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Chem 2202
Chapter 1: ALL the following is found in the TEXTBOOK. ALL notes follow the chapters for Chem 2202.
For the Midterm it is Ch1 to 4 and Ch 7. The book can be followed with these notes by turning pages in
the text. Remember the only difference between ordinary and extra ordinary is that little extra. These
notes summarize the chapters.
Matter has mass and takes up space. Einstein told us there was only Matter and Energy in the inverse,
so by studying matter, we are ½ way to understanding the universe!
There are two main ways of classifying matter.
1) Matter is classified first by its physical state as a solid, liquid, or gas.




Gas
low density
easy to
expand/compress
fills container




Liquid
high density
hard to compress
takes shape of
container



Solid
high density
hard to expand/compress takes
shape of container
Secondly matter is classed as an elements or compounds (ionic or molecular).
The Elements of the Periodic Table:
Periods are horizontal (across) the table. They correspond to the energy levels in Bohr Diagrams.
Basically, count the boxes across (each element) and that will be the electrons in that energy level.
Groups (or families) are columns on the table, They have similar chemical properties. Like Human
families, that have same last names and same “look” and way of acting, so to do chemical families.
FAMILIES : are Blocks are regions on the table.
1)
alkali metals (Group IA, first column )


soft, extremely reactive metals, to form hydrogen gas
form +1 ions
2) alkaline earth metals (Group IIA, second column):
 soft, reactive metals
 form +2 ions
3) halogens (Group VIIA, next-to-last column):



poisonous and extremely reactive nonmetals
all form -1 ions
4) noble gases ( last column)
 all are monatomic gases and inert which means unreactive
Other Important blocks:
1) metals, nonmetals, and metalloids (semimetals)
 metallic properties
 luster (shine)
 malleability: can be hammered into thin sheets
 ductility: can be drawn into wire
 conduct heat and electricity well
Non-metals, have the opposite properties.
Metalloids have combinations of both. Two metalloids of note are Carbon and Silicon
Mixtures (Characteristics)
- percentage composition varies from sample to sample
- components are chemically different and retain properties in a mixture
2 Types of Mixtures
- heterogeneous mixtures where components not uniformly mixed
- homogeneous mixture where the components uniformly mixed
Separating mixtures: a mixture's components have different properties. They can be separated from one
another based on their different densities, melting & boiling points, solubility, magnetism, or polarity.
Compounds:
Molecular (or covalent) Vs Ionic
Molecular compounds, are made up of 2 or more NON-METALS, valence electrons are SHARED.
examples: H2O, CO2, C6H12O6, NH3, CH4
Ionic Compounds are made of a metal TRASFERRING a valence electron to it’s non-metal partner(s)
They are made of positive ions (cations) and negative ions (anions). Cations combine with anions in just
the right numbers to give an electrically neutral compound. examples: NaCl, KBr, Na2S, MgBr2
Molecular compounds
Ionic compounds
Smallest things
molecules
ions
Elements present
Same side of line (fingers
together)
Fingers apart, opposite side of
diving line
Electrical conductivity
poor
good, when melted or dissolved
State at room
temperature
solid, liquid, or gas
ALL solid at room temp
Other names
covalent compounds
Salts
Atoms (History):
An atom is an extremely small particle of matter that retains its identity during chemical reactions.
An atom consists of a nucleus and electrons surrounding the nucleus. The nucleus, the core of the atom,
has the majority of the mass of the atom, and a positive charge.
An electron is a very light particle, which flies around the nucleus. It has a negative charge.
In an electrically neutral atom, the number of electrons equals the positive charge on the nucleus. In an
electrically neutral atom the positive charge of the protons, combined with the negative charge of the
electrons, would result in no charge.
A proton's mass, however, is a whopping 1836 times that of the electron. A neutron has a mass almost
identical to a proton's, but it has no electrical charge associated with it.
The atoms that are in a molecule are not just stuffed together without any order. The atoms are
chemically bonded to one another in order to form a definite arrangement. A structural formula is a
chemical formula which shows how the atoms are bonded to one another to form a molecule.
A molecule is a definite group of atoms that are chemically bonded together. They are tightly connected
by attractive forces.
A molecular formula is a chemical formula that gives the exact number of different types of atoms in a
molecule. Some simple molecular substances are carbon dioxide, CO2; ammonia, NH3; and water, H2O.
Chemical Properties :
Involves the chemical changing! Ex. determining the flammability of gasoline involves burning it,
producing carbon dioxide and water. The chemical changes into something else. Gasoline turns into
water ! Ok that is weird, no that is a reaction.
Physical properties such as melting/ boiling point or solubility with water, where the chemical does NOT
change H2O(s)  H2O(l)  H2O (g)
As you can see in a physical change , the CHEMICALS, do not change, water is melting or boiling.
Changing shape, sure, but that’s about it!
Diatomics: gaseous elements commonly occur as diatomic molecules (except for the noble gases),
memorize these as the halogens , oxygen, nitrogen , and hydrogen. The polyatomics are sulfur and
phosphorous.
Chapter 2 :
Moles, Molar Mass, and Average Atomic Number
The atomic number is the number of protons there are in the nucleus. Hydrogen's atomic number is 1.
Helium's atomic number is 2.
The mass number is the number of neutrons added to the number of protons.
The mass number of C 13 , is 13, clear enough, so this isotope of C has 7 neutrons:
# neutrons= (13, neutrons and protons)) – 6 (protons) = 7 Neutrons.
Elements are defined by the number of protons in an atom's nucleus. So, for example, an atom with 6
protons must be carbon and an atom with 92 protons must be uranium. If you change the protons, then
you change the element.
In addition to protons, the atoms of every element (except the simplest form of hydrogen) also contain
neutrons.
Isotopes occur when an element's atoms exist with different numbers of neutrons.
As a result of their having different numbers of neutrons, an element's isotopes differ in mass.
The stability of each atom's nucleus depends on the ratio of protons to neutrons. Many isotopes have a
ratio of protons to neutrons that renders them unstable and, as a result, they are radioactive.
Consider Carbon again, which exists naturally with 6, 7 or 8 neutrons. These carbon isotopes have
atomic masses of 12, 13 and 14. (atomic mass = mass protons + mass neutrons).
The isotopes are called carbon-12, carbon-13 and carbon-14. Alternatively, they may be written 12C, 13C
and 14C. Carbon-12 and carbon-13 are stable. Carbon-14 is unstable, decaying with a half-life of about
5,700 years. It is produced in earth's atmosphere by cosmic ray bombardment of nitrogen-14. This
knowledge enables us to Radioactively Carbon date fossilized life forma, well beyond 10, 000 years, so
science can prove life did NOT start 10,000 years ago as creationists propose.
This chapter is about average molar masses from relative abundances in nature. See pg 45 #1,2,3 and
the sample problem on the same page.
That page will explain the following, which is how to find average Atomic Mass or Molar mass from the
average (weighted / percentages) of the isotopes that make up the element in nature.
Ex: Chlorine-35 and Chlorine-37 are both isotopes of chlorine. All elements have different isotopes.
Chlorine consists of roughly 77.300 % Cl-35 and roughly 22.700 % Cl-37. What is the relative atomic
mass of chlorine?
We take an average of the two figures which is weighted towards the figure 35 to take into account the
fact that Chlorine-35 is three times more common than Chlorine-37. The relative atomic mass of chlorine
is usually quoted as 35.454 g / mol
Or
(35CL) 0.77300* 35 = 27.055g
(37CL) 0.22700 * 37 = 8.399g
= 35.454 g /mol
Mole:
Molar masses (M) of chemical compounds are equal to the sums of the molar masses of all the atoms in
one molecule of that compound.
If we have a chemical compound like NaCl, the molar mass will be equal to the molar mass of one atom
of sodium plus the molar mass of one atom of chlorine. If we write this as a calculation, it looks like this:
(1 atom x 23 grams/mole Na) + (1 atom x 35.5 grams/mole Cl) = 58.5 grams/mole NaCl
For other compounds, this might get a little bit more complicated. For example, take the example of zinc
nitrate, or Zn(NO3)2. In this compound, we have one atom of zinc, two atoms of nitrogen (one atom inside
the brackets multiplied by the subscript two) and six atoms of oxygen (three atoms in the brackets
multiplied by the subscript two). The molar mass of zinc nitrate will be equal to (1 atom x 65 grams/mole
of zinc) + (two atoms x 14 grams/mole of nitrogen) + (six atoms x 16 grams/mole of oxygen) = 189
grams/mole of zinc nitrate.
The mole concept applies to all kinds of particles: atoms, molecules, ions, formula units etc. The amount
of substance is measured in units of moles. The approximate value of Avogadro's constant, 6.02 x 10 23
To convert between moles and molecules you need to remember that one mole of any substance
contains 6.02 x 1023 particles (e.g., atoms or molecules).
When going from moles to molecules you multiply by 6.02 x 1023.
When going from molecules to moles you divide by 6.02 x 1023
__________________________________
Example: How many molecules of ethanol are in 89.0 g of ethanol?
Solution: 89.0g / 46.08 g/mol ethanol x 6.02 x 10
23
= 1.2 X 10 24 molecules.
To find molecules (or F.U. (formula units)) = n X 6.02 x 1023 ...Remember. To get molecules or (F.U’s)
FIRST GET MOLES ( see pg 53: #12,13,14,15).
Finding Moles:
In order to be able to ride the MOLE bus, we have to get molar masses to find moles. Adding molar
masses should be easy. Ex water = 18.02 g/mol, you should speed through this.
The general equation:
moles = mass / molar mass (n = m / M) .
So, to combine all so far , see the chart below:
Let us now continue with more solutions to the example problems above. Here are the same two
problems as before, but with gram replacing mole:
1. 0.450 gram of Fe contains how many atoms?
2. 0.200 gram of H2O contains how many molecules?
Look at the solution steps and you'll see we have to go from grams (on the left) across to the right
through moles and then to how many. So, for the first one it would be like this:
Step One: 0.450 g divided by 55.85 g/mol = 0.00806 mol
Step Two: 0.00806 mol x 6.022 x 1023 atoms/mol
and for the second, we have:
Step One: 0.200 g divided by 18.0 g/mol = 0.0111 mol
Step Two: 0.0111 mol x 6.022 x 1023 molecules/mol
Gases: Moles of a Gas @ STP:
The volume of any gas at STP (standard temperature and pressure, which is 0 ºC and standard pressure
is 1 atm.) is 22.4 L (equivalent is one mole = 22.4 L). We can use this relationship to convert between
moles of a gas at STP and volume of a gas. STP = 22.4 L in one mole of any gas. Basically GET MOLE
and MULTIPLY by 22.4 to get gas litres. Page 70 and 71 sample problem and pg 73 # 38 to 43.
Ex: The problem below shows how you would convert 5.50 moles of a gas at STP into volume.
how you would convert volume into moles:
Here's
V = n (22.4 L/mol)
= 5.50 mol X 22.4 L/mol
= 122.0 L
Chapter 3:
Percent Composition and Empirical Formula
Percent Composition:
Percent composition is the percent by mass of each element present in a compound.
Water, H2O, is the first example. One mole of water is 18.02 grams.
In that compound, there are two moles of H atoms and 2 x 1.01 = 2.02 grams. That's how many grams of
hydrogen are present in one mole of water.
There is also one mole of oxygen atoms weighing 16.00 grams in the mole of water.
To get the percentage of hydrogen, divide the 2.02 by 18.02 and multiply by 100, giving 11.19%.
For oxygen it is 16.00 ÷ 18.015 = 88.81%.
Empirical Formula: MOLES !!! and Foomp
It is the formula of a compound expressed as the smallest possible whole-number mole ratio of elements
in a compound. For example, CH3COOH has two carbons, four hydrogens and two oxygens. So we
could write the formula like this: C2H4O2 and so it reduces to CH2O. Here are the four formulas being
used as examples:
Molecular Formula Empirical Formula
H2O
CH3COOH
CH2O
C6H12O6
H2O
CH2O
CH2O
CH2O
Notice two things:
1. The molecular formula and the empirical formula can be identical.
2. You scale up from the empirical formula to the molecular formula by a whole number factor.
Steps to Calculating E.F.
Step 1: calculate moles, Step2 , foomp! ( divide all moles by lowest number of moles), if answers not
whole, multiply till they are.
Here's an example of how it works. A compound consists of 72.2% magnesium and 27.8% nitrogen by
mass. What is the empirical formula?
(1) Percent to mass: Assume 100 g of the substance, then 72.2% becomes 72.2 g Mg and 27.8 g N
(2) Mass to moles:
for Mg: 72.2 g Mg x (1 mol Mg/24.3 g Mg) = 2.97 mol Mg
for N: 27.8 g N x (1 mol N/14.0 g N) = 1.99 mol N
(3) Divide by small:
for Mg: 2.97 mol / l.99 mol = 1.49
for N: 1.99 mol / l.99 mol = 1.00
(4) Multiply 'til whole:
for Mg: 2 x 1.49 = 2.98 (i.e., 3)
for N: 2 x 1.00 = 2.00
and the formula of the compound is Mg3N2.
Chapter 4: Stoichiometry
The word stoichiometry derives from two Greek words: stoicheion (meaning "element") and metron
(meaning "measure").
Use this flow chart, or the one given in class, and you can solve all stoic prob’s.
Stoichiometry and Limiting Reagents: What can I say but the mole bus, be a mole to ride. To become a
mole you DIVIDE to RIDE. Whether you are give n your info in grams, litres of gas, or molecules, you
DIVIDE to ride (except if info is given in concentrations then n = c * v).
Example:
1. How many moles of O2 are produced when 7.5 moles of KClO3 decompose according to the
following equation?
2KClO3 --> 2KCl + 3O2
Limiting Reactant:
The first step in a limiting reactant question is to determine the limiting reactant. To do this, for all
reactants calculate the amount of product that would be produced if all of the reactant was used and
there was excess of all the other reactants. The reactant that has the least product is the limiting
reactant. The amount of products is due to the amount of the limiting reactant.
Limiting reagent problems explained on pg 130 (sample prob).Work through pg 131 #’s 23 to 26
An example:
Zn(s) + 2HCl(aq) --> ZnCl2(aq) + H2(g)
Which is the limiting reactant and how much H2 is produced if there is 0.30 mol of Zn and 0.52 moles of
HCl?
Step one:
Since HCl is the limiting reactant, the number of moles of H2 produced is .26 moles.
Chapter 7:
SOLUTIONS
Molarity
Molar Concentration or Molarity is defined as the moles of solute dissolved in one liter of solution.
Molarity =
moles of solute
liter of solution
Example:
0.0678 g of NaCl is placed in a 25.0 ml flask full of water. When the NaCl dissolves, what is the molarity
of the solution?
Moles (n) = m/M
C = n/v
0.0678g / 58.44g/mol = 0.00116 moles
0.00116 moles/0.0250 L = 0.0464 M NaCl
SOLUTIONS & THEIR CONCENTRATIONS
Solution:
a homogeneous mixture composed of a solute and a solvent
Solute:
the substance being dissolved
Solvent:
the substance that has other substances dissolved in it
NOTE:
solutes and solvents can be either gases, liquids or solids
Examples:
1.
Air
- a solution of gases (nitrogen, oxygen, carbon dioxide, etc.)
2.
Vinegar
- a solution of liquids (5% acetic acid + 95% water)
3.
Metal alloy
- a solution of solids (eg. Sterling silver is a mixt of silver & copper)
4.
NaCl(aq)
- a solid ionic compound dissolved in a liquid (water)
5.
NH3(aq)
- a gaseous molecular compound dissolved in a liquid (water)
IMPORTANCE OF SOLUTIONS
Solutions are important to non-living systems:

Eg: in order for chemical reactions to occur, reactants must often be dissolved in water
Solutions are important to living systems

Eg.
Blood plasma is mostly water with substances such as O2 and CO2
dissolved in it
B.
CLASSIFYING SOLUTIONS
1.
a)
Electrolytes and Non-electrolytes:
Electrolytes: conduct electricity
 an electric current involves the movement of charged particles, therefore solutions
containing ions will conduct electricity
 include acids, bases and neutral ionic compounds in solution
b)
Non-electrolytes: do not conduct electricity

include neutral molecular compounds in solution
2.
Concentrated and Diluted Solutions:
a)
Concentrated solution: has a relatively large quantity of solute in a given volume
b)
Diluted solution: has a relatively small quantity of solute in a given volume
Solubility:

the maximum amount of a solute that dissolves in a given quantity of solvent at a certain
temperature
 eg. the solubility of NaCl in water at 200C is 36 g per 100 mL

The ability of a solvent to dissolve a solute depends on the attractive forces between the
particles
 All ionic substances dissolve in water to varying degrees:
Insoluble:
(<0.01 mol/L)
Soluble:
High Solubility ( ≥0.1 mol/L)
Low Solubility (<0.1 mol/L but > 0.01 mol/L)
a)
Saturated solution



A solution that contains the maximum amount of solute at a given temperature
A saturated solution is formed when no more solute will dissolve in a solution and excess
solute is present
eg. 100 mL of a saturated solution of NaCl in H2O at 200C contains
36g of NaCl. The solution is saturated only with respect to NaCl.
Dynamic Equilibrium:


the balance that exists when 2 opposing processes occur at the same time
In a saturated solution, although no changes appear to occur, dissolving and re-crystallizing
continues. These 2 processes continue at the same rate (speed), therefore no net change in
concentration occurs and the solution is in “dynamic equilibrium”.
The rate of dissolving = the rate of recrystallization
b)
Supersaturated solution


Forms when a saturated solution is heated so that the excess solute present dissolves. If the
solution cools slowly the excess solute does not immediately crystallize.
A supersaturated solution holds a greater amount of solute than the equilibrium amount. It is in
an unstable state that persists as long as it is undisturbed.
c) Unsaturated solution
 A solution containing less than the maximum amount of solute
Low Solubility:
same state as pure form
Type of
Substance
Includes
Ionic
see solubility table
Molecular
Hydrocarbons
High Solubility:
Examples
Electrolytic Ability
AgCl(s)
Weak Electrolyte
CH4(g), C3H8(g), C6H6(l)
aqueous ions or molecules
Type of
Substance
Includes
Ionic
See solubility table
Behavior
Examples
Dissociate NaCl(s) Na+(aq)+Cl(aq)
Strong Acid
Non-electrolyte
HCl(g)  H+(aq) + Cl-
Electrolytic
Ability
Strong
Electrolyte
HCl, HBr, HI,
H2SO4, HNO3,
HClO4
Ionize
Weak Acid
all other acids
Disperse
HNO2(l)  HNO2(aq)
Weak
Electrolyte
Molecular
CH3OH(l),
Disperse
H2O2(l)  H2O2(aq)
Non-electrolyte
(aq)
Strong
Electrolyte
C2H5OH(l) H2O2(l),
C12H22O11(s)
IONIC EQUATIONS
1.
Nonionic Equations
 Elements and compounds are written in their molecular or formula unit forms
 Eg:
2 AgNO3(aq) + BaCl2(aq) → 2 AgCl(s) + Ba(NO3)2(aq)
2.
Total Ionic Equations
 Elements and compounds: written in the form in which they mainly occur:
 Electrolytes are written in ion form
 Nonelectrolytes, precipitates and gases are written in their molecular or formula unit forms
 Eg: 2Ag+(aq) + 2NO3-(aq) + Ba2+(aq) + 2Cl-(aq) → 2 AgCl(s) + Ba2+(aq) + 2NO3- (aq)
Net Ionic Equations
 Only those molecules, formula units or ions that have changed in form are included in the
equation = the predominant reacting species
 Ions or molecules that do not change, are omitted = spectator species
3.
Eg: 2Ag+(aq) + 2Cl-(aq) → 2 AgCl(s) which reduces to:
Ag+(aq) + Cl-(aq) → AgCl(s)
Summary of writing net ionic equations:
 Species that do not change or take part in the reaction are not shown.
 The equations must be balanced, both in atoms and in charge, using the lowest whole number ratio.
 Species are written in the form in which they actually exist.
Acids: Acid solutions are written in molecular form except for the 6 strong acids which
are written in ionic form:
Hydrochloric acid:
HCl(aq) → H+(aq) + Cl-(aq)
Hydrobromic acid: HBr(aq) → H+(aq) + Br-(aq)
Hydroiodic acid:
HI(aq) → H+(aq) + I-(aq)
Perchloric acid:
HClO4(aq) → H+(aq) + ClO4-(aq)
Nitric acid:
HNO3(aq) → H+(aq) + NO3-(aq)
Sulphuric acid:
H2SO4(aq) → H+(aq) + HSO4-(aq)
NOTE:
Only aqueous ionic compounds and strong acids are written as separate aqueous ions (in total
ionic or net ionic equations). All other species remain intact.
Expressing Concentration:
1. As a Percent = “parts per hundred”
ppm of solute = mass solute
X 106
mass of solution
a)
Volume/Volume Percent or percent (v/v)



Used when both solute and solvent are liquids
The volume of a solute dissolved in a volume of solution
Eg: white vinegar contains 5% acetic acid, by volume
 5 mL of pure acetic acid in 100 mL of solution
Volume/volume percent =
volume of solute (mL)
x
100
Volume of solution (mL)
b)
Mass/Volume Percent or percent (m/v)



The mass of solute dissolved in a volume of a solution
Equal to the number of grams of solute per 100 mL of solution
Eg: a solution containing 2.0 g of glucose in 100 mL of solution is 2.0% (m/v)
Mass/volume percent =
mass of solute (g)
x
100
Volume of solution (mL)
2.
Parts per million (ppm) & parts per billion (ppb)

Used when the amount of solute is very small, such as with trace impurities in water. Eg.
Water in a public swimming pool contains 1.0 ppm of chlorine
ppb of solute = mass solute
mass of solution
3.
Molar Concentration, C


The moles of solute in one liter of solution
Most common way of expressing concentration
C = moles of solute
liters of solution
C =n/V
X 109


Also called concentration (mol/L) or Molarity (M)
Eg: concentrated nitric acid has a concentration of 15.4 mol/L or 15.4 M
DILUTION
Most solutions used in laboratories are bought or are prepared in concentrated form. These stock
solutions can then be used to prepare more dilute solutions as needed.

When a solution is diluted, only the amount of solvent is increased. The number of moles of solute in
the initial (concentrated) solution is equal to the number of moles of solute in the final (diluted)
solution.
Thus:
ni =nf
Since
n = Cv
i = initial
then
f = final
Civi = Cfvf

Since the amount of solvent is increased, but not the solute, the molar concentration of the dilute
solution is less than the molar concentration of the concentrated solution.

Thus, Ci will always be greater than Cf and vf will always be greater than vi.
Chapter 8:
SOLUTION STOICHIOMETRY
General Steps:
1.
Write a balanced chemical equation, identify the required and given substances, and record the
required and given data.
2.
If not given moles, convert the given quantity (mass, or solution volume and concentration) to
number of moles using one of the formulas below:
3.
n=m
n = Cv
M
Use the mole ratio to calculate the moles of the required substance from the moles of the given
substance. (Every problem will involve this calculation).
nR = nG x ( R/G ) or BUS!
where nR = moles of required substance
nG = moles of given substance
R/G = mole ratio using the coefficients from the balanced equation
ie. moles of required/ moles of given
4.
If necessary, convert the moles of the required substance to the required quantity (mass, gas
volume, solution volume or concentration) using one of the formulas below:
m = nM
v = nV
v = n/C
C = n/V
Solutions and Solutions Stoichiometry: Read pg 237 to 241 for several TESTED definitions. Then read
pages 266 ( c = n/v) to pg 273.
Solution Stoich . Solubility chart pg 280 - 286 then Stoichiometry pg 299 to pg 306.