Integrated Physical Science Summer Assignment2016-2017 School Year

Welcome to Advanced Physical Science (PreAP)! We are eagerly anticipating a great year in science. In order to ensure the best start for each student next fall, we have prepared a Summer Assignment that reviews basic metrics, math review, chemistry and physics concepts previously studied. Most of the information in the packet will be familiar to you, but is designed to strengthen your foundation in chemistry and physics and ensure that all students are on a relatively even plane. It is important for everyone to come prepared and ready for the first day of class. While we will have a brief review of these concepts at the start of school, extensive remediation is not an option as the course only allows for a semester of each class.

The Summer Assignment is due the FIRST MONDAY OF SCHOOL – NO EXCEPTIONS! There will be a test covering these basic concepts included in this packet the first full week of school. Copying this assignment or simply trying to fill in the blanks will put you behind all semester!

We hope that you are gearing up for an exciting, challenging, and rewarding academic course. It will build your knowledge base, enhance your work habits and your organizational skills, and you will grow as an independent, critical thinking, learner. You are all certainly fine students, and along with motivation and hard work, you will find this class a successful and rewarding experience.

Finally, we recommend that you find a “study buddy” and spread out this assignment over time. Both Chemistry and Physics takes time to process ad grasp the level necessary for success. Taking this course requires dedication and is a great investment in your education, so prepare yourself and arrive ready to learn.

Video tutorials for Chemistry will be available on Blackboard by mid June. In order to view the videos, you must enroll in the summer Blackboard course named IPS summer assignment. The password needed to enroll is IPS2016. You also may contact us during the summer vial email or twitter. Email addresses are lhammond.oh@oxford.k12.al.us and rpoe.oh@oxford.k12.al.us You may sign up for Ms. Poe’s Remind 101 text message account by sending the message @ipschem to the number 81010.

Good websites for assistance: The Physics classroom, chemmybear.com, apluphysics.com, and khan academy.

Have a great summer!

Sincerely,

Ms. Poe and Mrs. Hammond

Unit 1: Nature of ChemistryContent Outline: Scientific Equipment and Safety (1.1)

Test Tube Rack – used to hold and dry test tubes.

Thermometer – used to measure how much heat energy is in an object.

Test tubes – used to hold small amounts of pre-measured substances.

Digital Balance – used to determine the mass of smaller, lighter objects.

Triple Beam Balance – used to determine the mass of larger, heavier

objects.

Scoopula – used for transferring dry chemicals.

Weigh Boat – used to weigh or transfer chemicals (usually dry).

Pipette – used to transfer liquid from one container to another. (Increments are very small – usually .25 mL per increment.)

Beaker – used to measure large volumes of liquid. (Increments are larger than a graduated cylinder – usually 25 -50 mL per increment.)

Graduated Cylinder – used to measure volume of liquids – more precise than a beaker. (Increments are smaller than a

beaker – usually 1 mL per increment.)

Bunsen Burner – used to heat liquids and substances in glassware or ceramics.

Ring Stand - Bunsen burner goes on the stand and a wire screen goes on top of the ring. The wire mesh is used to hold glassware and ceramics while heating.

Beaker Tongs – used to pick up hot objects and small glassware.

Test Tube Tongs – used to hold test tubes over flames or in beakers of hot fluid.

Heat Resistant Gloves – used to handle hot objects.

Normal Laboratory gloves – used in laboratory anytime chemicals are to be used.

Goggles – protects your eyes.

Meter stick- measures length.

Erlenmeyer Flask - glassware with a wide base and with sides that taper upward to a short vertical neck; allows contents to be mixed by swirling.

Volumetric Flask – used for making liquid solutions of precise volumes.

Microscope – used to magnify very small items for easier viewing.

Microscope Slides – used for holding specimens on the microscope.

Scientific Equipment and Safety- Continued

I. Safety Rules (10 Commandments of Safety) 1.  Know what equipment is being used and what it is used for. 2.   Never do anything in the lab without instructions OR without your teacher’s permission.3. Never eat or drink in the lab OR eat, drink, or sniff the lab chemicals.4.    Know your safety symbols and identify all possible dangers.5.    No open shoes in the lab. Pull long hair up when needed.6.    Make sure your lab area is clean and uncluttered.7.    Dispose of materials properly. 8.    In case of any accident, inform your teacher. 9.    Know where the nurse and emergency equipment are located.10. Only 1 person per group at a time may go get supplies or use equipment.

II. Safety Symbols

Unit 1: Nature of ChemistryContent Outline: Review of the Scientific Method (1.2)

I. The Scientific MethodA. Observation1. This observation of something in nature leads you to a question such as “Why or How

did that happen?” or “What if…?”2. Types of observations in science:a. Qualitative (Sounds like quality.)

i. These are qualities (descriptions) that an object possesses, such as color, shape, and texture.b. Quantitative (Sounds like quantity.)

i. These are numbers dealing with amounts of an object(s), such 9 bowling balls, ½ of a cake, 2 quarters and 3 dimes.

3. Areas of observation in science:a. System – a specific portion of matter (Anything with mass and takes up space.) in a

given region of space that has been selected for study during an experiment.i. Open system – this type interacts by exchanging matter or energy with the surroundings.

An example would be you (the open system) surrounded by the air and environment around you (surroundings).

ii. Closed system – this type does NOT interact with the surroundings. There is NO exchange of matter or energy.

An example of a closed system would be almost any lab exercise done in a controlled lab environment.

b. Surroundings – areas outside and surrounding the system.B. Research1. You look through textbooks, scientific journals, and maybe on the Internet to see if you

can find an acceptable and logical answer to your question.a. If you cannot find an acceptable answer, then you might design an experiment to test

your question and hopefully find an acceptable answer to your question.C. Formulating a Hypothesis1. A hypothesis is an educated (because you have performed some prior research) guess

about the possible outcome of an experiment, such as the one you developed.a. It may get proven, or it may not. If it is not proven, then you might need to redo or

modify your hypothesis and then retest it.b. It needs to be an “If … then” statement, such as “If water boils at 100° C, then we should

be able to heat and measure water on a stove to prove this.”i. The “If portion” is your initial question.ii. The “then portion” is your educated guess about the outcome of the experiment.

c. Your hypothesis must be testable.D. Procedure of the experiment1. You must have a step-by-step descriptive procedure for your experiment.a. You must list quantities of items, such as chemicals, temperatures, or time.b. You must also state all the equipment needed to perform the experiment in each step.E. Experimentation and Data Collection1. This is the actual performing of the experiment using the procedure you developed.2. You need to be making quantitative and qualitative observations the entire time your

experiment is being performed.a. To help keep the data in an organized, easy to understand format, you need to construct

data tables.

i. Data tables usually tell us information such as the Independent Variables (IV), Dependent Variables (DV), and constants.

. α Independent (controlled) Variable – this is the part of the experiment you are controlling, but modifying, to see if it has an effect on the outcome of your experiment. Some examples would possibly be: temperature, time, concentration.

b. Dependent (changing) Variable – this is the outcome you are measuring. It is dependent upon the outcome of your experiment. It may change as you modify the Independent Variable being tested. Some examples would be: number of bubbles produced, % change in decomposition, or change in color.

c. Constants – These are conditions that are the same (uniform) for all parts of the experiment. They are kept constant and unchanging.

b. Data tables allow you to measure your accuracy, precision, and percentage of error.i. Accuracy – the closeness of measurements to a correct or accepted value.ii. Precision – the closeness of measurements to the same quantity. The quantity may or may

not be the accurate value, though.3. You must perform the same exact experiment several times to ensure your accuracy as it

will be tested by your peers (other scientists) to see if you are telling the truth!F. Analysis1. This is where you will look over and think about your results.2. At this point you will begin making graphs and calculations from your data tables and

observations that you collected during the experiment.a. Percentage error – the mathematical difference (value) between what you

observed and what was expected. Measured using the below equation:

% Error = (VObserved - VExpected)/ VExpected X 100

Here, V represents any value that is being measured. The closer your percent error gets to zero; the more it becomes a perfect outcome, not a perfect experiment. The farther away from zero, the worse your results are.

Example Problems:1. Joshua uses his thermometer and finds the boiling point of ethyl alcohol to be 75o C.  He looks in a reference book and finds that the actual boiling point of ethyl alcohol is 80oC.  What is his percent error?

2. Betty Lou weighed an object on her balance and recorded a mass of 24.21 grams.Her teacher told her that there was obviously something wrong with her balance because it was giving her a reading which was 33.22% too high. What was the actual mass of the object?

3. The density of water at 4oC is known to be 1.00 g/mL. Lucille experimentally found the density of water to be 1.08 g/mL. What is her percent error?

4. You will begin theorizing (trying to prove) your hypothesis and backing it up descriptively using your collected data, graphs, calculations, and making models.a. Theory – a broad, generalized statement that attempts to explain a body of facts or phenomena.i. These can change over time as new data comes to light.b. Scientific Law – These “never” change, as the outcome is always the same. For example, Newton’s Law of Motion – an object remains in motion until acted upon by another object. c. Model – these are structures or formulas used for representing hard to see or hard to understand concepts. For example, it is hard to see the Solar system; but you probably made a model of it back in the 6th grade.G. Publishing1. This is were you will communicate the findings/outcomes of your experiment so that others canpeer review/reproduce (to see if you are accurate) your work or use your work to expand their work.

Unit 1: Nature of Chemistry

Content Outline: Scientific Measurement (1.3)

I. QuantityA. This term is used to describe something that has magnitude, size, or amount.B. This is not the same thing as measurement.

1. Measurement is a process that scientists perform to represent a specific unit of some object. For example, you measured the length of a piece of paper to be 11 inches, or you measured out 3 cups of salt.

2. A measurement nearly always has a number plus a unit.

II. The SI System of measurement used in science.A. SI stands for the French Le Système International d’Unités (International

System of Measurement) that was globally accepted in 1960 at the General Conference on Weights and Measures in Sèvres, France.1. It is used and recognized by all scientists around the world, despite the

reluctance of Americans to adopt the system over the old English system of measurement.

B. The SI system is based upon 7 Fundamental Units of Measurement. They are:1. Length (l)

a. Length is measured in meters (m).2. Mass (m)

a. Mass is measured in grams (g).i. Mass is measured using a scale or balance.

b. Mass is different from weight.i. Weight a measure of the gravitational pull on matter (an

object).ii. Weight is measured on a spring scale and measured in

Newtons, after the great scientist Isaac Newton, who worked with gravity.

c. Weight can change from location to location (earth vs. moon); but mass does not change.

3. Time (t)a. Time is measured in seconds (s).

4. Temperature (T)a. It is measured in Kelvin (K).

i. To convert degrees Celsius (°C) to Kelvin:

273 K + °C ; for example 273 + 27 °C = 300K

ii. To convert degrees Fahrenheit (°F) to degrees Celsius (°C):

(°F – 32) X 5/9; for example (78° F - 32) X 5/9 = 46 X 5/9 = 25.6°C

Example Problems:Convert the following to Kelvin:

1. 90o C ________

2. -20o C ________

Convert the following to Celsius:

3. 200 K ________

4. 273 K ________ 5. 85° F ________

5. Amount of a given substance (n)a. It is measured in moles (mol).b. A mole is a quantity equal to the Formula Weight of a molecule but

measured out in grams. 6. Electric Current (I)

a. Electric current is measured in Amps (A).7. Luminosity (IV)

a. Luminosity is measured in candelas (cd). (Sounds like candles.)b. You can see this one on light bulb packages in stores. The more

candelas…the brighter.C. Prefixes (Additions at the front of a word.) for Magnitude (greater than 1):

getting larger1. Deka (da) = 10; therefore 1 dekameter = 10 meters2. Hector (h) = 100; therefore 1 hectometer = 100 meters3. Kilo (k) = 1,000; 1therefore 1 kilometer = 1,000 meters4. Mega (M) = 1,000,000 ( 1 million); therefore 1 Megameter = 1,000,000

meters5. Giga (G) = 1,000,000,000 (1 billion); therefore 1 Gigameter =

1,000,000,000 meters6. Tera (T) = 1,000,000,000,000 (1 Trillion); therefore 1 Terameter = 1

Trillion metersD. Prefixes for portions (pieces) of a whole: getting smaller

1. deci (d) = 1/10; therefore 1 decimeter is 1/10 of 1 meter.2. Centi (c) = 1/100; therefore 1 centimeter is 1/100 of 1 meter.3. Milli (m) = 1/1,000; therefore 1 millimeter is 1/1,000 of 1 meter.4. Micro ( ) = 1/1,000,000 (millionth); therefore 1 micrometer is μ

1/1,000,000 of 1 meter5. Nano (n) = 1/1,000,000,000 (billionth); 1 nanometer is 1 billionth of 1

meter.E. Derived (made from the Fundamental) units:

1. Area (A)a. Derived by Length (m) X Width (m).b. Area is measured in square meters (m2).

2. Volume (V)a. Derived by Length X Width X Height.b. Volume is measured in cubic meters (m3).

3. Density (D)

a. Derived by Mass divided by Volume (m/m3).b. Density is measured in Grams per Meter cubed (g/m3).

4. Molar Mass (M)a. Derived by Mass divided by amount of a substance (mole) m/mol.b. Molar mass is measured in Grams per Mole (g/mol).

5. Molar Volume (VM)a. Derived by Volume divided by the amount of a substance (mole)

v/mol.b. Molar Volume is measured in cubic Meters per Mole (m3/mol).

6. Energy (E)a. Derived from Force X length (m).b. Energy is measured in Joules (J).

7. Pressure (P)a. Derived from mass divided by meter/second squared m/ms2 (

= multiplied by).b. Pressure is measured in Pascals (Pa) or Atmospheres (Atm.).

Unit 1: Nature of ChemistryContent Outline: Math – The Language of Science (1.4) – Part 1

I. Dimensional Analysis (Basically, changing one unit/dimension into another unit/dimension.)A. This is the mathematical expression of relationship between two different

sets of units, that are related; but in the form of a ratio/fraction. For example, you know than $1 is equal to 4 quarters (0.25) and that a quarter (0.25) is equal to 5 nickels (0.05). See the relationships below:

$1.00 = 4 Quarters; 1 quarter = 5 nickels ; therefore $1 = 20 nickels (4 x 5)

4 quarters(0.25) X 5 nickels (0.05) = 20 nickels (0.05)1 dollar (1.00) 1 quarter (0.25) 1 dollar (1.00)

Notice that each part of the fraction numerator and denominator add up to the to the same Quantity. So each ratio equals 1.

1. Numerator – this term refers to the top number in the fraction.2. Denominator – this term refers to the bottom number in the fraction.3. Always start by listing your units (They must have a direct relationship

or it will not work.) and then put in your numbers.

II. Conversion Factor A. This is a mathematical technique that allows you to use units to solve

problems involving measurements (remember measurements need units). The basic premise is this:

Unit given = Unit wanted = Unit wanted Unit given

This relationship allows us to cancel out the given unit of measurement and convert/replace it with the wanted unit of measurement. For example:

You know that your desk is 3 feet wide. How many inches is that?

3 feet (unit given) = 12 inches (unit wanted) = 36 inches (unit wanted) (3X 12 ) = 36

1 foot (unit given)1

Notice we used the conversion factor of 1 foot being equal to 12 inches; BUT notice the unit placement so that we could convert/replace the given unit with the wanted unit. The units must work in converting or you cannot solve the problem.

III. Significant Figures (These are important numbers in measurements.)A. Consists of all digits known with certainty and one final digit that is

estimated/uncertain.B. All digits 1 – 9 are considered significant. For example:

III.95 centimeters = 3 significant figures; 1 textbook = 1 significant figure

C. The significance of zeros depends on the location within the number. The four rules of zeros apply:1. Zeros appearing between non-zero digits are significant. For example:

305 grams = 3 significant figures; 40.06 meters = 4 significant figures

2. Zeros appearing at the end of a number and to the right of the decimal point are significant. For example:

25.00 liters = 4 significant figures; 3,000,000.00 milligrams = 9 significant figures

3. Zeros at the end of a number but to the left of the decimal point may or may not be significant . It depends on if a decimal point appears. For example:

4000 grams = 1 OR 2 OR 3 OR 4 significant figure…we do not know for sure; 4000.0 grams = 5 significant figures…we know for sure as indicated by the decimal being present.

The zeros in the first number act as “space holders” (they may or may not be known). The presence of the decimal point indicates that they are known, as in the second example.

4. Zeros appearing at the front of all non-zero digits are not significant, as they too are considered “place holders”. For example:

0.0086 meters = 2 significant figures; 0.0000002346 candelas = 4 significant figures

D. Addition or Subtraction involving significant figures:1. When adding OR subtracting decimals, the answer must have the same

number of digits to the right of the decimal point as there are in the measurement having the fewest digits to the right of the decimal point. For example:

14.3 g + 2.64 g = 16.9 g; 14.3 g – 2.645 g = 11.7 g

E. Multiplication or Division involving significant figures:1. When multiplying OR dividing, the answer must have no more significant

figures than are in the measurement with the fewest number of significant figures. For example:

25.6 g = 6.1g/mL (6.09 rounds up to 6.1) ; 25.6 g X 14.765 = 378 (377.9 rounds up to 378)

4.2 mL

F. The rules for D and E do not apply when using conversion factors as those conversions are considered as standards that do not change as they express an established relationship. For example:

2.065 g X 1000mg = 2065 mg 1 g

Example Problems:1. How many centimeters are in 3.25 km?

2. How many milliliters are in 5.67 L3. How many inches are in a 3.67 mile?

(1 mile = 5280 ft)

4. Convert 0.0448 m3 to mm3.

5. 0.000004 m = ________nm

6. 1000000 g = ________kg

7. A block of wood measures 20.5 cm x 4.60 cm x 1.60 cm. Its mass is 71.3 grams. Calculate the density of thee wood.

8. How many cubic centimeters are occupied by 75.0 grams of zinc if its density is 7.14 grams per milliliter?

Unit 1: Nature of ChemistryContent Outline: Math – The Language of Science (1.4) – Part 2

I. Scientific Notation A. This is essentially a way of writing numbers with large amounts of digits in a

condensed form.B. Only significant figures are written when using Scientific Notation.C. It is also based on the powers of 10; but as exponents.

1. Exponents are whole numbers written in superscript(above) to represent a specific number of places the decimal point has moved. (Seen as “Z” below in Part D.)a. If the exponent is a positive whole number, the decimal point has

been moved to the left. This would be a larger than 1 number.b. If the exponent is a negative whole number, the decimal point has

been moved to the right. This would be a smaller than 1 number.D. Numbers written in scientific notation have a basic format:

M.N X 10Z ; M = First Significant digit in the number. (Always followed by the decimal point.)

N = Second Significant digit in the number. Z = a whole number, representing the number of places the

decimal point has moved.

For example: 1,000,000.0 g = 1.0 X 106 g250.0 L = 2.5 X 102 L0.000465 m = 4.65 X 10-4 m

E. Addition and subtraction using Scientific Notation:1. These mathematical operations can only be performed if they possess

the same exponent value.For example:

2.4 X 106 + 5.3 X 106 = 7.7 X 106 OR 5.3 X 106 – 2.4 X 106 = 2.9 X 106

a. If they do not have the same exponent, then one of the numbers will need to be converted so that they do match.2.4 X 105 + 3.1 X 103 = 2.4 X 105 + 0.031 X 105 = 2.431 X 105 (The green value converted.)

OR

2.4 X 105 + 3.1 X 103 = 240.0 X 103 + 3.1 X 103 = 243.1 X 103 (The orange value converted.)Teachers please help students see that the outcome is numerically the same, just a difference in where the decimal point is located.

F. Multiplication using Scientific Notation:1. The significant digits, of each number, are multiplied first.2. Then the exponents are added together.

For example:

(2.4 X 105) X (3.6 X 103) = 8.64 X 108

G. Division using Scientific Notation:1. The significant digits are divided first.2. Then the exponents are subtracted.

For example:2.45 X 10 23 = 4.3 X 1010

5.65 X 1012

Step one: 2.45 /5.65 = 0.433 (round to 0.43)Step two: 23 – 12 = 11Step three: Move the decimal to the right to turn 0.43 into 4.3Step four: Since you had to move the decimal to the right, you need to

correct your exponent number to reflect that 11 becomes 10

*If you move the decimal to the right; then subtract that number of moves to the exponent.

* If you move the decimal to the left; the add that number of moves to the exponent

Example Problems:1. Write each of the following numbers in exponential notation:

a) 91,100 __________ b) 0.000000075 __________c) 6400 __________ d) 0.00165 __________e) 0.0816 __________ f) 935 __________

2. Convert each of the following numbers in exponential notation to conventional decimal form:a) 2.24 x 10-5 ______________ d) 2.95 x 10-3 ______________b) 9.3 x 102 ______________ e) 7.35 x 10-2 ______________c) 4.20 x 104 ______________ f) 8.18 x 10-12 ______________

3. Rewrite each number with a coefficient between one and ten (in other words, write each number in CORRECT scientific notation:a) 275 x 103 ______________ d) 43.9 x 10-1 ______________b) 92 x 10-4 ______________ e) 0.0165 x 10-2 ______________c) 0.611 x 105 ______________ f) 0.0641 x 106 ______________

4. How many significant figures are in each of the quantities listed below:a) 454 mg __________ d) 0.0680 km __________ g) 0.1536 g__________

b) 0.0353 L __________ e) 10.0 mL __________ h) 0.0060 g__________

c) 52.20 mL __________ f) 3 x 107 kg __________ i) 1.898 x 10-3 __________

j) 2500 m __________

5. Round off the given quantity 7.758064 to the number of significant figures indicated:

one ____________________ three _____________________ five _____________________two _____________________ four _____________________ six _____________________

6. Perform the indicated calculation and express the answer in the correct number of significant figures:

0 .370×843×0. 07040 . 0042×17 .10

=_____________________

7. Complete the following operations and round off the answers to the proper number of significant figures:

a) 18 .7 - 0 . 56 = _____________________ b) 210 .0+3.19+1765+2.64 =_____________________

8. Round off the quantity 4.106738 to the number of significant figures indicated:one _____________________ three _____________________five _____________________

two _____________________ four _____________________ six _____________________

9. Express each of the following in correct scientific notation with the correct number of significant figures:a) (5.68 x 10-4) x (6.52 x 104) _______________ b) (7.53 x 105) x (1.56 x 10-9) _______________

10. Express each of the following in correct scientific notation with the correct number of significant figures:a) (6.74 x 10-8) / (2.34 x 106) _______________ b) (1.45 x 109) / (9.35 x 10-7) _______________

Unit 1: Nature of ChemistryContent Outline: Classification of Matter (1.6)

I. Classification of MatterA. There are two main categories of matter: pure substances and

mixtures.1. Pure substance – this term refers to substances that have uniform

composition throughout all samples and uniform properties.a. Pure substance can either be elements, such as everything

contained on the Periodic Table, or compounds, such as water.i. The Periodic table shows the three categories of known

elements on Earth: metals, metalloids (like-metals), and Non-metals.

ii. The horizontal rows on the table are called periods.iii. The vertical rows on the table are called groups or

families. This term is used because the elements in a family all have similar properties.

iv. Metals – have the ability to conduct electricity.v. Metalloids – possesses some characteristics of metals

but also non-metals. (“oid” means “like”… It is like a metal.)

vi. Non-metals – these are poor conductors of electricity. Most are gases.

2. Mixtures – this term is used for samples of matter that do NOT possess uniform composition throughout all samples with a combination of properties.a. Mixtures can even be classified into two groups:

homogeneous and heterogeneous.i. Homogeneous mixtures – these are mixtures that

appear as uniform in composition. (“ Homo” means “same”.) For example, Kool-Aid.

. Homogeneous liquid mixtures can also be called αsolutions.b. Colloid – these are substances that appear uniform, but if left undisturbed will separate based upon density. For example, milk.

ii. Heterogeneous mixtures - these are mixtures that appear un-uniform in composition. (“Hetero” means “different”.) For example, a plate of salt and pepper mixed or a chunk of granite.

3. Pure substance compounds can only be decomposed by chemical means.

For example, providing electricity to water to convert it to Hydrogen gas and Oxygen gas. (This process is called electrolysis.)

4. Mixtures can be separated based upon the unique properties of the composing substances.a. Some various ways of separating mixtures include:

i. Filtration – pouring the sample through filter paper.ii. Chromatography – separating solutions by differing

densities.iii. Centrifuge – separating particles, in a liquid, by

differing densities.iv. Vaporization – heating off the liquid of solution.

Practice Problems: Classification of Matter on an ATOMIC LEVEL

Each of the following diagrams shows a sample of a substance viewed at the atomic level. Characterize the contents of each container in terms of the following categories: A. Solid, Liquid, Gas, or Combination of PhasesB. Pure Substance, Homogeneous Mixture, or Heterogeneous MixtureC. Element(s), Compound(s), or Both Elements and Compounds

Example:

A. Gas (totally random distribution of particles) B. Pure substance (there is a single type of matter in the container C. Element (it is a monatomic substance) Ex: Neon gas is an example

A. ____________________ B. ____________________ C. ____________________

A. ____________________ B. ____________________ C. ____________________

A. ____________________ B. ____________________ C. ____________________

A. ____________________ B. ____________________ C. ____________________

A. ____________________ B. ____________________ C. ____________________

A. ____________________ B. ____________________ C. ____________________