EXPONENTIAL NOTATION
Integrated Physical Science Summer Assignment
2015-2016 Academic year
Future Honors Physical Science Students,
Welcome to Honors Physical Science! We are eagerly anticipating a great year of Honors Physical Science. In order to ensure the best start for everyone next fall, we have prepared a Summer Assignment that reviews basic chemistry and physics concepts. Most of the material covered in the summer 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 will be important for everyone to come prepared to class on the first day. While we will review, extensive remediation is not an option as you only have a semester of each course. The Summer Assignment is DUE THE FIRST MONDAY OF SCHOOL! There will be a test covering the basic concepts included in the summer packet during the first full week of school.
We hope 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 learner. You are all certainly fine students, and with motivation and hard work, you should find Honors Physical Science a successful and rewarding experience.
Finally, we recommend that you find a “study buddy” and spread out the summer assignment. Please do not try to complete it all in one day. Chemistry and physics takes time to process and grasp at a 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.
Have a great summer!
Ms. Poe and Mrs. Hammond
. You may contact us throughout the summer via email or twitter. Ms. Poe’s email address is rpoe.oh@oxford.k12.al.us. My twitter account is @mspoes_chem. She will be available at various time throughout the summer is you are having difficulties with anything. Mrs. Hammond’s email address is lhammond.oh@oxford.k12.al.us.
. We will be posting video tutorials for the summer assignment notes on Blackboard. In order to view the tutorials you will need to enroll in the Blackboard course. The course name is IPS summer assignment. The password needed to enroll is IPS2014. Videos are posted on Blackboard for reference.
Good websites for assistance: khan academy, , the physics classroom and aplusphysics.
Significant Figures Notes
Accuracy –
Precision –
The number of digits reported for a measured value conveys the quality of the measurement and hence, the quality of the measuring device. It is important to use significant figures correctly when reporting a measurement so that it does not appear to be more (or less) precise than the equipment used to make the measurement allows. We can achieve this by controlling the number of digits, or significant figures, used to report the measurement.
In this course and in others, you must use correct significant figures in reporting your results. Laboratory measuring instruments have their limits, just as your senses have their limits. One of your tasks, in addition to learning how to use various measuring instruments properly, will be to correctly determine the precision of the measuring devices that you use and to report all measured and calculated values to the correct number of significant figures. Exact numbers do not affect the number of significant figures in the answer to a calculation.
Significant Figures -
RULES FOR COUNTING SIGNIFICANT FIGURES
1. All non-zero numbers are significant
2. Sandwich zeros are significant.
3. Leading zeros are NEVER significant.
4. Trailing zeros are significant if decimal point is present
5. Exact numbers have as many significant figures as needed.
EXAMPLE Determine the number of sig figs
3.82 L _______ 24 m ______ 0.0619 kg
2.1 x 104 g ___ 3.4610 km ___ 712,000 cm ___
EXAMPLE
[pic]
RULES FOR ROUNDING
1. 0-4 : Leave it alone
2. 5-9 : Raise it!
EXAMPLE Round the following number as instructed 21.409653
two three four
five six seven
EXAMPLE Round the following number to four significant figures
4.000574 x 105 = ________________ 375.6523 = __________________ 89,762,334 = ________________
CALCS WITH SIG FIGS
In mathematical operations involving significant figures, the answer is reported in a way that reflects the reliability of the least precise number. An answer is no more precise that the least precise number used to get the answer. Imagine a team race where you and your teammates must finish together at the same time. Who dictates the speed of the team? Of course, the slowest member of the team. Your answer cannot be MORE precise than the least precise measurement.
Careful!! -- The rules for addition/subtraction are different from those of multiplication/division
Another note – in chemistry, we report our answers with decimals, not fractions.
• Multiplication and Division
The answer can contain no more significant figures than the number with the fewest number of significant figures in the given problem.
• Addition and Subtraction
The answer must be rounded off to contain only as many decimal places as are in the value with the least number of decimal places (i.e. the least number of digits pas the decimal.)
When doing multi-step calculations, keep at least one more sig fig in intermediate results than needed in final answer. Round to correct sig figs at the very end.
EXAMPLE Perform the following calculations
239.1 x 46.23 x 0.00290 = 0.01699021 =
16.508 x 114.29
2342 x 11.3 = 26464.6 =
27.14 148.77
9224.8 -106.409
5.7567 42.361=
+ 948.84
10206.5367 =
SCIENTIFIC NOTATION
To change a number to exponential notation
• Relocate the decimal so it appears immediately behind the first nonzero digit.
• Count the number of places the decimal was moved.
• Write the exponential factor as 10 raised to the power equal to that number.
• If the decimal was moved to the left, the exponent is positive, and no sign is necessary.
• If the decimal was moved to the right, the exponent is negative, so a minus sign must be inserted.
REMEMBER!!
Large numbers have positive exponents.
Small numbers have negative exponents.
EXAMPLE: Write 0.0000000674 in exponential notation.
EXAMPLE: A calculated result is 0.000378 x 10-9
Rewrite the number in standard scientific notation.
Calculations in Exponential Notation
1. When multiplying numbers in scientific notation, the exponents are added.
2. When dividing numbers in scientific notation the exponents are subtracted.
3. If two numbers have the same exponents, the numbers can be added, and leave exponents the same.
EXAMPLE (9.6 x 10-4)(2.5 x 10-5)
EXAMPLE (2.25 x 103) / (6.00 x 10-6)
Dimensional Analysis
• Start with your given. If more than one given, use one that only has one unit.
• Identify the unit that you are looking for.
• Set up problem in a path that cancels out units not wanted and leaves unit looking for.
EXAMPLE How many minutes are in 2.25 weeks?
EXAMPLE A student has a part time job in which she works 4 hour shifts. Her pay scale is $5.25 hour. How many shifts must she work to earn $220. ?
EXAMPLE If you used 16.0 gallons when driving 376 miles, what was your gas mileage over that distance?
EXAMPLE At the above gas mileage, how far can you travel on a tank of 22.0 gallons?
EXAMPLE At the same gas mileage, how many dollars will you spend on gasoline during a 658 mile trip if the price of gasoline is $3.35 per gallon?
THE METRIC SYSTEM
The prefixes you must know are:
kilo = k = 103 centi = c = 10-2 milli = m = 10-3 micro = ( = 10-6
EXAMPLE How many centimeters are in 3.25 km?
EXAMPLE How many milliliters are in 5.67 L?
EXAMPLE How many inches are in a 3.67 mile?
(1 mile = 5280 ft)
EXAMPLE How many cm are in 3.67 mile?
(2.54 cm = 1 inch)
REMINDER: Base your sig figs on the given number, not on conversion factors.
Some useful conversion factors:
|Mass |Length |Volume |
|1 pound = 453.6 g |1 inch = 2.54 cm |1.057 quarts = 1 L |
|2.205 pounds = 1 kg |39.97 inches |1 gallon = 3.785 L |
|1 ounce = 28.35 g |1 mile = 1.609 km |1 ounce = 29.57 mL |
MORE ON DIMENSIONAL ANALYSIS
Doing cubic conversions
• use the normal factor conversions
• for the factor conversion, be sure to cube the unit and the number
EXAMPLE Convert 0.0448 m3 to mm3.
EXAMPLE An automobile engine has an internal engine volume of 286 in3. What is this in liters?
CALCULATIONS WITH PERCENTS
Any percentage is defined to be:
% = (part/whole) x 100
EXAMPLE How many grams of salt and how many grams of water would you use to prepare 85 g of a solution that is 12.0% salt?
EXAMPLE A mineral has a copper content of 25.5%. In a certain experiment, 561 grams of copper were recovered from a pure sample of the mineral. What was the mass of the sample?
Density Notes
• The density of a substance is its mass per unit volume.
• Units of density match the definition: mass units over volume units
• A useful formula for density is:
Density = mass
volume
EXAMPLE: 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 the wood.
EXAMPLE: How many cubic centimeters are occupied by 75.0 grams of zinc if its density is 7.14 grams per milliliter?
Graphing Density……
When plotting a line, it is known that the slope of a line is rise or change in Y
run change in X
Since Density = mass and slope = change in Y
volume change in X
then plotting mass on the y axis and volume on the x axis will give you a graph of density of a substance.
WE WILL GO OVER GRAPHING MORE IN DEPTH AT THE BEGINNING OF THE SCHOOL YEAR.
Classification of Matter Notes
(You are responsible for everything in this outline)
I. Matter – anything that takes up space and has mass. Material refers to a specific kind of matter.
A. Pure substances – matter that contains only 1 type of material
1. Elements – substances composed of a single type of atom.
a. Smallest possible unit that can’t be broken down any further and still retain the same properties or characteristics.
b. An atom is the smallest particle of an element.
c. It is possible to have a molecule of an element if more than 1 atom of the same element is chemically bonded together.
d. Examples – anything on the periodic table of elements – duh
2. Compounds – substance composed of 2 or more elements
a. Can only be separated by chemical means
b. A molecule is the smallest particle of a compound.
c. Examples
i. Inorganic compounds – do not contain carbon (ie. NaCl)
ii. Organic compounds – contains carbon (ie. CH4, C6H12O6)
B. Mixtures – matter that contains 2 or more types of materials
1. Homogeneous (aka Solutions) – mixture w/ only one phase; each part looks identical
a. Examples: air, antifreeze, kool-aid, ice tea, glass, jello
2. Heterogeneous – mixture composed of more than identifiable component differences can be seen with the naked eye
a. There is an interface boundary between phases.
b. Examples – jello w/ fruit, oil & water, rocky road ice cream
3. Separating Mixtures
a. Distillation (separates heterogeneous & homogenous mixtures)
b. Evaporation (separates heterogeneous & homogenous mixtures)
c. Filtration (separates heterogeneous mixtures)
d. Decantation (separates heterogeneous mixtures)
[pic]
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 Phases
B. Pure Substance, Homogeneous Mixture, or Heterogeneous Mixture
C. Element(s), Compound(s), or Both Elements and Compounds
[pic]
1.[pic] 2.[pic] 3.[pic]
A. _____________________ A. _____________________ A.____________________
B. ______________________ B. _____________________ B. ____________________
C. ______________________ C. _____________________ C. ____________________
4.[pic] 5.[pic] 6.[pic]
A. ______________________ A. _____________________ A. ____________________
B. ______________________ B. _____________________ B. ____________________
C. ______________________ C. _____________________ C. ____________________
Chemical and Physical Properties Notes
(You are responsible for everything in this outline)
I. Changes in Properties
a. Physical Changes
1. Definition – same substance remains. Basic chemical composition doesn’t change. Can be observed and measured w/o changing the matter itself.
2.Physical Properties – describes the substance
e. Extensive properties – depend on amount of matter present (i.e. Mass, length, volume)
f. Intensive properties – does not depend on amount of matter present (color, density, freezing point, melting point, malleability, ductility, conductivity)
3.Examples of Physical changes – changes in state, dissolving, tearing, pounding
B. Chemical Changes
1. Definition – Different substance formed. Describes how a certain type of matter will react and change (or not react and change) when it is in the presence of other kinds of matter.
2.Chemical Properties – describes behavior of substances undergoing chemical change
3.Examples of Chemical changes –reaction w/ oxygen, corrosion, flammability
4. How to Identify a Chemical Reaction
1. Color change
2. Gas formation
3. Formation of a precipitate (insoluble solid)
4. Energy change (light, heat, cold, sound, etc)
PHYSICAL VS. CHEMICAL
PROPERTIES
A physical property is observed with the senses and can be determined without destroying the object. For example, color, shape, mass, length and odor are all examples of physical properties.
A chemical property indicates how a substance reacts with something else. The original substance is fundamentally changed in observing a chemical property. For example, the ability of iron to rust is a chemical property. The iron has reacted with oxygen, and the original iron metal is changed. It now exists as iron oxide, a different substance.
Classify the following properties as either chemical or physical by putting a check in the appropriate column.
| |Physical Property |Chemical Property |
|1. blue color | | |
|2. density | | |
|3. flammability | | |
|4. solubility | | |
|5. reacts with acid to form H2 | | |
|6. supports combustion | | |
|7. sour taste | | |
|8. melting point | | |
|9. reacts with water to form a gas | | |
|10. reacts with a base to form water | | |
|11. hardness | | |
|12. boiling point | | |
|13. can neutralize a base | | |
|14. luster | | |
|15. odor | | |
CHANGES
In a physical change, the original substance still exists, it has only changed in form. In a chemical change, a new substance is produced. Energy changes always accompany chemical changes.
Classify the following as being a physical or chemical change.
1. ________ Sodium hydroxide dissolves in water.
2. ________ Hydrochloric acid reacts with potassium hydroxide to produce a salt, water, and heat.
3. ________ A pellet of sodium is sliced in two.
4. ________ Water is heated and changed to steam.
5. ________ Potassium chlorate decomposes to potassium chloride and oxygen gas.
6. ________ Iron rusts
7. ________ When placed in water, a sodium pellet catches on fire as hydrogen gas is liberated and sodium hydroxide forms.
8. ________ Evaporation.
9. ________ Ice melting.
10. ________ Milk sours.
11. ________ Sugar dissolves in water.
12. ________ Wood rotting.
13. ________ Pancakes cooking on griddle.
14. ________ Grass growing in the lawn.
15. ________ A tire inflated with air.
16. ________ Food is digested in the stomach.
17. ________ Water is absorbed by a paper towel.
Nuclear Composition/Nuclide Composition
I. Components of the Atom
| |Symbol |Mass |Charge |Location |
|Proton | | | | |
|Neutron | | | | |
|Electron | | | | |
Note: 1 amu = 1.66 x 10-24
II. Nuclide symbols
Atomic number (Z)
Mass number (A)
Condensed form:
[pic] Or element-__
Example:
|Isotope Symbol |Atomic Number |Mass Number |Number of Protons |Number of Electrons |Number of Neutrons |
|__ __ | | | | | |
|__ Ne | |22 | |10 | |
| | | | | | |
| |24 |52 | |22 | |
| | | | | | |
| | | |35 |36 |46 |
III. Isotopes --
Examples)
[pic] [pic] [pic]
hydrogen deuterium tritium
You try!
carbon-12 carbon-13
Notes – Mass Number vs. Atomic Weight
Mass number -
Atomic weight -
Why is the actual weight of an isotope not a whole number like the mass number?
(It is always slightly less)
How do you calculate an atomic weight?
Example: Three isotopes of magnesium occur in nature. Their abundances and masses are listed below. Use this information to determine the atomic weight of magnesium.
Isotope % abundance mass(amu)
[pic] 78.99 23.98504
[pic] 10.00 24.98584
[pic] 11.01 25.98259
Example: Calculate the atomic weight of zinc using the data in the table given below:
Isotope % abundance mass(amu)
[pic] 48.6 63.929
[pic] 27.9 65.9260
[pic] 4.7 66.9271
[pic] 18.8 67.9298
Example: The atomic weight of gallium is 69.72 amu. The masses of two naturally occurring isotopes are 68.9257 amu for [pic] and 70.9249 amu for [pic]. Calculate the % abundance for each isotope.
Example: The atomic weight of chlorine is 35.453 amu. There are only two isotopes of chlorine: chlorine-35 (34.96885 amu) and chlorine-37 (36.9658). Calculate the % composition of each isotope of chlorine.
Atomic History
400 BC [pic]Democritus – Greek philosopher
Thought that all matter was composed of small, indivisible particles. Beliefs based on reasoning.
1803 John Dalton – British school teacher
Based on data from chemical reactions and the behavior of gases, he formed the Atomic Theory:
*1. Elements are composed of atoms that cannot be further broken down
2. Atoms are neither destroyed nor created in chemical reactions
*3. All atoms of given element weigh the same; different elements weigh different amounts
4. Elements can be ordered according to weight
* Not totally correct; later revised
Dalton’s atomic theory was accepted because it successfully explained several laws and observations. Dalton’s atom was a solid, indivisible sphere.
His findings made it possible to analyze compounds and use atomic masses to calculate formulas.
1811 Amadeo Avogadro –
All gases having the same volume and temperature have same number of particles. Distinguished difference between atoms and molecules.
1869 Dimitri Mendeleev – Russian chemist
Published arrangements of known elements. Classification was based primarily on chemical properties of the elements and increased according to atomic weight.
1896 Henri Becquerel
Found that radiation was emitted from uranium
1897 Marie and Pierre Curie –
Discovered 2 new radioactive elements – radium and polonium, and identified 3 types of radiation emitted by radioactive elements.
1897 JJ Thomson – British physicist
Devised a system with fluorescent material that glowed when struck by charged particles which provided a way to measure the deflections of the beam in a CRT (or Crooke’s tube).
[pic] [pic]
Factors that affect the moving, charged particle by the magnet or charged plates: 1) mass of particle, 2) velocity of particle, 3) electric charge of particle, 4) strength of magnet, 5) amount of charge on plates.
Showed mass / charge ratio for negative particles was always the same, regardless of metal used for cathodes or gas used in tube.
[pic]
1909 Robert Millikan – American physicist
Used oil drop experiment to determine the electron’s charge.
[pic]
1911 Ernest Rutherford – New Zealand physicist
He wanted to prove that atoms are mostly empty space, so he fired alpha (+) particles at gold leaf foil:
-most went straight through or were slightly deflected,
-while some bounced straight back.
[pic]
1913 Niels Bohr
[pic]
1926 Erwin Schrodinger
Proposed a model of the atom based on the mathematical probability of locating an electron within the atom. At any given time, there is a high probability that the electron exists within the electron cloud.
[pic]
Introduction to the Periodic Table
I. Founders of the Periodic Table
A. Dmitri Mendeleev (1869) – published first periodic table; organized elements in vertical columns by similar
properties; arranged column so that the elements were in horizontal rows by increasing atomic mass.
(Grandfather of the Modern Periodic Table)
B. Henry Moseley (1913) – arranged elements in order of increasing atomic number (kept elements arranged so
that elements with similar chemical properties were in the same vertical column. (Father of the Modern Periodic
Table)
II. Organization of the Periodic Table
A. Metals vs. Nonmetals – divided by zig-zag line
1. Metals – found to the left of the zig-zag line; the majority of the periodic table
2. Nonmetals – found to the right of the zig-zag line
3. Metalloids – found between the metals and nonmetals; have properties of both; have two sides on zig-zag line
but no sides on the outside edge of the table.
4. Color coding on the wall Periodic Table
a. gas = red
b. liquid = blue
c. solid = black
d. manmade = outline
B. Groups and Periods
1. Groups – vertical columns, also called families; elements in the same group or family have similar chemical properties
a. IUPAC Numbering System – number from left to right (1 –18)
b. Traditional System – Roman numerals and letters
2. Some groups have special names
a. Group 1 (IA) – alkali metals
b. Group 2 (IIA) – alkaline earth metals
c. Group 14 (IVA) – Carbon group
d. Group 15 (VA) – Nitrogen group
e. Group 16 (VI A) - chalcogens
f. Group 17 (VIIA) – halogens
g. Group 18 (VIIIA) – noble gases (Why are they “noble”?)
3. Periods – horizontal rows
a. All periods designated by a number (1 – 7)
b. Note that periods 6 & 7 have had some elements placed at the bottom of the periodic table.
C. Representative vs. Transition Elements
1. Representative Elements (“A” groups in traditional system, Groups 1, 2, 13 – 18 in IUPAC system)
2. Transition Elements (“B” groups in traditional system, Groups 3 – 12 in IUPAC system)
3. Inner Transition Elements – at the bottom of the Periodic Table, no group designation
a. Lanthanide Series (elements 57 – 71)
b. Actinide Series (elements 89 – 103)
[pic]
-----------------------
Inexact numbers:
Any measured value:
Use of a 1m mL graduated cylinder might give a volume of 8.81 mL (3 sig figs)
A 100 mL graduated cylinder might give a less precise volume of 8.7 mL (2 sig figs)
Exact Numbers
There are exactly 12 eggs in a dozen.
Most people have exactly 10 fingers and toes.
1 meter = 100 centimeters
1 yard = 36 inches
1 dollar = 100 cents = 4 quarters
1 kilometer = 100 centimeters
A. What measuring device is used in the picture to the right?
B. How should you report the volume of a liquid?
C. How many significant figures did you report?
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
Concluded that all cathode rays were composed of identical negatively charged particles that he named electrons.
By measuring the amount and direction of deflection, he was able to calculate the ratio of the electron’s charge to its mass.
Created “Plum Pudding Model”.
Showed that an electron has the smallest possible negative charge.
All other negative charges were whole number multiples of the charge of an electron.
Combined this with Thomson’s work to calculate the mass of an electron.
Conclusions:
1. Most of the atom is empty space,
2. The atom has a dense, tiny positively charged core (nucleus),
3. The electrons (-) whirl around the outside of the nucleus.
EFG
.
9
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HIJbcˆŽ/]^˜š¯ÀÁñ
He measured the diameter of the nucleus to be 1 / 100,000 the diameter of the atom.
Proposed that electrons are arranged in concentric circular paths or orbits around the nucleus. Electrons in a particular path have a fixed energy, thus they do not lose energy and fall into the nucleus.
Experiments were based on the hydrogen atom. Results did not hold true for other elements.
Electrons have certain wave-like properties
Light has certain particle-like properties
Electrons and light have a dual wave-particle nature.
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