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Chapter 1Matter,Measurement, and Problem

Solving

Classification of Matter

- matter is anything that has mass and occupies

space - we can classify matter based on whether its

solid, liquid, or gas

- Fixed keeps shape when placed in a container
- Indefinite takes the shape of the container

Solids

- the particles in a solid are packed close

together and are fixed in position - though they may vibrate
- incompressible
- the inability of the particles to move around

results in solids retaining their shape and

volume when placed in a new container, and

prevents the particles from flowing

Crystalline vs. Amorphous solids

- some solids have their particles arranged in an

orderly geometric pattern we call these

crystalline solids - salt and diamonds

- some solids have their particles randomly

distributed without any long-range pattern we

call these amorphous solids - plastic
- glass
- charcoal

Liquids

- the particles in a liquid are closely packed, but

they have some ability to move around - Incompressible
- take the shape of their container and to flow
- however, they dont have enough freedom to escape

and expand to fill the container

Gases

- in the gas state, the particles have complete

freedom from each other - the particles are constantly flying around,

bumping into each other and the container - Compressible

- because there is a lot of empty space, the

particles can be squeezed closer together

therefore gases are compressible - because the particles are not held in close

contact and are moving freely, gases expand to

fill and take the shape of their container, and

will flow

Classification of Matterby Composition

- made of multiple types of particles
- samples may show different intensive properties

- made of one type of particle
- all samples show the same intensive properties

Classification of Pure Substances

- made of one type of molecule, or array of ions
- molecules contain 2 or more different kinds of

atoms

- made of one type of atom (some elements found as

multi-atom molecules in nature) - combine together to make compounds

Classification of Mixtures

- made of multiple substances, whose presence can

be seen - portions of a sample have different composition

and properties

- made of multiple substances, but appears to be

one substance - all portions of a sample have the same

composition and properties

Separation of Mixtures

- separate mixtures based on different physical

properties of the components - Physical change

Changes in Matter

- changes that alter the state or appearance of the

matter without altering the composition are

called physical changes - state changes
- boiling / condensing
- melting / freezing
- subliming

Physical change vs. Chemical change

- changes that alter the composition of the matter

are called chemical changes - during the chemical change, the atoms that are

present rearrange into new molecules, but all of

the original atoms are still present - rusting
- processes that release lots of energy
- burning

Burning propane gas

Energy

- changes in matter, both physical and chemical,

result in the matter either gaining or releasing

energy - energy is the capacity to do work
- work is the action of a force applied across a

distance - a force is a push or a pull on an object
- electrostatic force is the push or pull on

objects that have an electrical charge

Energy

- kinetic energy is energy of motion
- motion of the atoms, molecules, and subatomic

particles - potential energy is energy that is stored in the

matter - due to the composition of the matter and its

position in the universe - chemical potential energy arises from

electrostatic forces between atoms, molecules,

and subatomic particles - Law of Conservation of Energy
- energy is neither created nor destroyed. It is

converted from one form to another

Temperature

- measure of the average amount of kinetic energy
- higher temperature larger average kinetic

energy - heat flows from the matter that has high thermal

energy into matter that has low thermal energy - until they reach the same temperature
- heat is exchanged through molecular collisions

between the two materials

Temperature Scales

- Fahrenheit Scale, F
- used in the U.S.
- Celsius Scale, C
- used in all other countries
- Kelvin Scale, K
- absolute scale
- no negative numbers
- directly proportional to average amount of

kinetic energy - 0 K absolute zero

Fahrenheit vs. Celsius

- a Celsius degree is 1.8 times larger than a

Fahrenheit degree - the standard used for 0 on the Fahrenheit scale

is a lower temperature than the standard used for

0 on the Celsius scale - the size of a degree on the Kelvin scale is the

same as on the Celsius scale - so 1 kelvin is 1.8 times larger than 1F
- the 0 standard on the Kelvin scale is a much

lower temperature than on the Celsius scale

TC (TF 32)

TF 1.8 TC 32

1.8

K C 273.15

Example

- The melting point of gallium is 85.6oF. What is

this temperature on - Celsius scale
- Kelvin scale

The Standard Units

- Scientists have agreed on a set of international

standard units for comparing all our measurements

called the SI units - Système International International System

Quantity Unit Symbol

length meter m

mass kilogram kg

time second s

temperature kelvin K

Common Prefix Multipliers in the SI System

All units in the SI system are related to the

standard unit by a power of 10 The power of 10 is

indicated by a prefix multiplier The prefix

multipliers are always the same, regardless of

the standard unit Report measurements with a unit

that is close to the size of the quantity being

measured

Prefix Symbol Decimal Equivalent Power of 10

mega- M 1,000,000 Base x 106

kilo- k 1,000 Base x 103

deci- d 0.1 Base x 10-1

centi- c 0.01 Base x 10-2

milli- m 0.001 Base x 10-3

micro- m or mc 0.000 001 Base x 10-6

nano- n 0.000 000 001 Base x 10-9

pico p 0.000 000 000 001 Base x 10-12

Common Units and Their Equivalents

Mass

1 kilogram (km) 2.205 pounds (lb)

1 pound (lb) 453.59 grams (g)

1 ounce (oz) 28.35 grams (g)

Volume

1 liter (L) 1000 milliliters (mL)

1 liter (L) 1000 cubic centimeters (cm3)

1 liter (L) 1.057 quarts (qt)

1 U.S. gallon (gal) 3.785 liters (L)

Length

1 kilometer (km) 0.6214 mile (mi)

1 meter (m) 39.37 inches (in.)

1 meter (m) 1.094 yards (yd)

1 foot (ft) 30.48 centimeters (cm)

1 inch (in.) 2.54 centimeters (cm) exactly

Volume

- Derived unit
- any length unit cubed
- Measure of the amount of space occupied
- SI unit cubic meter (m3)
- Commonly measure solid volume in cubic

centimeters (cm3) - Commonly measure liquid or gas volume in

milliliters (mL)

Mass Volume

- two main physical properties of matter
- mass and volume are extensive properties
- the value depends on the quantity of matter
- extensive properties cannot be used to identify

what type of matter something is - Large iceberg and small ice cube
- even though mass and volume are individual

properties, for a given type of matter they are

related to each other!

Density

- Ratio of massvolume is an intensive property
- value independent of the quantity of matter
- Solids g/cm3
- 1 cm3 1 mL
- Liquids g/mL
- Gases g/L
- Volume of a solid can be determined by water

displacement Archimedes Principle - Density solids gt liquids gtgtgt gases
- except ice is less dense than liquid water!

- For equal volumes, denser object has larger mass
- For equal masses, denser object has smaller

volume - Heating an object generally causes it to expand,

therefore the density changes with temperature

A Measurement

- the unit tells you what standard you are

comparing your object to - the number tells you
- what multiple of the standard the object

measures - the uncertainty in the measurement
- scientific measurements are reported so that

every digit written is certain, except the last

one which is estimated - If the length is reported as 3.26 cm,
- the digits 3 and 2 are certain (known).
- the final digit, 6, is estimated (uncertain).
- all three digits (2, 7, and 6) are significant,

including the estimated digit.

Known Estimated Digits

For the following volume readings, what would be

measured values?

E.g l8. . . . l . . . . l9. . . . l . . . .

l10. . cm What is the length of the line? 1)

9.2 cm 2) 9.13 cm 3) 9.19 cm

Uncertainty in Measured Numbers

- accuracy is an indication of how close a

measurement comes to the actual value of the

quantity - precision is an indication of how reproducible a

measurement is

Accuracy, Precision, and Significant Figures

- Significant figures The number of meaningful

digits in a measured or calculated quantity. They

come from uncertainty in any measurement. - Generally the last digit in a reported

measurement is uncertain (estimated). - Exact numbers and relationships (7 days in a

week, 30 students in a class, etc.) effectively

have an infinite number of significant figures.

Examples

- Classify each of the following as (1) exact or

(2) measured - numbers.
- A.__Gold melts at 1064 C.
- B.__1 yard 3 feet
- C.__The diameter of a red blood cell is 6 x 10-4

cm. - D.__There are 6 hats on the shelf.
- E.__A can of soda contains 355 mL of soda.

Accuracy, Precision, and Significant Figures

- Rules for counting significant figures

(left-to-right) - Zeros in the middle of a number are like any

other digit they are always significant. - 4.803 cm 4 sf
- 2. Rules for counting significant figures

(left-to-right) - Zero at the beginning of a number are not

significant (placeholders). - 0.00661 g 3 sf or 6.61 x 10-3 g
- 3. Zeros at the end of a number and after the

decimal point are always significant. - 55.220 K 5 sf
- 4. Zeros at the end of a number without a written

decimal point are ambiguous and should be avoided

by using scientific notation - if 150 has 2 sig. figs. then 1.5 x 102
- but if 150 has 3 sig. figs. then 1.50 x 102

Rounding Numbers

- If the first digit you remove is 4 or less, it

and all following digits are dropped from the

number - 5.664 425 5.664 (4 s.f)
- If the digit you remove is 5 or greater, the

last digit of the number is increases by 1 - 5.664 525 5.665 (4 s.f)
- Sometimes, a calculator displays a small whole

number. To give an answer with the correct

number of significant figures, significant zeros

may need to be written after the calculator

result. - E.g 8.00 2.00 4 ?

4.00 - 3 s.f 3 s.f calculator

2 zeros are needed -

result to give 3 s.f

Significant figures in calculation

- When multiplying or dividing
- the final answer must have the same number of

significant figures as the measurement with the

fewest significant figures. - Example
- 110.5 x 0.048 5.304 5.3

(rounded) - 4 SF 2 SF calculator

2 SF - When adding or subtracting
- the final answer must have the same number of

decimal places as the measurement with the fewest

decimal places. - 25.2 one decimal place
- 1.34 two decimal places
- 26.54 calculated answer
- 26.5 final answer with one decimal

place

Both Multiplication/Division and

Addition/Subtraction with Significant Figures

- when doing different kinds of operations with

measurements with significant figures, do

whatever is in parentheses first, evaluate the

significant figures in the intermediate answer,

then do the remaining steps - 3.489 (5.67 2.3)
- 2 dp 1 dp
- 3.489 3.37 12
- 4 sf 1 dp 2 sf 2 sf

Metric Equalities

- An equality
- states the same measurement in two different

units. - can be written using the relationships between

two metric units. - Example 1 meter is the same as 100 cm and 1000

mm. - 1 m 100 cm
- 10-2 m 1cm
- 1 m 1000 mm
- 10-3m 1 mm

Conversion Factors

- A conversion factor is
- obtained from an equality.
- E.g Metric U.S system
- Equality 1 in. 2.54 cm
- written as a fraction (ratio) with a numerator

and denominator. - inverted to give two conversion factors for every

equality. - 1 in. 1 2.54 cm
- 2.54 cm 1 in.

Arrange conversion factors so given unit

cancels Arrange conversion factor so given unit

is on the bottom of the conversion factor

Using Two or More Factors

- Often, two or more conversion factors are

required to obtain the unit needed for the

answer. - Unit 1 Unit 2 Unit 3
- Additional conversion factors are placed in the

setup problem to cancel each preceding unit. - Given unit x factor 1 x factor 2

needed unit - Unit 1 x Unit 2 x Unit 3

Unit 3 - Unit 1 Unit 2

Example

- If a ski pole is 3.0 feet in length, how long is

the ski pole in m? - Convert 288.0 cm to yard
- Convert 9255 cm3 to gallons

1.8 Density

- Density
- compares the mass of an object to its volume.
- is the mass of a substance divided by its volume.
- Density Expression
- Density mass g or g g/cm3

- volume mL cm3
- Note 1 mL 1 cm3
- Can we use density as a conversion factor to

calculate mass or volume?

Examples

- What is the density (g/cm3) of 48.0 g of a metal

if the level of water in a graduated cylinder

rises from 25.0 mL to 33.0 mL after the metal is

added? - A drop of gasoline has a mass of 22.0 mg and a

density of 0.754 g/cm3. What is its volume in

Liters?

33.0 mL

25.0 mL

object