Final Exam Review – Fall 2015
1. Scientific Method
a. Read the following story and then answer the questions.
A scientist wants to find out why sea water freezes at a lower temperature
than fresh water. The scientist goes to the library and reads a number of
articles about the physical properties of solutions. The scientist also reads
about the composition of sea water. The scientist travels to a nearby beach
and observes the conditions there. The scientist notes the taste of sea
water and other factors such as waves, air pressure, temperature, and
humidity. After considering all this information, the scientist sits at a desk
and writes, “IF sea water has salt in it, it will freeze ata a lower temperature
than fresh water.” The scientist goes to the laboratory and does the
following:
i. Fills each of two beakers with 1 liter of fresh water
ii. Dissolves 35 grams of table salt in one of the beakers
iii. Places both beakers in the freezer at a temperature of -1°C
iv. Leaves the beakers in a freezer for 24 hours.
After 24 hours, the scientist examines both beakers and finds the fresh water
to be frozen. The salt water is still a liquid. The scientist writes in a notebook,
“It appears that salt water freezes at a lower temperature than salt water.” The
scientist continues, “I suggest that the reason sea water freezes at a lower
temperature is that sea water contains dissolved salts, while fresh water does
not.”
v. Questions
1. Which statement(s) contain conclusions?
2. Which statement(s) contains a hypothesis?
3. Which statement(s) contain observations?
4. Which statement(s) describe an experiment?
5. In which statement is the problem described?
6. Which statement(s) contain data?
7. Which is the independent variable in the experiment?
8. Which is the dependent variable in the experiment?
2. Lab Safety – Be aware of the lab safety rules on the Flynn Scientific Lab Safety Contract
you signed in August. No, you do not have to memorize them, just be aware of them and
use your common sense.
3. Lab Equipment – Be able to identify the following pieces of lab equipment.
4. Scientific Notation – IS NOT A MATH PROBLEM, IT IS JUST A METHOD FOR
WRITING REALLY BIG OR SMALL NUMBERS! You need to be able to put numbers into
scientific notation and solve problems, using a calculator.
a. Writing numbers in scientific notation.
i. Put the decimal place right after the first number.
ii. Add × 10
iii. Add the exponent- it’s how many places you moved the decimal point. If
the original number was a decimal (smaller than one) make the exponent
negative.
Examples
1. Mass of the International Space Station
417,289,000 g. =>
4.17289 × 108 g
2. Mass of a flea
0.00008704 mg => 8.704 × 10-5 mg
b. Put these numbers into scientific notation.
i. Speed of light:
299792458 m/s
ii. Number of seconds in a day: 86400 s
iii. Mean radius of Earth:
6378 km
iv. Density of oxygen gas at 0°C: 0.00142 g/mL
v. Radius of an argon atom:
0.000000000098 m
c. Use your calculator (Remember the EE/EXP/× 10n button on your
calculator) to solve these problems.
i. (4.3 × 108) × (2.0 × 106) =
ii. 7.8 × 103 / 1.2 × 104 =
iii. (6.02 × 1023) × (3.00 × 10-5) =
iv. 8.64 × 10-6 / 7.53 × 10-7 =
5. Significant Figures
a. Rules for what is or is not considered significant.
i. Any nonzero number is significant.
ii. Any zeroes between two numbers is significant.
iii. Any zeroes before a number is NOT significant.
iv. Any zeroes after numbers is significant ONLY if that number has a
decimal point in it.
1. Examples – The underlined numbers are significant
a. 234.56 g
= 5 s.f.
b. 6070809.06 cm = 9 s.f.
c. 0.000465 mL
= 3 s.f.
d. 5500 J
= 2 s.f.
e. 44400.00 J
= 7 s.f.
b. Adding & Subtracting using significant figures: Your answer should have the
same number of significant figures as the original number with the LEAST
AMOUNT of significant figures AFTER the decimal point.
i. Examples
1. 199.71 g + 11.2602 g = 210.9702 g => 210.97 g
2. 428.6 mL – 50.83 mL = 377.77 mL => 377.8 mL
c. Multiplying & Dividing using significant figures: You answer should have the
same amount of significant figures as the original number with the LEAST
AMOUNT of significant figures total.
i. Example
1. 0.0456023 g/mL × 0.5402 mL = 0.02463436 g => 0.02453 g
ii. Significant Figures when doing dimensional analysis: Since dimensional
analysis is really just a conversion, anything you multiple or divide by has
infinite significant figures. Therefore, your answer should have as many
significant figures as your original number does.
iii. Examples
1. 15.5 years × (365.25 days) × (24 hours) = 135873 hours =
(1 year )
(1 day)
136000 hours
iv. Solve these problems.
1. 1.942 g – 1.926 g =
2. 33.43 g / 7.14 mL =
3. .417 cm × .8030 cm =
6. Nuclear Chemistry
a. Nuclear Fission - breaking down of a larger atom into smaller particles + energy
i. Produces a lot of energy
ii. Used in nuclear power plants and nuclear bombs
1. Creates radioactive waste that must be contained and disposed
b. Nuclear Fusion – combining of two or more smaller particles into a larger atom +
energy
i. Produces a lot more energy that fission but only works at really high
temperatures
ii. Produced all the elements heavier than helium in stars
7. States of Matter
a.
Solid
Liquid
Gas
Plasma
Particles
Very tightly
Closely
Far apart
Very far apart
packed
packed
Motion
Vibrate in
Slide past
Constant,
Constant,
position
each other
random
random, very
fast
Energy
Least
Most
Shape
Definite
Not Definite
Not Definite
Not Definite
Volume
Definite
Definite
Not Definite
Not Definite
8. Pure Substances vs. Mixtures
a. Determine if the following is an element (E), Compound (C), Homogeneous
Mixture (Ho), or Heterogneneous Mixture (He).
i. Iron filings (Fe)
ii. Limestone (CaCO3)
iii. Orange juice (juice w/ pulp)
iv. Ocean water
v. Air
vi. Brass
vii. Pure water
viii. Chromium
ix. Pizza
9. Physical vs Chemical Properties & Changes
a. Physical Properties & Changes
i. Physical Properties and Changes – can be seen through direct
observation, changes do NOT result in a new substance forming.
ii. Chemical Properties and Changes – can only be observed when one
substance changes into a new substance.
iii. Determine if the following are physical properties/changes or
chemical properties/changes.
1. Flammability
2. Boiling point
3. Reactivity
4. Density
5. Malleability
6. Glass breaking
7. Water evaporating
8. Burning leaves
9. Melting ice
10. A rusting bicycle
11. Frying an egg
12. Crushing a can
10. Nomenclature – Feel free to use the attached flow charts to assist you when work on
these problems.
a. Type I Ionic Compounds – regular metal + nonmetal or a polyatomic ion
i. Formulas: Criss – cross charges
ii. Names: Change the ending of the 2nd element to –ide (do not so that for
polyatomic ions, however.)
b. Type II Ionic Compounds – transition metal + nonmeatl or a polyatomic ion
i. Formulas: Criss-cross charges
ii. Names: same as for type I BUT you must write the original charge of the
metal as a Roman Numeral in parentheses!
1. Roman Numerals
a. One  I
b. Two  II
c. Three  III
d. Four  IV
e. Five  V
f. Six  VI
g. Seven  VII
h. Eight  VIII
i. Nine  IX
j. Ten  X
c. Type III Covalent Compounds – 2 nonmetals
i. Formulas: No Criss-crossing at all! The prefix becomes the subscript.
ii. Names: Use the prefixes for the subscripts!
1. Prefixes:
a. One  monob. Two  dic. Three  trid. Four  tetrae. Five  pentaf. Six  hexag. Seven  heptah. Eight  octai. Nine  nonaj. Ten  decad. Acids – Always have H+ first in the formula or acid in the name!!!
i. Formulas: Write H+ and then the 2nd element (if binary) or the polyatomic
ion (if an oxyacid). Criss-cross charges.
ii. Names:
1. Binary (no oxygen): Write hydro + 2nd element + ic acid
2. Oxyacid (w/oxygen)
a. Comes from –ate? Change –ate to –ic acid.
b. Comes from –ite? Change –ite to –ous acid.
e. Write the names or formulas for these compounds:
i. BaSO3
xi. Hydrobromic acid
ii. PBr3
xii. Chromium (III) carbonate
iii. CaO
xiii. Lead (II) nitrate
iv. H3PO4
xiv. Sodium oxalate
v. SO3
xv. Iron (III) sulfate
vi. HClO
xvi. Silicon dioxide
vii. MnO
xvii. Calcium chloride
viii. Al2(CrO4)3
xviii. Ammonium hydroxide
ix. Sn(NO3)4
xix. Carbonic acid
x. NaC2H3O2
xx. Magnesium phosphate
11. Chemical Reactions
a. Evidence of a chemical reaction
i. A permanent color change
ii. Formation of gas bubbles
iii. A precipitate (solid)forms
iv. Change in energy (temperature)
v. Formation of water
b. Balancing Chemical Equations – you use coefficients (numbers that go in front of
a formula & multiply through every atom in that formula) to make sure that you
have the same number of atoms of each element on both sides of the reaction.
This satisfies the Law of Conservation of Mass!
i. Balance these chemical equations.
1. GaF3 + Na3PO4  GaPO4 + NaF
2. CaBr2 + Mg(NO3)2  Ca(NO3)2 + MgBr2
3. Na2CO3 + H3PO4  Na3PO4 + H2O + CO2
4. Mg + Co(NO2)3  Co + Mg(NO2)2
5. Sn + P4  Sn3P4
6. C4H8O + O2  CO2 + H2O
7.
8.
9.
10.
c.
C5H12 + O2  CO2 + H2O
O3  O2
S8 + O2  SO2
H2O2  H2O + O2
Types of Chemical Reactions
i. Identify the type of reaction for the following equations.
1. NaBr + H2SO4  Na2SO4 + HBr
2. Mg + Fe2O3  Fe + MgO
3. PbSO4  PbSO3 + O2
4. C2H4 + O2  CO2 + H2O
5. H2O + SO3  H2SO4
12. Moles
a. Molar Mass – the sum of the atomic mass of each atom of each element in a
compound.
i. Example
H2CO3
H: 2 × 1.01 g/mol
= 2.02 g/mol
C: 1 × 12.00 g/mol
= 12.01 g/mol
O: 3 × 16.00 g/mol
= 48.00 g/mol
62.03 g/mol
ii. Find the molar masses of these compounds.
1. H2SO3
2. FeCl3
3. C6H12O6
4. Ca(ClO4)2
5. Al2(C2O4)3
b. % Composition –
i. % elememt = mass element × 100
molar mass
ii. Find the % of each element in the following compounds.
1. H2SO3
2. FeCl3
3. C6H12O6
c. Empirical & Molecular Formulas
i. Empirical formula – is the formula of a compound that has had its
subscripts reduced to the lowest possible ratio.
ii. Molecular formula – the true formula for a compound.
iii. Determine if the following compounds are empirical or molecular
formulas. If a formula is molecular, write its empirical formula.
1. C6H6
2. H2O
3. N2O4
4. C6H12O6
5. C12H22O11
6. H2C2O4
7. H2CO3
d. Mole Relationships – Use the mole map to help you solve the problems.
e. Solve these problems.
i. How many moles are in 62.00 g of N2O?
ii. How many atoms are in 0.777 mol of Au?
iii. How many moles are in 20.00 L of O2 gas at STP?
iv. What is the mass, in grams, of 1.05 L of CO2 gas at STP?
v. How many formula units are in 25.00 g of Mg(NO3)2?
vi. What is the volume, in L, of 5.12 × 1022 molecules of SO3 gas at
STP?
vii. What is the mass, in grams, of 3.00 × 1023 formula units of NaCl?
13. Stoichiometry
What is stoichiometry?
One time I was making sandwiches for some of my son’s friends who had inexplicably been invited over to
my house on a “playdate”. All I had was crackers and cheese in the house, and the kids all decided that the
proper way to eat them was to put one piece of cheese between two crackers to make a little sandwich. It’s
a miracle I didn’t kill any of them. Anyway, I looked in the fridge and found about a zillion pieces of
cheese, and the package of crackers had a sleeve of 20 remaining. The question: How many cracker
sandwiches can I make?
Here’s the math:

If I have 20 crackers and assume that I have infinite quantities of cheese, I can make 10 cracker
sandwiches.
Because I run out of crackers at 10 cracker sandwiches, that’s the maximum quantity I can make. And
that’s what I made them.
Believe it or not, this story actually answers the question of what stoichiometry is. Here’s a more explicit
version for those of you who didn’t like the story:
Stoichiometry is a set of calculations you perform to figure out how much stuff you can make in a
reaction, or how much stuff you will need to make the reaction occur.
In other words, stoichiometry is used to figure out if you’ve got enough crackers to make 30 sandwiches, or
how much cheese you’ll need to make 15 sandwiches. Of course, since chemistry uses fancy symbols,
we’ll deal with all of that in a second. However, that’s the basic idea.
How to do stoichiometry
Before we do anything, we’re going to make a modified version of the diagram we saw back when we were
doing mole calculations:
Let’s see what it all means using the following example:
Example: Using the equation 2 H2 + O2 →2 H2O, determine how many grams of water can be
formed from 45.0 grams of oxygen and an excess of hydrogen gas.
So, where do we begin? We begin by figuring out what that diagram above means:

The box that says “grams of what you’ve got” refers to the number of grams that you’ve been
given in the problem. In our example, we literally see “45.0 grams of oxygen”, so that’s where we
start.



The box that says “moles of what you’ve got” means that before we even start talking about water,
we’ve got to figure out how many moles of oxygen we have. Since you already know how to do
mole calculations (using the molar mass of what you’ve got, shown above), you should be OK.
The box that says “moles of what you want” refers to the fact that, using the equation for this
reaction, you can convert “moles of oxygen” to “moles of water.” We do this using the mole ratio,
which literally just consists of the numbers written down in the equation. We’ll get back to that in
a sec.
The box that says “grams of what you want” refers to what is likely your desired answer. To get
this value, convert the moles of water to grams of water using water’s molar mass. When you’re
finished with this, you’re done!
Let’s just go ahead and do this example, using the methods you’ve seen before to do conversions: The Tchart method:
Step 1: Draw a t
There it is!
Step 2: Put whatever the problem tells you in the top left of the t.
In this case, the problem tells you that you have 45.0 grams of oxygen, so write “45.0 grams of oxygen” in
the top left of this t.
Step 3: Write the units of whatever was in the top left at the bottom right.
Since “grams of oxygen” was written at the top left, write “grams of oxygen” at the bottom right.
Step 4: Write the units of whatever the next step is on the top right.
In the first step of this calculation we use our table to see that we’re converting from grams of oxygen to
moles of oxygen. As a result, write “moles of oxygen” in the top right:
Step 5: Put numbers before each blank on the right side of the t, corresponding to the conversion
factors you need.
This is exactly the same as grams/moles conversions, except that we’ll do more later. What this means is
that we’ll put “1” in front of “moles” (because we always do during mole calculations) and the molar mass
of O2 in front of “grams” (it’s 32.0 g for those of you playing at home):
Step 6: Repeat these steps until you’re done.
You’ll get the hang of what to do before long, but I’ll keep going through all of these steps in this example
to make sure you’re comfortable with the calculations.
Step 7: Add another section to the t, and write the units of the thing in the top left on the bottom
right:
Step 8: Write the units of the thing you want to find in this step in the top right.
We’re converting from moles of oxygen to moles of water here, so write “moles of water” in the top right:
Step 9: Add the conversion factors in the blanks on the right.
Now, given that we have “moles” on both the top and the bottom, it doesn’t really make sense to put “1” in
each spot as we usually do. Instead, realizing that the equation gives us a ratio of the number of moles of
oxygen to number of moles of water (these are the coefficients in the equation), we’ll put these numbers in
front of each number. This ratio is called the “mole ratio”, because it’s a ratio of moles.
Step 10: Do the last conversion from moles of water to grams of water, using the standard t-chart
method.
Step 11: Do the math:
The whole t-chart thing you just did is just a big bunch of fractions being multiplied together, so think of it
like this:
And that’s how you do stoichiometry!
a. Solve these stoichiometry problems. The balanced chemical reaction for all
the problems is below:
6 CO2 (g) + 6 H2O (l)  C6H12O6 (l) + 6 O2 (g)
i. How many moles of C6H12O6 are produced from the reaction of 18.0
moles of CO2 with excess H2O?
ii. How many moles of O2 gas are produced from the reaction of 54.06
g of water?
iii. What mass of glucose (C6H12O6) are produced from the reaction of
88.02g of carbon dioxide with excess water?
iv. What mass of glucose (C6H12O6) are produced from the reaction of
9.01 g of water with excess carbon dioxide?
b. Limiting Reactants
i. The limiting reactant is the reactant that controls how much product can
be produced in a chemical reaction because you run out of that reactant
first.
ii. The excess reactant is the other reactant(s) in the reaction.
iii. Solving limiting reactant problems
1. In a limiting reactant problem, you must be given the amounts of
all the reactants in the reaction. You have to complete a
stoichiometry equation for each reactant, solving for the same
product each time. Whichever reactant gives the smallest
amount of product, that’s the limiting reactant!
a. By the way, that smallest amount of product made is
called the theoretical yield.
2. Example
2 K + Cl2  2 KCl
If 40.00g of K reacts with 85.00 g of Cl2, which reactant is the
limiting reactant and what is the theoretical yield of KCl?
40.00 g K × (1 mol K) × (2 mol KCl) = 1.023 mol KCl
(39.10 g K) (2 mol K )
85.00 g Cl2 × (1 mol Cl2 ) × (2 mol KCl) = 2.128 mol KCl
(79.90 g Cl2) (1 mol Cl2)
Since 40.00 g K made the least amount of KCl, that makes K the
limiting reactant.
To find the theoretical yield, take the 1.023 mol of KCl and
convert it to grams.
1.023 mol KCl × (74.55 g KCl) = 76.27 g KCl
1 mol KCl
3. Solve this limiting reactant problem.
2 NaOH + H2S  Na2S + 2 H2O
If 20.00 g of NaOH reacts with 20.00 g of H2S, what is the
limiting reactant and the theoretical yield of H2O?
14. Collision Theory
15. Kinetic Molecular Theory
16. Heat
a. heat is the flow of energy due to a temperature difference, from a higher
temperature to a lower temperature. (There is no such thing as cold!)
b. temperature is a measure of the average kinetic energy a substance’s molecules
have
c. q = mCΔT
i. q = heat, measured in joules (J)
ii. m = mass, measured in grams (g)
iii. C = specific heat capacity, measured in joules/gram degree Celsius (J/g
°C)
iv. ΔT = Tf – Ti difference in temperature, measured in degrees Celsius (°C)
v. Endothermic processes – heat is added to the system, ending up at a
higher temperature than before (q is positive)
vi. Exothermic processes – heat is released by the system to the
surroundings, ending up at a lower temperature than before (q is
negative)
vii. Solve these problems
1. Find the amount of heat needed to raise the temperature of
5.00g of a substance from 20.0°C to 30.0°C if the heat
capacity if 2.01 J/g°C.
2. The temperature of a 250.0 g ball of iron increases from
19.0°C to 32.0°C. How much heat did the ball gain if the heat
capacity of iron is 0.450 J/g°C.
3. 10.0 g of gold is cooled from 25.0°C to 15.0°C. The heat
capacity of gold is 0.13 J/g°C, how much heat is given off?
17. Equilibrium
a. Definition – occurs when the rate that reactants become products (forward
reaction) EQUALS the rate that products become reactants (reverse reaction).
b. LaChâtelier’s Principle – states that when a stress is added to a system in
equilibrium, the system will shift in order to relieve the stress.
i. It can shift to the right (forward reaction) or to the left (reverse reaction).
ii. Stresses
1. Catalysts – only increase the rates of the forward or reverse
reaction but does not have any effect on the equilibrium.
2. Temperature –
a. Adding heat/increasing temperature will shift equilibrium
away from where the heat is in the reaction (endothermic
– reactant side or exothermic – products side)
b. Removing heat/decreasing temperature will shift the
equilibrium toward where the heat is in the reaction.
3. Concentration –
a. Increasing concentration/adding a substance will shift
equilibrium away from that substance.
b. Decreasing concentration/removing a substance will shift
equilibrium toward that substance.
4. Pressure – gas systems only
a. Increasing pressure will shift equilibrium away from the
side with more moles of substances.
b. Decreasing pressure will shift equilibrium toward the side
with more moles of substances.
iii. Answer these questions.
1. 2SO2 (g) + O2(g)  2SO3 (g) + heat
a. Which way will equilibrium shift if the temperature
increases?
b. Which way will equilibrium shift if the pressure
increases?
c. What will happen to the equilibrium if the
concentration of SO2 decreases?
d. What will happen to the equilibrium if the
concentration of O2 increases?
2. N2 (g) + O2 (g) + heat  2NO (g)
a. Which way will equilibrium shift if the pressure
decreases?
b. Which way will equilibrium shift if the concentration
of N2 decreases?
c. What will happen to the equilibrium if the
concentration of O2 increases?
d. What will happen to the equilibrium if the
temperature decreases?
18. Bonding
a. Electrons are the subatomic particle involved in all bonding between atoms.
b. 3 Types of Bonds
i. ionic bonds
1. occur between a metal and a nonmetal
2. transfer electrons from the metal to the nonmetal
3. strong bond with high melting points, form crystals
ii. covalent bonds
1. occurs between 2 nonmetals
2. share electrons a
a. pure covalent – equal sharing
b. polar covalent – unequal sharing
3. strong bond with lower melting points, can be liquids, solids, or
gases at room temperature
iii. metallic bonds
1. occurs between 2 metals
2. sea of delocalized electrons free to move across the alloy.
19. Atomic Structure
a. Atomic number (Z) = # of protons in the nucleus (also = # electrons in the atom)
b. Mass number (A) = # protons + # neutrons in the nucleus.
c.
Isotopic Symbol
A X
Z
d. Fill in the chart with the appropriate information.
Isotopic
Atomic #
Mass #
# protons
# electrons
Symbol
3
1H
11
24
26
35
17
# neutrons
30
20. Electrons and Light
Electrons can absorb energy and move from the ground state (lowest, stable state) to an
excited state where they are unstable!
a. Electron configurations- shows the energy levels and sublevels filled by the
electrons in an atom.
i. Each sublevel has an associated shape
1. s sublevels are spheres
2. p sublevels resemble infinity signs, or 2 teardrops touching at the
pointed ends
3. d & f sublevel shapes are more complicated.
s = 2e-
7s
7p
6s
6p
6d
5s
5p
5d
5f
d = 10e-
4s
4p
4d
4f
d = 14e-
3s
3p
3d
2s
2p
p = 6e-
1s
•nucleus
b. Write electron configurations for the following elements:
i. C
ii. Mg
iii. Cl
iv. Fe
v. As
21. Periodic Table
a. Arranged in columns and rows of increasing atomic number
b. There are
vertical columns, called groups.
c. There are
horizontal rows, called periods.
d. Label the periodic table below correctly with the following:
Alkali metals
alkaline earth metals
transition metals
Metalloids
halogens
noble gases
Inner transition metals
e. Periodic trends
i. Atomic and Ionic Radius
1. Down a group – radius increases
2. Across a period – radius decreases
ii. Ionization energy and Electronegativity
1. Down a group – trend decreases
2. Across a period – trend increases
iii. Order the elements: Pd, Sb, and Y
1. Increasing atomic radius
2. Decreasing ionization energy
3. Increasing electronegativity
4. Decreasing ionic radius
iv. Order the elements: Ga, B, and In
1. Increasing atomic radius
2. Decreasing ionization energy
3. Increasing electronegativity
4. Decreasing ionic radius
22. Reaction Rates – the pace that a reaction occurs
a. Several factors can effect the rate of a reaction.
i. Temperature – the higher the temperature, the
the reaction because there are more / less collisions.
ii. Concentration - the higher the concentration, the
iii. The reaction because there are more / less particles to collide and
react.
iv. Surface area – The higher the surface area, the more / less surface
there is for collisions.
v. Catalysts – are substances that speed up / slow down the reaction
by increasing / decreasing the amount of activation energy needed
for the reaction.
b. Draw and label potential energy diagrams for an endothermic and an
exothermic reaction.
c. Define entropy.
23. Solutions
a.
b.
i. unsaturated solution – contains less than the maximum amount of solute
at a specific temperature
ii. saturated solution – contains the maximum amount of solute at a specific
temperature
iii. supersaturated solution – contains more than the maximum amount of
solute at a specific temperature, very unstable
iv. Ways to increase the solubility of a substance
1. Increase the surface area – more surface of solute available to
dissolve
2. Agitation (Stirring) – brings solute and solvent into contact more
often so dissolving can occur much faster
3. Temperature
a. For solids – increasing temperature will increase how
fast a substance dissolve (think about how sweet tea is
made)
b. For gases – decreasing temperature will increase the
amount of gas that can dissolve (keeping the CO2
bubbles in your soda)
c. Measuring concentration – Molarity
i. Molarity is moles of solute/ liter of solution
ii. M = n/V
1. Solve these molarity problems
a. What is the molarity of a 0.30 L solution containing
0.50 moles of NaCl?
b. Calculate the molarity of 0.289 moles of FeCl3
dissolved in 120 mL of solution.
c. How many moles of NaCl are present in 600. mL of a
1.55 M NaCl solution?
d. What is the molarity in a 650.0mL solution
containing 63.3 g of NaCl?
e. How many grams of Ca(OH)2 are needed to produce
0.500L of a 1.66 M Ca(OH)2 solution?
24. Acids and Bases
a.
b. Amphoteric subtstance can act as
c. Determine the Bronsted-Lowry acid (A), Bronsted-Lowry base (B),
conjugate acid (CA) and conjugate base (CB) in the equations below.
i. HCO31- + H2O  H2CO3 + OH1ii. NH41+ + Br1-  NH3 + HBr
d. pH
pH = -log[H+]
pOH = -log[OH-]
.
pH + pOH =14
[H+] = 10-pH
[OH-] = 10-pOH
[H+] × [OH-] = 1.00 × 10-14 M2
i. Solve these pH problems.
1. What is the pH of a solution of HNO3 with a concentration of
4.00 × 10-3 M?
2. What is the pOH of a solution of Ca(OH)2 with a
concentration of 4.00 × 10-9 M?
3. What is the pH of the solution in #2.
4. Calculate the [H+] of a solution of HBr with a pH of 3.45.
5. Calculate the [H+] of a solution of NaOH with a pOH of 9.99.
e. Neutralization Problems
MAVA = MBVB
i. What volume of 0.456 M LiOH is needed to neutralize 150.0 mL of
0.234 M H2SO4?