I, 3 % 128I
______________
2. 15 % 55Fe, 85 % 56Fe
_______________
3. 99% 1H, 0.8 % 2H, 0.2 % 3H
_______________
4. 95 % 14N, 3 % 15N, 2% 16N
________________
4
EC Version
5. 98 % 12 C, 2 % 14 C
_______________
3. Al
Electron Configuration
Electrons are distributed in orbitals with principal quantum numbers
of 1 – 7, sublevels, (s,p,d,f), orientations (s has 1, p has 3, d has 5 and
f and 7) and spin (two electrons are allowed in each orbital). Use
your periodic table or the diagonal rule to place electrons into orbitals
Draw the orbital notation for sodium:
Sodium has 11 electrons:
Orbital notation: __ __
1s
__ __ __
2s
2p
__
3s
Electron configuration: 1s22s22p63s1
* remember Hund’s Law
4. O
Valence Electrons
Valence electrons are defined as the electrons in the largest energy
level. They are always the electrons in the s and p sublevels, so the
number of valence electrons can never exceed 8. You can determine
the number of valence electrons from the electron configuration or by
looking at the element’s placement on the periodic table.
Carbon: 1s22s22p2
Largest energy level is 2 where a total of 4 electrons
reside (look at superscripts). Therefore the valence is 4
Determine the number of valence each of the following has.
Element
Draw both the orbital notation and write the electron configurations
for:
# Valence e-
Element
Fluorine
Lithium
Phosphorus
Zinc
Calcium
Carbon
Nitrogen
Iodine
Iron
Oxygen
# Valence e-
1. Cl
2. N
Review Part 2: Chapters 4 -7, 25
5
EC Version
Argon
Barium
Potassium
Aluminum
Helium
Hydrogen
Magnesium
Xenon
Sulfur
Copper
Review Part 2: Chapters 4 -7, 25
6
EC Version
Nuclear Refresher
problem asks for the number of grams that remain, multiply the
starting amount by the fraction remaining.
Perform the following half-life calculations:
1. 10-g of a radioactive element decays by beta emission. What is
the half life of the element if 2.5 g of the undecayed element
remains after 2 days?
Nuclear decay occurs when a nucleus is unstable or radioactive. It
becomes unstable when the neutron to proton ratio is too high or
too low. In an attempt to become stable, a particle and/or energy
are released, frequently resulting in the identity of the element
changing. Some examples of decay are alpha, beta, gamma, or a
neutron. Radioactive isotopes can also undergo fusion (joining
together of nuclei) or fission (a single nucleus separating)
What are the symbols for alpha particles:
and beta particles:
, neutrons:
,
?
2. The half-life of protactinium-234 is 6.75 hours. What fraction of
a given sample will remain after 27 hours?
For the following decay reactions, fill in the missing part.
1.
2.
214
Bi 
239
4
Np 
He + ________
239
Pu + _____
Es + 4He  1n + ________
3.
253
4.
238
5.
75
6.
6
7.
252
U +
12
C  _____
Se 
0
e + _____
Li +
63
Cf +
-1
3. The half-life of radium-226 is 1600 years. How many grams of a
0.25g sample will remain after 4800 years?
Ni  _____
10
B  _____
Half-Lives
The half-life for an unstable element is the amount of time required
for half of the element to transmutate into something else. Thus, after
1 half-life ½ of the original remains undecayed, after 2 half-lives:
½ x ½ = ¼ of the original remains, 3 half-lives: 1/8 remains and so
on.
To find the # of half-lives divide the total time by the time for one
half life. The fraction remaining will be (½ )^# half-lives. If the
Review Part 2: Chapters 4 -7, 25
4.
Sodium-24 has a half-life of 15 hours. How much sodium-24
will remain in an 18.0g sample after 60 hours?
5. The half-life of tritium (3H) is 12.3 years. If 48.0 mg of tritium is
released from a nuclear power plant during the course of a
mishap, what mass of the nuclide will remain after 49.2 years?
7
EC Version
Periodic Table
The Periodic Table is organized in groups and periods. The elements
in a group are placed together because they have many similarities.
Also from the periodic table, one can compare various trends like
atomic radius, electronegativity, reactivity and ionization energy
1. In what group are the most active metals located?
2. In what group are the most active nonmetals located?
3. As you go from left to right across a period, the atomic size
(increases / decreases)
4. As you travel down a group, the atomic size (decreases /
Bonding
A bond forms when an atom tries to become more stable. It wants to satisfy
the octet rule – have eight valence electrons.
There are a variety of ways an atom can bond. They are:
1. Metallic – between two metals resulting in a sea of mobile electrons
2. Ionic – between a metal and a nonmetal or polyatomic ion. Results
from a transfer of electrons from metal to nonmetal. Form positive
and negative ions.
3. Covalent – between two nonmetals due to the sharing of electrons.
There are two types of covalent bonds:
A. Polar – an uneven sharing – usually two different nonmetals
create dipole
B. Nonpolar – equal sharing of electrons – usually between two
identical nonmetals.
increases)? why
5. In what group are the most electronegative elements found?
6. In what group are the least electronegative elements found?
7. Element of group 2 are called
8. As you go from left to right across the periodic table, the
elements go from (metals/ nonmetals) to (metals/ nonmetals).
9. Group 18 elements are called
Identify the following as being metallic, ionic, polar covalent, or
nonpolar covalent bonds.
1. Cu – Cu
2. Na – O
3. Li – Cl
10. What sublevel are filling in across the transition elements?
4. I - I
11. Which element in group 2 is most likely to lose an electron?
5. C – H
12. Elements within a group have the same number of
13. The majority of elements on the periodic table are
(metals/nonmetals).
6. B – F
7. Zn – Zn
14. Which element in group 16 is most likely to gain an electron?
8. CrF3
9. O2
Review Part 2: Chapters 4 -7, 25
8
EC Version
10. NaNO3
4. F2
8. BCl3
Lewis Dot Diagrams
A Lewis dot diagram is used to show how the electrons in a bond are
distributed. If the bond is an ionic bond, the Lewis Dot Structures shows the
electron transfer and resulting ions. If the bond is covalent the Lewis Dot
Structure shows the electron sharing. For covalent bonds the rules are:
1. Draw a skeleton of the structure (identify the center atom
2. Count the total number of valence electrons (Use PT)
3. Distribute electrons so that each atom has eight dots around it.
4. If you run out of dots and every atom is not satisfying the
octet rule, try double (sharing 4 electrons) or even triple bonds
(sharing 6 electrons between two atoms)
5. Exceptions: H is happy with 2 electrons and B is happy with 6
electrons
Complete the Lewis Dot Structures for the following covalent
molecules. Also predict the shape as linear, bent, trigonal planar,
trigonal pyramidal or tetrahedral.
1. NF3
2. SiI4
Overall Polarity
The bonding that occurs between two atoms has an impact on the overall
polarity of a substance. The overall polarity impacts the substance’s
properties like phase, solubility in water (“like dissolves like”), ability to
conduct electricity etc. To determine the overall polarity use these guidelines:
1. If the bonding is metallic its overall polarity is a metal.
2. If the bonding is ionic its overall polarity is an ionic salt.
3. If the bonding is nonpolar it is considered nonpolar overall.
4. If the bonding is polar covalent, then you must look at the
symmetry of the molecule (lone pairs around center atom).
A. If symmetrical – no lone pairs - (linear, trigonal planar
and tetrahedral shapes) then it is nonpolar overall.
B. If the shape is asymmetrical - has lone pairs - (bent or
trigonal pyramidal shapes) then it is polar overall.
Predict the overall polarity of the following substances.
1. NF3
2. CaO
2. CO2
4. SO2
3. MgBr2
4. SiI4
3. O2
6. OF2
5. F2
6. O2
Review Part 2: Chapters 4 -7, 25
9
EC Version
7. OF2
5. N2O
_______________________________
8. Cu –Cu
6. NO
_______________________________
7. CaCO3
_______________________________
8. KCl
_______________________________
9. FeSO4
_______________________________
10. LiBr
_______________________________
9. BCl3
10. Zn – Zn
Nomenclature
11. MgCl2
12. FeCl3
Writing formulas and naming compounds depends entirely on the
type of bonding present in the substance. The first question that must
be asked is: is the substance ionic, covalent or an acid?
13. Zn3(PO4)2
14. NH4NO3
15. HNO3
Naming Compounds:
Covalents: write first nonmetal name with prefix if there is a subscript,
write second nonmetal name always with prefix and ide ending.
Acids: If two nonmetals, write hydro – nonmetal stem – ic acid
If polyatomic ion (no hydro) change polyatomic ion ending:
ate  ic
ite ous
Ionics: write metal name then either nonmetal name with ide ending or
polyatomic ion name. Decide if metal needs Roman numeral
Write the correct name for the following substances. Decide whether
you need to follow covalent, ionic or acid rules.
1. CO2
_______________________________
2. CO
_______________________________
3. SO2
_______________________________
4. SO3
_______________________________
16. HCl
17. H2SO4
18. H2SO3
19. HC2H3O2
To write formulas correctly, again you have to ask yourself the
question, what type of bonding exists? Is it ionic (starts with a meta),
covalent (two nonmetals), or and acid ( has acid in the name)
Writing Formulas
Covalent:
write first nonmetal symbol using prefix for
subscript. Then write second nonmetal symbol
using its prefix for its subscript.
Ionic:
write metal symbol. Write second part of formula
using the ending as a guide. It is ends in ide write
a nonmetal symbol. If it ends in ate or ite write a
polyatomic symbol. The swap and drop oxidation
numbers to get subscripts.
Review Part 2: Chapters 4 -7, 25
Acids:
write H. If hydro in name, write a nonmetal
10
EC Version
Assigning Oxidation Numbers
Write the correct formulas for the following chemical names.
1. ammonium phosphate
2. iron(II) oxide
An oxidation number can be assigned to each element in a formula.
Here are some of the rules for assigning oxidation numbers.
3. iron(III) oxide
4. carbon monoxide
5. calcium chloride
6. potassium nitrate
7. magnesium hydroxide
8. aluminum sulfate
9.copper(II) sulfate
10. lead(IV) chromate
Rules
1. Elements in their elemental state are 0
2. All the charges must add up to 0 or the superscript
3. Fluorine is always -1 in a compound
4. Oxygen is always -2 in a compound
5. Group 1 metals are +1 in a compound
6. Group 2 elements are +2 in a compound
Example: determine the oxidation number of Cr in Cr2O72–
Let x = ox# for chromium  2x + 7(-2) = –2
x = +6  Cr+6
11. diphosphorus pentoxide
Assign oxidation numbers to each element in the following:
12. carbonic acid
1. HCl
2. KNO3
3. OH-
4. Mg3N2
5. KClO3
6. Al(NO3)3
7. S8
8. PbO2
9. NaHSO4
10. H2SO3
11. H2SO4
12. KMnO4
13. LiH
14. MnO2
15. SO3
16. NH3
17. Na
18. LiClO2
13. potassium permanganate
14. sodium carbonate
15. hydrobromic acid
16. zinc nitrate
17. aluminum sulfite
18. dinitrogen trioxide
19. phosphorous pentachloride
20. aluminum nitrate
21. sulfuric acid
22. acetic acid
23. hydrosulfuric acid
Review Part 3: Chapters 8 – 9, 20
11
EC Version
19. Mg(NO3)2
20. SO4
-2
21. S2O3
-2
Redox Reactions
In a redox reaction one substance is reduced and one substance is oxidized.
You must have both things occurring! To decide what is oxidized and
reduced you must look at what is happening to the oxidation number from
the reactant to the product side:
Reduction:
The oxidation number goes down (reduces) from reactant
to product side.
This occurs because electrons are gained: GER.
 X+2 + 2e-  X
Oxidation:
The oxidation number goes up (increases) from reactant
to product side.
This occurs because electrons are lost: LEO
 X  X+2 + 2e-
1. Examine the reaction H2 + 2 NO  2H2O + N2
a. What is the oxidation half reaction?
b. What is the reduction half reaction?
2. Examine the reaction Cu +H2SO4  CuSO4 + H2O + SO2
a. What is the oxidation half reaction?
b. What is the reduction half reaction?
3. Examine the reaction: KClO3  KCl + O2
a. What is the oxidation half reaction?
Review Part 3: Chapters 8 – 9, 20
b. What is the reduction half reaction?
Gram Formula or Molar Mass
Molar mass is the mass of one mole of a substance. If it is a single
element, it is equal to the mass off the periodic table measured in
grams. If the substance is a formula, you must dissect the formula
and add everything up off the periodic table.
Determine the molar mass of the following:
1. KMnO4 _______________
2. KCl ________________
3. Na2SO4 _______________
4. Ca(NO3)2 ___________
5. Al2(SO4)3 _____________
6. (NH4)3PO4 ___________
Percent Composition
Percent Composition is an analysis of a chemical’s content
using masses. Here is how it is calculated.
% Comp. = mass element
x 100
mass of compound
Determine the % composition of the element specified in the
following:
1. % K in KMnO4
__________
2. % P in (NH4)3PO4
__________
3. % H2O in CuSO4  5H2O
__________
4. % N in Mg (NO3)2
__________
5. % O in Al2(SO4)3
__________
12
Mole Conversions
1 mole = Molar Mass (off PT in g) = 6.02 x 10 23
atoms or molecules = 22.4 L (of gas at STP)
1. Read the problem and identify the unit you are starting
with and the unit you are trying to convert to.
2. Set up factor label and put given in the upper left.
3. Choose the two parts of the above conversion factor and
place them so that units cancel.
4. Multiple across the top and divide by anything on the
bottom.
EC Version
6. Determine the mass of 0.50 moles Al2(SO4)3
Molarity
Molarity is one of about a dozen ways to measure the concentration
or strength of a solution. Other ways include molality, mole fraction,
% by volume, % by mass, or normality. However, molarity seems to
be the most commonly used method.
Molarity = Moles Solute
Liter of Solution


1. Convert 25 g NaCl into moles.
__________

2. Convert 4.3 moles CO2 to molecules
__________
3. Determine the mass of 33.6 L of O2 at STP
__________
__________
The solute must be converted to moles using the molar
mass of the periodic table
The volume of solution or solvent must be in liters. Use
“King Henry…” (KHDBdcm) to complete the metric
conversion
Molarity is abbreviated with a capital M
1. What is the molarity when 107 g of NaCl are dissolved in 1.0 L
of solution?
2. What is the molarity of a solution in which 100.g of AgNO3 are
dissolved in 500. mL of solution?
4. What is the mass of 7.21 x 10 24 atoms P?
__________
3. How many grams of KNO3 should be used to prepare 2.00 L of a
0.500 M solution?
5. Convert 0.25 moles Na2SO4 to grams
Review Part 5: Chapters 14 – 16, 19
__________
13
EC Version
Empirical Formulas
An empirical formula is a formula where the subscripts are reduced.
Most of the ionic formulas that we work with are also empirical
formulas.
Molecular (True) Formulas
A molecular or true formula is one where the subscripts are not
reduced or simplified. Many organic formulas are molecular, like
glucose C6H12O6.
Steps
Steps
1.
2.
3.
4.
Cross out % and make grams if given %
Convert grams to moles for each element
Divide each moles answer by the smallest
The whole numbers obtained in step 3 are your subscripts. If a
.5 number is obtained in step 3 multiply all answers by 2.
5. What is the empirical formula for a substance that is 75 %
carbon and 25 % hydrogen?
6. What is the empirical formula for a substance that is 52.7 %
potassium and 47.3 % chlorine?
1. You must have the empirical formula first.
2. Divide the mass given in the problem by the
molar mass of the empirical formula.
3. Use the whole number answer from number 3 to
multiply the subscripts.
1. The empirical formula of a compound is NO2. Its molecular
molar mass is 92 g/mol. What is the molecular formula?
2. The empirical formula of a compound is CH2. Its molecular
molar mass is 70 g/mol. Determine the molecular formula.
7. What is the empirical formula for a substance that is 22.1 %
aluminum, 25.4 % phosphorus and 52.5 % oxygen?
8. What is the empirical formula for a substance that is 13 %
magnesium and 87 % bromine?
3. A compound is found to be 40.0 % carbon, 6.7 % hydrogen,
and 53.5 % oxygen. Its molecular mass is 60.0 g/mol. What
are the empirical and molecular formulas?
9. What is the empirical formula for a substance that is 32.4 %
sodium, 22.5 % sulfur, and 45.1 % oxygen?
Review Part 5: Chapters 14 – 16, 19
14
Balancing Equations
An equation is balanced when it adheres to the Law of Conservation of
Matter or Mass. In other words:
Mass reactants = Mass products
# atoms of each element of reactant = # atoms of each element of product
Coefficients (whole numbers in front of symbols or formulas) are used to
balance equations
EC Version
Translating Equations
Predict the products of the reactions below. In order to complete
the problem you may want to follow the steps below.
Steps
1. Translate from words to formulas.
2. When writing a formula that is ionic be sure to criss-cross
the oxidation numbers. Watch for Diatomics!
(H2,N2,O2,F2,Cl2,Br2,I2)
3. Balance the equation.
Subscripts (the small numbers in the formulas) tell you how many of the
preceding atom are in each formula unit or molecule.
1. aluminum bromide + chlorine yield aluminum chloride +
bromine
1. _____N2 + _____H2  _____ NH3
2. _____ NaCl + _____F2  _____ NaF + _____Cl2
3. _____AgNO3 + ___MgCl2  ___ AgCl + ___Mg(NO3)2
4. _____CH4 + _____O2  _____CO2
+ _____ H2O
2. potassium chlorate when heated yields potassium chloride and
oxygen gas
5. _____ C8H18 + _____O2  _____ CO2 + ____H2O
6. _____P + _____O2  _____ P2O5
7. _____S8 + _____O2  _____ SO3
8. ____CO2 + _____H2O  _____C6H12O6
3. calcium hydroxide + phosphoric acid yield calcium
phosphate and water
+ _____O2
9. ___CaCO3 + ___HCl  ___CaCl2 + ___H2O + ___CO2
10. ___Na2SO4 + ___AlCl3  ___Al2(SO4)3 + ___ NaCl
4. hydrogen + nitrogen monoxide yield water + nitrogen
11. ___H3PO4 + ___Mg(OH)2  ___ H2O + ___Mg3(PO4)2
Review Part 5: Chapters 14 – 16, 19
15
Types of Reactions
A reaction occurs when bonds are broken or formed. Indicators of a
chemical reaction include: color change, heat exchange, precipitate
formation or gas production (bubbling).
Predicting Products
Predict the products of the reactions below. In order to complete
the problem you may want to follow the steps below.
5 Types
5.
6.
7.
8.
9.
Synthesis – A + B  AB
Decomposition – AB  A + B
Combustion – CxHy + O2  H2O + CO2
Single Replacement A + BC  AC + B (look at activity series)
Double Replacement – AB + CD  AD + CB (look for precipitate)
EC Version
Steps
1. Identify the reaction pattern (type) to follow.
2. Write the products by criss-crossing oxidation numbers for
compounds. Watch for Diatomics! (H2,N2,O2,F2,Cl2,Br2,I2)
3. Balance the equation.
4. If the reaction will not occur, write NR after the arrow.
Identify the type of reaction represented in each equation below.
1. 2H2 + O2  2 H2O
_____________
1. ___Al + ___FeCl3 
2. 2H2O  2 H2 + O2
_____________
2. ___I2 + ___NaCl 
3. Zn + H2SO4  ZnSO4 + H2
_____________
3. ___AgNO3 + ___ZnCl2 
4. 2 CO + O2  2CO2
_____________
4. ___H2 + ___AuCl3 
5. 2HgO  2Hg + O2
_____________
5. ___Sn + ___H2O 
6. 2KBr + Cl2  2KCl + Br2
_____________
6. ___(NH4)2CO3 + ___KOH 
7. CaO + H2O  Ca(OH)2
_____________
7. ___Cl2 + ___NaI 
8. AgNO3 + NaCl  AgCl + NaNO3
_____________
8. ___Zn + ___HCl 
9. 2H2O2  2H2O + O2
_____________
9. ___Ba + ___Mg(OH)2 
10. Ca(OH)2 + H2SO4  CaSO4 + 2H2O
_____________
10. ___(NH4)2CO3 + ___SrCl2 
Review Part 5: Chapters 14 – 16, 19
16
EC Version
Stoichiometry
In a stochiometry problem, you are given the amount of one reactant
or product and you are assigned the task of determining the amount
of any other reactant or product. The key to relating two different
chemicals in an equation is the mole ratio which comes from the
COEFFICIENTS. The problems may be mass-mass, mole-mole,
mass-mole, or mass-volume. Below are the steps to completing any
stoichiometry problem.
12. How many moles of HCl are required to react with 125 g of Zn?
Zn + 2 HCl  ZnCl2 + H2
13. How many molecules of water are produced if 2.0 g of sodium
sulfate are produced in the equation below?
H2SO4 + 2 NaOH  2 H2O + Na2SO4
Stoichiometry Steps
1.
2.
3.
4.
Get a balanced equation.
Write the given and X above the equation with units
Write mole ratio (coefficients) below equation
Do a mole conversion to convert the bottom units to match the
top units of the equation. If the unit on top is….
in grams – use molar mass off PT.
in atoms or molecules – use 6.02 x 10 23
in L at STP - use 22.4 L
in moles – you do not need to do anything
in kJ – write the heat value below the equation
5. Cross multiply and solve for X
14. How many grams of NH3 are produced by reacting 50.0 g of
nitrogen?
N2 + 3 H2  2 NH3
15. What volume of hydrogen (at STP) is required to react with the
nitrogen in the above reaction? N2 + 3 H2  2 NH3
10. How many moles are required to complete react with two moles
of nitrogen? N2 + 3 H2  2 NH3
16. What mass of phosphorus must react with 32.5 g of oxygen?
2P + 5O2  2P2O5
11. What volume of oxygen gas is produced at STP by decomposing
6.0 moles of potassium chlorate? 2KClO3  2KCl + 3O2
17. Determine the amount of heat produced when 100. g of CH4
combust.
CH4 + 2O2  CO2 + 2H2O + 1063 kJ
Review Part 5: Chapters 14 – 16, 19
17
EC Version
Limiting Reactants
The limiting reactant is the reactant that runs out first. When it is
gone the reaction stops. Almost always, when there are two
reactants, one is a limiting and one is the excess reactant. You will
be given two numbers in these problems. Sometimes the problem
will just ask you to identify the limiting reactant. Other times it will
ask for how much product will be formed.
3. What mass of water is produced in the reaction in question #2?
4. How many liters of CO2 are produced when 2 L of C3H7OH react
with 4 L of oxygen at 50ºC under 1 atm of pressure?
2 C3H7OH(g) + 9 O2(g)  6 CO2(g) + 8 H2O(g)
Steps
1. Make sure the reaction is properly balanced.
2. Determine the number of moles of product that each
reactant amount will give.
3. The limiting reactant is the one that produces the
SMALLER number of moles of product.
4. Cross out the molar amount of product calculated
from the reactant in EXCESS.
5. Finish calculating the number of grams of product
using the remaining number of moles of product.
1. Determine the limiting reactant when 10 g of Li react with 10 g of
H2O.
2Li + 2H2O  2 LiOH + H2
2. Determine the limiting reactant when 4.7 moles of oxygen react
with 100. g of C3H7OH.
2C3H7OH + 9O2  6 CO2 + 8 H2O
Review Part 5: Chapters 14 – 16, 19
Pressure Conversions
The reference packet identifies standard pressure as being: 1.00 atm =
101.3 kPa = 760. mmHg = 760. Torr You can use this information to
convert from one pressure unit to another. For example, convert
1atm
 0.992atm
754 Torr to atm 754 Torr x
760Torr
Complete the following unit conversions:
1. 54 mmHg =
atm
8. 900 Torr =
atm
2. 780 mmHg =
atm
9. 120 kPa =
atm
3. 95 kPa =
atm
10. 76 Torr =
atm
4. 100 mL =
L
11. 5 x 10–2 mL =
L
5. 5 mL =
L
12. 1.5 x 105 mL =
L
6. 25 C =
K
13. 100 C =
K
7. -60 C =
K
14. Standard temp =
K
18
EC Version
Combined Gas Law
5. Determine the volume at STP of a gas originally occupying 7.3 L
at 25 C and 70 kPa.
The combined gas law is Boyle’s, Charles’, and Gay-Lussac’s gas
laws all put together. Within the one combined gas law you can see
all three of the individual gas laws.
6. A gas at 33 C and 100 kPa is heated to 75 C. What is the new
P1V1
T1
= P2V2
T2
P1 = Initial pressure
P2 = new pressure
V1 = Initial volume
V2= new volume
T1 = Initial temperature in K
T2 = new temperature
Use this formula for changing conditions. Look for words
that mean change – increase, decrease, rise, lowers new,
original
Temperature must be in Kelvin K =  C + 273
pressure?
Graham’s Law of Diffusion
Graham’s Law of Diffusion basically says that a light gas travels
faster than a heavier gas. Use the molar masses off the Periodic
Table to determine the heavy versus the lighter gas.
1. Determine the volume if 1.5 atm of gas at 20 C in a 3.0 L vessel
are heated to 30 C at a pressure of 2.5 atm.
2. A sample of oxygen gas occupies a volume of 250 mL at 740 torr
of pressure. What will the volume be when the pressure is
increased to 800 torr?
# times faster = heavy/light
Heavy = molar mass of heavy gas
Light = molar mass of lighter gas
There are no units on the answer
1. Which gas is faster - Carbon tetrachloride or Bromine gas?
2. How much faster is hydrogen gas than sulfur trioxide gas?
3. A sample of nitrogen occupies a volume at 250 mL at 25 C.
What will it occupy at 95 C?
3. If helium is placed in one end of an evacuated tube and dinitrogen
monoxide is placed in the other end, where will the two gases
meet?
4. What is the temperature in Celsius when 2.5 L of gas at 22 C and
600 mm of Hg are transferred to a 1.8 L vessel at 760 mm of Hg?
Review Part 5: Chapters 14 – 16, 19
4.
How much faster is helium gas than xenon gas?
19
EC Version
Ideal Gas Law
The ideal gas law is the only gas law that relates the four measurable
aspects of gases – moles, volume, pressure and temperature. It is
NOT used for changing conditions like the combined is. It is the one
law where you must be picky about the units.
PV = nRT
P = pressure(must be in atm)
V = volume (must be in L)
n = moles (may need to convert grams to moles)
R = universal gas constant = 0.0821 if pressure is in atm
T = temperature (must be in Kelvin)
K =  C + 273
Ptotal = P1 + P2 + P3
Patmospheric = Pdry gas + PH2O
P1, P2, and P3 = the pressure of each individual gas
and must have the same units
PH2O = pressure of water vapor left behind during
water displacement. It depends on the temperature
and is looked up on a table
1. What is the total pressure of a mixture of gases if the
oxygen’s pressure is 75 kPa, the nitrogen’s pressure is 120
kPa and the carbon dioxide’s pressure is 33 kPa? __________
Is this mixture at standard pressure? ________
4. How many moles of oxygen will occupy a volume of 2.5 L at 1.2
atm and 25 C?
5. What pressure will be exerted by 25 g of CO2 at a temperature of
25 C and a volume of 500 mL?
6. What mass of argon occupies 4.3 L at 70 kPa and 20 C?
Dalton’s Law of Partial Pressures
Dalton’s Law of Partial Pressures is a relatively easy law that
says that the sum of the pressures of the individual gases in a
mixture is equal to the total pressure of that mixture. The key
word to this law is mixture. However, there is a special
circumstance when this law is used and it is when a gas is
produced in lab via water displacement or “collected over water”.
Review Part 5: Chapters 14 – 16, 19
2. What is the pressure of the carbon dioxide gas if a mixture is
at standard pressure and the oxygen is at 0.37 atm and the
nitrogen is at 450 mm Hg? ___________________________
3. What is the pressure of oxygen gas that is collected over
water at 25 C at an atmospheric pressure of 800 mm of Hg?
___________________________________
4.
A 32 mL sample of hydrogen is collected over water at 20 C
and 750 mm of Hg. What is the volume of the dry gas at
STP? (Vapor pressure of water at 20 C = 17.5 mm of Hg)
_______________________
20
Solubility Curves
EC Version
12. What happens when the solution in #11 cools down to room
temperature (~25ºC)?
Answer the following questions about the solubility curve located at
right.
1. What salt is least soluble in water at 20C_________
2. How many grams of potassium chloride can be dissolved in 200 g
of water at 80C? _____________________
3. At 40C, how much potassium nitrate can be dissolved in 300g of
water? ____________________________
4. Which salt shows the least change in solubility from 0 – 100C?
_____________________
5. At 30C, 90 g of sodium nitrate is dissolved in 100 g of water. Is
this solution saturated, unsaturated, or supersaturated?
____________________________
6. A saturated solution of potassium chlorate is formed in one
hundred grams of water. If the saturated solution is cooled from
80 C to 50 C, how many grams of precipitate are formed?
________________________
7. What compounds show a decrease in solubility for 0 – 100 C?
_____________________why_______________
8. Which salt is most soluble at 10 C? _____________
9. Which salt is least soluble at 50 C? ____________
10. Which salt is least soluble at 90 C? _____________
11. How can a saturated solution of ammonium chloride at 50ºC be
prepared?
Review Part 5: Chapters 14 – 16, 19
21
EC Version
Colligative Properties
Colligative properties are properties of a solvent that are changed
when a solute is added, and depend ONLY on the number
(moles) of particles that are present. They include boiling point
elevation, vapor pressure lowering and freezing point depression.
Nonelectrolytes are solutes that DO NOT break apart into ions.
The number of particles that can be obtained from an electrolyte
(an ionic solute) can be determined by adding up the subscripts in
the formula.
Which of the following will have the lowest freezing point/highest
boiling point (circle)?
1. 0.1 M sucrose or 0.1 M NaCl
2. 0.1 M NaCl or 0.05 M Al(NO3)3
3. 0.1 M NaCl or 0.1 M Al2(CO3)2 (check your solubility rules!)
4. A student measures the conductivity (with a light bulb and a
battery) and freezing point of the following solutions. What can
she conclude about the solubility of her solutes based on her data?
Solution
Freezing point (ºC)
Conductivity
1 M NaCl
-3
Bright
1 M MgCl2
-6
Bright
1M Al(OH)3
-0.2
Very dim
Electrolytes and
Electrolytes are substances that when placed in water dissociate
(separate) into ions which allow the solution to conduct electricity.
Substances that contain ionic bonds are electrolytes along with acids
and bases. Nonelectrolytes contain covalent bonds. Classify the
following as either an electrolyte or nonelectrolyte.
Nonelectrolytes
Review Part 5: Chapters 14 – 16, 19
Compound
Electrolyte
Nonelectrolyte
NaCl
CH3OH (methanol)
C3H5(OH)3 (glycerol)
C6H12O6
HCl
NaOH
Like Dissolves Like
Liquids can dissolve substances that contain the same type of overall
polarity. However, if two substances do not have the same type of
overall polarity, they will not be soluble in each other or, in other
words, will be immiscible. An example of two immiscible substances
are oil and water because water is a polar substance overall and oil is
nonpolar overall. Water is special because and is called the universal
solvent because, although it is polar, it can dissolve both polar and
ionic substances.
Substance
Dissolve in Water
Dissolve in Oil
22
EC Version
CCl4
Heat
NaCl
Heat is a form of energy. Like all forms of energy its quantity can be
measured in joules or calories. When no reaction is occurring, adding
or removing heat can change the temperature of an object or change its
phase, but it cannot do both at the same time.
Changing Temperature:
Changing Phase:
Ca(NO3)2
I2
NH3
q= mCpT
q = m Hfus or q = mHvap
q = heat, m = mass, Cp = specific heat
Predicting Properties
Once the overall polarity is known, then the general properties of a
substance can be determined.
1. Ionic Salts – have high melting points, dissolve to dissociate into
ions and create a solution that conducts electricity, dissolve in
water, and are all solids.
2. Metals – very high melting points, conduct electricity as solids, are
all solids (except mercury), do not dissolve in water
3. Nonpolar – lowest melting points, do not dissolve in water, never
conduct electricity, frequently gases.
4. Polar – low melting points, dissolve in water, never conduct
electricity, and can be any phase.
Fill in, completely, the following table for each substance.
Substance Melting Point Dissolve Ability to Conduct
(high/low)
in water? Electricity
CaO
F2
(melting)
(boiling)
Complete the following heat calculations. First decide if it is a
changing temperature problem or a changing phase problem. Then
select the correct formula and find the constants on the reference
table.
1. Determine the heat needed to warm 25.3 g of copper from 22 C
to 39 C.
2. What quantity of heat is required to boil 150 g of water?
3. What metal is present in a 250 g sample if 4041 J of heat were
released to cool the sample from 41  C to 5  C?
Phase
4. How much heat is needed to raise the temperature of 100 g of
water from 10ºC to 50ºC?
OF2
Zn- Zn
SiI4
NF3
MgBr2
Review Part 5: Chapters 14 – 16, 19
Acids and Bases
23
EC Version
Identify whether each of the following is a property of an acid (A) a
base (B) or both (AB)
____ 1. Feel slippery
____ 6. pH < 7
____ 2. Turn blue litmus red
____ 7. Turn red litmus blue
____ 3. corrosive
____ 8. electrolytes
____ 4. Taste bitter
____ 9. Taste sour
____ 5. pH > 7
____ 10. Cause indicators to
change colors
2. Boiling can occur at 50  C and at about ___________ atm.
3. What is the phase at 75 C and 0.01 atm? ____________
4. All three phases can exist at _________ C and ______ atm.
These conditions are called the _______________________.
5. What is the phase at – 5 C and 0.5 atm? ______________
6. What is the phase at 300 C and 224 atm? _____________
Phase Diagrams
Reaction Energy Diagrams
Identify the quantities represented by letters a – e in Figures 1 and 2.
Include numerical values where appropriate.
1. On the diagram, label where condensation, freezing and
sublimation each occur.
Review Part 5: Chapters 14 – 16, 19
24
EC Version
Figure 1
Figure 2
(a)
(a)
(b)
(b)
(c)
(c)
(d)
(d)
(e)
(e)
(f) Is this an endothermic or exothermic reaction?
(f) Is this an endothermic or exothermic reaction?
(g) How will adding a catalyst affect this diagram?
(g) How will adding a catalyst affect this diagram?
Review Part 5: Chapters 14 – 16, 19
25
Chemical Equilibria
EC Version
A reaction is in equilibrium when the:
Rate of the forward reaction = rate of the reverse reaction
The equilibrium constant, K, is the ratio of the molar concentrations
of the products divided by the molar concentrations of the reactants
(solids and liquids do NOT appear in the equilibrium constant
expression). Weak acids and weak bases are designated by Ka and Kb,
respectively. The equilibrium that involves a partially soluble solute
in equilibrium is designated with Ksp.
Stressing a reaction that is at equilibrium causes it to increase the rate
of either the forward or reverse reaction in order to restore
equilibrium. This response is called a Le Châtelier shift.
Write equilibrium constant expressions for the following reactions:
1. N2(g) + 3 H2(g)  2 NH3 (g)
2. CaCO3(s)  Ca2+(aq) + CO32–(aq)
3. H2O(l)  H2O(g)
4. How will increasing the pressure affect reactions #1 and 3?
5. What kind of shift will be observed if more H2 is added to
reaction #1?
6. What kind of shift will be observed if more Ca2+ is added to
reaction #2?
7. What kind of shift will be observed if the [H+] concentration is
lowered in reaction #3 by adding NaOH?
Review Part 5: Chapters 14 – 16, 19
26