Matter, Energy, and Measurement

Home Work-1.16, 1.17, 1.19, 1.21, 1.25, 1.27,1.33, 1.37, 1.38, 1.39, 1.41, 1.45, 1.47, 1.48, 1.49, 1.51, 1.57, 1.59, 1.69, 1.73

Chemistry

• Why?

• The universe consists of three things: Matter, Energy, Empty Space •

Matter

- is anything that has mass and takes up space.

Chemistry

•

Chemistry

- is the science that deals with matter: the structure and properties of matter and the transformations from one form of matter to another.

• Matter can undergo two types of changes

Changes

•

Chemical Change:

also called a chemical reaction, substances are used up (they disappear) and others are formed to take their place.

• Examples:

Changes

•

Physical Change-

changes in which the identity of a substance remains unchanged. (usually involves changes in state and/or appearance) • Examples:

Properties of Matter

• There are two types of properties: •

Chemical properties:

the chemical reactions a substance undergoes •

Physical Properties:

properties that do not involve chemical reactions such as: density, color, melting point, physical state

The Scientific Method

• The scientific method establishes a process that provides a foundation of evidence to back up all scientific information!!

• It has four parts!!!

The Scientific Method

•

Fact-

is a statement based on direct experience. It is a consistent and reproducible observation.

•

Hypothesis-

is a statement that is proposed without actual proof, to explain the

Fact

and/or relationships betweens different

Facts.

The Scientific Method

•

Tests-

Designed experiments or observations used to determine the validity of the

Hypothesis.

•

Theory-

the formulation of an apparent relationship of certain observed phenomena, which has been verified to some extent. (A

Hypothesis

with evidence from the

Tests

to support it!!)

Serendipity

•

Serendipity-

is chance observation. Accidental discovery.

Experimental Notation

• This system provides an easy way to express very large and/or very small numbers • It is based on the power of tens system • Examples:

Adding and Subtracting in EN

• Appendix 1 page A-1 • Numbers must have the same exponent • Add/subtract coefficients • Leave exponent as is • Examples:

Multiply and Divide in EN

• First multiply/divide Coefficients • Then

Add Subtract

exponents for multiplication, exponents for division.

• Examples:

Significant Figures

• Appendix 2 page A-5 • Defined as: The number of digits of a measured number that have uncertainty only in the last digit.

• Examples:

Rules

1) Nonzero digits are Always significant 2) Zeros at the beginning of a number are Never significant 3) Zeros between nonzero digits are Always significant

Rules (cont)

4) Zeros at the end of a number that has a decimal point are Always significant 5) Zeros at the end of a number with no decimal point May or May Not be Significant.

Sig. Figs. In E.N

• In E.N., the EN number must contain the same number of Sig. Figs. as the original number.

Sig. Figs in Functions

• Multiplication/Division- answer must have the same number of

sig figs

as the one with the

Fewest

sig figs.

Sig. Figs in Functions (cont)

• Addition/Subtraction- sig fig not relevant. The answer must contain the same number of

decimal places

the

Fewest

.

as the one with

Rounding

• If the digit to be dropped is 5,6,7,8 or 9, we round up • Otherwise, we simply drop the digit.

Defined/Counted Numbers

• Numbers that are defined, or counted, are treated as though they have infinite significant figures.

• Example:

Measurements

• A measurement consists of

TWO

parts: a number and a unit.

The units must always be present.

English Units

• Mass- pounds • Length- miles, inch, feet, etc • Volume- gallons, pints, quarts, etc • Time- Seconds

Metric Units

• Mass- gram, kilogram • Length- meter, kilometer • Volume- Liter, milliliter • Time- Seconds SI Units are typically the same as metric units, just more specific!

Metric System

• Establishes a base unit, and other units are related to that base unit by powers of 10.

Length

• English • 12 inches = 1 foot • 3 feet = 1 yard • 1760 yards = 1 mile • Metric – Base unit is the meter • 1 kilometer = 1000 meters • 1 millimeter = 0.001 meters

Metric Prefixes

Prefix giga mega kilo BASE UNIT deci centi milli micro nano d c m  n Symbol G M k 10 9 10 6 10 3 Value 1 billion 1 million 1 thousand 10 -1 10 -2 one-tenth one- hundreth 10 -3 one-thousandth 10 -6 one-millionth 10 -9 one-billionth

Volume

• Volume is space. • The volume of a substance is the amount of space it occupies • Base unit= Liter

Mass

• Mass is the quantity of matter in an object • Base Unit= Gram • There is a difference between Mass and Weight!!!!!

Time

• The base unit for time is Seconds.

• This is the same in all 3 systems!!

Temperature

• Base unit= Celsius or centigrade ( o C) • English system uses Fahrenheit ( o F) • The following can be used to convert between the two:

o F = (9/5) o C + 32 o C = (5/9) x ( o F – 32)

The SI temperature unit uses the Kelvin (K) K= o C + 273 0K = Absolute zero!!!

Comparisons

Length 1 in = 2.54 cm 1 m = 39.37 in 1 mile = 1.609 km Mass 1 oz = 28.35 g 1 lb = 453.6 g 1 kg = 2.205 lb

Comparisons

Volume 1 qt = 0.946 L 1 gal = 3.785 L Temperature 0 K = -459 o F = -273 o C 233K = -40 o F= -40 o C 273K = 32 o F= 0 o C 1 L = 33.81 fl oz 1 fl oz = 29.57 mL 1L = 1.057 qt 1 mL = 1 cc = 1 cm 3 310K = 98.6

o F= 37 o C 373K = 212 o F = 100 O C

Unit Conversions

• Factor-Label Method- we multiply and divide units with numbers using conversion factor.

• Using a conversion factor is the same thing as multiply by 1, only the units cancel out!!!!

Examples

• Convert 2,750 L into kL.

• Convert 120 lbs to grams.

More Examples

• Convert 3.5 miles to meters • Convert 900 g/ml to lbs/qt

States of Matter

• Matter can exist in three states- Solid, Liquid and Gas • Gases- no definite shape or volume, highly compressible • Liquids- no definite shape, but do have definite volume, only slightly compressible • Solids- Definite shape and volume, essentially noncompressible

Density

•

Density

- the mass of a substance per unit of volume.

– All states of matter have a density.

– When two liquids are mixed and one does not dissolve in the other, the one with the lower density floats on top!!

– Density is calculated by dividing the mass of a substance by its volume d= m/v

Density

• Density is a physical property and always has the same value at a given temperature • Density usually decreases as temperature increases because the mass remains the same while the volume increases • EXCEPTION: WATER!!! From 4-100 o C density increases, but from 0-4 o C it actually decreases!!!

Specific Gravity

• Numerically, it is the same as density, only it has no units.

• It is the density of a substance compared to water.

Energy

• •

Energy-

the capacity to due work

Kinetic Energy

is energy of motion • KE increases when either an object moves faster or a heavier object is moving.

Energy

•

Potential Energy (PE)

is stored energy • The PE possessed by an object arises from its capacity to move or cause motion.

Forms of Energy

• Mechanical, light, heat, and electrical energy – Kinetic energies possessed by all moving objects • Chemical energy and Nuclear energy – Potential energies

Chemical Energy

• The energy stored within chemical substances and given off when they take part in a chemical reaction.

• Examples:

Energy

• Various forms of energy can be converted from one to another • Example •

Law of Conservation of Energy-

Energy can neither be created nor destroyed.

Heat

•

Heat

is the form of energy that most frequently accompanies chemical reactions • HEAT AND TEMPERATURE ARE DIFFERENT!!!

• Heat is a form of energy, temperature is a measurement.

Heat (cont)

• The Heat unit is usually a calorie • calorie- the amount of heat necessary to raise the temperature of 1 gram of water by 1 o C.

• This is a small unit so kilocalories is typically used (1 kcal = 1000 calories

• Nutritionists use the word Calorie to mean the same thing as kilocalorie, so: 1 Cal = 1000 cal = 1 kcal • The SI unit is the Joule (J) 1 cal = 4.184 J

Specific Heat

• The amount of heat necessary to raise the temperature of 1 g of any substance by 1 o C.

• Each substance has its own specific heat • Is this a physical property or chemical property?

• Specific Heat can be used to calculate the amount of heat needed or used Amt of Heat Used = SH x m x (T 2 -T 1 )

Fun with SH!!!

Aluminum has a SH of .22, Iron has a SH of .11. How much heat is required to raise the temperature of 100 grams of each from 30 o C to 100 o C?