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Chemistry-1 CP
Chapter 3, Scientific Measurements (pp. 62-97)
Ms. Wack
Introduction: In this chapter, you will apply the scientific method to various problems and use experiments to prove
hypotheses. You will also learn the basic mathematical skills needed to succeed in chemistry.
Objectives (what you should know/be able to do by the end of the chapter):
 Define chemistry
 Apply the steps of the scientific method to solve a problem
 Set up an experiment to test a hypothesis, including the variables, the control and the constants.
 Graph data and recognize a trend
 Know the names and symbols of the SI Base Units and Non-SI Units
 Know how to derive a unit of measurement
 Define and calculate density
 Know the metric system prefixes
 Know the differences between precision and accuracy
 Know the correct number of figures to use when making measurements on various instrumentation
 Know how to round to the appropriate number of significant figures when performing mathematical calculations
 Know how to convert numbers between regular notation and scientific notation
 Know how to calculate percents, percent error and ratios
 Know how to convert between units using dimensional analysis
Chapter 1
Chemistry and You
Chemistry in
the Real World
Scientific
Method
Mathematics in Chemistry
SI Units
Uncertainty
Graphs
Hypothesis
Dimensional
Analysis
Sig Figs
Experiments
Variables
Data
Control
Theory
Constants
Law
Scientific
Notation
Percents/
Ratios
Percent
Error
Density
Chemistry:
o
Also known as:
o
Examples of Chemistry in Life:
The Scientific Method:

Steps of the Scientific Method
1. State the Problem
2. Make Observations
3. Develop a Hypothesis
Hypothesis:
Examples: Write a hypothesis to explain each of the following observations.
1) The static on your radio increases right before it thunders during a storm.
2) People who smoke cough more than people who don’t smoke.
3) You sneeze every time you visit your best friend’s house.
4) On a cold morning, the air pressure in the tires of your car measures 34 psi. After several
hours of high-speed driving, the pressure measures 38 psi.
4. Design an Experiment
Experiment:

Variable:
o
Independent Variable:
o
Dependent Variable:

Control:

Constant:
Example: What experiment could be done to prove/disprove the following hypothesis:
“Clean” laundry detergent causes skin rash.
Independent Variable:
Dependent Variable:
Control:
Constant:

Data:
Qualitative Data:
Quantitative Data:
o
Graph
 X-axis:

Y-axis:

numbering:

labeling:

best-fit line:
o
Inversely Proportional:
o
Directly Proportional:
Example: Create a line graph of the following data:
Mass (g)
25
30
40
50
54
Volume (cm3)
100
115
134
160
163
5. Draw Conclusions
 Theory:

Law:
o
What’s the difference between a theory and a law?
UNCERTAINTY IN MEASUREMENTS
 Why are measurements uncertain?

What does “uncertainty in measurement” mean when reading the measurements on instrumentation?

Precision:
o

If measurements are precise then all the measurements are:
Accuracy:
o
If measurements are accurate then all the measurements are:
Percents:
Example: If 20 students in a class of 33 students scores an A on a test, what percentage of students got an
A?

Percent Error:
o
To calculate percent error:
Example: You measure the classroom temperature to be 23°C. The actual classroom temperature is 20°C.
What is your percent error?
Rounding:
Example: Round the following numbers to the number of digits designated in parentheses.
a) 2.3344 (1) ____________
b) 1.029 (3) _____________
Significant Figures:
o
Rules to determining Significant Figures:
1. All nonzero numbers are significant.
Example: 123 has _________________ significant figures.
2. All zeroes at the beginning are not significant.

Example: 0.0025 has ______________ significant figures.
3. Zeroes between 2 nonzero digits are significant.

Example: 5007 has ______________ significant figures.
c) 0.00234 (2) ____________
4. Zeroes at the end of a number are only significant if the number contains a decimal point.

Example: 470 has ________________ significant figures.

Example: 470.0 has ______________ significant figures.

Example: 0.00470 has ___________ significant figures.
5. In scientific notation, all numbers in the coefficient are significant.

o
Example: 2.020 x 104 has ___________ significant figures.
Easier Rule: To count significant figures, if there is a decimal, count all digits including and after the first
non zero number. If there is not a decimal, start counting at the first non zero number but do not count
zeroes at the end of the number.
Example—How many significant figures is in each of the following measurements?

a) 3.3333
_____________________
d) 2000.0 _________________
b) 3023
_____________________
e) 0.216 __________________
c) 72800
_____________________
f) 0.009030 ________________
Significant Figures in Calculations
o Multiplication/Division:
o
Addition/Subtraction:
Example: Perform the following calculations and round off to the correct number of significant figures.
a) 0.3287 g x 45.2 g = _________________________
b) 125.5 kg + 52.68 kg + 2.1 kg = __________________________
c) 0.258 mL  0.36105 mL = _____________________________
d) 68.32 ns – 1.001 ns  0.00367 ns = ________________________

Metric System:
o
International System of Units (SI):
Base Units:
Physical Quantity
Mass
Unit Name
Unit Symbol
Measured Using a…
Length
Time
Quantity
Temperature
Electric Current
Luminous Intensity

Non-SI Units Commonly Used in Chemistry
Physical Quantity
Volume
Unit Name
Unit Symbol
Pressure
Temperature
Energy

Derived Units:
o
Area:

o

Derived SI Unit for Area:
Volume:


Derived SI Unit for Volume:

To calculate density:

Derived SI Unit for Density:
Density:
Examples:
a) Calculate the density of a piece of iron with a mass of 1.23 kg and a volume of 156 cm 3.
b) An unknown liquid has a mass of 30.6 g and a volume of 53.3 mL. What is the density of the
liquid?
c) Iron has a density of 7.86 g/cm3. Could a block of metal with a mass of 18.2 g and a volume
of 2.56 cm3 be iron?
d) If aluminum has a density of 2.70 g/cm3, what is the mass of a piece of aluminum that has a
volume of 35 cm3?
e) Iron has a density of 7.86 g/cm3. What volume will 160.0 g of iron occupy?
Temperature
-Measured with:
Temperature Scales
-Fahrenheit Scale
Example
Freezing
Water
Boiling
Water
Normal
Body
Temp.
-Celsius Scale
-Kelvin Scale
-Absolute Zero:
Celsius
Kelvin
Examples:
1) Pure gold melts at 1064C. What is the melting point of gold in Kelvins?
2) 4120 K = ____________ C
3) Tyrone uses a Celsius thermometer to measure the temperature of boiling water in the lab.
He then records the temperature in his lab manual as 273 K. What is the error in his
measurements?

Metric Prefixes:
Prefix
megakilodekaBASE UNIT
decicentimillimicronanopico-

In 1 base unit there are:
Dimensional Analysis:
o
Unit Equality:
o
Conversion Factor:
o
What is a conversion factor equal to?
o
How do you use conversion factors?
o
Steps to Dimensional Analysis
1.
Example: How many cm in 2.30 meters?
2.
3.
4.
Example 1: How many grams in 33.2 dekagrams?
Example 2: How many kilometers in 22.3 meters?
Example 3: How many milliliters in 0.0234 kilometers?
Example 4: How many nanoseconds are in 5000 kiloseconds?
SCIENTIFIC NOTATION

Scientific Notation:
o
o
Exponent:
To convert a number to scientific notation:
o
To convert a number from scientific notation to regular notation:
Example 1: Express each of the following numbers in scientific notation.
a) 8960 = ___________________________
c) 0.00023 = ________________________
b) 36,000,000 = _____________________
d) 0.000 000 025 3 = _________________
Example 2: Express each of the following numbers in regular notation.
a) 4.563 x 107 = ___________________________
b) 2.53 x 10-3 = _____________________
c) 6.805 x 108 = ___________________________

d) 1.33450 x 10-7 = __________________
Calculating in Scientific Notation (all calculations can be done using your calculator)
o Multiplication

Multiply the coefficients and add the exponents. Write the answer in scientific notation.
Example: (1.9 x 10-3) x (2.2 x 106) =
Multiple the coefficients: 1.9 x 2.2 = 4.2 (new coefficient)
Add the exponents: -3 + 6 = 3 (new exponent)
Write in scientific notation: 4.2 x 103
o
Division

Divide the coefficients and subtract the exponents. Write the answer in scientific notation.

Example: (1.9 x 10-3) =
(2.2 x 106)
Divide the coefficients: 1.9/2.2 = 0.86
Subtract the exponents: -3-(6) = -9
Write in scientific notation: 0.86 x 10-9
The new coefficient is still not in scientific notation (it is not a # between 1 & 10. So now what?
Convert 0.86 to scientific notation: 8.6 x 10-1
Substitute: 0.86 x 10-9 is the same as (8.6 x 10-1) x 10-9
Simplify: Now it is a multiplication problem again.
Add the exponents: -1 + -9 = -10
Therefore, 0.86 x 10-9 is the same as 8.6 x 10-10
Final answer; 8.6 x 10-10.
o
Scientific Notation on the Calculator

If your calculator has a ^ key: This is your raise to a power key. To enter 10 -3, you would type 10^-3.

If your calculator has a yx key or 10x key: These are your raise to a power keys. To enter 10 -3, you would
type either 10yx-3 or -3 10x.

When calculating using scientific notation on your calculator, it is easiest to put each separate # into
parentheses. Otherwise, your calculator may take things out of order.
Practice Problems:
Without Calculator
(5.5 x 106) x (1.111 x 10-1) =
(9.896 x 10-34)  (3.311 x 10-24) =
With Calculator