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1
Welcome to CHEM 1101
Instructor: Dr. Muhannad Amer
Office Location: 44 staff Bldng
2
• Beyond the chemistry theory of this class taking this class will enable you to:
•Apply knowledge to solve new problems.
• Analyze information you have gathered.
• Work with and delegate responsibility to others.
• Have confidence in yourself and your work.
• Be organized in your thoughts and actions.
• Ask and answer questions.
3
•The scientific method provides the method by which
scientists solve problems.
• Chemists use this method to understand matter at the
atomic or molecular level.
observation
hypothesis
prediction
experiment
Scientific Method
(explanation of
observation)
Carrying out experiment
A will prove the hypothesis
by giving result B
4
TheoryTheory
• When consistency is obtained, When consistency is obtained, hypotheseshypotheses become become a theorya theory
• Typically a fact of nature, often a math constant/number and unit.– Law of Conservation of Mass— “In a chemical
reaction matter is neither created nor destroyed.”– Speed of Light, E = mc2, Dalton’s Gas Law,
Universal Gas Constant, etc…
5
Theories
• Explains how nature behaves.– Newton’s Gravitational Theory: how an apple falls– Dalton’s Atomic Theory: atoms look like…– Darwin’s Theory of Evolution: we always change– Einstein's Theory of Relativity: light is constant
• Used to predict future observations.
66
What’s the Difference Between aLaw and a Theory?
• Laws: Very specific, “What will happen” often expressed in mathematical equations.
• Theories: Very general, “Why it will happen,” often includes many “Laws”
7
•Observations
•Observations can be quantitative(which involve numbers.)which involve numbers.)
• or qualitative (changes in color and physical state)changes in color and physical state)
•All measurements MUST consist of• a number and a unit!
•Example: charge of an electron is 1.60 x 10-19 coulombs
•Scientific notation
•1.60 x 10-19 = 0.000000000000000000160
8
Scientific Notation
number x 10n
1-9
integer
1.60 x 10 = 1.60 x 1 = 1.60
1.60 x 101 = 16.0
1.60 x 10-1 = 0.160
1.60 or 1.6 or 1.600 can be used
0
9
Are Units of Measurement that Important?
July 23rd, 1983: Gimli Glider, an Air Canada aircraft ran out of
fuel
Needed for trip: 22,300 kg of fuel
Used to fill plane: 22,300 pounds of fuel (10,115 kg !)
Not enough fuel!
10
Important SI (International system) base units
Quantity SI Base Unit
Length meter (m)
Mass kilogram (kg)
Time second (s)
Temperature Kelvin (K)
Amount mole (mol)
Volume = length3
1L = 1 dm3 = 1000 cm3 = 10-3 m3 = 1000 ml
1cm3 = 1ml
11
Common Prefixes used to adjust the size
of Base Units
Prefix Meaning Abbreviation
Exponential
Notation
deci- tenth of d 10-1
Mega- million M 106
kilo- thousand k 103
centi- hundredths of c 10-2
milli- thousandths of m 10-3
micro- millionths of µ 10-6
nano- billionths of n 10-9
pico- trillionths of p 10-12
12
The number obtained in measurement
is obtained using a measuring device that
introduces some degree of uncertainty
to this measurement and this must be
indicated.
Uncertainty in Measurement
Uncertainty in the measurement lies in the
last digit and is assumed to be +1 or -1
Recorded measurement of 0.0508 g
= 5.07 x 10-2 or 5.09 x 10-2 g
Actual mass is 0.0507 g or 0.0509 g
13
1. Digits from 1-9 are always significant.
Example: 26981 has 5 significant figures
Significant Figures
2. Zeros between two other significant digits are always
significant. Example: 1023 has 4 significant figures
3. One or more additional zeros to the right of both the
decimal place and another significant digit are significant.
Example: 5.00 and 500. both have 3 significant
figures
The recorded certain and the first uncertain digit or
estimated number of a measurement are called its
significant figures.Rules for Significant Figures
14
Significant Figures
4. Zeros used solely for spacing the decimal point
(placeholders)
are not significant.
Example: 0.000231 has 3 significant figures 5. The absence of a decimal point means terminal zeros
are NOT significant.
Example: 600 has 1 significant figure
6. Exact numbers have an infinite number of significant
figures. They are obtained via counting, e.g. 1 dozen
eggs, or by definition, e.g. the 2 in 2r. When used in
calculations, exact numbers do not limit the number of
significant figures.
15
How many significant figures are in each of the following measurements?
24 mL 2 significant figures
3001 g 4 significant figures
0.0320 m3 3 significant figures
6.4 x 104 molecules 2 significant figures
560 kg 2 significant figures
1.8
1616
Practice—Write the Following in Scientific Notation, Continued
123.4 = 1.234 x 102
145000 = 1.45 x 105
25.25 = 2.525 x 101
1.45 = 1.45 x 100
8.0012 = 8.0012 x 100
0.00234 = 2.34 x 10-3
0.0123 = 1.23 x 10-2
0.000 008706 = 8.706 x 10-6
17Tro's "Introductory Chemistry",
Chapter 2
17
Practice—Write the Following in Standard Form, Continued
2.1 x 103 = 2100
9.66 x 10-4 = 0.000966
6.04 x 10-2 = 0.0604
4.02 x 100 = 4.02
3.3 x 101 = 33
1.2 x 100 = 1.2
1818
Determine the Number of Significant Figures,
• 12000
• 120.
• 12.00
• 1.20 x 103
• 0.0012
• 0.00120
• 1201
• 1201000
2
3
4
3
2
3
4
4
19
How many sig figs?
45.8736
.000239
.00023900
48000.
48000
3.982106
1.00040
6
3
5
5
2
4
6
•All digits count
•Leading 0’s don’t
•Trailing 0’s do
•0’s count in decimal form
•0’s don’t count w/o decimal
•All digits count
•0’s between digits count as well as trailing in decimal form
20
Significant Figures
1.8
Addition or Subtraction
The answer cannot have more digits to the right of the decimalpoint than any of the original numbers.
89.3321.1+
90.432 round off to 90.4
one significant figure after decimal point
3.70-2.91330.7867
two significant figures after decimal point
round off to 0.79
21
Significant Figures
1.8
Multiplication or Division
The number of significant figures in the result is set by the original number that has the smallest number of significant figures
4.51 x 3.6666 = 16.536366 = 16.5
3 sig figs round to3 sig figs
6.8 ÷ 112.04 = 0.0606926
2 sig figs round to2 sig figs
= 0.061
22
Significant Figures
1.8
Exact Numbers
Numbers from definitions or numbers of objects are consideredto have an infinite number of significant figures
The average of three measured lengths; 6.64, 6.68 and 6.70?
6.64 + 6.68 + 6.703
= 6.67333 = 6.67
Because 3 is an exact number
= 7
23
Multiplying and Dividing Significant Figures
22.37 x 3.10 x 85.75
4 sig. figs 3 sig. figs 4 sig. figs
Least number of significant
figures dictates the number
of significant figures to be
stated in the calculated answer
= 5946.50525 Seen on
calculator
but not to be
recorded as
the answer
= 5950
5946.505259 sig. figs
59503 sig. figs
Rounding 5 round up < 5 round down
Calculated results are never more reliable than
the measurements they are obtained from.
24
Adding and Subtracting Significant Figures
3.76 + 14.83 + 2.1
2 dec. places
2 dec. places
1 dec. place
= 20.69 Seen on calculator but
not to be recorded as
the answer.
Least number of decimal places
dictates the number of decimal
places to be stated in the
calculated answer.
20.692 dec. places
20.71 dec. place
= 20.7
Rounding to one dec. place
Calculated results are never more reliable than the
measurements they are obtained from.
25
Addition (subtraction) with Multiplication (Division)
732.11 + 6.3
760.00do addition (subtraction) first
732.11 + 6.3 =2 decimalplace
1 decimalplace
738.41 NEVER round
intermediate results for
multistep calculationsdo division (multiplication) last
738.4
760.00
4 sig fig
5 sig fig
738.41
760.00= 0.971592105
(4 sig fig)Answer: 0.9716
(Not to be recorded as the answer)
(738.4)
26
Examples of RoundingFor example you want a 4 Sig Fig number
4965.03
780,582
1999.5
0 is dropped, it is <5
8 is dropped, it is >5; Note you must include the 0’s
5 is dropped it is = 5; note you need a 4 Sig Fig
4965
780,600
2000.
27
Precision vs. Accuracy of Calculated Results
28
Precision = reproducibility
How much of a clone are you?
Standard values
Sugar content: 54 grams
pH: 2.6
Accuracy = Closeness of measured value to standard value
How do you measure
up?
29
Dimensional Analysis
A problem-solving method that uses the fact that any
number or expression can be multiplied by one without
changing its value.
Unit factors may be made from any two terms that
describe the same or equivalent "amounts" of what we
are interested in. 1 inch = 2.54 centimeters
Unit factors
30
Steps for Using Dimensional Analysis
Steps:
1.Identify what units are required, what units have
been given.
2. State the equivalent of these units.
3. Multiply the given data and its units by the
appropriate unit factors so that only the
desired units are present at the end.
31
Notice that the unit factor was chosen that allowed the
units required to remain while the other cancels
during the calculation.
1 inch = 2.54 centimeters
Unit factors
Example: How many centimeters are in 6.00 inches?
Units required: centimeters
Units given : inches
32
Kelvin ( K ) - The “Absolute temperature scale”
At absolute zero and only has positive values.
Celsius ( oC ) - Commonly used scale around the
world
and in laboratories.
Fahrenheit ( oF ) - Commonly used scale in
America for
weather reports.
Temperature Scales and Interconversions
T (K) =T (oC) + 273.15
T (oC) = T (K) − 273.15
T (oF) = 9/5 T (oC) + 32
T (oC) = 5/9 T (oF) - 32
33
Density
Density is the mass per unit volume of a substance and has
compound units of grams per cubic centimeter (g/cm3)
Example: Calculate the density of an object that has a
volume of 64 cm3 and a mass of 34g.
Density = mass volume
Solution:
Density = 34g 64cm3
= 0.53g/cm3
34
What is the mass, in grams, of 1.00 gallon of
water ?
The density of water is 1 g/mL (1 ml of water = 1g)
1.00 gal x
=
1 gal = 4 qt
4 qts
1 galx 1 L
1.057 qts
1 g = 1 mL1 L = 1000 ml
x 1000 mL1 L
x 1 g
1 mL3 sig figs
= 3.78 x 103 g3784.295
Solution
Units given: gallon, g/ml Units required: g
All equivalent values are EXACT numbers and do not limit
the
number of significant figures in the answer.
given required
calculator
= 3780 g3 sig figs
1.057 qt = 1 L
35
Quiz 1
• Perform the following mathematical operations and express the result to the correct number of sf.
• The volume of a diamonds is found to be 2.8ml . What is the mass of the diamond in carats ?
1 carat = 0.200g . The density of diamond is 3.51 g/cm3 .
01.1
73.20821.0102.0