Transcript Data: I am writing out the question and underlining it.

Not the
real Mr. Cooper
“Good morning, and welcome to
introduction to chemistry.”
Info
• Class: Chemistry
• Instructor: Mr. Cooper
• Office: A112 - I’m pretty much in my
classroom before and after school
• E-mail: [email protected]
• WebPage: lsw.lps.org click on teachers
and find my name.
Class Format
• 1. Daily Quizzes (D.Q.): Each day will start
with a quiz over the previous days material.
Clear your desk, have out a piece of paper
and be ready at the beginning of class.
Recommendation: Write out the question and
answer and keep a running list of the quizzes
on the same sheet of paper. At the end of the
term you will have created a review sheet. I
will not be giving you a review sheet at the
end of the term. Quizzes will not be picked up
but your score will be recorded on your selfevaluation sheet.
Class Format
• 2. Lecture: Lecture notes are posted on the
web. You can print them off before class. This
allows you to listen and formulate questions
and take your own notes rather than just
copying down my outline. Printing them
ahead of time is not required; but if all of the
class has notes ahead of time, we can spend
more time on lab work, book work and
individual help rather than copying notes.
Class Format
• 3. Laboratory:. One typed lab report of your choice
will be submitted per unit. You choose the lab you
want to do the report on. It is suggested that you
choose one at the beginning of the unit so that you
can get it done ahead of time. It will be due day of
test. Leave 5 minutes at the end of each period for
clean up. Labs are not expected to be homework.
If you work diligently in class, you should get them
done. If not, the labs are on-line for you to finish at
home. Lab may be replaced by worksheets, group
work, a video, or demonstrations
Class Format
• 4. If there is time left at the end of class
you are expected to be working on book
problems, which are assigned on a daily
basis. See unit outline.
• This is also time for you to get individual
help.
Grading
• Lab Books and Lab Quizzes: A simple
one-subject notebook is recommended.
The day before each unit test will be a
unit lab quiz. It is open lab book. See
ppt notes for specifics on lab book
expectations.
Grading
• Quizzes and tests will be announced in
advance. Tests and quizzes will be closed
note and closed book. Exception: lab quizzes
are open lab book. Your textbook is your first
resource; so read it!!! All materials for this
course are based off of your text book. Also,
calculators will be allowed on tests and
quizzes. It is your responsibility to provide a
calculator.
Grading
• Tests and quizzes are m.c., short
answer, problem solving, make you
think exams; not memorizing
exams(although you will need to have
some things memorized.)
Grading
• Test retakes will be offered during plc
time and Sat. school in the media
center. You must sign up for a retake so
I can get it to the media center. It is your
responsibility to rearrange your
schedule if you wish to take advantage
of retesting.
Grading
•
•
•
•
•
The grading scale is as follows:
A= 90.0-100
B+= 85.0-89.9
B= 80.0-84.9
C+=75.0-79.9
C= 70.0-74.9
D+= 65.0-69.9
D= 60.0-64.9
F= Below 60.0
Misc.
• Any assignments or test missed for
truancy results in 60.0% of earned
grade. This is district policy. Missed labs
need to be done ASAP. Late work is
accepted 1 day late for 1/2 credit.
Misc
• Tardies - Building policy is followed.
• Cell Phones - If I see it or hear it; I can take it
for the rest of the day, turn it into security, or
write a referral.
• iPods - iPods are not to be used during
instructional time or lab time. You may use
them during individual work time at your desk.
I reserve the right to revoke privileges.
Student Expectations
• Do your job as a student which means:
• 1. Bring all needed materials to class.
• ex) books, notebooks, writing utensils, brain, good
attitude, etc.
• 2. Respect each students right to learn and their
property.
• 3. Listen carefully and follow instructions given.
• 4. READ, STUDY, PAY ATTENTION TO DETAIL
• 5. No food or drink.
• 6. Use class time to work.
• 7. ASK QUESTIONS in class or see me after school
for help.
Misc.
• “I do not feel obliged to believe that the same
God who has endowed us with sense,
reason, and intellect has intended us to forgo
their use.” - Galileo Galilei (astronomer and physicist)
• Remember, I am working hard for you. I
expect that you will work hard for me.
• I find it offense when at the end of the term
you are begging me to round or expecting me
to do you some extra credit favor when you
didn’t give me your best to begin with.
Mr. Cooper
Equipment Use Review
• What lab equipment is used for handling a hot
beaker?
• What lab equipment would be used to hold a piece of
metal in a flame?
• What piece of lab equipment is used to measure
volume?
• A BEAKER OR FLASK IS NEVER USED AS A
MEASUREMENT DEVICE!!!
Tirrill (Bunsen)
Burner
How the parts work.
Turning the barrel
Controls type of flame
(orange or blue) by
opening and closing
the air vent.
Always use blue flame
(open vent);
however, vent
should not be wide
open for initial
igniting.
How the parts work.
Gas Flow Control
• Controls the height
of the flame through
controlling the
amount of gas
flowing.
• Use appropriate
flame height. NO
TORCHES.
Operation of the Tirrill (Bunsen)
Burner
• Hook the hose to the gas
inlet and gas jet
• Place spark ignitor next to
top of barrel
• Turn on gas and ignite with
sparker
• Make barrel and gas flow
control adjustments for
proper flame
Trouble Shooting
• You should be able
to hear the gas
flowing. If not:
• Check if gas flow
control valve is open
• Check if jet valve is
clogged. If so see
your teacher.
Troubleshooting
• Gas attempts to light
but goes out.
Possible cause is:
• Air vent is too far
open. Turn the
barrel down.
Formatting a Lab Report
• Title: The word “title” is written and underlined; followed then by the
name of the lab.
• Purpose: The word “purpose” is written and underlined; followed by the
purpose of the lab.
• Procedure: Usually extremely detailed. You can summarize. Just a
couple of sentences is fine. Procedure questions will be on quizzes.
• Data: The word “data” is written and underlined. For this section you
will either be filling out charts or questions will be asked to help you
gather data. Write out the question, underline it, leave a space, then
answer the question.
Formatting a Lab Report
• Conclusion: The word “conclusion” is written and underlined. For
this section you will be asked questions. Write out the question,
underline it, leave a space, then answer the question.
• Application: Where is this concept used in the real world or in
the scientific community? How does this affect your life or why is
this important to have this knowledge for society or other real
world application or future predicted use?
– Minimum 3 sentences and Maximum 5 sentences
– Must have one source to accompany this section. If you use
a website please make sure you do not have a typo in the
address.
– No opinions. I am not your source nor are you a source. This
is a research component. Do some research and quote your
source or than your text. Do not use any “I” statements.
Sample lab report
Title: Place title here.
Purpose: Place the purpose here.
Procedure: A couple of general sentences summarizing lab steps.
Data:
1.
I am writing out the question and underlining it.
A space was left and question 1 is answered.
2. Another question is written out and underlined.
A space was left and question 2 was answered.
Sample lab report
Conclusion:
1.
I am writing out the question and underlining it.
A space was left and question 1 is answered.
2. Another question is written out and underlined.
Do you see a pattern here?
Application: Use or Application
Source: WWW.SCIENCERULES.COM
Keep answers clear and concise. Length is not
important. I care about content and good
communication.
Bad Student Example - Very Bad
• Application:
• After doing this lab, I sat and wondered how I would apply what I learned to
something that expands outside of our classroom. Thinking about the candle burning
sent me into deep contemplation. And then, out of complete randomness, I started
thinking about our environment and the things that we burn which pollute it. I then
thought of where all the statistics we hear about come from, and how the claims are
substantiated. How do scientists know exactly what percent our ozone layer has
deteriorated, and what percent of our atmosphere is made up harmful pollutants?
Well when fossil fuels are burned, or maybe even things like wood or who knows,
scientists most likely calculate the molecules given off so they can come up with
these statistics. Well maybe they deal with moles or liters of gas at STP, who knows,
but I’m sure somewhere in there scientists will have to convert from moles to
molecules, or grams to moles, or grams to molecules, and in a sense that is what we
have done in this lab. We found out how many moles of wax were burned over 3
minutes, and if we know what wax is made of then we can figure out what exactly
was released into the atmosphere.
• Source: My brain .. no seriously .. my brain.
Acronyms
•
•
•
•
•
•
•
•
KISS
Keep It Simple Stupid
SOP
Standard Operating Procedure
HUA
Heard Understood Acknowledged
WAG
Wild Ass Guess
Metric Conversions
grams (g) is used for mass (weight)
liters (l) is used for volume
meters (m) is used for distance
kilo
k
hecto
h
deka
dk
SI
g
l
m
deci centi milli
d
c
m
Metric Conversions
kilo
hecto
deka
SI
deci centi milli
• We will use a problem solving process
called dimensional analysis (tracks).
• Example 25.0 cg = _______ g
• 25.0 cg 1 g
100 cg
= 0.250 g
Example 0.351 hl = _______ ml
0.351 hl 100,000 ml
1 hl
= 35,100 ml
Metric Conversions
kilo
hecto
deka
SI
deci centi milli
• Your turns - convert the following:
• 15.72 g = ______ mg 15.72 g = _____ kg
Density = m  v
D = m/v
• intensive property
– density is the same no matter size
• 50 grams of gold has the same density
as 150 grams of gold.
• Important density to remember
– water is 1.0 g/ml at 4 oC
– 1 cm3 = 1 ml
Density Problem
• A substance has a volume of 1.74 ml
and a mass of 20.0 grams. What is the
density?
• What is the substance? Use page 96 of
your text.
Density Problem
• One more to test your algebra:
• What is the volume of ice in a container
if the density is 0.920 g/ml and the mass
is 58.39 g?
• 0.920 g/ml = 58.39 g
V
V = 63.5 ml
Properties of Matter
Definitions
• Matter - anything that has mass and
takes up space
• Mass - amount of matter an object
contains
• Substance (pure) - matter that has a
uniform composition
– Ex. Sugar - C12H22O11
– Lemonade is not a pure substance
Properties of Matter
States of Matter (Solid)
• Definite shape
• Definite volume
• Is incompressible (atom or molecules
can not be pushed closer together)
• Examples – coal, sugar, ice, etc
Properties of Matter
States of Matter (Liquid)
•
•
•
•
•
Matter that flows
Has a fixed volume
Takes the shape of its container
Incompressible
Examples
– Water, milk, blood, etc
Properties of Matter
States of Matter (Gas)
• Matter that takes the shape and volume
of its container
• Easily compressed
• Examples
– Oxygen, nitrogen, helium, etc
States of Matter Video
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Properties of Matter
Physical Property
•
•
•
•
•
•
•
•
•
•
•
•
•
An observed condition of the substance
Physical properties help identify substances
Examples include:
Color
Solubility
Odor
Density
Hardness
Melting point (m.p.)
Boiling point (b.p.)
Malleability
Ductility
Luster
Properties of Matter
Physical Change
• A change which alters a given material
without changing its composition
• Nothing new is made
• Example
– Ice melting - new state of matter but
substance is still H2O
– Vapor - a substance that is in a gaseous
state but liquid at room temp
Change of State - a physical change
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Properties of Matter
Chemical Property
• The ability or inability of a substance to
rearrange its atoms.
• Example
– Gasoline has the ability to react violently
with oxygen
Physical and Chemical Properties
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Properties of Matter
Chemical Change
• The actual rearrangement of atoms
• Example
– The combustion of gasoline to make
carbon monoxide, carbon dioxide, carbon,
water (this produces a great amount of
energy)
Classifying a physical or
chemical change
• Ask yourself these questions:
• 1. Has something new been made?
– If yes than a chemical change occurred
– Indicators - color change, formation of precipitate,
absorption or release of energy, formation of a gas
• 2. What does it take to get back to the original
form?
– If a physical process can revert it back than the
change was physical.
A chemical change
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Classify the following as a
physical or chemical change
•
•
•
•
•
•
•
•
A sidewalk cracking
Blood clotting
Getting a tan
Making Kool-Aid
Making a hard boiled egg
Plastic melting in the sun
Autumn leaf colors
Digestion of food
• The ripening of a
banana
• Making ice cubes
• Milk curdling
• Turning on the
television
• Making toast
• Mowing the grass
• Paint fading
• Grey hair
Categorizing our environment
Classifying Matter
Pure Substances
Element
Compound
Mixtures
Heterogeneous
Homogeneous
aka - solution
Classifying Matter
Mixtures
• A physical blend of two or more
substances.
• Examples:
– Beef stew, air - mixture of gases
Classifying Matter
Mixture (Heterogeneous)
• Not uniform in composition
• One portion of the mixture is different
from the composition of another portion
• Example:
– Soil - sand, silt, clay, decayed material
Classifying matter
Mixture ( homogeneous)
• Completely uniform composition
• Components are evenly distributed
throughout the sample
• Example
– Alloys - mixture of metals (brass, steel)
• AKA - solution
– Example - ammonia, alloys, kool-aid
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Classifying Matter
Mixtures can be separated by physical means
• Examples
– A spoon can separate beef stew
– Sulfur and iron can be separated with a magnet
• Tap water is a mixture that can be separated by
distillation.
• Distillation - a separation techniques based on the
physical property of boiling points.
– Liquid is boiled to produce a vapor
– Then condensed to a liquid leaving impurities behind
Categorizing our environment
Classifying Matter
Pure Substances
Element
Compound
Mixtures
Heterogeneous
Examples
river water
milk
beef stew
Homogeneous
aka - solution
Examples
pop
steel
kool-aid
Classifying Matter
Elements
• Simplest form of matter
• Cannot be broken down into anything
else
• Building blocks for all other substances
• Examples
– Hydrogen, oxygen, carbon
Classifying Matter
Compounds
• Two or more elements combined
through a chemical bond
• Can only be separated into simpler
substances by chemical reactions
• Example
– Sugar - C12H22O11
Chemical and physical properties of compounds
are different from their constituent elements.
• Examples
• Sugar
– Carbon is black
– Hydrogen is a gas
– Oxygen is a gas
• Salt - NaCl
– Na (sodium) soft metal that explodes with water
– Cl (Chlorine) pale yellow-green poisonous gas
Classifying matter review
Remember
• Substance - all of one kind of matter
– Examples: element or compound
• Mixture - has more than one kind of
material
– Examples - two or more compounds or
elements that are mixed, not chemically
combined
Categorizing our environment
Classifying Matter
Pure Substances
Mixtures
Element
Compound
Heterogeneous
Examples
Carbon (C)
Gold (Au)
Neon (Ne)
Examples
Sodium Chloride (NaCl)
Sugar (C12H22O11)
Dihydrogen Monoxide (H20)
Examples
river water
milk
beef stew
Homogeneous
aka - solution
Examples
pop
steel
kool-aid
The anatomy of the periodic table
• Get out your periodic tables
• Know where the following are on your
periodic table (p.t)
• Group A (representative elements)
• Group B
• Metals
• Nonmetals
• Metalloids (Semimetals)
– Note - aluminum is not considered a metalloid
The anatomy of the periodic table
• Know where the following are on your periodic table
(p.t) continued
• Transition metals
• Inner transition metals
• Alkali metals
• Alkaline metals
• Halogens
• Noble gases
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Naming Compounds
Ionic
Molecular
i.e. covalent
Naming Molecular
Compounds
• Molecules are made up of nonmetals
• Prefixes are used to represent numbers of
atoms. See your text for prefixes
• Binary compounds end in -ide
• Examples
• Name? - Cl2O8 and OF2
• Formula for? - dinitrogen tetroxide
• Answers -
Naming Molecular
Compounds
• Your turn. Try these.
• Name or write the formula for:
– Boron trichloride
– Dinitrogen tetrahydride
– N2O5
PF5
S4N2
• Answers
CCl4
SO3
H2O
Take ten minutes and work a
few problems on the “Naming
covalent compounds” side of
your worksheet.
Ions
• An atom that carries a charge
• The charge on the ion is called the
Oxidation state or Oxidation Number
• Cation - positively charged atom
– Metals form cations
– CATions are PAWsitive
• Anion - negatively charged atom
– Nonmetals form anions
Naming Cations
Name the metal followed by the word
ion
• Example
– Na - sodium - neutral element
– Na1+- sodium ion - cation of the element
• Another example:
– Mg - magnesium
Mg2+ - Magnesium ion
Naming Anions
• Ending changes are used for Anions
• Elemental anions will end in -ide
• Example
– Cl2 - chlorine - neutral element
– Cl1- - chloride - anion of the element
• Another example
– O2 - oxygen
O2- Oxide
Writing Formulas for Binary Ionic
Compounds
• The periodic table tells you the charge for
group A (aka - the representative elements)
• Group 1A - 1+ Group 2A - 2+
• Group 3A - 3+ Group 4 - depends
• Group 5A - 3- Group 6A - 2• Group 7A - 1- Group 8A or (0) - does
not form ions
Naming
• Your turn:
– Name or write the symbol for the following:
•
•
•
•
Aluminum
Calcium Ion
Ga3+
K
Phosphide
Iodine
Nitrogen
Sulfide
Naming Binary Ionic Compounds
Name the metal then the nonmetal with
the ending changing to -ide
– The -ide tells the person it is a binary
compound and the anion portion.
• Examples: MgCl2
K2S
• Magnesium Chloride Potassium
Sulfide
Writing Formulas for Binary Ionic
Compounds
• All compounds are electrically neutral
• To write the formula, figure out how many
cations and anions are needed so that the
number of positives and negatives are equal.
Find the least common multiple to figure
out the total number of +’s and -’s. Then
divide by the charge to find out how many
of each atom is needed!
• If X1+ and Y2-, what would be the formula?
• X2Y - Charges total 2 +’s and 2 -’s
Writing Formulas for Binary Ionic
Compounds
• If X3+ and Y2-, what would be the
formula?
• X2Y3 - Charges total 6 +’s and 6 -’s
• Find the formula for the following pairs
of ions:
– Na1+ , P3-
• Answers:
Sr2+ , N3-
• Now:
– Finish side 1 of worksheet
– Work sections 1 - 4 on back of worksheet
– Work homework problems
Writing formulas for multivalent ionic
compounds
• Transition metals have the ability to
form more than one cation
• Therefore, a roman numeral is placed in
the name to signify the charge on the
cation
• Example:
– Iron (III) Chloride
• Write the formula?
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Writing formulas for mulitvalent ionic
compounds
• Write formulas for the following:
• Copper (I) Oxide
• Copper (II) Oxide
Answers -
Naming compounds with multivalent
metals
• If the metal is in group B it requires a
roman numeral in the name.
• You will have to deduce the roman
numeral based on the formula.
• Example
– Name CoI2
• Answer -
Naming compounds with multivalent
metals
• Deducing the roman numeral
• Multiply the charge on the anion by the number of
anions and then divide by the number of cations to
get the roman numeral.
• Write the names for Fe2S3 SnO2
• Answers -
• Take ten minutes and work on sections
5 and 6 on the back side of your
worksheet.
Polyatomic Ions
• A group of atoms that carry a charge
• Examples:
– SO42-
NO31-
• Names of polyatomic ions that contain oxygen will end
in -ate or -ite
• -ite is one less oxygen then ate
• Example
– Sulfate is SO42- Sulfite is SO32– Chlorate is ClO31- Chlorite is ClO21-
• Other polyatomic ions
– NH41+ Ammonium
– OH1- Hydroxide
CN1- cyanide
Writing formulas using polyatomic ions
• The polyatomic ion is treated as one unit.
• Balance the charges
• Place parenthesis around the polyatomic ion
if there is more than one
• Example
– Write the formula for Iron (II) Nitrate
Naming using Polyatomic ions
• Name the metal then name the
polyatomic ion. If you need a roman
numeral; include it.
• Treat the polyatomic ion as one unit (as
if it were one atom)
• Example - Name CuSO4
Exceptions for roman numerals
• Silver, Cadmium and Zinc do not get
roman numerals.
• Ag is always +1, Cadmium and Zinc are
always +2
• Tin and Lead need roman numerals.
They are multivalent (multiple oxidation
states)
Naming Acids
•
•
•
•
•
•
Memorize
HCl - Hydrochloric Acid
H2SO4 - Sulfuric Acid
HNO3- Nitric Acid
H3PO4 - Phosphoric Acid
Note - Acids give off H1+ (Hydrogen ions) and
bases give off OH1- ions
• What do you get when an acid and base
combine?
Naming Compounds
Is there a metal?
Yes
No
Ionic
Molecular
Does the compound contain a
multivalent ion?
aka - transition metal or
group B element
Use prefixes to represent
the number of atoms.
Example: H2O Dihydrogen Monoxide
CO2 Cabon Dioxide
No
Yes
Name the cation first
then name the anion
Example: Lithium Fluoride
Magnesium Carbonate
Name the cation first
Place a roman numeral
Name the anion
Example: Iron (II) Sulfate
Check for understanding
• Name or write the formula for:
–
–
–
–
Potassium Sulfate
Chromium (III) Cyanide
Fe(ClO3)3
CuCl
• Answers
• Now finish your worksheet and work on your
homework.
• Get help
• Make sure and check your answers. You will be
writing formulas all year and doing math based on
these formulas. You get the formula wrong you get
the math wrong.
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Helpful hints for balancing
chemical equations
•
•
•
•
•
Balance hydrogens second to last
Balance oxygens last
Check for lowest ratio
Coefficients must be whole numbers
Don’t break up your compounds with coefficients
– NaCl cannot become Na6Cl
• Do not change your subscripts
• Balance the polyatomic ions as one unit (if it didn’t
break apart)
• Perform a final check
Balance the following
•
•
C2H6 +
O2 -->
Na3PO4 +
Mg3(PO4)2
CO2 +
H2O
Mg(NO3)2 --> NaNO3 +
Types of Reactions
Including reaction prediction
Generals about writing Equations
• Reactants on the left and products on the right
• Symbols - see text for symbols that are
included in equations.
– Ex: g for gas, l for liquid, s for solid
– Downward arrow for precipitate, aq for aqueous
• Catalyst goes above the arrow
•
KI
– Ex H2O2(aq) ---> H2O(l) + O2(g)
• Diatomic Molecules - BrINClHOF
– Elemental state - Br2I2N2Cl2H2O2F2
1. Synthesis (Combination)
• Two or more substances react to form a
single substance
• R + S --> RS
• Ex) SO3(g) + H2O(l) --> H2SO4(aq)
• Usually gives off energy when forming bonds
• Example: Write the balanced equation for:
magnesium ribbon reacting with oxygen
• Mg(s) + O2(g) ---> MgO(s)
• 2 Mg(s) + O2(g) --> 2MgO(s)
1. Synthesis (Combination)
• Your turn. Write balanced equations for
the following:
– Aluminum (s) reacts with oxygen (g)
– Hydrogen (g) reacts with oxygen (g)
• Answers:
2. Decomposition
• A single compound is broken down into simpler
products
• RS --> R + S
• Ex) BCl3 --> B + Cl2
• Requires energy to break chemical bonds (heat, light,
electricity)
• Example - Write the balanced equation for mercury
(II) oxide decomposing;
• HgO --> Hg + O2
• 2HgO --> 2Hg + O2
2. Decomposition
• Your turn. Write balanced equations for
the following:
• The decomposition of water
• The decomposition of lead (IV) oxide
• Answers
3. Single Replacement Reactions
• An element replaces an element of a
compound
• T + RS --> TS + R
• Ex) Zn(s) + H2SO4(aq) --> ZnSO4(aq) + H2(g)
• A metal may replace a metal or a nonmetal
may replace a nonmetal
• Activity Series - list of metal in order of
decreasing activity
• Nonmetals reactivity decreases as you go
down the periodic table
• This is limited to the halogens -group 7A
3. Single replacement reactions
• Ex) Write the balanced equation when
aluminum reacts with sulfuric acid
• Al(s) + H2SO4(aq) --> Al2(SO4)3(s) + H2(g)
• 2Al(s)+ 3H2SO4(aq) --> Al2(SO4)3(s) + 3H2(g)
3. Single replacement reactions
• Your turn. Write balanced equations for the following:
• When chlorine reacts with potassium iodide
• When copper (assume Cu2+) is added to Iron (II)
Sulfate
• Answers
–
4. Double Replacement
• Exchange of positive ions between two
compounds. Just swap the positive ions and
write the new formula.
• R+S- + T+U- --> R+U- + T+S• Ex) FeS(s) + 2HCl(aq) --> H2S(g) + FeCl2(aq)
• Ex) Write the balanced equation for barium
chloride added to potassium carbonate
• BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + KCl(aq)
• BaCl2(aq) + K2CO3(aq) --> BaCO3(s) + 2 KCl(aq)
4. Double Replacement
• Your turn. Write balanced equations for
the following.
• Iron (III) Sulfide reacting with
hydrochloric acid
• Answer
5. Combustion Reactions
• Oxygen reacts with another substance, often
producing heat and light
• Often involve hydrocarbons
– Compounds of hydrogen and carbon
• Combustion of hydrocarbons produces a lot of
energy, therefore, hydrocarbons are used as fuels.
• Examples: methane, propane, butane, octane
5. Combustion Reactions
•
•
•
•
•
Two types of combustion
1. Complete combustion
CxHy + O2(g) --> CO2(g) + H2O(g) + energy
2. Incomplete combustion
Two more products: CO and C
• CxHy + O2(g) --> CO2(g) + H2O(g) + CO(g) + C(s) + energy
5. Combustion Reactions
• Ex) Write a balanced equation for the
complete combustion of C3H8.
• C3H8(g) + O2(g) --> CO2(g) + H2O(g) + energy
• C3H8(g) + 5 O2(g) --> 3 CO2(g) + 4H2O(g) + energy
• Your turn: Write a balanced equation for the complete
combustion of C8H18.
• Answer
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
QuickTime™ and a
YUV420 codec decompressor
are needed to see this picture.
Precipitation Reactions
• Most ionic compounds dissociate into
cations and anions when dissolved in
water.
• A complete ionic equation (basically a
double replacement reaction) shows
ionic compounds as free ions.
• In other words, write in the charges.
Precipitation reactions
Predicting the precipitate
• Use the chart on the back of your periodic table.
• Which of the following compounds are not soluble
–
–
–
–
–
Calcium Sulfate
Sodium Acetate
Silver Chloride
Aluminum Hydroxide
Potassium Phosphate
Precipitation Reactions
Complete Ionic Equations
• In an aqueous solution, substances exist as
free ions. The equation shows this.
• Example for AgNO3(aq) + NaCl(aq)
Ag+1(aq) + NO31-(aq) + Na1+(aq) + Cl1-(aq) --> AgCl(s) + Na+1(aq) + NO31-(aq)
Precipitation Reactions
Net Ionic Equation
• A net ionic equation indicates those ions
that took part in the reaction.
• Net ionic equation for the reaction from
the previous slide is:
• Ag1+(aq) + Cl1-(aq) --> AgCl(s)
Precipitation Reaction
• Example: Write a complete and net ionic
equation for the reaction of aqueous
solutions of iron (III) nitrate and sodium
hydroxide.
Fe3+(aq) + NO31-(aq) + Na+1(aq) + OH-(aq) --> Fe(OH)3(s) + Na+1(aq) + NO31-(aq)
• Fe3+(aq) + OH1-(aq) --> Fe(OH)3(s)
Precipitation Reactions
• Your turn. Write a complete ionic equation
and a net ionic equation for the reaction of
aqueous solutions of silver nitrate and
potassium sulfate.
• Answer
I. Molar Conversions
The Mole
• 1 mole of hockey pucks would
equal the mass of the moon!
• 1 mole of basketballs would fill a
bag the size of the earth!
• 1 mole of pennies would cover the
Earth 1/4 mile deep!
Molar Conversions
Converting from moles to grams to representative particles
and vice versa. Use the following conversion factor:
1 mole = 6.02 x 1023 representative units = molar mass (g)
or formula weight
Representative units a. ionic compounds are called formula units
b. molecular compounds are called molecules
c. atoms are called atoms.
Example of representative units 6.02 x 1023 atoms Cu
6.02 x 1023 molecules O2
6.02 x 1023 units NaCl
Molar Conversion Examples
• How many moles of carbon are in
26.0 g of carbon?
26.0 g C 1 mol C
12.0 g C
= 2.17 mol C
Molar Conversion Examples
• How many molecules are in 2.50
moles of C12H22O11?
2.50 mol 6.02  1023
C12H22O11 molecules
= 1.51  1024
1 mol
molecules
C12H22O11 C H O
12
22
11
Molar Conversion Examples
• Find the mass of 2.1  1024
formula units of NaHCO3.
2.1  1024
units NaHCO3
84.0 g NaHCO3
6.02  1023
units NaHCO3
= 290 g NaHCO3
Molar Conversion Examples
• Find the number of units of Iron
(III) Chlorate in 98.6 g of Iron (III)
Chlorate.
98.6 g
Fe(ClO3)3
6.02 x 1023
units Fe(ClO3)3
306.3 g
Fe(ClO3)3
= 1.94 x 1023 units Fe(ClO3)3
Moles in a Gas
• 1 mole of gas takes up 22.4 L of space at
standard temperature and pressure.
• Conversion factor - 1 mole = 22.4 L
– Remember this is for a gas only
• Standard Temperature and Pressure (STP)
– Temp = 0oC
– Pressure = 1 atm (atmosphere)
– 1 atmosphere is defined as the amount of
pressure the earth’s atmosphere places on you at
sea level
Calculations w/ molar volume
• Determine the volume, in liters, of 0.60
mol SO2 gas at STP.
• Answer – 0.60 mol SO2 22.4 L SO2
1 mol SO2
= 13 L SO2
Calculations w/ molar volume
Your Turn
• How many atoms of He are contained in your
party balloon if the balloon takes up 4.2 L of
space? Of course, this is one cold party, as it
would be held at STP.
• Answer -
Molarity
• Unit of Concentration
– There are many units of concentration
• Molarity is most useful to the chemist
moles
of
solute
M=
Liters of solution
Liters of solution means the total volume of
water and solute.
If I want a liter of solution I will not use a liter of
water.
Molarity Problems
You work them.
• A saline solution contains 0.90 g NaCl in exactly 100
ml of solution. What is the molarity of the solution?
Molarity Problems
You work them.
• How many moles of solute are present 1.5 L of 0.24
M Na2SO4?
Preparing a solution
• How would you make 500.0 ml of a 0.25 M solution of
copper (II) chloride?
• 0.25 M = mol/0.5000 L - change ml to liters and solve
for moles.
• You need 0.13 moles of CuCl2. Converting to grams
equals 17 grams.
• Final answer
– Take 17 grams of CuCl2 and dissolve in enough water to make
500.0 ml of solution.
• Dissolve the 17 grams in say 400 ml of water. Once the
CuCl2 is dissolved add water up to 500.0 ml.
Preparing a solution
your turn
• How would you prepare a solution of 0.40 M KCl? If a
volume is not given assume 1 L.
Making Dilutions
• Making dilutions from known
concentrations:
• M1 x V1 = M2 x V2
• Volume can be in liters or mL as long as
the same units are used.
Dilution Problems
• How would you prepare 1.00x102 mL of
0.40 M MgSO4 from a stock solution of
2.0 M MgSO4?
• 0.40 M x 100 mL= 2.0 M x V2
• V2 = 20 mL
• Answer - Take 20 mL of 2.0 M MgSO4
and dilute with enough water to make
100 mL of solution.
Dilution Problems
Your Turn
• How would you prepare 90.0 mL of 2.0
M H2SO4 from 18 M stock solution?
• Answer
Dilution Problems
Your Turn - 1 more
• If 250 mL of a 12.0 M HNO3 is diluted to
1 L, what is the molarity of the final
solution?
• Answer -
Percent Composition
QuickTime™ and a
Sorenson Video 3 decompressor
are needed to see this picture.
% composition
your turn
• Hydroxide makes up what percent of
Calcium Hydroxide?
• Answer
Hydrates
• Hydrates are substances that contain
water within the crystalline structure of
the compound.
• The water is not chemically bound; it is
trapped within the crystal.
• Ex. FeSO4 . 7H2O
Empirical vs. Molecular
Formula
Calculating Empirical Formulas
– Lowest whole-number ratio of the atoms of
the elements in a compound
• C6H12O6 (glucose)
• The ratio that glucose normally has for
carbon:hydrogen:oxygen is 6:12:6.
• The lowest ratio that glucose has for
carbon:hydrogen:oxygen is 1:2:1 (each
number can be divided by the smallest
number in the ratio which is 6).
• The empirical formula for glucose is
CH2O since this is the lowest wholenumber ratio of atoms for that
compound.
– May or may not be the same as the normal
molecular formula of a compound
• Next - Calculating empirical formulas
What is the empirical formula of a
compound that is 25.9% nitrogen and
74.1% oxygen?
• If 25.9% of the compound is nitrogen and 74.1% of
the compound is oxygen, then a compound with a
mass of 100 g has 25.9 g of nitrogen and 74.1 g of
oxygen.
• To calculate the empirical formula, we need to relate
the moles of each atom in the compound, so we need
to convert the masses of the elements to moles.
25.9 g N
1 mol N
mol N 

 1.85 mol N
1
14.0067 g N
74.1 g O
1 mol O
mol O 

 4.63 mol O
1
15.9994 g O
This would mean that the ratio of nitrogen to oxygen
is N1.85O4.63.
We can divide each number in the ratio of N1.85O4.63 by
1.85 to get N1O2.50.
Since we cannot have 2.50 atoms of
oxygen, we must multiply through each number by 2 to
even it out, getting N2O5 as our empirical formula.
• In calculating empirical formulas, remember
that the number of atoms is a whole number.
If the number of atoms for an element is close
to a whole number (i.e., 2.1 or 2.2 or 2.8, or
2.9), you can usually round up or down to get
a whole number.
• If you should get a number of atoms closer to
2.33 or 2.5, multiply each number in the
formula by a number that gets that to a whole
number. For example, if you calculated 2.33,
you would multiply this by 3 to get a value of
7 for that number.
• Give it a try
• Determine the empirical formula for a
compound containing 7.8% carbon and
92.2% chlorine.
Empirical vs. Molecular
Formula
Calculating Molecular Formulas
• Although sometimes a molecular
formula may be the same as a
molecule’s empirical formula, like in
carbon dioxide (CO2), we have seen
that the empirical formula for glucose is
not the same as its molecular formula.
• One can determine the molecular
formula of a compound by knowing its
empirical formula and its mass.
• Next - Example
Calculate the molecular formula of the compound
whose molar mass is 180.1583 g and empirical formula
is CH2O.
We know that the molecular formula will have a molar
mass of 180.1583 g. We also know, by calculating the
gmm of CH2O, that CH2O has an empirical formula
mass (efm) = 30.0264 g CH2O.
Now, in order to figure out what we must multiply each
number in the empirical formula by, we must figure out by
what number we must multiply the empirical formula mass
to get the molecular formula mass.
To get from 30.0264 to 180.1583, 180.1583
30.0264
6
• Therefore, we must multiply each number of atoms in
CH2O by 6 to get the molecular formula of C6H12O6.
• You can double-check your answer by recalculating
the molar mass of C6H12O6.
• gmm C6H12O6 = 6 x 12.0111 g + 12 x 1.00794 g + 6 x
15.9994 g = 180.1583 g C6H12O6
• This agrees with the molar mass we were given, so
the molecular formula we calculated is correct.
• Give it a try
• Determine the molecular formula of a
compound that is 40.0% C, 6.6% H, and
53.4% O and the molar mass is 120.0g.
Example (toughy)
• 1.00 g of menthol on combustion yields
1.161 g of H2O and 2.818 g of CO2.
What is the empirical formula?
• Solution:
Stoichiometry
Calculations of quantities in chemical
reactions.
The use of ratios to calculate quantities
The five step process
• 1. Start with the balanced equation
• 2. Set up the problem - put down the tracks
• 3. Convert to moles if needed. This means you would
be given grams, representative units or liters.
• 4. Convert to moles of what you want. You will use
the mole ratio from the balanced equation.
• 5. Convert to what you are trying to find (grams,
liters, representative units) if needed.
Stoichiometry
Example Problem #1
• How many moles of ammonia are
produced when 0.60 mol of hydrogen
reacts with nitrogen?
Stoichiometry
Example Problem #2
• Your Turn
• How many moles of aluminum sulfide
are produced when 1.2 moles of
aluminum reacts with sulfur?
Stoichiometry
Example Problem #2
• Answer
Stoichiometry
Example Problem #3
• How many grams of ammonia will be
produced by reacting 5.40 g of
hydrogen with nitrogen?
Stoichiometry
Example Problem #4
• Your Turn
• How many grams of aluminum are
needed to react with 2.45 g of copper
(II) chloride?
Stoichiometry
Example Problem #4
• Answer
% Yield
• Definitions
• 1. Theoretical Yield
• The maximum amount of product that
can be formed from a given amount of
reactants
• In other words, the calculated amount
predicted through stoichiometry
% Yield
• Definitions
• Actual Yield
• The amount that is actually formed
when the reaction is carried out in the
laboratory.
% Yield =
actual yield
X 100
theoretical yield
% yield will never be over 100%
Most likely it will never even be 100%
Why will % yield never be 100%
• Advantageous to add an excess of an inexpensive reagent
to ensure that all of the more expensive reagents reacts
• Reactant may not be 100% pure
• Materials are lost during the reaction
– If a reactions takes place in a solution it may be
impossible to get all of the reactants or products out of
the solution
• If the reactions takes place at a high temperature,
materials may be vaporized and escape into the air
• Side reactions may occur
– Example Mg burned in air. Some of Mg reacts with
nitrogen reducing the amount of MgO produced.
• Loss of product when filtering or transferring
• If reactants are not carefully measured
% yield example problem
• In a reaction between barium chloride
and potassium sulfate, 3.89 g of barium
sulfate is produced from 3.75 g of
barium chloride. What is the percent
yield?
% yield example problem
answer
• BaCl2 + K2SO4 --> BaSO4 + 2 KCl
3. 75 g BaCl2 1 mol BaCl2 1 mol BaSO4 233.4 g BaSO4
208.3 g BaCl2 1 mol BaCl2 1 mol BaSO4
= 4.20 g BaSO4
3.89 g BaSO4 x 100 = 92.6 %
4.20 g BaSO4
% yield example problem
your turn
13.35 grams of magnesium hydroxide is
produced when 42.50 grams of
magnesium nitrate reacts with an
excess of aluminum hydroxide. What is
the percent yield?
% yield example problem
answer
Limiting Reagent
• 1. Limits or determines the amount of product that
can be formed
• 2. The reagent that is not used up is therefore the
excess reagent
• These types of problems require 2 sets of tracks.
Quantities of both reagents will be given. Therefore,
you need to find out which one is the limiting reagent.
Limiting Reagent
• One track to determine limiting reagent
• A second track to determine product
Limiting Reagent Example problem
• How many grams of copper (I) Sulfide can be produced when 80.0
grams of Cu reacts with 25.0 grams of sulfur?
• 2Cu + S --> Cu2S
• Pick a reactant and calculate how much of the other reactant is
needed.
80.0g Cu 1mol Cu 1mol S 32.1g S
63.5g Cu 2mol Cu 1mol S
= 20.2g S
So, 20.2 g of S is needed; 25.0g is supplied
Plenty of S; therefore, Cu is limiting reagent.
Use Cu to solve the problem
80.0g Cu 1mol Cu 1mol Cu2S 159.1g Cu2S
63.5g Cu 2mol Cu
1mol Cu2S
= 1.00x102 g Cu2S
Limiting Reagent Example Problem - Your Turn
• How many grams of hydrogen can be produced when 5.00g of Mg is
added to 6.00 g of HCl?
Limiting Reagent Example problem- Your Turn
• Acetylene (C2H2) will burn in the presence of oxygen. How many grams
of water can be produced by the reaction of 2.40 mol of acetylene with
7.4 mol of oxygen?