welcome to advanced placement...

Welcome to Advanced Placement Chemistry! I am excited that you have decided to take, or are at least thinking about taking AP Chemistry. Just to let you know, this class is a lot of work, but you will gain a better understanding of chemistry because of it. So that we can get a head start and have plenty of time to cover all of the material on the AP exam, you will have a summer assignment to complete. The AP Chemistry curriculum changed last year so now I have a better understanding of what needs to be covered. The new curriculum requires a little less memorization and math. However, you will be expected to truly understand everything that has not been omitted at a much deeper level. Several new topics have been added, as well. The summer assignment is designed to keep basic chemistry concepts fresh in your brain so we can jump right in to new material when school starts up again. Your summer assignment will consist of reviewing basic chemistry concepts which are covered in the first two chapters of the textbook as well as reviewing and adding to your knowledge over gases. Summer assignments covering chapter one, book work from page 2 and the pages 5‐6 worksheet, and chapter two, book work from page 7 and the pages 11‐13 worksheet, will be due at the school by Monday, July 14th. Summer assignments covering naming, the pages 15‐16 worksheet, and chapter five, book work from page 17 and the pages 21‐23 worksheet, will be due at the school by Monday, August 4th. The following links are to podcasts I have created that help to support some of these topics: http://bit.ly/18qKLEq : Matter http://bit.ly/17IY7dO : Properties, Units and Scientific Notation http://bit.ly/18qKQYP : Significant Figures and Dimensional Analysis (Fence and Rail) http://bit.ly/18qKRMo : Atoms, Ions, and Nomenclature (Naming) The first week of school I will have returned your work and we will be briefly reviewing the material on the summer assignments. At the end of the week, we will have an exam covering all summer topics. Then, be ready to dive in and learn some chemistry! As always, feel free to email me any time if you have questions at [email protected]. I am looking forward to the upcoming year, and I hope you are, too. Mr. Burkett

Study guide for AP test on Chapter 1 Matter & Measurement

In order to be fully prepared you should seek help if required, refer to the relevant chapter in the textbook

and review ALL relevant notes, homeworks, worksheets, classwork and other materials

ALL students should;

Recall a definition of chemistry Understand the process and stages of scientific (logical) problem solving Recall the three states of matter, their general properties and the methods for their

interconversion Understand and recall definitions for physical and chemical change Know the difference between elements, mixtures and compounds including the difference

between heterogeneous and homogeneous mixtures Understand and be able to use scientific notation (standard form) Recall and use SI units and prefixes Be able to convert between units Understand the concept of derived units and use relationships relating to density Recall the meaning of uncertainty and understand and be able to use the rules for determining

significant figures and rounding off Understand the differences between, and be able to apply, the concepts of accuracy and

precision Learn, and be able to use, formula for the conversion of the three different temperature units

studied in TOPIC 1

End of Chapter Homework Problems

** Density, Conversions and Sig Figs 1.21, 1.25, 1.26, 1.29, 1.31, 1.33, 1.35, 1.57, 1.59 & 1.60 ** Dimensional Analysis 1.37, 1.39, 1.42 and 1.43

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Chapter 1AP Notes

I. Classification of Matter – anything that has mass and occupies space

A. Mixtures – combinations of 2 or more substances that are not chemically combined and can be separated by physical means (combined in no certain proportions).

1. heterogeneous – unevenly distributed composition 2. homogeneous – evenly distributed composition

B. Pure substances – matter that cannot be physically separated into simpler parts (combined in only certain proportions) 1. compounds – 2 or more chemically combined elements; can be chemically separated 2. elements – simplest form of matter, cannot be chemically separated into simpler form

II. Properties of Matter A. Physical – can be determined without changing a substance’s chemical makeup/identity

1. intensive – does not depend on sample size (used to identify matter) 2. extensive – depends on sample size

B. Chemical – characteristic ability of matter to react/change its chemical identity C. Energy – ability to do work

1. kinetic – energy of motion 2. potential – stored energy (in chemical bonds) or energy due to position

III. Changes of Matter A. Physical – change in physical property, but not chemical identity B. Chemical – change in a substances chemical identity; evidenced by:

1. change in energy (heat/light) 2. formation of precipitate 3. gas production 4. color change?

C. Energy – changes are always accompanied by a change in energy 1. exothermic – energy released 2. endothermic – energy absorbed

V. Quantitative Measurements – SI (metric) system – useful measurements include A. Factor - number

1. significant digits a) all non-zero digits are significant b) zeros between sig. figs. are significant c) zeros at the end of a number AND after a decimal are significant d) zeros that act as placeholders for a decimal (before and after) are not significant e) all counting numbers have an infinite number of sig. figs.

2. rounding numbers – answers to calculations must be rounded to the correct number of sig. figs. so that they do not express more accurate measurements than the numbers from which they were calculated. a) adding/subtracting – answer cannot have sig. figs. further toward right than any numbers in the calculation b) multiplying/dividing – answer must have same number of sig. figs. as number in calculation with the fewest sig. figs.

3. scientific notation – abbreviated way of expressing very large/small numbers 602,000,000,000,000,000,000,000 = 6.02 x 1023 0.00000456 = 4.56 x 10-6

B. Units 1. fundamental units – based on actual physical standards

a) mass – measures amount of matter (kilograms or grams) 1 kg cylinder weight – pull of gravity on an object b) length – (meter) 1/10 millionth of dist. From equator to N. pole c) temperature – measure of avg. kinetic energy (Kelvins = oC + 273) boiling/freezing point of water.

2. derived units – from combinations of fundamental units

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a) volume – length x height x width (1 dm3 = 1 L = 1000 mL = 1000 cm3) b) density – mass/volume (g/mL)

3. prefixes- mega- M 1,000,000 106 (million) kilo- k 1,000 103 (thousand) BASE UNIT 1 1 deci- d 0.1 10-1 (tenth) centi- c 0.01 10-2 (hundredth) milli- m 0.001 10-3 (thousandth) micro- 0.000 000 1 10-6 (millionth) nano- n 0.000 000 000 1 10-9 (billionth)

4. unit conversion – dimensional analysis (fence-and-rail) C. Indications of Uncertainty

1. accuaracy – measure of nearness to an accepted value % error = (experimental value – accepted value) x 100 accepted value

2. precision – measure of nearness to other measurements of the same thing a) deviation from the mean – how far apart one measurement is from the average measurement b) standard deviation from the mean – +/- average of the absolute values of all deviations; expresses a range around avg. in which true value should be Group Mass Abs. Dev. 1 3.60 0.06 2 3.70 0.04 3 3.80 0.14 4 3.70 0.04 5 3.50 0.16 Total 18.30 0.44

Avg. 3.66 +/- 0.09

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© Adrian Dingle’s Chemistry Pages 2004, 2005, 2006. All rights reserved. These materials may NOT be copied or redistributed in any way, except for individual class instruction.

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AP WORKSHEET 1a: Matter & Measurement 1. Classify each of the following as either, an element , a compound or a mixture . If you classify something as a mixture then also state whether it is a homogeneous or a heterogeneous mixture. (10)

(i) Helium

(ii) Nitrogen

(iii) Pure water

(iv) Pure table salt (sodium chloride)

(v) Coke

(vi) Air

(vii) Fruit cake 2. Classify the following as either chemical or physical changes. (3)

(i) Ice melting

(ii) Gasoline burning

(iii) Evaporation of perfume from an open bottle 3. Mercury is a liquid metal that has a density of 13.58 g/mL. Calculate the volume of mercury that must be poured out in order to obtain 0.5000 g of Mercury. (2)

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4. Classify the following as either quantitative or qualitative observations. (4)

(i) My eyes are brown

(ii) My neck size is 17 inches

(iii) My average grade last year was 79%

(iv) Physics is a difficult subject 5. Give an example of a natural law (other than the law of conservation of mass). (1) 6. Convert these numbers to scientific notation. (2)

(i) 35800000000000

(ii) 0.00000000821 7. Round the following numbers to four figures. (6)

(i) 2.16347 x 105

(ii) 4.000574 x 106

(iii) 3.682417

(iv) 7.2518

(v) 375.6523

(vi) 21.860051

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Study guide for AP test on Chapter 2 Ions & Nomenclature

In order to be fully prepared you should seek help if required, refer to the relevant chapter in the textbook

and review ALL relevant notes, homeworks, worksheets, classwork and other materials


Recall a very brief history of Atomic Theory Know and understand the five main aspects of Dalton's Atomic Theory Recall some of the experiments that led to the identification of sub-atomic particles Know the three particles that make up the atom and their relative charges, masses and positions

in the atom Be able to use the Atomic # and Mass # of an isotope to calculate the numbers of protons,

neutrons and electrons present Know what the term isotope means and be able to perform simple calculations relating to isotopic

data Know the approximate locations of metals, non-metals and metalloids on the periodic table Understand the meaning of the terms Molecule and Ion Learn the lists of common anions and cations (including polyatomic ions) studied in TOPIC 3 Know how to combine those anions and cations in the correct proportions to form ionic

compounds with no net charge Be able to name binary ionic compounds of a metal and a non-metal Be able to name binary molecular compounds of two non-metals Be able to name simple binary acids Be able to name ionic compounds containing polyatomic anions Be able to name oxoacids and compounds containing oxoanions Be able to name hydrated salts

End of Chapter Homework Problems

**Chemical Formulas 2.35, 2.38, 2.39, 2.40, 2.45 & 2.46

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Chapter 2 AP Notes I. History

A. Democritus (400 B.C.) – world is made of two things 1. empty space 2. atomos – smallest “indivislible” particles 3. matter made of different types of atomos

B. Aristotle (300 B.C.) – matter is uniform and continuous 1. hyle – can be divided into infinitely smaller particles 2. matter made of different combinations of earth, air, wind and fire

C. Robert Boyle (1627-1691) – “first chemist” 1. evidence for atomic nature of matter 2. defined element – substance that can’t be chemically divided 3. predicted existence of several different elements

D. Joseph Priestly (1733-1804; 1774) – isolated the element oxygen E. Antoine Lavoisier (1743-1794)

1. proved oxygen is a key component of combustion 2. Law of conservation of mass (mass of combustion products = mass of combustion reactants)

F. Joseph Proust (1754-1826) 1. Law of Definite Proportions – chemical compounds contain elements in specific mass ratios

G. John Dalton (1766-1844; 1808) Atomic Theory Postulate: Explains: Exceptions:

1. All elements are composed of very small particles called atoms.

1. Democritus 2. Boyle 3. Priestly

None

2. All atoms of a given element are identical in properties (mass), but different than atoms of other elements.

1. Democritus 2. Boyle 3. Priestly 4. Proust

1. Isotopes – atoms of same element with different masses (numbers of neutrons).

3. The identity of atoms does not change in chem. rxns. (atoms are combined, separated or rearranged).

1. Lavoisier 1. Nuclear rxns. (fission or fussion) 2. E = mc2 (energy __> mass)

4. Compounds are formed when atoms of more than one type combine, always in certain small whole-number ratios.

1. Proust 1. Law of Multiple Proportions (Dalton) – same elements can combine in different ratios to make different compounds (CO and CO2, H2O and H2O2)

II. Structure A. “Plum-pudding” Model

1. J.J. Thomson (1897) – cathode ray tube experiment a. all matter contains electrons (small negatively charged particles) b. charge to mass ratio of e- = 1.76 x 108 Coulmbs/gram

2. James Millikan (1909) – oil drop experiment a. measured charge of e- = - 1.60 x 10-19C b. calculated mass of e- 1.76 x 108 C/g = 1.60 x 10-19C mass = 9.10 x 10-28 g mass (g)

3. “Plum-Pudding” Model = e- embedded in a positively charged solid sphere B. Nuclear Model

1. Rutherford (Marsden & Geiger) (1910) – gold foil experiment (5 x 10-4 cm/17,500 atoms thick) a. most of atom is empty space b. all of (+) charge and most of the mass are in small dense center (nucleus) - size of nucleus = “pea in a stadium” (approx. 10-10 m)

- density = 1 Tbs. Is approx. 2.5 billion tons (1013 to 1014 g/mL)

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2. subatomic particles Particle: Charge: Mass: Location:

C relative u = 10-10 g relative Proton + 1.60 x 10-19C +1 (p+) 1.0073 1 Nucleus

Neutron none 0 (n0) 1.0087 1 Nucleus Electron - 1.60 x 10-19C -1 (e-) 5.486 x 10-4 0 Empty space

3. isotopes – atoms of same element (same # protons) with different masses ( different # neutrons)

a. atomic number (Z) – # of p+ in nucleus (# e- in empty space); determines element’s identity b. mass number (A) = # of p+ + # of n0

Isotopes of Hydrogen Name: Symbol: Z A # of p+ # of n0 # of e-

Hydrogen-1 11H 1 1 1 0 1

Hydrogen-2 21H 1 2 1 1 1

Hydrogen-3 31H 1 3 1 2 1

4. atomic masses

a. relative atomic mass – mass (in u; atomic mass units) compared to carbon-12 (12.0 u) b. average atomic mass – takes into account relative abundance of isotopes of an element ex. 69.17% of Cu is copper-63; relative atomic mass = 62.939589 u 30.83% of Cu is copper-65; relative atomic mass = 64.927793 u copper-63: (62.939589 u x 0.6917) = 43.54 u copper-65: (64.927793 u x 0.3038) = 20.02 u average atomic mass = 63.57 u

III. Ionic Compounds – formed due to electrostatic attractions between oppositely charged ions A. Ions

1. monatomic – single atoms with (+) or (-) charge shown by superscripts; formed when atoms gain/lose e- to attain octet in valence shell (noble gas config.) a. cations – atoms that lose e-; have (+) charge Na ___> Na+1 + e- Na+1 = sodium ion * some have more than one possible charge; shown in name with Roman numerals Cu+1 = copper (I) ion Cu+2 = copper (II) ion b. anions – atoms that gain e-; have (-) charge Cl + e- ___> Cl-1 Cl-1 = chloride ion

2. polyatomic - group of covalently bonded atoms with a (+) or (-) charge most are oxyanions – contain oxygen + some other element(s) and have (-) charge ex. ClO2

-1 = chlorite ion (-ite = one less O than –ate) ClO3

-1 = chlorate ion (-ate = most common form of ion) ClO4

-1 = hyperchlorate ion (hyper- = one more O than –ate) B. Formulas & bonding – ions always combine so that compound is neutral, (+) = (-)

1. Write formulas for ions – cation (least EN ion) first Mg+2 Cl-1 2. Cross charges over to find subscripts Mg1 Cl2 3. Write empirical formula (subscripts in lowest whole-# ratio MgCl2 Mg+2 + PO4

-3 ___> Mg3(PO4)2 C. Nomenclature

1. binary ionic – monatomic cation + monatomic anion name = cation’s name + anion’s name MgCl2 = magnesium chloride CuCl2 = copper (I) chloride CuCl = copper (II) chloride

2. ternary ionic – include one or more polyatomic ions name = cation’s name + anion’s name Mg3(PO4)2 = magnesium phosphate

3. hydrates – water molecule bound in ionic crystal structure name = cation’s name + anion’s name + prefix + “-hydrate” Na2SO4

. 10 H2O = sodium sulfate decahydrate

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4. acids a. binary – hydrogen + monatomic anion name = “hydro-“ + root of monatomic anion + “-ic” + acid HCl = hydrochloric acid b) ternary (oxyacids) – hydrogen + polyatomic anion (oxygen and some other nonmetal) name = root of polyatomic ion + suffix + acid

suffix: “-ite” on polyatomic ion ___> “-ous” “-ate” on polyatomic ion ___> “-ic” HNO2 = nitrous acid HNO3 = nitric acid

IV. Covalent bonds – formed by sharing of e- by two atoms; shared on overlapping orbitals, area of high e- density attracts nuclei and holds atoms together

A. Nomenclature 1. Binary molecular cmpds. – two nonmetals name = prefix +”least EN” nonmetal + prefix + root of 2nd nonmetal + “-ide” prefixes: one = mono- six = hexa- two = di- seven = hepta- three = tri- eight = octa- four = tetra- nine = nona- five = penta- ten = deca- CO = carbon monoxide CO2 = carbon dioxide P2O5 = diphosphorus pentoxide

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AP WORKSHEET 2a: Atomic Structure 1. Using the periodic table, complete the following table. (13)

Isotope Symbol Atomic Number Mass Number # of Protons # of Neutrons # of Electrons

9 18

24Na11

35 44

2. What is the charge on a sodium atom ? (1) 3. What is the charge on a sodium nucleus ? (1) 4. What is the atomic number of potassium? (1) 5. How many protons in the nucleus of a potassium atom? (1) 6. How many electrons in the potassium nucleus ? (1)

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7. Define the term isotope . (2) 8. Using the periodic table; (a) Complete the following table. (22) Isotope Symbol Atomic # # Protons # Neutrons Mass #

13C

17 18

26 56

17 37

2 3

52 128

50 70

(b) Consider the second and fourth row in the table above. What do they have in common? (2) (c) Consider the second and fourth row in the table above. What are the differences? (2) 9. Naturally occurring Ni is found to have the following approximate isotopic abundance. 58Ni 68%, 60Ni 26%, 62Ni 4% and 61Ni 2%. Calculate the average relative atomic mass of Ni from the data. (2)

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10. The results taken from a mass spectrum of chlorine gas show peaks at m/e 35.00 and m/e 37.00 (The mass spectrum identifies the different isotopes of an element that are present in a sample). (a) Given that the relative abundances of Cl 35.00 and Cl 37.00 are 77.50% and 22.50% respectively, calculate the average relative atomic mass of chlorine atoms to four significant figures . (2) (b) Suggest all the possible masses of CI2 molecules , that are made when two chlorine atoms bond together. (3) (c) Which of the molecules you have suggested in (b) will be the most abundant? Explain your answer. (2)

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Polyatomic IonsName Symbol Charge

ammonium NH4 +1acetate C2H3O2 -1bromate BrO3 -1chlorate ClO3 -1chlorite ClO2 -1cyanide CN -1

dihydrogen phosphate H2PO4 -1hypochlorite ClO -1

hydrogencarbonate(bicarbonate) HCO3 -1hydrogen sulfate (bisulfate) HSO4 -1hydrogen sulfite (bisulfite) HSO3 -1

hydroxide OH -1iodate IO3 -1nitrate NO3 -1nitrite NO2 -1

perchlorate ClO4 -1permanganate MnO4 -1

thiocyanate SCN -1carbonate CO3 -2chromate CrO4 -2

dichromate Cr2O7 -2oxalate C2O4 -2selenate SeO4 -2silicate SiO3 -2sulfate SO4 -2sulfite SO3 -2

phosphate PO4 -3phosphite PO3 -3

Rules for Naming Ionic Compounds1. Balance Charges (charges should equal zero)2. Cation is always written first ( in name and in formula)3. Change the ending of the anion to -ide

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AP WORKSHEET 2f: Inorganic Nomenclature II Add either a name or a formula to complete each table.

1. Potassium dichromate

2. Lithium sulfide

3. Potassium bromide

4. Cesium iodide

5. Calcium phosphide

6. Sodium fluoride

7. Strontium oxide

8. Beryllium sulfide

9. Magnesium bromide

10. Lithium oxide

11. Strontium chloride

12. Barium bromide

13. Magnesium sulfide

14. Magnesium iodide

15. Hydrogen fluoride (Hydrogen monofluoride)

16. Barium phosphide

17. Sodium hydrogen phosphate

18. Potassium chloride

19. Lithium nitride

20. Calcium sulfide

21. Rubidium oxide

22. Strontium nitride

23. Cesium phosphide

24. Magnesium carbonate

25. Beryllium sulfate

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51. ScCl3

52. HCl

53. PtO2

54. Sb(ClO3)5

55. GeS2

56. ZnO

57. VSO4

58. CuCl2

59. TiO2

60. NiN

61. Ni3(PO4)2

62. CoF3

63. Au2O3

64. Zn3P2

65. Cr(NO3)6

66. NaIO2

67. NaIO3

68. NaI

69. H2SO3

70. H2CO3

71. AlN

72. AlH3

73. Li3AsO4

74. NaCN

75. Na2O2

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Study guide for AP test on TOPIC 15 Gases Chapter 5

In order to be fully prepared you should seek help if required, refer to the relevant chapter in the

textbook and review ALL relevant notes, homeworks, worksheets, classwork and other materials


Be able to convert between different units of pressure Be able to convert between different units of temperature Recall and be able to use Boyle's law in calculations Recall and be able to use Charles' law in calculations Recall and be able to use Gay-Lussac's law in calculations Recall and be able to use Avogadro's law in calculations Recall and be able to use the Combined gas law and the General gas law in calculations Recall and be able to use the Ideal gas law in calculations Understand and be able to use the van der Waals equation (modified ideal gas law) in

calculations Recall and be able to use Dalton's law of partial pressures in calculations Recall the conditions that are used as standard in calculations Be able to use molar gas volume in calculations Understand the Kinetic theory as applied to gases Understand the concept of, and be able to perform calculations involving, the root-mean-

square-speed of gases Understand the terms effusion and diffusion and be able to perform calculations relating

to those concepts

End of Chapter Problems Pressure and Gas Laws 5.14, 5.20, 5.24 & 5.26 Ideal Gas Equation 5.42, 5.46, 5.47 & 5.49 Gas Stoichiometry 5.56 & 5.60 (these are more difficult so I attached the solution, but try them first) Dalton’s Law of Partial Pressures 5.63 & 5.64 Kinetic Molecular Theory of Gases 5.77 & 5.78 Additional 5.94

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Chapter 5 AP Notes I. Variables

A. Pressure (P) = force/area = (m x a)/area = (kg x m/s2)/m2 = (kg/m x s2) = Pascals air P on 1 m2 at sea level at 0oC = (10,300 kg x 9.8 m/s2) = 101,300 Pa = 101.3 kPa 1 m2 1. barometer (atmospheric pressure)

101,300 Pa = 101.3 kPa = 760 mm Hg = 760 torr = 1 atm = 1.10 bar 2. manometer (gas pressure in a closed container - closed: PGAS = ht. (PH) - open: PATM > PGAS: PGAS = PATM - PH PATM < PGAS: PGAS = PATM + PH B. Volume (V) 1 dm3 = 1 L = 1000 mL = 1 cm3

A. Temperature (T) Kelvins = oC + 273 B. Moles (n) C. Standard Temperature & Pressure

1. T = 0oC = 273 K 2. P = 1 atm…

D. Standard Molar Volume – volume of 1 mole of any gas at STP = 22.4 L Avogadro’s hypothesis – if two gases have same n, P and T then they have the same V

II. Gas Laws A. Combined Gas Law P1V1 = P2V2 (* T in Kelvins) T1 T2

1. Boyle’s Law – V is inversely proportional to P (when T and n are constant) V 1/P ___> V = k (1/P) ___> P1V1 = k = P2V2

___> P1V1 = P2V2 (k = constant)

2. Charles’ Law – V is directly proportional to T (when P and n are constant) V T ___> V = k (T) ___> V1/ T1 = k = V2/ T2

___> V1/ T1 = V2/ T2 (T in Kelvins)

3. Gay-Lussac’s Law - P is directly proportional to T (when V and n are constant) P T ___> P = k (T) ___> P1/ T1 = k = P2/ T2

___> P1/ T1 = P2/ T2 (T in Kelvins)

4. Avogadro’s Law V is directly proportional to n (when P and T are constant) V n ___> V = k (n) ___> V1/ n1 = k = V2/ n2

___> V1/ n1 = V2/ n2 B. Ideal Gas Law PV = nRT R = ideal gas constant R = PV/nT = (1 atm)(22.4 L) = 0.0821 atm L/mol K (1 mol)(273 K) (* P = atm; V = L; n = moles; T = Kelvins) 1. Molar mass (M) = mass (g)/mol ___> mol (n) = m / M PV = nRT ___> PV = mRT ___> M = mRT M PV 2. Density (D) = mass/vol. = m/V M = mRT ___> m = MP PV V RT

A. Dalton’s Law of Partial Pressures – total pressure of a gas mixture = sum of partial pressures PT = P1 + P2 + P3…+ Pn

The fraction of P1/PT = mole fraction of n1/nT B. Graham’s Law of Effusion (escape of gas through a small hole in a container)/Diffusion (mvmt. of

gas from areas of high ___> low conentration) 1. Avg. molecular speed (rms – root mean square) = u = _ / 3RT (R = 8.31 kg m2/s2 K mol) \/ M (M = kg / mol)

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2. Comparisons of 2 different gases at the same T (same KE) KEA = KEB

½ MavA2 = ½ MBvB

2 MavA

2 = MBvB2

vA2 = MB

vB2 MA

vA = _ / MB vB \/ MA

ex. Compare the rates of effuison/diffusion of H2(g) and O2(g) at the same temp. vH2 = _ / 32.0 g = _ / 16.0 = 4 H2 is 4.0 x faster than O2

vO2 \/ 2.0 g \/ 1 III. Kinetic Molecular Theory (KMT) for Ideal Gases A. Types of Motion

1. translational (pt. A ___> pt. B) 2. rotational 3. vibrational

A. Postulates 1. molecules are point masses – have mass with negligible volume 2. molecules move randomly in all directions at various speeds 3. attractive/repulsive intermolecular forces are negligible 4. collisions are elastic – no loss in avg. KE (avg. velocity) KE = ½ mv2 5. Avg. KE Temp.

C. Ideal Gases (imaginary) - follow all postulates of KMT under all conditions A. Real Gases – follow KMT except when gas molecules are very close (at high P, low T = low V)

1. molecules will have significant volumes under these conditions (postulate # 1) 2. IMF become significant under these conditions (postulate # 3) and

- inhibit movement (postulates # 2 and 4) - collisions are not elastic (postulate # 4) - reduce P

3. Van der Waals equation – accounts for P and V deviations from KMT for real gas due to IMF (P + [a(n)2/V2]) (V – bn) = nRT

IV. Gas Stoichiometry – volumes of gaseous reactants and products are related in the same way at moles of reactants and products.

N2(g) + 3 H2(g) ___> 2 NH3(g)

1mol + 3 mol 2 mol 1 L + 3 L 2 L

Converting moles volume - at STP use standard molar volume (1 mole = 22.4L) - at non STP use ideal gas eqn. vol. ___> mole: n = PV/RT mole ___> vol.: V = nRT/P ex. 50 g NH3 = ? L H2 at STP 50 g NH3 | 1 mol NH3 | 3 mol H2 | 22.4 L H2 = 98.8 L H2 | 17 g NH3 | 2 mol NH3 | 1 mol H2 ex. 50 g NH3 = ? L H2 at 250 K and 1.6 atm 50 g NH3 | 1 mol NH3 | 3 mol H2 = 4.4 mol H2 | 17 g NH3 | 2 mol NH3 V = (4.4 mol)(0.0821)(250 K) = 56.4 L H2 (1.6 atm)

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AP WORKSHEET 6b: Gas Laws In Questions 1 through 3 the temperature and the amount of gas are both constant. 1. Which gas law relates volume and pressure? Express the law mathematically? (2) 2. Calculate the new pressure if a 2.45 L of a gas at a pressure of 1.01 atm is contracted to a volume of 2.29 L. (1) 3. Calculate the new volume if a 13.3 L of a gas initially at a pressure of 2.51 atm is subjected to an increase in pressure equivalent to 65.0 mmHg. (2) In Questions 4 through 6 the pressure and the amount of gas are both constant. 4. Which gas law can relates volume and temperature? Express the law mathematically. (2) 5. Calculate the new volume of a particular gas if 1.23 L of it, initially at a temperature of 32.0 oC is subjected to a drop in temperature of 19.0 degrees Celsius. (1) 6. Calculate the new volume of a gas if a 12.78 L of it, initially at a temperature of –50.00 oC is heated to a temperature of 28.00 oC. (1)

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In Questions 7 and 8, the pressure and temperature of gas are both constant. 7. Which gas law can relates volume and number of moles? Express the law mathematically. (2) 8. What mass of nitrogen gas occupies a volume of 11.2 L, if 4.20 g of nitrogen occupies 100. L? (2) In questions 9 and 10, assume the gas behaves ideally. 9. A sample of a group I bromide weighing 2.000 g was converted to a gas at 504.0 oC and 1.000 atm pressure. The resulting vapor occupied a volume of 1238 mL. Identify the group I metal present in the compound. (3) R = 0.0821 (atm L K-1 mol-1) 10. What volume does 1.24 g of Fluorine gas occupy under conditions of 5.20 oC and 2.04 atm? (1) 11. If 5.0 g of nitrogen gas and 5.0 g of chlorine gas are injected in to a 2.0 L vessel at a temperature of 65 oC, what will the partial pressure of each gas be? What will the total pressure in the container be? (3) 12. Which gas law allows the simple calculation of the total pressure in question #11? (1)

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In questions 13 and 14, assume 1.000 mol of any gas occupies 22.40 L at standard temperature and pressure (s.t.p). 13. What volume of hydrogen gas is obtained when 23.00 g of zinc metal reacts with an excess of dilute sulfuric acid at s.t.p? (2) Zn(s) + H2SO4(aq) � ZnSO4(aq) + H2(g) 14. What volume of oxygen, at s.t.p, is required to burn exactly 11.60 L of methane (CH4(g)), according to the reaction below? (2) CH4(g) + 2O2(g) � CO2(g) + 2H2O(g)

15. What is the Molar Mass (RMM) of a gas that diffuses at 51

the rate that hydrogen diffuses at?

(1) 16. How much faster does Helium gas escape through a porous container than SO2? (1) 17. Calculate and compare the urms for Helium and Nitrogen at 20 oC. Comment on the two values. (4)

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