Chapter 1web.gccaz.edu/~jaszi38221/2014/Fall/CHM 151... · 2 Scientific Method •A systematic way to conduct research – Define the problem – Perform experiments and make observations

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Chapter 1

Chemistry: Matter and Measurement

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Why Study Chemistry? aka Why are you here?

Huh?

Why Study Chemistry? aka Why are you here?

• What’s the difference?

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Why Study Chemistry? aka Why are you here?

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Natural Non-Natural

Why Study Chemistry? aka Why are you here?

• Medicine

• Energy sources

• Materials

• Technology

• Food

• Agriculture

• Cooking

• Cars, clothes, computers, sporting goods

2

Scientific Method

• A systematic way to conduct research

– Define the problem

– Perform experiments and make observations

– Record data from experiment

– Interpret data

– Formulate/test hypothesis

– Summarize data

• (devise theory or law to explain phenomena)

The Periodic Table

On the periodic table, elements are located adjacent to other elements with similar general physical properties

– Metals – usually solid at room temperature, lustrous,

conduct electricity, malleable

– Non-metals – solid, liquid, or gas at room temperature.

Solids are brittle. Do not conduct electricity.

– Metalloids or Semi-metals – shares properties of both

metals and non-metals. Lustrous like metals, but brittle like non-metals

Antimony. Looks like a metal, but

is not a good conductor of

electricity

The Periodic Table

(Semi-metals)

+1 +2 -1-2-3

Elements and the Periodic Table

Groups (Families): columns.

Divides elements into more closely related properties

Periods: rows (we’ll talk more

about this relationship in later chapters)

Elements and the Periodic Table

Alkaline earth metals (less reactive than alkali metals)

Alkali metals (react violently with water)

Noble gases (very unreactive)

Halogens colorful, corrosive gases, liquids and solids

Elements and the Periodic Table

Transition Metals

“Other Stuff”

Inner Transition Metals

3

Relative Abundances of elements

In the Human Body

On Earth

Chemistry and the Elements

Elements are not distributed evenly within the universe.

• H is 75% of mass of universe

• O, Si are 75% of the earth’s crust

• O, C, H are more than 90% of the human body

– Fats

– Proteins (with N)

– Carbohydrates

– All contain O, C, H……..

Some Chemical Properties of the Elements

• Intensive properties: do not depend on amount of material (freezing point, boiling point)

• Extensive properties: depend on amount of material (volume, mass)

• Physical properties: color, temperature, texture, odor, boiling point, melting point, conductance

• Chemical properties: reactivity; how an element reacts with other elements. (flammability, enthalpy of combustion)

Physical and Chemical Changes

Physical Changes: changing from one state to another. How elements are arranged does not change. How atoms are chemically bonded does not change.

Chemical Changes: changes how elements are arranged with in matter. How atoms are chemically bonded does change.

States of Matter

17

Physical or chemical change?

Physical and Chemical Changes

18

Physical or chemical change?

What has changed?

4

Which elements should I know?Need to know names and symbols:

• Elements 1-36

• Groups 1, 2, 17, 18• Ag, Cd, Pt, W, Au, Hg, Sn, Pb, Sb, Bi

SI Base Units

Remove the “kilo” when other

prefixes are used

Each physical quantity measured has its own SI base unit

Volume meters cubed m3

liters l or L

These are NOT the same

m3 = 1000L, cm3 = mL

SI Prefixes units

Need to be memorized

Base unit

Measurements

• Mass

– kg, g, mg, mm

• Length

– m, km, cm, mm, mm, nm

• Temperature

– K, oC, oF

– K = oC + 273.15

– F = (9/5 * oC) + 32oF

– C = 5/9 * (oF – 32)

• Volume

– L = dm3, mL = cm3

Note: “micro” is a Greek mu

Converting Prefixes

s10

1ks

1Ms

s10 1Ms

3

6

Many times, we need to convert from one prefix to another.

For example, how many kiloseconds (ks) are in 1 megasecond (Ms)?

mega = 106 kilo = 103

The “fool-proof” method

ks10 s10

1ks

1Ms

s10 1Ms 3

3

6

Set-up the equation & cancel units

Converting Prefixes

Short Cut: Move-the-decimal method

mega = 106 kilo = 103

Subtract the exponents. This

determines how many places to move the decimal

6-3 = move 3 spaces

Larger to smaller unit → move to the right

Smaller to larger unit → move to the left

1 ks = 1 0 0 0 ks1Ms =

= 1000 ks = 103 ks

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Uncertainty in Measurement

• There is always uncertainty in measurement!

• Measurements are never exact.

• Different measuring devices have differing levels of precision and accuracy.

How close repeated

measurements are in respect to one

another. Reproducibility.

How close measurements are

to the “actual” value. We use

reference standards.

Precision vs. Accuracy

Describe each in terms of accuracy and precision

X

X

X

X

X

X

X

XX

XXXXXX

XXXX X

Uncertainty in Measurement

• The temperature outside is 71.21oF. Several thermometers made by one manufacturer record the temperature as 67.8, 68.2, 67.2, 67.6, and 68.0oF.

• How would you describe this data in terms of accuracy and precision?

• Would the price of the thermometers change your answer?

– 99¢ vs. $999

Significant Figures

• What are significant figures?

• Why are they important?

• When do you need to worry about them?

• How many decimal places can you use on the rulers below (shown in cm)?

28http://webphysics.iupui.edu/webscience/courses/chem101/chem101/images/ruler.10.gif

Rules for Significant Figures

• All non-zero digits are significant (335 cm).

• Zeroes in the middle of a number are significant (3406 mg).

• Zeroes at the beginning of a number are NOT significant. Called leading zeros. (0.000345 km).

• Zeroes at the end of a number and after the decimal point are significant. Called following zeros. (43.21000 g).

• Zeroes at the end of a number and before the decimal point may or may not be significant (6890 ft). You will have to look at the measurement to determine this.Ambiguous. Not sure if the measurement to the 10’s or 1’s place.

Rules for Significant Figures

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Scientific Notation and Powers of tensSometimes certain numbers, especially large or small numbers, are awkward to

write

0.000000000000013cmA better way to write this is by using scientific notation

How to write using scientific notation

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1.3

•Remove all leading zeros

•Write all numbers after the first non-zero number

•Place the decimal after the first non-zero number

•Add “x10” after the number

•Count how many spaces the decimal moves from its original location

•If the decimal point moves to the right, the exponent will be negative

•If the decimal point moves to the left, the

exponent will be positive

1.3 x 10

1.3 x 10-14

Scientific Notation and Powers of tens

Write the following numbers in scientific notation and

place in order of increasing value:

32

•1 x 10-6

•3 x 10-5

•8 x 105

•700000

•10

•0.001

•0.00002

•1 x 104

Scientific Notation Practice

• Addition and Subtraction– Combine numbers with same exponent and add numbers

– 7.4 x 103 + 2.1 x 103 = 9.5 x 103

• Multiplication– Add exponents and multiply numbers

– 8.0 x 104 * 5.0 x 102 = 40 x 106 = 4.0 x 107

• Division– Subtract exponents and divide numbers

– 6.9 x 107 / 3.0 x 10-5 = 6.9/3.0 x 107-(-5) = 2.3 x 1012

How to Use Your Calculator

• To enter 1.00 x 104 in your calculator, DO NOT enter “x” or “10”.

• Instead, use the exponent key (“EXP” or “EE”).

• Press: 1.00 “EXP” (or “EE”) 4

• Rule of thumb, “EE” and “EXP” basically mean x10 within scientific notation.

Significant Figures

•1.45

•0.38

•0.0670

•301.9

•072.8

•1.0

•44.20

•278

•1098.40

•0.00041560

•98.76

•100

•190

•1.90 x 103

•1063

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How many significant figures are in the following assuming they are measurements?

Significant Figures in Calculations

• Don’t round for sig. figs. until the END of all calculations. Keep extra sig figs between steps.

• Multiplication and division: report to the least number of significant figures.

– Ex: 2.8 x 4.5039 2 sig. figs. in answer

= 12.61092 13

• Addition and subtraction: report to the least number of decimal places.

– Ex: 2.097 – 0.12 2 digits after decimal

= 1.977 1.98

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Significant Figures Practice

• Calculate the following using the correct number of sig figs:

1.67890 x 56.32

• 94.56

9.0210 + 856.1

• 865.1

(6.02 + .5) x (3.14 + 2.579)

Be careful when you must use both rules

in one calculation.

Use one rule at a time

Dimensional Analysis

•1 dozen eggs = 12 eggs

•1 inch = 2.54 cm

•3 feet = 1 yard

•1 Mm = 1x106 m

1dozen 1

eggs 12

38

1eggs 12

dozen 1

1cm 2.54

inch 1

1yard 1

feet 3

11inch

cm 2.54

1feet 3

yard 1

or

or

or

We use conversion factors convert from one unit to another unit.

1m 10

Mm 16

1Mm 1

m 106

or

Dimensional Analysis

in 12

ft 1in 12.77

How many feet are in 12.77 in?

1.064 ft

How many feet are in 562.0 cm?

in 12

ft 1

cm 2.54

in 1 cm 0.562 18.44 ft

How many inches are in 52 km?

More than one possible way to solve this problem:

• km → m → cm → in

• km → mi → ft → in

Dimensional Analysis

• A cop clocks you going 4.47 x 10-2 km/s. Were you speeding? If so, are you going to jail?

• 1 mi = 1.61 km

• 60 s = 1 min

• 60 min = 1 hr

Density

V

m density

•If a steel ball bearing weighs 54.2 grams and has a volume of 6.94 cm3, what is its density?

•If a steel beam is measured to have a volume of 94390 cm3, how much does it weigh?

Is density intensive or extensive?

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Volumes (cubed) and Areas (squared)

If 1 in = 2.54 cm…does 1 in3 = 2.54 cm3 ?

So….(1 in)3 = (2.54 cm)3

(1)3(in)3 = (2.54)3(cm)3

1 in3 = 16.387 cm3

1 in or 2.54 cm