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Welcome to CHEM 1101

Instructor: Dr. Muhannad Amer

Office Location: 44 staff Bldng

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• 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.

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•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

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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…

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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.

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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”

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•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

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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

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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!

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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

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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

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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

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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

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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.

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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

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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

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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

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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

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How many sig figs?

45.8736

.000239

.00023900

48000.

48000

3.982106

1.00040

6

3

5

5

2

4

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•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

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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

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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

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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

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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.

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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.

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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)

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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.

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Precision vs. Accuracy of Calculated Results

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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?

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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

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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.

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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

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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

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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

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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

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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 .

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73.20821.0102.0