1. Ch. 5 Notes---Scientific Measurement
Qualitative vs. Quantitative
Qualitative measurements give results in a descriptive nonnumeric form. (The result of a measurement is an _____________ describing the object.)‏
*Examples: ___________, ___________, long, __________...
Quantitative measurements give results in numeric form. (The results of a measurement contain a _____________.)‏
*Examples: 4’6”, __________, 22 meters, __________...
Accuracy vs. Precision
Accuracy is how close a ___________ measurement is to the ________ __________ of whatever is being measured.
Precision is how close ___________ measurements are to _________ ___________.
adjective
short
heavy
cold
number
600 lbs.
5 ºC
single
true
value
several
each
other

2. Practice Problem: Describe the shots for the targets.
Bad Accuracy & Bad Precision
Good Accuracy & Bad Precision
Bad Accuracy & Good Precision
Good Accuracy & Good Precision

3. Significant Figures
Significant figures are used to determine the ______________ of a measurement. (It is a way of indicating how __________ a measurement is.)‏
*Example: A scale may read a person’s weight as 135 lbs. Another scale may read the person’s weight as lbs. The ___________ scale is more precise. It also has ______ significant figures in the measurement.
Whenever you are measuring a value, (such as the length of an object with a ruler), it must be recorded with the correct number of sig. figs.
Record ______ the numbers of the measurement known for sure.
Record one last digit for the measurement that is estimated. (This means that you will be ________________________________ __________ of the device and taking a __________ at what the next number is.)‏
precision
precise
second
more
ALL
reading in between the
marks
guess

4. Significant Figures
Practice Problems: What is the length recorded to the correct number of significant figures?
length = ________cm
11.65
(cm)
length = ________cm 58

5. For Example
Lets say you are finding the average mass of beans. You would count how many beans you had and then find the mass of the beans.
26 beans have a mass of grams.
44.56 grams ÷26 = grams
So then what should your written answer be?
How many decimal points did you have in
your measurement?
Rounded answer = 2
1.71 grams

6. The SI System (The Metric System)‏
Here is a list of common units of measure used in science:
Standard Metric Unit Quantity Measured
kilogram, (gram) ______________
meter ______________
cubic meter, (liter) ______________
seconds ______________
Kelvin, (˚Celsius) _____________
The following are common approximations used to convert from our English system of units to the metric system:
1 m ≈ _________ kg ≈ _______ L ≈ 1.06 quarts
1.609 km ≈ 1 mile gram ≈ ______________________
1mL ≈ _____________ volume 1mm ≈ thickness of a _______
mass
length
volume
time
temperature
1 yard
2.2 lbs.
mass of a small paper clip
sugar cube’s
dime

7. The SI System (The Metric System)‏

10. kilo- hecto- deka- deci- centi- milli-
Metric Conversions
The metric system prefixes are based on factors of _______. Here is a list of the common prefixes used in chemistry:
kilo- hecto- deka deci- centi- milli-
The box in the middle represents the standard unit of measure such as grams, liters, or meters.
Moving from one prefix to another involves a factor of 10.
*Example: 1000 millimeters = 100 ____ = 10 _____ = 1 _____
The prefixes are abbreviated as follows:
k h da g, L, m d c m
*Examples of measurements: 5 km 2 dL 27 dag 3 m 45 mm
mass cm dm m
grams Liters meters

11. 380 km = ______________m 1.45 mm = _________m
Metric Conversions
To convert from one prefix to another, simply count how many places you move on the scale above, and that is the same # of places the decimal point will move in the same direction.
Practice Problems:
380 km = ______________m mm = _________m
461 mL = ____________dL cg = ____________ dag
0.26 g =_____________ mg ,000 m = _______km
Other Metric Equivalents
1 mL = 1 cm L = 1 dm3
For water only:
1 L = 1 dm3 = 1 kg of water or mL = 1 cm3 = 1 g of water
(1) How many liters of water are there in 300 cm3 ? ___________
(2) How many kg of water are there in 500 dL? _____________
kilo- hecto- deka deci- centi- milli-
380,000
4.61
0.0004
260
230
0.3 L
50 kg

12. Metric Volume: Cubic Meter (m3)
10 cm x 10 cm x 10 cm = Liter

18. Ch. 4 Problem Solving in Chemistry
Dimensional Analysis
Used in _______________ problems.
*Example: How many seconds are there in 3 weeks?
A method of keeping track of the_____________.
Conversion Factor
A ________ of units that are _________________ to one another.
*Examples: 1 min/ ___ sec (or ___ sec/ 1 min)
___ days/ 1 week (or 1 week/ ___ days)
1000 m/ ___ km (or ___ km/ 1000 m)
Conversion factors need to be set up so that when multiplied, the unit of the “Given” cancel out and you are left with the “Unknown” unit.
In other words, the “Unknown” unit will go on _____ and the “Given” unit will go on the ___________ of the ratio.
conversion
units
ratio
equivalent 60 60 7 7 1 1
top
bottom

19. If your units did not ________ ______ correctly, you’ve messed up!
How to Use Dimensional Analysis to Solve Conversion Problems
Step 1: Identify the “________”. This is typically the only number given in the problem. This is your starting point. Write it down! Then write “x _________”. This will be the first conversion factor ratio.
Step 2: Identify the “____________”. This is what are you trying to figure out.
Step 3: Identify the ____________ _________. Sometimes you will simply be given them in the problem ahead of time.
Step 4: By using these conversion factors, begin planning a solution to convert from the given to the unknown.
Step 5: When your conversion factors are set up, __________ all the numbers on top of your ratios, and ____________ by all the numbers on bottom.
Given
Unknown
conversion factors
multiply
divide
If your units did not ________ ______ correctly, you’ve messed up!
cancel out

20. How many hours are there in 3.25 days?
Practice Problems:
How many hours are there in 3.25 days?
(2) How many yards are there in 504 inches?
(3) How many days are there in 26,748 seconds?
24 hrs
3.25 days
78 hrs x =
1 day
1 ft
1 yard
504 in.
14 yards x x =
12 in.
3 ft
1 min
1 hr
1 day
26,748 sec
days x x x =
60 sec
60 min
24 hrs

21. Converting Complex Units
A complex unit is a measurement with a unit in the _____________ and ______________.
*Example: 55 miles/hour meters/sec g/mL
To convert complex units, simply follow the same procedure as before by converting the units on ______ first. Then convert the units on __________ next.
Practice Problems: (1) The speed of sound is about 330 meters/sec. What is the speed of sound in units of miles/hour? (1609 m = 1 mile)
(2) The density of water is 1.0 g/mL. What is the density of water in units of lbs/gallon? (2.2 lbs = 1 kg) (3.78 L = 1 gal)
numerator
denominator
top
bottom
330m
1 mile
3600 sec
738 miles/hr x x =
sec
1609 m
1 hr
1.0 g
1 kg
2.2 lbs
1000 mL
3.78 L
8.3 lbs/gal x x x x = mL
1000 g
1 kg
1 L
1 gal

22. Ch. 6 Notes -- Chemical Composition
What is a mole?

23. Ch 6 – Chemical Quantities
The Mole!!!
A counting unit
Similar to a dozen, except instead of 12, it’s
602 billion trillion…
(602,000,000,000,000,000,000,000)
___________ (in scientific notation)
This number is named in honor of Amedeo _________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present…6.02 x 1023
6.02 x 10 23
Avogadro

24. Just How Big is a Mole?
Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles.
If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles.
If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

25. The Mole 1 dozen cookies = ___ cookies12
1 dozen cookies = ___ cookies
1 mole of cookies = ___________ cookies
1 dozen cars = ___ cars
1 mole of cars = __________ cars
1 dozen Al atoms = ___ Al atoms
1 mole of Al atoms = __________ atoms
Note that the NUMBER is always the same, but the ______ is very different!
Mole is abbreviated ______ .
6.02 X 1023 12
6.02 X 1023 12
6.02 X 1023
MASS
mol

26. The Mole and Mass
Mass in grams of 1 mole equal to __________ of the atomic masses.
Practice problem:
Calculate the mass of 1 mole of CaCl2.
Ca = 1 x ________ g/mol =
Cl = 2 x ________ g/mol =
= __________ g/mol CaCl2
1 mole of CaCl2 =
the sum
40.1
40.1 g/mol
35.5
71.0 g/mol
40.1 g/mol
71.0 g/mol
111.1
111.1 g/mol

27. Mole Conversion Factors that you will need to know!
1 mol = __________ atoms/molecules/etc.
1 mol = _____________ grams
1 mol = ________ Liters of gas at STP
(STP is Standard Temp. and Pressure… we will talk about what this means later!)
6.02 x 1023
? ( molar mass)
22.4

28. Ch. 6 Notes -- Chemical Composition
Practice Problems:
(1) How many atoms of hydrogen are there in each compound?
a) Ca(OH)2 ___ b) C3H8O___ c) (NH4)2HPO4 ___ d) HC2H3O2 ___
(2) Calculate the formula mass of each compound. (Add up all the
atomic masses for each atom from the Periodic Table.)
a) CaCO3 b) (NH4)2SO4
c) C3H6O d) Br2
e) H3PO4 f) N2O5 2 8 9 4
2 N’s = 2 x 14.0 = 28.0
8 H’s = 8 x 1.0 = 8.0
S = 32.1
4 O’s = 4 x 16.0 = 64.0
Ca = 40.1
C = 12.0
3 O’s =3 x 16.0 = 48.0
Add them up!
132.1 g/mol
Add them up!
100.1 g/mol
159.8 g/mol
C = 3 x 12.0 = 36.0
H = 6 x 1.0 = 6.0
O =16.0
2 Br’s = 2 x 79.9 =
Add them up!
58.0 g/mol
2 N’s = 2 x 14.0 = 28.0
5 O’s = 5 x 16.0 = 80.0
3 H’s = 3 x 1.0 = 3.0
P = 31.0
4 O’s = 4 x 16.0 =64.0
Add them up!
98.0 g/mol
Add them up!
108.0 g/mol

29. 3) Convert 835 grams of SO3 to moles.
4) How many molecules of CH4 are there in 18 moles?
5) How many grams of helium are there in 5.6 x 1023 atoms of helium?
6) How many molecules are there in 3.7 grams of H2O?
1 mole SO3
835 g SO3
10.4 moles of SO3 x =
80.1 g SO3
6.02 x 1023 molecules CH4
18 moles CH4 x =
1.08 x 1025 molecules CH4
1 mole CH4
4.0 grams He
5.6 x 1023 atoms He x
3.72 grams He =
6.02 x 1023 atoms He
6.02 x 1023 molecules H2O
3.7 grams H2O
1.24 x 1023 molecules H2O x =
18.0 grams H2O

30. Calculating Percent Composition by Mass
Step 1: Find the formula mass of the compound by adding the individual masses of the elements together.
Step 2: Divide each of the individual masses of the elements by the formula mass of the compound.
Step 3: Convert the decimal to a % by multiplying by 100.
Practice Problems:
(1) Find the % composition of the elements in each compound.
a) Na3PO b) SnCl4
= = 42.1%
3 Na’s = 3 x 23.0 = 69.0
P = 31.0
4 O’s = 4 x 16.0 =
÷ 164
Sn = 118.7
4 Cl’s = 4 x 35.5 =
÷ 260.7
= 45.5%
= = 18.9%
÷ 164 +
÷ 260.7
= 54.5%
260.7
= = 39.0% +
÷ 164
164

31. Elements in the Universe: % Composition by Mass

32. Earth’s Crust: % Composition by Mass

33. Entire Earth (Including Atmosphere): % Composition by Mass

34. Human Body: % Composition by Mass

35. Determining the Empirical Formula for a Compound
The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound.
Examples: H2O is the empirical formula for water.
_______ is the empirical formula for glucose, C6H12O6.
Practice Problems: What is the empirical formula for the following compounds? a) C6H6= ________
b) C8H14O2 = ________
c) C10H14O2 = _________
d) Ca5Br10 = ________
e) N3O9 = ________
whole
ratio
C1H2O1 CH
C4H7O
C5H7O
CaBr2
NO3

36. Determining the Molecular Formula for a Compound
The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule.
Examples: 1) H2O is the empirical & molecular formula for water ) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!)
Practice Problems: (1) If the empirical formula for a compound is CH2, which of the following is a possible molecular formula for the compound? a) C8H16 b) C8H8 c) C4H2 d) C3H9
(2) If the empirical formula for a compound is C2H3, which of the following is a possible molecular formula for the compound? a) C2H6 b) C10H c) C6H12 d) C8H14
whole # multiple 6

37. Determining the Molecular Formula for a Compound
Find the molecular formula for C2H7 if the molecular mass of the compound is 93.0 g/mol.
Find the molecular formula for P2O5 if the molecular mass of the compound is g/mol.
C2H7 = 31.0 g/mol
93.0 g/mol
= 3
31.0 g/mol
C2H7 x 3
= C6H21
g/mol
P2O5 = g/mol
= 2
g/mol
P2O5 x 2
= P4O10