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AP Chem Summer Assignment:

Independent Study Chapters 1 and 2 of textbook

Read Chapters 1 and 2 of text book, these are a review of Chem 1

We will not be discussing these chapters in class.

Answer the questions in this packet to accompany Chapters 1 and 2

The answers to the questions in the packet will be posted on my website

during the last week of August. You can check your answers then

Do the following questions from your text: These will be turned in and

graded on Friday, September 7th

.

Pages 31-35 : # 5,15,16,24,25,28,29,35,39,41,45,57,60,65,70

Pages 71-76 : # 1,9,27,29,31,33,39,41,59,61,63,64,65,68,69,71,73,75,

84,94,96

There will be a test on Chapters 1 and 2 on Friday, September 7.

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AP Chemistry Unit I Chemical Foundations

Chapters 1 Matter and Measurement , Chapter 2-Atoms Molecules and Ions

Objectives

Chapter 1 Matter and Measurement

Recall a definition of chemistry

Understand the stages and process of scientific problem solving

List the states of matter, defining and listing characteristics of the common states

Compare and contrast physical and chemical changes

Differentiate among elements, mixtures and compounds

Differentiate between homogeneous and heterogeneous mixtures

Understand and use scientific notation

Define and convert among SI units and prefixes

Understand the concept of derived units and use relationships relating to density

Define uncertainty in a measurement and recall the rules for significant digits in

calculations

Compare and contrast accuracy and precision

Interconvert between Celsius, Kelvin and Fahrenheit temperature scales

To state the fundamental chemical Laws of Conservation of Mass, Definite Proportions,

and Multiple Proportions

Chapter 2 Atoms, Molecules and Ions

To Summarize Dalton’s Atomic Theory

To explain how the Experiments of Thomson, Millikan, and Rutherford led to the

Modern view of Atomic Structure

Define isotope, and describe how mass spectrometers determine the mass of an isotope

Calculate the Average atomic mass

Distinguish among terms chemical formula, empirical formula, molecular formula and

structural formula

Determine the oxidation numbers of monatomic ions

State the formula and oxidation numbers of a variety of polyatomic ions

Distinguish among three types of inorganic substances: ionic compounds, acids, and

binary molecular compounds

Name and write formulas for some simple Organic compounds

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Classifying

Matter

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Classify the following as element, compound, homogeneous mixture , heterogeneous

mixture:

1. Aqueous solution of sodium chloride ________________________

2. Air_____________________________________________________

3. Granite_________________________________________________

4. Sodium chloride _________________________________________

5. White Gold (gold and palladium in differing amounts) ____________

States of matter

Properties described on the macroscopic level:

•gas (vapor)_________________________________________.

•liquid: _______________________________________________

•solid: ________________________________________________

Properties described on the molecular level:

•gas: ___________________________________________________.

•liquid: ___________________________________________________

•solid: ___________________________________________________

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, a completely different substance.

Intensive properties do not depend on the amount of substance (for example, alcohol is flammable).

Extensive properties do depend on the amount (for example, the mass of the alcohol).

Classify the following properties as either chemical or physical by putting a check in the appropriate

column. Also indicate if it is an intensive or extensive property.

Property

Ph

ysi

cal

Ch

emic

al

Inte

nsi

ve

Ex

ten

siv

e

1. weight

2. blue color

3. density

4. flammability

5. volume

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Physical vs Chemical 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 by writing a C or a P.

1. When placed in water, a sodium pellet catches on fire as hydrogen gas is liberated

and sodium hydroxide forms Sodium hydroxide dissolves in water.

2. Sugar dissolves in water

3. Hydrochloric acid reacts with potassium hydroxide to produce salt, water, and hat.

4. Leaves change color.

5. Potassium chlorate decomposes into potassium chloride and oxygen gas.

6. Iron rusts.

7. Distillation of a liquid

8. Ice melting.

9. Milk sours.

10. evaporation

Counting Significant Digits

Quantities in chemistry are of two types:

Exact numbers – These result from counting objects such as desks (there are 24

desks in this room), occur as defined values (there are 100 cm in 1 m), or as

numbers in formulas (area of a right triangle = ½ B x H). They (24, 100, 1, and ½

for these examples) all have an infinite(∞) number of significant digits. B and H

are measurements and do not have an infinite number of digits.

Inexact numbers – These are obtained from measurements and require judgment.

Uncertainties exist in their values.

When making any measurement, always estimate one place past what is actually known.

For example, if a meter stick has calibrations to the 0.1 cm, the measurement must be

estimated to the 0.01 cm. When making a measurement with a digital readout, simply

write down the measurement. The last digit is the estimated digit.

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Significant digits are all digits in a number which are known with certainty plus one

uncertain digit. The following rules can be used when determining the number of

significant digits in a number:

Rule Example Sig Digs

1. All nonzero numbers are significant. 132.54 g 5

2. All zeros between nonzero numbers are significant. 130.0054 m 7

3. Zeros to the right of a nonzero digit but to the left of an

understood decimal point are not significant unless shown

by placing a decimal point at the end of the number.

190 000 mL

190 000. mL

2

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4. All zeros to the right of a decimal point but to the left of

a nonzero digit are NOT significant.

0.000 572 mg 3

5. All zeros to the right of a decimal point and to the right

of a nonzero digit are significant.

460.000 dm 6

A good way to remember which side to start on is:

decimal point present, start at the Pacific

decimal point absent, start at the Atlantic

How many significant digits do each of the following numbers have?

1. 2.3000 x 106_____________________ 6. 500. _____________________

2. 45.1 _____________________ 7. 970 _____________________

3. 800000.103 _________________ 8. 0.002 _____________________

4 0.000000001500_________________ 9. 0.007 80 _____________________

5. 600 _____________________ 10. 145.55 _____________________

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Rounding Rules Calculators often give answers with too many significant digits. It is often necessary to round off the answers to the correct number of

significant digits. The last significant digit that you want to retain should be rounded up if the digit immediately to the right of it is (Each of the examples are being rounded to four sig digs):

Rule Example 4 sig digs

….. greater than 5 532.79 532.8

….. 5, followed by a nonzero digit 17.255 1 17.26

….. 5, not followed by a nonzero, but has an odd digit directly in front of it. 3 213.5 3214

The last significant digit that you want to retain should stay the same if the digit immediately to the right of it is:

Rule Example 4 sig digs

….. less than 5 5 454.33 5 454

….. 5, not followed by a nonzero digit, but has an even digit directly in front of it. 0.007 85 0.007 8

Round the following numbers to 3 sig digs.

1. 279.3 ___________ 3. 32.395 ___________ 5. 18.29 ___________

2. 42.353 ___________ 4. 32.25 ___________ 86 5 001 ___________

Applying significant digits to arithmetic operations

Addition and Subtraction – Look at the numbers being added or subtracted and identify which one has the lowest number

of decimal places. Calculate the answer. Round the answer to the lowest number of decimal places.

14.565 + 7.32 = 21.885 7.32 has only 2 decimal places, so the answer should be rounded to 21.88

143.52 – 100.6 = 42.92 100.6 has only 1 decimal place, so the answer should be rounded to 42.9

Multiplication and Division – Look at the numbers being multiplied or divided and identify which one has the lowest number of significant digits. Calculate the answer. Round the answer to the lowest number of significant digits.

172.6 x 24.1 = 4159.66

24.1 has only 3 significant digits, so the answer should be rounded to 4160

172.6 ÷ 24.1 = 7.161 82

234.1 only has 3 significant digits, so the answer should be rounded to 7.16

Express each answer with the correct number of significant digits.

1. 320 x 24.9

0.080

2. 432.7 x 6.5 x .002300

62 x 0.103

3. 32.44 + 4.9- 0.304

82.94

4. 8.002 + 0 .3040

13.4 - 0.066 + 1.02

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Temperature is defined as the average kinetic energy of the particles in a sample of

matter. The units for this are oC and Kelvin (K). Note that there is no degree symbol for

Kelvin. Absolute zero is ) Kelvin. This is the coldest possible temperature where all

molecular motion ceases

Heat is a measurement of the total kinetic energy of the particles in a sample of matter.

The units for this are the calorie (cal) and the Joule (J).

The following equation can be used to convert temperatures from Celsius (t) to Kelvin

(T) scales and fahrenheit to celsius and celsius to fahrenheit

T(K) = t(oC) + 273.15

t(oC) = T(K) - 273.15

You are simply subtracting 273.15 from your Kelvin temperature.

Convert the following:

1. 335C = _______________K

2 50.0 K = _______________C

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SI Fundamental Units

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SI Derived Units- calculated from fundamental units

http://cornellchem.wikispaces.com/Unit+2+-+Measurements+and+Calculations

SI Prefixes

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Using Dimensional Analysis to convert among metric units

1. Convert 1.09 ng to g

2. Convert 90.1 ms to s

3. Convert 145 pm to m

4. Convert 1.55 kg/m3 to g/L

5. Determine the number of L in a box that measures 4cm x 8cm x 10cm.

(1 ml= 1 cm3)

6. If a car has an EPA mileage rating of 30 miles per gallon, what is this rating in

km/L?

7.

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Accuracy versus Precision

Percent Error – Measures accuracy of measurement.

Calculate the percent error for the following problems. Show your work and use the

correct number of significant digits in your answer.

1. A student measures a rectangular piece of metal. According to the student, it has

a length of 1.98 cm, a width of 5.15 cm, and a height of 1.03 cm and a mass of

77.89 g. The metal is iron, which has a density of 7.80 g/cm3. What is the

percent error?

2. A technician is making a 0.500 M salt solution for an experiment. The

experiment will work as long as the solution is within 2.00 % of the accepted

0.500 M. In actuality, the solution was 0.522 M. Will the experiment work?

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Density

The formula for density is d=m/v

The density of wáter at 22 C is 1 g/ml

1 cm3= 1 ml

1. An irregularly shaped piece of metal with a mass of 125g is placed into a

graduated cylinder that contains 25.00mL of water. This raises the water level to

56.00mL. What is the density of the metal?

2. Calculate the mass of a block of iron (density = 7.86g/ml) with dimensions of

528mm x 67.4mm x 37.3mm.

3. The volume of a red blood cell is about 9.00x10-11

cm3. Assuming that red blood

cells are spherical, what is the diameter of a red blood cell in inches?

4. A fertilizer contains 21% nitrogen by mass. What mass of this fertilizer, in

kilograms, is required for an application of 775g of nitrogen?

5. A column of mercury is contained in a cylindrical tube. The tube has a diameter

of 8.0 mm, and the height of the mercury column is 1.20 m. Given that the

density of mercury is 13.6 g/mL and that the volume of mercury in the tube can

be calculated by the relation V = πr2h (r = the radius of the tube and h = the height

of the mercury column), calculate the mass of mercury present in the cylindrical

tube.

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Fundamental Chemical Laws and Dalton’s Atomic Theory

What Laws do Dalton’s atomic theory explain?

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

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Thomson’s Cathode Ray Tube

Experiments

Describe Thomson’s Cathode ray tube experiment and describe the major conclusions of

this experiment:

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

Describe this picture in terms of Thomson’s model

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Millikan’s Oil Drop Experiment

What was Milikan able to determine from his oil drop experiment?

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

Radioactivity

Explain the behavior of alpha, beta and gamma rays as illustrated in the diagram above

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

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Rutherford’s Gold Foil Experiment

Describe Rutherford’s Gold Foil Experiment and describe the major conclusions of this

experiment:

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

__________________________________________________________________

Atomic theory Summary

Scientist Experiment Knowledge gained Relating to

Thomson

Millikan

Rutherford

Chadwick

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Modern view of Atomic Structure

Atomic number & the number of electrons, protons, and neutrons.

XAZ

is the general symbol for an isotope where A is the mass number (number of

protons + number of neutrons), X is the element’s symbol, and Z is the atomic

number (number of protons). This information can also be found on the periodic

table. Use this information to fill in the chart for the following elements

Symbol Atomic

Number

Atomic

Mass Protons Electrons Neutrons Element Name

Gold

copper

nickel

lead

zinc

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Isotopes and Mass Spectroscopy

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Average Atomic Mass

Directions:

1. Write the isotopes.

2. Write the masses.

3. Multiply the masses by the decimal equivalent of the relative abundance.

4. Add the products. This is the average atomic mass.

Example:

A sample of cesium is 20.0 % 132

Cs ,75.00 % 133

Cs, and 5.0 % 134

Cs

132

Cs = 132 x 0.200 = 26.4 133

Cs = 133 x 0.7500 = 99.75 134

Cs = 134 x 0.050 = _6.7_

132.8

Determine the average atomic mass of the following mixtures of isotopes:

1. 80.0 % I-127, 17.0 % I-126, and 3.00 % I-128.

2. 50.0 % Au-197, 50.0 % Au-198.

3. Naturally occurring chlorine molecules (chlorine is diatomic) have masses of

70, 72 and 74. They occur in the percentages of 56.25%, 37.50% and

6.250%. What is the average atomic mass of chlorine atoms?

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Chemical Nomenclature Flowchart

Case 1: Ionic compounds containing monatomic ions (i.e. ions that can only have one charge)

Name of Compound = name of metal + name of non-metal w/ide suffix or name of polyatomic ion. No

prefixes are used!

e.g. NaF = sodium fluoride; Na3PO4 = sodium phosphate; (NH4)3PO4 = ammonium phosphate

Case 2: Ionic compounds containing a metal that can form more than one ion

Name of Compound = name of metal, followed by charge of metal in Roman numerals in parentheses,

followed by name of non-metal w/ -ide suffix or name of polyatomic ion. No prefixes are used!

e.g. PbCl2 = Lead (II) chloride; Cu(NO3)2 = copper (II) nitrate

Case 3: Binary molecular compounds:

Name of Compound = name of first element + name of second element with -ide suffix.

Use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, nona-, deca-) to indicate the number of atoms.

The mono prefix is not used with the first element.

e.g. CO = carbon monoxide; NO2 = nitrogen dioxide; N2O = dinitrogen monoxide; P2O5 = diphosphorus

pentoxide

Case 4: Binary acid solutions (i.e. binary acids dissolved in water = binary acids in aqueous solution)

Name of Compound = hydro + name of halogen w/ -ic suffix e.g. HF(aq) = hydrofluoric acid; HCl(aq) =

hydrochloric acid

Unless stated otherwise assume the formula of a binary acid is for the acid dissolved in water. E.g.

assume HCl = HCl(aq)

Naming Oxoacids (i.e. compound with the general formula HxMOy, where M = nonmetal)

The name of an oxoacid is based on the name of the polyatomic ion from which the acid is derived.

Case 5: -ate -ic

If the name of the polyatomic ion ends in “-ate,” the name of the corresponding acid ends in “-ic acid.”

Polyatomic ion (-ate) Acid (-ic)

sulfate = SO42- H2SO4 = sulfuric acid

Chlorate = ClO31- HClO3 = chloric acid

Case 6: -ite -ous

If the name of the polyatomic ion ends in “-ite,” the name of the corresponding acid ends in “-ous acid.”

Polyatomic ion (-ite) Acid (-ous)

sulfite = SO32- H2SO3 = sulfurous acid

Chlorite = ClO21- HClO2 = chlorous acid

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Write the formula.

1. Ammonium Nitrate

2. Calcium Bicarbonate or Calcium

Hydrogen Carbonate

3. Barium Chlorate

4. Dimercury (I) Iodide

5. Hydronitric acid

6. Plumbic Oxide or Lead (IV) Oxide

7. Potassium Thiocyanide

8. Perchloric acid

9. Sulfur hexafluoride

10. Zinc Hydroxide

11. Potassium Sulfite

12. Copper (I) Sulfide

13. Potassium bisulfate

14. Zinc Bromide

15. Ferric Chromate

16. Sodium Perchlorate

17. Potassium Hypochlorite

18. Magnesium Nitride

19. Sodium Permanganate

20. Potassium Permanganate

Write the name.

21. Hg2(CN)2

22. H2SO4

23. Fe(C2H3O2)2

24. KClO3

25. PbF2

26. HBr

27. N2O4

28. HgCrO4

29. Ag3PO4

30. K2Cr2O7

31. Ba3N2

32. NF3

33. CuBr2

34. (NH4)2S

35. Ca(NO3)2

36. Zn(OH)2

37. NaHCO3

38. PbO2

39. KClO4

40. Hg2I2

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Mixed Practice

Write the correct formula, Write I, A, or M to the left to state what type of compound it is

1. ______calcium fluoride _________________________

2. _______ferrous carbonate _________________________

3. _______aluminum hydroxide __________________________

4. _______dinitrogen pentaoxide ___________________________

5. _______chromic acid ___________________________

6. _______nitrogen monoxide ___________________________

7. _______plumbic sulfate ___________________________

8. _______hydrophosphoric acid _____________________________

9.________tetraphosphorus decoxide ____________________________

Write the correct name. IF THERE ARE TWO WAYS TO NAME A COMPOUND

USE BOTH!!

9. ______ LiS ___________________________________________

10. ______NH4Cl _______________________________________

11. _______SnSO4 ____________________________________________

12. _______ H3PO3 _____________________________________________

13. ________ P2O5 ____________________________________________

14. ________ CO ______________________________________________

15. _______H3N ______________________________________________

16. _______ CuS ________________________________________________

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Naming Acids

“ate” becomes “____ic acid” I ate something icky

“ite” becomes “____ous acid”

“ide” becomes hydro ___ic acid

**in sulfur acids add “ur” in phosphorous acids add “or”

Write formula and oxidation number of anion, write correct formula

for acid, name the acid:

anion Anion

formula

/ox #

Acid

formula

Acid name

bromate

periodate

carbonate

sulfide

chloride

chloric

thiosulfate

dichromate

chromate

permanganate

bromate

iodite

iodide

perchloride

phosphate

hyporchlorite

phosphide

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Simple Organic Compounds

Alkanes

Most basic type of hydrocarbon

Carbón bonded to 4 other atoms

General fromula is CnH2n+2

End in suffix –ane

Name derived from number of carbons in hydrocarbon chain. KNow the

following:

Part of name derived from # carbons # carbons in hydrocarbon chain

Meth- 1

Eth- 2

Prop- 3

But- 4

Pent- 5

Hex- 6

Hept- 7

Oct- 8

Non- 9

Dec- 10

Write the molecular formula for the following:

1. Propane___________________

2. Hexane ____________________

3. Octane _____________________

4. Methane _____________________

5. Butane _______________________

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Alcohols

Some hydrogen atoms in an alkane chain replaced with hydroxyl -OH

Name is derived from its parent alkane

Name ends with –ol

WHen there is more tan one carbón upon which the hydroxyl group can be

present, a number is placed in front of the name to inducate which carbón in the

hydrocarbon chain has the –OH group attached

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