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 exponents for multiplication, Subtract 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 as the one with the Fewest.

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 Symbol Valuegiga G 109 1 billion

mega M 106 1 million

kilo k 103 1 thousand

BASE UNIT

deci d 10-1 one-tenth

centi c 10-2 one- hundreth

milli m 10-3 one-thousandth

micro 10-6 one-millionth

nano n 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 (oC)

• English system uses Fahrenheit (oF)

• The following can be used to convert between the two:

oF = (9/5)oC + 32oC = (5/9) x (oF – 32)

The SI temperature unit uses the Kelvin (K)

K= oC + 273

0K = Absolute zero!!!

Comparisons

Length Mass

1 in = 2.54 cm 1 oz = 28.35 g

1 m = 39.37 in 1 lb = 453.6 g

1 mile = 1.609 km 1 kg = 2.205 lb

Comparisons

Volume Temperature

1 qt = 0.946 L 0 K = -459oF = -273oC

1 gal = 3.785 L 233K = -40 oF= -40oC

1 L = 33.81 fl oz 273K = 32oF= 0oC

1 fl oz = 29.57 mL 310K = 98.6oF= 37oC

1L = 1.057 qt 373K = 212oF = 100OC

1 mL = 1 cc = 1 cm3

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 4-100oC density increases, but from 0-4oC 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 1oC.

• 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 1oC.

• 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 (T2-T1)

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 30oC to 100oC?

Fun with SH!!!