Chapter 3 & A Li’l Bit
About “Moles”
Plus some problem-solving techniques from Chapter 4. We aren’t
going to do all of chapter 4, but if YOU are having difficulty with
problems, I would review this chapter in detail on your own.
Objective A
(remember, the objective refer to the Study Guides)
http://www.magazine-agent.com/officials-logic-problems/magazine
 Look at the following numbers. How easy
would it be to memorize this list and
regenerate it in a week or so on a quiz?
 1, 3, 5, 11, 21, 43, 85, 171, 341, 683, 1365,
2731, 5461, 10923, 21845
 Pretty hard? Pretty Easy? Why?
 Easy huh? OK, what’s the next
number in the sequence? How
about the next 5 numbers?
How bout this?
http://shop.crackerbarrel.com/online/shopping/default.asp
 What if I never mentioned the list again? What if
at the end of the semester…the day before the
SOL…I asked you to remember those numbers?
 Could you do that?
Cracker Barrel
“Brain Teaser”
Game
 If so, you’re going to do really, really well when I
asked you to memorize some “polyatomic ions” on
page 147 of your textbook.
 My guess is you wouldn’t remember though.
Can you make sense out of
NONSENSE?
 Force
Hour
Upon
Neigh
 Koran
Force
Cis
Shun
 Heaven
Fodder
Count
 Ye
Brat
Anent
 Sago
Farce
Anew
 No luck?
The Gettysburg Address
http://www.old-picture.com/defining-moments/pictures/Abraham-Lincoln-Antietam-Battlefield.jpg
http://en.wikipedia.org/wiki/Gettysburg_Address
 “Four score and seven years ago our
fathers brought forth on this continent a
new nation, conceived in Liberty, and
dedicated to the proposition that all men
are created equal.”
Yes, Ms Rackley, I
know this was
Antietam, but it’s a
nice picture.
 Abraham Lincoln delivered these words
on Nov 19, 1863. (Go back a slide and read
down instead of just left to right).
 Being able to look at something new, and
make sense of it is an important skill for
the rest of this course.
1, 3, 5, 11, 21, 43, 85, 171, 341, 683,
1365, 2731, 5461, 10923, 21845
http://www.antiqueradiomuseum.org/RR%20Rule%20Book%20for%20CMSP&P%20RR.jpg
 The numbers don’t seem to follow
any kind of pattern.
 But, what if you knew a rule? It
wouldn’t just be memorization then.
 OK, good point! The rule is:
You just wish it
was that EASY!




Start with 1.
Double and add 1.
Double and subtract 1.
And so on.
Let’s look at the numbers again
http://equintconsulting.com/wp-content/uploads/2008/10/istock_000005164183small.jpg
 1, 3, 5, 11, 21, 43, 85, 171, 341, 683,
1365, 2731, 5461, 10923, 21845
 Start with 1.
 Double and add 1. Doubling 1 gives me 2
and adding one gives me 3.
 Double and subtract 1. Doubling 3 gives
me 6 and subtracting 1 gives me 5.
Learn HOW to do the problem;
don’t just memorize the answer!
http://familyfun.go.com/Resources/printable-previews/previews/beaver_memorygame_august2.jpg
Don’t
 Doesn’t “knowing the rule” make
the list easier to memorize?
 You don’t have to memorize a
bunch of unrelated numbers.
 IF YOU KNOW the rule, you can
generate the number list on the test
easily. Anytime! Anywhere!
OK, does this relate to
Chemistry at all?
 What makes Chemistry hard for most students, is
you DON’T know the rule. You didn’t read the
book. You didn’t look at the sample problems.
You don’t know the RULE!
 Every problem looks like it’s a totally brand new
problem. But you’re not solving for x or y.
You’re finding the “density” or some other REAL
PROPERTY of matter.
 Every problem looks totally unrelated to anything
we’ve done before.
Objective A
http://www.52shows.com/wp-content/uploads/2009/02/huh.jpeg
 If we can understand the rules, it makes the
problems much, much easier. And if you figure
out how to do one problem, you should be able
to figure out other problems just like it (like on
the test).
 Let’s start with a simple problem. I’ll give you
two numbers, and you tell me the answer.
 The numbers are 2 and 3.
 What’s the answer?
Objective A
http://www.platformnation.com/wp-content/uploads/2009/06/shrugging.jpg
 What are you supposed to do
with those numbers?
What?
 It’s like that in Chemistry. If you
don’t have a clue, anything you
do is potentially just as valid as
anything else.
 But usually it’s pretty easy to
figure out what to do.
Objective A
http://www.tvgasm.com/newsgasm/Bill-nye.jpg
What
would
Bill Nye
do?
 In Chemistry, when we have
numbers, much of the time we do
one of three things:
 Multiple 2 x 3
 Divide 2 / 3
 Divide 3 / 2
 If you can figure out which
operation to do when, you can
actually solve the problem. That’s
our goal.
Bill Nye is the Guy!
Objective A
http://francisanderson.files.wordpress.com/2009/01/billy-mays.jpg
 We don’t use abstract or
imaginary or hypothetical
quantities in chemistry.
But wait!
There’s more!
 We use real amounts WITH
UNITS.
 5 grams
 14.7 milliliters
 3.6 x 10-3 moles (don’t worry about moles for
a while…we’ll get to these)
 6.02 x 1023 molecules
Note the
cool red
LED
numbers
Objective A
http://www.vintagecalculators.com/html/texas_insturments_ti_58.html
 Notice on the last slide, most of the the
numbers aren’t as nice as 2 or 3.
 But so what. You have a calculator.
Calculators don’t care how hairy the numbers
are…they just add ‘em up.
 You need to get in the habit of always using
units. If the units in your answer don’t work
out right, the problem is WRONG. Try again.
Second calculator I ever got (1977). I got my first one in 1975 and
all it did was add, subtract, multiply, divide AND do square
roots. Quite an upgrade!
Units TELL you the right
ANSWER!
 If your units don’t work out right, your
answer is wrong.
 Do the problem over. Read the problem
carefully before you start, and check
your work when you get done.
 Don’t just ignore wrong units. It is a big
clue for you. This is very important.
 LISTEN to the units!
Qualitative and
Quantitative Data
 Qualitative (think quality)
 This is more a description. These are observations usually.
 It’s blue. It’s hot. It’s cool. It’s smells like rotten eggs.
 Remember to look for qualitative date when you do your lab
experiments.
 Quantitative (think quantity or amount)
 This is telling me how MUCH of something you have.
 5.0 g of baking soda. 2L of diet Coke. 4.5 moles of CO2.
 Remember to record quantitative data in your lab notebook
when you do experiments too.
Math Alert!
Objective B
 Scientific notation is used to express very large and
very small numbers.
 Two very important numbers in chemistry are
 6.02 x 1023 = 602,000,000,000,000,000,000,000 (Avogadro’s Number)
 6.6 x 10-34 = 0.00000000000000000000000000000000066 (Planck’s
Constant)
 One is a VERY LARGE number. One is a very small number.
Objective B
http://phoenix.fanster.com/suns/files/2009/05/pile-of-money.png
 6.02 x 1023 is very large.
 Let’s say you had that much
money.
 $602,000,000,000,000,000,000,000.00
 If you spent a billion dollars
every second of every day,
how long do you think the
money would last?
Objective B
http://musikality.net/wp-content/uploads/2009/02/shocked.jpg
http://hebrewandgreekreader.files.wordpress.com/2009/05/waynes_world_15b15d.jpg
 Answer:
More than 19
million years!
Way!
No Way!
Power Point Assignment
 Another one for everyone…we are going to study
“moles” in Chap 7. However, I want to introduce
moles much earlier than that, and Chap 3 seems
like a great time to do so.
 Research moles online and write a 150 word
summary about moles. (Note: mole is a chemical
UNIT and not a little creature that burrows into
the ground.)
 Be prepared for a one question quiz on Chap 3!
Scientific Notation
(we now resume our regularly scheduled power point…)
 We use scientific notation to more easily represent
very large and/or very small numbers.
 We represent numbers as a “factor” times a power
of 10.
 Let’s look at 500, as an example.
Scientific Notation
 500 = 5 x 10 x 10, right?
 But 10 x 10 = 100 and as you know, 100 = 1 x 102.
 So we can rewrite 500 as 5 x (1 x 102).
 But 5 x 1 just equals 5, so we usually ignore the “1”
and just write 500 as 5 x 102.
 That’s scientific notation. I’m sure this is just a
review for all of you.
Scientific Notation
 How about 999? The decimal point moves from left to right.
 999 = 9.99 x 10 x 10
 So, 999 = 9.99 x 102.
 When you put a large number into scientific notation, the
exponent will be positive.
Scientific Notation
 How about 0.0014?
 Well, 1.4 divided by 10 = 0.14.
 0.14 divided by 10 = 0.014.
 0.014 divided by 10 = 0.0014.
 So we’re dividing by 10 and doing it 3 times.
Scientific Notation
 So 0.0014 = 1.4 ÷ 10 ÷ 10÷ 10
 Or 0.0014 = 1.4 ÷ 103 since we are dividing by
10 three times.
 1.4
------103
= 1.4 x 10-3
Scientific Notation
 So 0.0014 = 1.4 x 10-3.
 The decimal point moved from left to right.
 When you put a small number into scientific notation, the
exponent will be negative.
 Remember that these are exponents. Negative exponents
don’t mean negative numbers, they mean very small numbers
(between 0 and 1).
Scientific Notation
 103 = 1,000
Go up,
multiply by 10
 102 = 100
 101 = 10
 100 = 1 (NOT 0!!)
Go down,
divide by 10
 10-1 = 0.1
 10-2 = 0.01
 10-3 = 0.001
How do we add or subtract
using scientific notation?
 Make the exponents the same, if necessary, and then just
add.
 5.4 x 103 + 2.6 x 102 = ?
 Let’s change the second number so that the exponents
match up…
 5.4 x 103 + 0.26 x 103 = 5.66 x 103
How do we add or subtract
using scientific notation?
 Does that make sense?
 Well 5.4 x 103 = 5,400
 And 2.6 x 102 = 260
Math Alert!
 So 5,400 + 260 = 5,660. Putting that back in scientific
notations gives us 5.66 x 103.
 Either way, your calculator will do it for you without
any problems. Maybe!! Use parentheses with
scientific notation on TI-83 and TI-84 graphing
calculators or you will get the wrong answer.
How do we multiply or divide
using scientific notation?
 Even simpler
 To multiply, you multiply the numbers and add the
exponents.
×
 4 x 107 × 2 x 10-3 =
8 x 104 (80,000)
 To divide, you divide the numbers and subtract the
exponents.
÷
 4 x 107 ÷ 2 x 10-3 =
2 x 1010
(20,000,000,000)
Sample Problems on
Scientific Notation
 A 91.43
 B 0.000000000154
 C 6,378,000
 D 0.000008
 E 149,600,000,000
 F 8934.8
Objective c…Accuracy,
Precision and Error
http://www.edupics.com/en-coloring-pictures-pages-photo-dartboard-p9574.jpg
 Think of a dartboard. You
are trying to hit the bullseye.
 Accuracy is hitting what
you are aiming for.
 Precision is hitting the same
Is this accurate, precise, or
spot over and over.
both, or neither?
Objective c…Accuracy,
Precision and Error
http://comps.fotosearch.com/comp/IGS/IGS170/dartboard-darts_~IS028-013.jpg
 Think of a dartboard. You are
trying to hit the bullseye.
 If you hit the edge of the
dartboard, you are not accurate.
Both accurate
AND precise!
 If your darts are all over the place,
you are not precise.
Objective c…Accuracy,
Precision and Error
http://www.durhamtech.edu/graphics/programs/univtransf/chemlab1lg.jpg
Take your time. Know what you
are doing BEFORE you do it.
Pay attention to detail. No one in
this group is not involved.
You’ll get good results.
Everyone working together safely
and observing what’s happening.
 In Chemistry labs, we will take measurements.
 Accuracy is how close you are to the TRUE value.
 Precision is how close all of your measurements are to each other.
Objective c…Accuracy,
Precision and Error
Yes, accuracy is
compared to a “true”
value.
No, precision is
“closeness” to all other
measurements
 Can you be accurate (or have
accuracy) with only 1
measurement?
 Can you be precise (or have
precision) with only 1
measurement?
Objective c…Accuracy,
Precision and Error
Accurate, Not Precise
Accurate, Precise
Not Accurate, Precise
Not Accurate, Not Precise
 Student A had 3 measurements:
Average = 85.0
 Student B had 3 measurements:
85.1
Average = 85.0
 Student C had 3 measurements:
82.3
Average = 82.2
 Student D had 3 measurements:
94.2
Average = 82.2
80, 85, 90
84.9, 85.0,
82.1, 82.2,
70.2, 82.2,
 The accepted value = 84.9. Who is
accurate? Who is precise?
Objective C…Error
 I put the formula for error and % error in your study
guide.
 The error is the difference between the experimental value
and the actual or true value. We take the “absolute value”
because there’s no such thing as “negative error.”
 Being 5% high is the same as being 5% low. You still have
5% error in either case.
Objective C…Error
 If the actual value is 10, it doesn’t matter if you get 9
or you get 11. You are still off by a unit of “1”
 % error = Error / True Value x 100
 Using our example, % error = 1/10 x 100 or 10%.
Objective C…Error
 We will discuss error primarily in our labs. Be sure
to include a discussion of any errors that happened
in your experiment in your “Analysis of Data”
section of your lab report.
 What kind of errors can you have in your labs?
 Two kinds:
 Random
 Systemic
Objective C…Error
http://www.budapesthotels.com/sitepic/error_button.png
http://images.intomobile.com/wp-content/uploads/2009/06/easy-button.jpg
Easy Button
 Random errors are mistakes. You can
take care to reduce or eliminate random
errors. These usually come from being
unprepared (not reading the lab prior
to doing it), rushing to get done, and
careless errors (like forgetting to do a
step).
Error Button
 Systemic errors are errors that are
present in your system. You can’t do
anything about these. If your data is
precise, but not accurate this might
indicate that you had systemic error.
Objective d…Significant Figures
http://www.musicdirect.com/shared/images/products/large/aayremyrtle.
http://becauseican.co.za/wp-content/uploads/2008/04/ruler_0_10.jpg
http://mrsdlovesscience.com/meniscusirr.jpg
 Measure the block of
wood using the ruler.
How long is it?
 How much water is in
the graduated cylinder?
23.0 mL
?? mL
Objective d…Significant Figures
 When you take measurements in Chemistry class or in lab,
you have to worry about how many significant figures
(usually abbreviated as sig figs) you have.
 What are sig figs?
 All the digits you can read and the first one you can estimate.
Between 7.1 and 7.2, so record as 7.15. All 3 digits are significant,
because you read the first 2 and estimated the last 1.
Objective d…Significant Figures
http://www.freefoto.com/images/2000/98/2000_98_1---Number-Zero_web.jpg
Significant?
Maybe or
maybe not!
 However, if you didn’t measure it,
there are rules for figuring out
how many sig figs something has.
 RULE #1. All non-zero digits are
significant.
 So that means that the only thing
you have to determine is whether
or not the zeros are significant.
Objective d…Significant Figures
 Rules are in the study guide.
 2. Zeros between non zero digits are significant. So,
2.003 has 4 sig figs.
 3. Zeros at the end of a number AND to the right of
the decimal point are significant. So, 1.000 has 4 sig
figs.
 4. Zeros at the beginning of a number are never
significant. So, 0.00034 has 2 sig figs.
 5. Exact measurements or exact quantities have an
unlimited amount of sig figs. Example 1 hour = 60
min. “1” and “60” both have an unlimited number
of sig figs.
Objective d…Significant Figures
 Adding/Subtracting
 Answer cannot have more sig figs AFTER the decimal point than the
number with the lowest number.
 Ex: 2.1 + 2.22 + 2.345 = 6.665
 Answer can only have 1 digit after the decimal
 Correct answer = 6.7
 Multiplying/Dividing
 Answer cannot have more sig figs than the number with the lowest
number of sig figs.
 Ex: 2.22 x 2.345 = 5.2059
 2.22 has 3 sig figs and 2.345 has 4 sig figs. Answer can only have 3 sig
figs.
 Correct answer = 5.21
±1 sig figs off is usually OK. If answer should have 4
sig figs and you have 3, 4 or 5, I won’t mark it wrong.
Objective d…Significant Figures
http://www.tvguide.com/celebrities/stuart-scott/214001
 But isn’t 50 and 50.0 and 50.00
and even 50.000 the exact same
number?
50 ≠ 50?
Wha-haphappen?
 Maybe, outside of chemistry
class, but not in here.
 Let me explain.
Objective d…Significant Figures
 50 means that your measurement is somewhere
between 49 and 51.
 50.0 means that your measurement is
somewhere between 49.9 and 50.1
 50.00 means that your measurement is
somewhere between 49.99 and 50.01
 50.000 means that your measurement is
somewhere between 49.999 and 50.001
The more sig figs in your measurement, the more
confidence you have that it’s “exactly” 50 mL or g.
Objective e…SI Units AKA
The Metric System
http://www.boston.com/ae/celebrity/more_names/blog/KG.JPG
 You need to know THESE:





Meters for length (m and cm and nm)
Kilograms for mass (kg and g)
Kelvin for temperature (K)
cm3 or liters for volume (L and mL)
Kilopascals for pressure (KPa and atm)
1.0
Objective e…SI Units AKA
The Metric System
 You need to know
Certified SI Genius
Kilo means 1000 times (1 g = 1 kg)
Centi means 1/100th (100 cm = 1 m)
Milli means 1/1000th (1,000 mm = 1 m)
Micro means 1/1,000,000th (1000 μg = 1
mg or 1,000,000 μg = 1 g)
 Nano means 1/1,000,000,000th (1 billion
nm = 1 m)
 Å = Angstrom means
1/10,000,000,000th (10 billion Å = 1 m)




μ (mu) =
micro
Objective f…Density
http://www.tungsten-spheres.com/density_model.jpg
 Density = mass / volume
 Usually has units of g / cm3
 Density of water = 1.000 g / cm3
 You should memorize this number!
 Density is the mass in grams of 1
cubic centimeter of volume.
Densely packed
Objective f…Density
http://www.stevespanglerscience.com/img/cache/bcb9b8db117ee64376aedaf7af3595ca/sevenlayer-251908.jpg
Higher density layers on
the bottom and lower
density layers on the top.
Is oil more or less dense
than water?
D=m/V
Objective g…Ice Floats
http://www.dharma.org/ims/images/pi_ice_on_pond.jpg
 Why does ice float?
 Why is this a good
thing?
 Ice is actually kind of
strange. Actually water
is a very unique
compound.
 Most substances are
more dense as a solid
than they are as a
liquid.
 Ice is an exception.
Objective h…Take your
Temperature!
 We don’t use Fahrenheit in this
course.
 We use Celcius (°C)
 We also use Kelvin.
View outside my little cuz’s
high school Chem class in
Miami. Probably a balmy 35°C
(308K)
 K = °C + 273
 0 K is called “absolute zero.”
Absolute zero is the temperature at
which all molecular motion stops.
Objective h…Take your
Temperature!
http://www.avogadro.co.uk/miscellany/t-and-p/thermometers.gif
 Kelvin temperature was defined to
be a measure of the kinetic energy of
the sample of matter.
 A sample of matter at 300K has twice
as much kinetic energy as a sample
of matter at 150K. Temperature is
directly proportional to kinetic
energy.
 You must remember how to convert
from °C to K and vice versa. It
WILL be on the test.
The Mole
http://www.naturephoto-cz.com/photos/others/european-mole-22725.jpg
http://www.waukeganschools.org/cspchem/stories/storyReader$12
 No, not this guy.
 A mole is a “counting unit.”
 For example, a dozen is 12 of anything.
 A “mole” is 6.02 x 1023 of anything.
 1 mole of carbon has 6.02 x 1023 carbon
atoms (or particles) and has a mass of 12.0 g
(hint, hint: look on the periodic table and
find the mass of C in the top left corner).
The Mole
Section 7.1, pg 170-181
we will do Section 2 and 3 in Chapter 7 later.
http://i.ehow.com/images/GlobalPhoto/Articles/2100705/PaperSkyscraperwrittenbyDr.Mom-main_Full.jpg
 Suppose I ask you to count this stack
of paper. I need to know how many
sheets there are for a class project.
How did
she do it?
 You come back in about a minute and
tell me, “hey, Mr. Schwartz, there are
1245 pieces of paper there.
 I don’t believe you could’ve counted
them that fast, but you assure me the
total is correct ± a couple of pages
either way.
The Mole
 When I asked for an explanation, the student said,
it’s simple.
 I weighed 1 piece of paper. I recorded that weight
and then weighed the entire stack.
 I divided the total weight by the weight of 1 sheet
and that gave me 1245.
 That’s essentially what we do to count atoms.
The Mole
 We can’t see atoms. We can’t possibly count them,
because they are too small.
 But we know how much 6.02 x 1023 of them should
weigh.
 Let’s say we have 2.12 g of carbon. How many
carbon atoms is that. We know that 1 mole of
carbon has a mass of 12.0 g and 1 mole of C contains
6.02 x 1023 atoms.
2.12 g C = 1.06 x
 2.12g C
23
10
1 mole
6.02 x 1023 atoms
12 g C
1 mole
atoms
Real simple: multiply everything on top.
Divid everything on bottom. Cancel out units
where you can.
2.12 g C = 1.06 x
 2.12g C
23
10
1 mole
6.02 x 1023 atoms
12 g C
1 mole
atoms
Notice that g of C cancels out and mole cancels out,
leaving us with the units of ATOMS for our answer.
2.12 g C = 1.06 x
 2.12g C
23
10
1 mole
6.02 x 1023 atoms
12 g C
1 mole
atoms
We really can’t “count” 1.06 x 1023 atoms. They are too
small. But by doing it this way, we can calculate how
many we have.
Mass of an Element
“gram atomic mass” or just atomic mass
 To find the mass of an element, you look it up on
the Periodic Table.
 For simpicity, let’s round all the elments to 1
decimal point.
 What is the mass of sodium? Iron? Krypton?
Tungsten? Uranium?
Mass of a Compound
“gram formula mass” (for ionic compounds) or “gram molecular mass” (for
molecular compounds) or just a generic “Molar Mass.”
 There are no compounds on the Periodic Table.
 So you take the mass of each element, multiply that
times the subscript for that element and then add
everything together to get the mass of the
compound.
 Ex: mass of CO2 = 12 + 16 x 2 = 12 + 32 = 44
 What is the mass of water (H2O)? Rust (Fe2O3)?
Table Sugar (sucrose = C12H22O11)? Salt (NaCl)?
Power Point Assignment
 If your last name ends in A-L
 Pg 181, 12-14 (a and b only)
 Pg 198, 46 (b and d), 48, 50 (b, d and f)
 If your last name ends in M-Z
 Pg 181, 12-14 (b and c for 12; c and d for 13 and 14)
 Pg 198, 46 (a and c), 48, 50 (a, c and e)
 Both groups be prepared for a one-question quiz on
Chapter 7 material in this powerpoint.
The End
What is next?
Well, Chapter 5, 28, 13, and 14.
If you are in advanced Chemistry, we are going to do all of those in
Unit 2 (5&28) and 3 (13&14)
If you are in Chemistry, it’s mostly going to be Chapter 5, with a little
bit of all the others thrown in. We are going to skip quite a lot in 28, 13
and 14, however.