Chemistry….

The study of matter and the changes it undergoes

5 Major Areas of Chemistry

Analytical Chemistry- composition of substances.

Inorganic Chemistry- substances without carbon

Organic Chemistry- most substances containing carbon

Biochemistry- Chemistry of living things

Physical Chemistry- describes the behavior of chemicals (ex. stretching)

1.1 Chemistry’s impact on society:

Health & Medicine

Biotechnology

Energy & Environment

Materials & Technology

Agriculture- world’s food supply

The Environment- both risks and benefits involved in discoveries

Astronomy and Space Exploration

1.3 The Scientific Method

A logical approach to solving problems or answering questions.

Starts with observation- noting and recording facts

hypothesis- an educated guess as to the cause of the problem, or a proposed explanation

Scientific Method

“controlled” experiment- designed to test the hypothesis

only two possible answers

hypothesis is right

hypothesis is wrong

Generates data & observations from experiments.

Modify hypothesis - repeat the cycle

Observations

Hypothesis

Experiment

Cycle repeats many times.

The hypothesis gets more and more certain.

Becomes a theory

A thoroughly tested model that explains whythings behave a certain way.

Data Collection:

Qualitative vs. Quantitative

Examples??

1.4 What is Matter?

Matter is anything that takes up space and has mass. Everything, but energy!

Mass- amount of material or “stuff” in an object

Weight is due to gravity, and changes from location to location; mass is always constant.

Types of Matter

Substance- a particular kind of matter - pure; is uniform (all the same) and has a definite composition (examples are elements & compounds)

water; gold

Mixture- more than one kind of matter; has a variable composition

Substances

Elements- simplest kind of matter cannot be broken down any simpler

all one kind of atom.

Compounds are substances that can be broken down only by chemical methods When broken down, the pieces have completely

different properties than the original compound.

Made of two or more atoms, chemically combined (not physical blend!)

10 Most Common Elements

Mixtures

Physical blend of at least two substances; variable composition

Heterogeneous- mixture is not uniform in composition Chocolate chip cookie, gravel, soil.

Homogeneous- same composition throughout; called “solutions” Kool-aid, air, salt water

Every part keeps its own properties.

Compound or MixtureCompound Mixture

Made of one kind

of material

Made of more than

one kind of material

Made by a

chemical change

Made by a

physical change

Definite

composition

Variable

composition

Classification of Matter

Which is it?

ElementCompoundMixture

1.5 States of Matter

Solid- matter that can not flow, vibrational movement, low kinetic energy

Liquid- flows, more kinetic energy than solid,

Gas- flows, independent molecular motion, fills container, highest kinetic energy

Vapor- a substance that is currently a gas, but normally is a liquid or solid at room temperature. (water vapor)

States of Matter

Solid

Liquid

Gas

Definite

Volume?

YES

YES

NO

Definite

Shape?

YES

NO

NO

Temp.

increase

Small

variation

Small

variation.

Large

Variation

Com-

pressible?

NO

NO

YES

Solid Liquid Gas

Melt Evaporate

or Boil

CondenseFreeze

1.6 Properties

Words that describe matter (adjectives)

Physical Properties- a property that can be observed and measured without changing the composition. Examples- color, hardness, m.p., b.p.

Chemical Properties- a property that can only be observed by changing the composition of the material. Examples-volatile, flammable

Types of Properties:

• Extensive Properties: Dependent on quantity of matter

ex: mass, volume

• Intensive Properties: Independent of quantity

ex: density, boiling point

Physical Changes

A change that changes appearances, without changing the composition.

Ex. Boil, melt, cut, bend, split, crack

Boiled water is still water.

Chemical changes - a change where a new form of matter is formed.

Ex. Rust, burn, decompose, ferment

Filtration: Physical Separation

Distillation: Physical Separation

1.7 Measurement:International System of Units

The number is only part of the answer; it also need UNITS

The standards of measurement used in science are those of the Metric System

International System of Units

Metric system is now revised as the International System of Units (SI), as of 1960

Simplicity and based multiples of 10

10 base units (Know them…p.17 Table 1.3)

International System of Units

Sometimes, non-SI units are used

Liter, Celsius, calorie

Some are derived units

Made by joining other units

Speed (miles/hour)

Density (grams/mL)

Most Commonly Used Prefixes

Volume

The space occupied by any sample of matter

Calculated for a solid by multiplying the length x width x height

SI unit = cubic meter (m3)

Everyday unit = Liter (L), which is non-SI

Solid Volume Calculations

1 cm3 = 1 mL

1 dm3 = 1000 mL = 1 L

1 m3 = 1,000,000 mL = 1,000 L

Volume Measuring Instruments

Graduated cylinders

Graduated Pipet

Buret

Volumetric Flask

Syringe

Volume from Water Displacement

Units of Mass

Mass is a measure of the quantity of matter

Weight is a force that measures the pull by gravity- it changes with location

Mass is constant, regardless of location

Working with Mass

The SI unit of mass is the kilogram (kg), even though a more convenient unit is the gram

Measuring instrument is the balance or scale

Density

The formula for density is:

mass

volume

• Common units are g/mL, or possibly g/cm3, (or g/L for gas)

• Density is an intensive, physical property

Density =

Density and Temperature

What happens to density as the temperature increases?

Mass remains the same

Most substances increase in volume as temperature increases

Thus, density generally decreases as the temperature increases

Density and water

Water is an important exception

Over certain temperatures, the volume of water increases as the temperature decreases

Does ice float in liquid water?

Why?

Temperature

Heat moves from warmer object to the cooler object

Remember that most substances expand with a temp. increase

Basis for thermometers

Temperature scales

Celsius scale- named after a Swedish astronomer

Uses the freezing point(0 oC) and boiling point (100 oC) of water as references

Divided into 100 equal intervals, or degrees Celsius

Temperature scales

Kelvin scale (or absolute scale)

Named after Lord Kelvin

K = oC + 273

A change of one degree Kelvin is the same as a change of one degree Celsius

No degree sign is used

3 most common temp. scales

0 K is called absolute zero, and equals –273 0C

Conversion Formulas

K = 0C + 273

0F = 9/5(0C) + 32

0C = 5/9 (0F – 32)

1.8 Handling Numbers

How do you know what number to round your calculation answers to?

Significant figures: Determining which numbers are meaningful in a measurement or calculated quantity.

Working with Scientific Notation

Regardless of magnitude: all numbers can be expressed in formula

N x 10n

N = number between 1 and 9.9

N = positive or negative whole #

Ex: Express 0.000000456 in scientific notation

Answer: 4.56 X 10-7

Working with Scientific Notation

Do you know how to use your calculator when you have numbers in scientific notation????

Significant Figures

Significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated or uncertain.

**This is what you did when you read the volume from the glassware in lab.

Significant Figures

Rules: (page 24: Know 1-5)1.) Any non-zero digit is significantEx: 1.2345 = 5 sig. figs2.) Zeros between non-zero digits are significantEx: 1.2340567 = 8 sig. figs3.) Zeros to the left of the first non-zero digit

are not significantEx: 0.00123456 = 6 sig. figs

Significant Figures

4.) All zeros to the right of the decimal point are significant if they follow a non-zero number

Ex: 4.560000 = 7 sig. figs0.00100 = 3 sig. figs

5.) #’s without decimals present ambiguous info. Always use scientific notation to clear up problems.

Ex: 56,700 = ? sig. figs5.670 x 104 = 4 sig. figs

Significant Figures

Try These:

How many sig. figs?

A.) 1,245

B.) 1,245,000.15

C.) 0.00001

D.) 0.0004560

E.) 5.090 x 10-5

Sig. fig. calculations

Addition and Subtraction

The answer should be rounded to the same number of decimal placesas the least number in the problem

Addition & Subtraction

Examples:

1.) 4.56 cm + 3.1 cm=

Answer: 7.7 cm

2.) 0.4567 L –0.00654 L =

Answer: 0.4502 L

3.) 450 g + 1.04 g=

Answer: 451 g

Sig. Fig. calculations

Multiplication and Division

Round the answer to the same number of significant figures as the least number in the measurement

Multiply & Divide

Examples:

1.) 451 x 3.2 =

Answer: 1,400 or 1.4 x 102

2.) 0.0345/5.60 =

Answer: 0.00616 or 6.16 x 10-3

3.) 0.2030 x 12 =

Answer: 2.4

Significant Figures

**Exact numbers obtained from definition or by counting of objects can be considered to have an infinite # of sig. figs. They are not considered in the calculation. Only use measurements!!

Ex: 12 eggs in a dozen

Uncertainty in Measurements

Need to make reliable measurements in the lab

Accuracy – how close a measurement is to the true value

Precision – how close the measurements are to each other (reproducibility)

**Do multiple trials for experiments!!

Uncertainty in Measurements

Accepted value – correct value based on reliable references

Experimental value – the value measured in the lab

Error – the difference between the accepted and experimental values

Uncertainty in Measurements Error = accepted – experimental

Can be positive or negative

Percent error = the absolute value of the error divided by the accepted value, times 100%

| error |

accepted valuex 100%% error =

% Error

Can you think of an easier way to calculate this???

1.9 Dimensional Analysis

A way to analyze and solve problems by using units (or dimensions) of the measurement

Based on conversion factors

Conversion factors are fractions that are equal to one. Both the top and bottom measurements are identical; they just use different units.

Examples: 1ft/12 in

5,280 ft/1 mi

Dimensional Analysis

Give me some more examples!

Dimensional AnalysisExample Problems

A ruler is 12.0 inches long. How long is it in cm? ( 1 inch = 2.54 cm)

in meters?

A race is 10.0 km long. How far is this in miles?

Pikes peak is 14,110 ft. above sea level. What is this in meters?

Dimensional Analysis

Another measuring system has different units of measure:

6 ft = 1 fathom 100 fathoms = 1 cable length10 cable lengths = 1 nautical mile

3 nautical miles = 1 league

Problem: Jules Verne wrote a book 20,000 leagues under the sea. How far is this in feet?

Problem solving

1. ANALYZE

a) Identify the unknown

Both in words and what units it will be measured in. Write it down!

May need to read the question several times.

Problem Solving

b) Identify what is given (the “known”)

Write it down!

Unnecessary information may also be given

Problem solving

c) Plan a solution

• Break it down into steps.

• Look up needed information:

*Tables

*Formulas

*Constants, or conversion factors

*Choose an equation

Problem solving

2. CALCULATE

• doing the arithmetic

• use a calculator

Problem Solving

3. EVALUATE• Round off to proper # of sig. figs.

• Proper units? Need Scientific Notation?

• Check your work!

• Reread the question, did you answer it?

• Is it reasonable?

• Estimate an approximate answer

Converting Complex Units

Units expressed as a ratio or raised to a power

speed is: miles/hour

gas mileage is: miles/gallon

density is: g/cm3

Volume is: cm3, dm3, m3

Examples:

The density of silver is 10.5 g/cm3. Convert the density to kg/m3.

Answer: 1.05 x 104 kg/m3

The density of the lightest metal, lithium, is 5.34 x 102 kg/m3. Convert the density to g/cm3.

Answer: 0.534 g/cm3

Lastly….

What makes you perfect…or close to it?

Practice

Practice

& More Practice!!!