Basic Concepts of Matter

Chapter 1

2

Classifying Matter

Matter– Anything that has mass and occupies

space

Mass vs. Weight

Kinetic-Molecular TheoryAll matter consists of extremely tiny particles in constant motion

3

States of Matter• Solid

– -Closely packed together with a definite ridged shape– -Vibrate back and forth in a confined space – -the particles are not able to move past one another

• Liquid– -arranged randomly with a definite volume – -“fluid” – -the particles are not confined in space and can move

past one another

• Gas – -no definite shape or volume – -“fluid” – -the particles are far apart and move very rapidly

colliding with other particles and the container walls

4

Categorizing Matter

• Elements– -cannot be decomposed into simpler

form via chemical reactions– -found on periodic chart– -atoms are the smallest particle that

retains the characteristic properties of the elements

• Pure Substance– -consists of all the same substance (pure

gold, distilled water, etc)– -have a set of unique properties that

identifies it

5

Categorizing Matter

• Chemical Compounds– -two or more elements in a definite ratio

by mass with unique properties that separate them from the individual elements

– -can be decomposed into the constituent elements by chemical reactions

– -chemical compounds are held together by a chemical bound

•Water– hydrogen and oxygen•Carbon dioxide – carbon and oxygen

6

Categorizing Matter

• Mixtures– two or more pure substances in the

same container

homogeneous mixtures (solution) • -uniform composition throughout• -single phase• -cannot be separated easily

heterogeneous mixtures• -nonuniform composition thoughout

• -easily separated

7

Physical and Chemical Changes

Physical changes • changes in physical properties

• -melting, boiling, and cutting

Chemical changes

changing one or more substances into one or more different substances (chemical reaction)

• 2H2 + O2 -> 2H2O

8

Chemical and Physical Properties

Chemical Properties observed during a chemical reaction (change in chemical composition)– -rusting, oxidation, burning…– -chemical reactions

Physical Propertiesobserved without changing the substance’s composition– -allow for identification and classification– -density, color, solubility, melting point…

9

Classification of Physical Properties

Extensive Properties

depend on the amount of substance present

-mass or volume

Intensive Properties do not depend on the amount of substance present

-melting point, boiling point, density…

10

Density• Describes how compact a substance is

• Who “discovered” density?

• Density = mass/volume or D = m/V

11

Density

• Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

• 1cm3 = 1mL => 97.3cm3 = 97.3mL

12

Density

• Example: You need 125 g of a corrosive liquid for a reaction. What volume do you need?

• – liquid’s density = 1.32 g/mL

13

Units of Measure• Qualitative measures

– Nonnumerical experimental observations describing the identity of a substance in a sample

• Quantitative measures– Numerical experimental observations

describing how much of a particular substance is in a sample

System International d’Unites (SI)measurement system used in the sciences based on the metric system

14

Math Review and Measurements

• We make measurements to understand our environment:

– Human senses: sight, taste, smell, hearing…• Our senses have limits and are biased

– Instruments: an extension of our senses meter sticks, thermometers, balances• These are more accurate and precise

– All measurements have units• METRIC SYSTEM vs. British System

15

SI units

QuantityQuantity UnitUnit SymbolSymbol length meter m mass kilogram kg time second s current ampere A temperature Kelvin K amt. substance mole mol

16

Measurements in Chemistry

NameName SymbolSymbol MultiplierMultiplier• mega- M 106 (1,000,000)• kilo- k 103 (1,000)• deka- da 10• deci- d 10-1 (0.1)• centi- c 10-2 (0.01)

17

Measurements in Chemistry

NameName SymbolSymbol MultiplierMultiplier• Milli- m 10-3(0.001)• Micro- 10-6(0.000001)• Nano- n 10-9

• Pico- p 10-12

• Femto- f 10-15

18

Units of Measurement

Length– Measure of space in any direction– -derived unit cm– -standard length is a meter (m)

19

Units of Measurement

• Volume– Amount of space occupied by matter– -derived unit: mL or cm3 (cc)– -liter (L) is the standard unit

20

Units of Measurement

• Time (t)– Interval or duration of forward events– -standard unit is the second (s)

• Mass (m)

– measure of the quantity of matter in a body

• Weight (W)– measure of the gravitational – attraction (g) for a body (w=m x g)

1 kg = 1000g 1 kg = 2.2 lbs

1 g = 1000mg

21

Heat and Temperature

• Heat (q) vs. Temperature (T)

• 3 common temperature scales: • all use water as a reference • -Fahrenheit (F)• -Celsius (C)• -Kelvin (K)

22

Temperature Reference Points

Melting Point Boiling Point of water of water• 32 oF 212 oF• 0.0 oC 100 cC• 273 K 373 K

• Body temperature 37.0 oC or 98.6 oF– 37.2 oC and greater—sick– 41 oC and greater, convulsions– <28.5 oC hypothermia

Fahrenheit Celsius

Kelvin

23

Temperature Scales

273.15KC

or

15.273 C K

ipsRelationsh Centigrade andKelvin

o

o

24

32-F9

5

1.8

32FC

or

32Cx5

9 32C 1.8F

ipsRelationsh Centigrade and Fahrenheit

oo

o

ooo

Example: Express 548 K in Celsius degrees.

•Example: Convert 211 oF to degrees Celsius.

Fahrenheit to Centigrade Relationships

Temperature Scales

25

Precision and Accuracy

Precise

Accurate

Neither

Both

Accuracy how closely measured values agree with the correct value

Precisionhow closely individual measurements agree with each other

26

Mathematics in Chemistry

• Exact numbers (counted numbers)– 1 dozen = 12 things

• Measured Numbers– Use rules for significant figures– Use scientific notation when possible

• Significant figures– digits in a measured quantity that reflect

the accuracy of the measurement– -in other words, digits believed to be correct

by the person making the measurement– Exact numbers have an infinite number of

significant figures12.000000000000000 = 1 dozen

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Significant figures (numbers/digits)

Why use significant numbers?• -Calculators give 8+ numbers• -People estimate numbers differently• -Dictated by the precision (graduation) on

your measuring device• -In the lab, the last significant digit is the

digit you (the scientist) estimate

Scientists have develop rules to help determine which digits are “significant”

28

Rules for Significant Figures

1. All Nonzero numbers are significant!!!2. Leading zeroes are never significant

– 0.000357 3. Imbedded zeroes are always significant

– 3.0604

4. Trailing zeroes may be significant- You must specify significance by how the number

is determined or even written

– 1300 nails - counted or weighed?– 1.30000 –How many significant figures?

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• Multiplication & Division rule:

• The product retains the number of significant figures that corresponds to the multiplier with the smallest number of significant figure (sig. fig.)

5.22 tooff round

21766.5

31.2x

224.4

3.9 tooff round

89648.3

41.x

2783.2

Significant Figures

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Significant Figures

• Addition & Subtraction rule:

– Answer retains the smallest decimal place value of the addends.

6.95 tooff round

9463.6

20.2

423.1

3692.3

16.671 tooff round

6707.16

312.2

7793.18

31

Scientific Notation Express answers as powers of 10 by

moving the decimal place right (-) or left (+)

• Use of scientific notation is to remove doubt in the Significant Figures:

2000 2 x 103

15000 1.5 x 10?

0.004 4 x 10-3

0.000053 __.__ x 10?

In scientific notation, zeros are given if they are significant!!!

1.000 x 103 has 4 significant figures2.40 x 103 has ? significant figures

Key to Sig. Figs…Locating the decimal and deciding when to count the zeros!!!

32

Review #2• Units of Measure

– -length– -volume– -time– -mass– -weight

• Heat vs. Temperature– -three temperature scales– -temperature conversions

• Precision vs. Accuracy• Significant Figures• Scientific Notation

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Conversion Factors

• Length – 1 m = 39.37 inches– 2.54 cm = 1 inch

• Volume– 1 liter = 1.06 qt – 1 qt = 0.946 liter

• See Text for more conversion factors

34

Conversion Factors

Why do conversions?• -Scientists often must convert between

units

Conversion factors can be made for any relationship of units-Use known equivalence to make a fraction that can be used to “convert” from one unit to the other

35

Dimensional Analysis

• 1 inch = 2.54 cm

– Use the ratio to perform a calculation so the units will “divide out”

Example: Convert 60 inches to centimeters

in

cmor

cm

in

1

54.2

54.2

1

36

Dimensional Analysis

• Example: Express 9.32 yards in millimeters.

3 ft = 1 yard1 ft = 12 in or1 in = 2.54 cm100 cm = 1 m1000 mm= 1 m

in 12

ft 1ft 1

in 12

37

Dimensional Analysis

• Example: Express 627 milliliters in gallons. 1 liter = 1.06 qt

1 qt = 0.946 liter

38

Practice on your Own

1kg = 2.20 lbs

– Convert 25 g to lbs

– Convert 1 mL to Liters

– Convert 20 meters to cm

39

Area = length x width

Area is two dimensional thus units must be in squared terms:

• Express: 2.61 x 104 cm2 in ft2

Dimensional Analysis

40

Volume =length x width x height

• Volume is three dimensional thus units must be in cubic terms

Express: 2.61 ft3 in cm3

– this volume is used in medical measurements--cc

Dimensional Analysis

41

Percentage

• Percentage is parts per hundred of a sample

• % = x100

• Example: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?

g of substancetotal g of sample