Welcome to AP Chemistry!€¦ · AP Chemistry Summer Assignment 2011‐2012 Page 1 ... so that...

AP Chemistry Summer Assignment 2011‐2012

Welcome to AP Chemistry!

AP Chemistry is a challenging course. The following summer assignment is designed to help us make the best use of our class time. You need to refresh what you learned in Chemistry Honors so that you’re ready to move forward from there. There will be a quiz on Chapters 1 and 2 and a test on chapters 3 and 4 during the first few days of school In addition to the actual summer assignment there are three unit pre-assignments that will not be due until just before we start those actual units during the school year. If you have a demanding schedule this fall or struggled with those units last year, I recommend that you work your way through the pre-assignments this summer.

The paper and pen information and exercises are designed to be read and completed as you read the textbook. After you have worked the simpler pencil and paper exercises, move on to the electronic homework (CALM) assignment. Jumping right into the CALM assignment can lead to frustration. Trust me. It’s happened.

The on-line homework system we’ll be using is called CALM. The website is CALM.indiana.edu. Ensure that you can get into the system and complete the trial problems BEFORE leaving school for the year. Your user name is your Firstclass login (eg sbuchanan12) and your password is your student id #. Most problems in CALM care about significant digits, and scientific notation is expressed in the format 3.42E-5.

Please complete the following tasks: Section Description Submission Date Done? Check on CALM Pre-summer assignment – CALM login June 13!!!!! Chapter 1 - intro Read Chapter 1 in BLB Chapter 1 paper and pencil September 7 CALM Chapter 1 August 14 Chapter 2 - molecules

Read Chapter 2 in BLB

Chapter 2 paper and pencil September 7 CALM Chapter 2 August 21 Chapter 3 - stoich Read Chapter 3 in BLB Chapter 3 paper and pencil September 7 CALM Chapter 3 August 28 Chapter 4 - reactions Read Chapter 4 in BLB Chapter 4 paper and pencil September 7 CALM Chapter 4 September 3 Chapter 10 - Gases Read Chapter 10 in BLB Chapter 10 paper and pencil pre-

assignment ~ September 16

Chapter 10 CALM Section 1 ~ September 16 Chapter 5 - Thermo Read Chapter 5 in BLB Chapter 5 paper and pencil pre-


Chapter 6 – QM Read Chapter 6 in BLB Chapter 6 paper and pencil pre-



Name ____________________________ Fill in the answers to the questions/prompts that follow as you read your textbook. Turn in filled in papers for the summer assignments (Ch 1-4) on the first day of school. After completing the paper assignments, go on to the CALM on-line assignments. Pre-assignments for Chapters 10, 5, and 6 will be turned in just before we start those chapters. Chapter 1: Introduction: Matter and Measurement The Study of Chemistry: State the difference between atoms, elements, and molecules: What is matter? The Classification of Matter: Define and give examples of the three States of Matter:

1.

2.

3. What is the difference between a compound and a mixture? Properties of Matter: Explain the difference between intensive and extensive properties: Explain both physical and chemical change Physical: Chemical: Is a change in the state of matter a physical or chemical change? List three or more techniques for separating mixtures, and the physical property on which that technique is based (ex: distillation, boiling points):

1. 2. 3. 4.


Units of Measurement: SI Units:

Physical Quantity Name of Unit Abbreviation Define what a derived unit is and give 2 examples: Uncertainty in Measurement: What is the difference between precision and accuracy? How many significant figures? 7.9152 ______ 0.00039 ______ 0.0400 ______ 5.470x103 ______ Do calculations (be sure to leave answers with correct number of significant figures): 2.45 x 5.427 = ________ 0.00230 x 2.1 = ________ 7.355 + 23.1 = ________ 435.423 – 237.91 = ________ Dimensional Analysis: Do calculations (be careful with units): The speed of a molecule in air is 137 m/s. Convert this speed to kilometers per hour using dimensional analysis. Show all work. How many moles do you have if you have a 54.3 g sample of Neon? (molar mass of Neon: 20.18 g/mol) What is the mass in grams of 7.00 in3 of solid zirconium? (density of Zr = 6.52 g/cm3 , 1 in. = 2.54 cm)


Summer Assignment Chapter 2 Vocabulary

Isotopes Atomic number Mass number Nuclide Proton Neutron Electron Atoms Molecular formula Empirical formula Ionic compound

a. Atoms of the same elements that differ in mass number.

b. Negatively charged particles, first subatomic particles discovered

c. Chemical formula that indicates the actual numbers and types of atoms in a molecule

d. Smallest particle of an element that retains he elements chemical identity

e. Neutrally charged particle, found in the nucleus in an atom

f. Number of protons in an atom g. This formula gives the relative numbers of

each type of atom in a molecule h. Electrically neutral, containing both

cations and anions, usually contain metallic and nonmetallic elements

i. Positively charged particles found in the nucleus

j. The total of protons plus neutrons in an atom

k. A nucleus of a specific isotope of an element

Fill in the Notes 2.1 Total mass of materials present after a chemical reaction is the same as the mass before the reaction 2.2 Is a fast-moving, positive charge, larger than β. High speed electrons, charge of -1 High-energy radiation, no charge, and does not consist of particles 2.5 Are located on the left side and middle of the periodic table Are elements in a column that have similar properties. Fall in line between metals and nonmetals and have properties of both.


2.6 Is a molecule made up of 2 atoms. Some examples are oxygen, nitrogen, and halogens. 2.7 ______ atoms tend to lose electrons to form cations are combinations of metals and nonmetals are generally composed of nonmetals only are compounds that are not combinations of C, H, O, N, S Naming & Formulas – the following represent examples. More practice is available in your book. You do not need to memorize all the names for polyatomic ions, you do need to understand the logic of the naming rules Formulas of Ions: Names of Ions: NH4

+

Hydronium ion

SO42-

Perchlorate ion

Nitrate ion

Carbonate ion

Ionic compounds: learn predictable (no roman numeral needed) vs unpredictable rules (Go Roman!) Formula Name

Cu(ClO4)2

CaCl2

Iron (II) sulfate

Cu2SO4

Aluminum Hydroxide

Acids: learn ic/ate and ous/ite relationships Acid and name: conjugate base and name: HClO4

HClO3

HClO2

HClO

HCl

HSO4-

HSO3-

HS-


Molecular compounds (go Greek!):

Name Formula

Dinitrogen tetroxide

Tetraphosphorus decasulfide

SO2

CO

H2O

Additional Notes for organic chemistry

Single bonds Double bond Triple Bond Mikey 1 carbon Methane NA! NA Eats 2 carbons Ethane ethene ethyne Peanut 3 carbons Propane propene propyne Butter 4 carbons Butane butene butyne After that it all makes sense… Pent- 5 carbons Pentane etc. Hex- 6 carbons etc. Hept- 7 carbons Oct- 8 carbons Non- 9 carbons Dec- 10 carbons Write the molecular and draw the structural formula for: Ethane Ethene Ethanol Propane Propene Propanol Propyne


Chapter 3: Stoichiometry: Calculations with Chemical Formulas and Equations

Terms and Concepts • Balancing equations • Law of conservation of mass • Identifying types of reactions: combustion, decomposition, combination • Calculating percentage composition:

o %element = (number of atoms)(atomic weight of element) x 100% Formula weight of compound

• Mole = Avogadro’s number = 6.02 x 1023 • Determining molar mass • Determining empirical formulas • Understanding quantitative information about balanced equations (ex: how many moles or grams

produced) • Limiting reagents • Calculating theoretical and percent yield.

Balance the following equations. (3.1)

1. __ Na + __ Cl2 __ NaCl

2. __ (NH4)3PO4 + __ Pb(NO3)4 __ Pb3(PO4)4 + __ NH4NO3

3. __ MnO2 + __ ClO3- __ MnO4

- + __Cl2

4. __ HCl + __Na2CO3 __NaCl + __CO2 + __H2O Label each reaction as a combination reaction, decomposition reaction, or combustion reaction. Note that some reactions fall into more than one category (3.2)

1. CH4 + 2O2 CO2 + 2H2O _____________________

2. H2SO3 H2O + SO2 _____________________

3. NH4OH NH3 + H2O _____________________

4. 2CO + O2 2CO2 _____________________

5. CaO + H2O Ca(OH)2 _____________________

6. C10H8 + 12O2 10CO2 + 4H2O _____________________

7. 8Fe + S8 8FeS _____________________

8. 2H2O 2H2 + O2 _____________________


Calculate the formula weight in grams to the nearest tenth decimal. (3.3) 1. KH ___________

2. Glucose (C6H12O6) ___________

3. Baking Soda (NaHCO3) ___________

4. NH3 ___________

5. NaC3H5O3 ___________

Calculate the percentage composition to the appropriate number of significant digits. (3.3)

1. KMnO4 K = _____, Mn = ______, O = _____

2. N2O5 N = _____, O = _______ Calculate the number of moles or mass using significant digit rules. (3.4)

1. 5.9 g of H2 ______ moles

2. 24.8 g of CO2 ______ moles

3. 0.1245 moles of water ______ grams

Determine the empirical formula (3.5) 1. 32.13% aluminum; 67.87% fluorine ______

2. Na 29.1 %, S 40.5% and O 30.4% ______

3. K 42.4%, Fe 15.2%, C 19.5% and N 22.8% ______

Determine the numbers of moles or grams of the product produced in the reaction (3.6)

2C4H10 + 9O2 4CO2 + 10H2O 1. If 2.5 moles of C4H10 is burned in excess oxygen, how many moles of water are produced?

______

2. How many moles of carbon dioxide are formed if 4.0 moles of C4H10 is burned in excess oxygen? ______

3. If 3.0 moles of carbon dioxide was formed, how many grams of butane were burned? ______


Calculate the number of moles or grams of the product produced in a limiting reagent reaction. (3.7) Mg(OH)2 + 2HCl MgCl2 +2H2O

1. If you have 10.0 g of Mg(OH)2 and 45.0 g of HCl, how many moles of MgCl2 will be produced? ______

2. How many grams of HCl will remain from the previous problem? ______

3. How many grams of Mg(OH)2 is needed to react completely with 20.0 g of HCl? ______

Determine the percent yield. (3.7) 1. A student adds 200.0 g of C7H6O3 to an excess of C4H6O3. The products are C9H8O4 and C2H4O2.

Calculate the percent yield if 231 g of aspirin (C9H8O4) is produced. ___________

2. Calculate the percentage yield if 550.0 g of C7H8 is added to an excess of nitric acid and 305 g of C7H7NO2 is formed. What is the percentage yield? ____________

3. Calculate the percent yield if 384 grams of S6 (s) is reacted with exess oxygen and 680 grams of sulfur dioxide is produced. ____________


Chapter 4 Aqueous Reactions and Solution Stoichiometry Vocabulary – do a mental check of your understanding of the following words, look them up in your book’s glossary, and write the definitions for any words you did not correctly define Aqueous solution- Solvent- Solute- Electrolyte- Nonelectrolyte- Chemical Equilibrium- Spectator Ions- Acids- Bases- Neutralization- Oxidation-reduction- Oxidation number- Equivalence Point- Indicators- 4.1 – General Properties of Aqueous Solutions ___________ electrolytes are solutes that exist in a solution completely or nearly completely as ions. Ex: ___________ electrolytes are those that exist in a solution as molecules. Ex: ___________ electrolytes are those that exist in a solution mostly as molecules. Ex: _____________________ (Ionic/covalent) compounds are strong electrolytes. List the Seven Strong Acids – all strong electrolytes List at least four Weak Acids– weak electrolyte List at least four Strong Bases - strong electrolyte List at least four Weak Bases – weak electrolyte


4.2 Precipitation Reactions The following is the prompt for the most important solubility rules. Write out what the ions are, and make sense of the acronym Sodium Chlorate Ammonium Acetate Potassium Nitrate If an ionic compound contains one or more ions from SAP CAN, then ___________________________ Write the molecular, complete ionic equation and the net ionic equation for the reaction between aqueous solutions of magnesium nitrate and sodium hydroxide. Molecular: Complete ionic: Net ionic: What are the spectator ions? 4.4 Oxidation-Reduction Reactions What is the oxidation number of sulfur in sodium thiosulfate, Na2S2O3? What are the oxidation numbers, and what is the process (oxidation or reduction) when Cl2 → 2Cl- 4.5 Concentrations of Solutions How many grams of Nickel (I) Sulfate are needed to make 82 mL of a 0.35M solution? How many milliliters of 2.0 M of CuSO4 are needed to make 75 mL of 0.10 M CuSO4? 35 mL of a 0.25M solution contains how many grams of Lead (II) Nitrate? 4.6 Solution Stoichiometry and Chemical Analysis What is the purpose of a titration? What is the difference between an endpoint and an equivalence point? How many milliliters of 0.113 M HCl are needed to neutralize 75 mL of 0.210 M Ba(OH)2 solution? If 56.7 mL of 0.107 M HCl is needed to neutralize a solution of Ca(OH)2, how many grams of Ca(OH)2 are in this solution?


Chapter 10 Gases Pre‐assignment (due on day Ch 10 started) Name _____________________ I. Definitions: Review, check, and write down for those you got wrong.

Term Definition Example/Equation Vapor * Homogeneous Gas mixture * Pressure * Force * Barometer Manometer Newton * Pascal * STP Boyle’s Law * Charles’s Law * Avogadro’s Law * Ideal Gas Equation * Dalton’s Law of Partial Pressures * Mole fraction * Equations to know: 10.2: Pressure=Force/Area 10.2: 1 atm=760mm Hg=760 torr=1.01325 x 105 Pa=101.325 kPa 22.4L=1 mol at STP. Conditions of STP = ___________________________? 10.5: With M as the molar mass, derive the formula for the density of a gas in terms of M, P, R, and T. Density =

II. Calculations and Questions: Use your rules for significant digits!!

1) 10.2: 265.23 kPa= torr= ___________ atm=______________mm Hg 2) 10.3: As pressure increases, volume (increases/decreases) 3) 10.3: As temperature decreases, volume (increases/decreases) 4) 10.3: As the number of moles of a gas increase, volume (increases/decreases) 5) 10.4: Given the following sets of values, calculate the unknown quantity: a) P = 1.01 atm, n = 0.00831 mol, T = 25°C V = ?


b) V= 0.602 L n = 0.00801 mol T = 311 K P = ? 6) 10.4: At what temperature would 2.10 moles of N2 gas have a pressure of 1.25 atm in a 25.0 L tank? 7) Propane is burned in excess oxygen.

(a) Write the balanced equation for this reaction

(b) If a container has a volume of 57 L and is to be filled with carbon dioxide at a pressure of 2.03 atm at a temperature of 32°C, how many moles of CO2 are to be generated? How many grams of propane must be burned?

8) 10.6: What is the total pressure exerted by a mixture of 6.00 g of oxygen and 2.00 g

nitrogen at 273K in an 8.0‐L vessel? 9) 10.6: What is the mole fraction of each component in a solution in which 3.57 g of oxygen is

dissolved in 25.0 g of carbon dioxide? III. Collecting Gases over water

• Ptotal = Pgas + PH2O • fill a test tube with water and invert it under the water so that the entire tube stays filled

with water when it is held upright. The gas you are collecting is bubbled into the test tube and the gas displaces the water.

• For video example, see http://www.kentchemistry.com/moviesfiles/Units/GasLaws/gasoverwater.htm

10) Write the equation for magnesium reacting with HCl in a single replacement reaction.

If 35.0 mL of H2 gas are collected over water when the temperature is 293K, the external pressure is 758 Torr and the vapor pressure of water at that temperature is 20 Torr, how many moles of H2 were generated? How many grams of Mg were used in the reaction?

http://www.kentchemistry.com/moviesfiles/Units/GasLaws/gasoverwater.htm


Pre‐assignment for Chapter 5 – Thermodynamics. Specific Heat and Heat of Combustion Name ________________________ Hint: If you write down all your units as you do your calculations, you will save yourself many errors!

Specific Heat – how much heat does it take to warm up a substance?

For this you will use one formula over and over, turned inside out and upside down: q = SmΔT. q = heat added = (specific heat of the substance being heated)(mass of substance heated)(change in temperature). The units for S are Joules/gram oC, for change in temperature are oC, and for mass is grams. Multiply those all together and you see that the unit for heat is Joules. One calorie is equal to 4.18 Joules. Just to make things confusing, the calories you see on nutritional labels are actually kcal, so one nutritional calorie is equal to about 4.18 kJ.

1. Calculate the amount of heat needed to increase the temperature of 75 g of water from 23oC to 95oC. The specific heat of water is 4.18 J/goC

2. Calculate the heat needed to increase the temperature of 35 g of Copper from 23oC to 95oC. The specific heat of copper is 0.387 J/goC.

Sometimes you need to calculate the heat, other times you know the heat added and want to calculate the specific heat, the mass of the substance, or the temperature change. Rearrange q = SmΔT three times, making three new equations where you have solved for each of S, m, and ΔT.

3. 480 J of heat is added to 15g of water. By how many degrees will the temperature of the water increase?

4. 300 g of Aluminum requires 25kJ to raise its temperature from 25oC to 118oC. What is the specific heat of aluminum?

5. 570kJ of heat is available to heat 2kg of water. If the water is starting out at 25oC, how high can you heat it?


Heat of Combustion – how much heat do you get from burning a fuel? Write the equation for the burning of propane, including heat as either a reactant or product as appropriate:

6. If you burn 5.00 g of propane and all the thermal energy goes into warming 1000g of water, the temperature of the water will rise by 60oC. How much thermal energy is the water absorbing?

7. The thermal energy to warm the water in problem 6 came from burning the 5.00 g of propane. How much heat would burning one gram of propane generate?

8. How much heat would burning one mole of propane generate?

9. A mole of heptane will generate 4817 kJ of heat. How much heat is generated by burning a gram of heptane?

Putting it all together….

10. The molar heat of combustion for butane is 2859 kJ/mol. Your camping stove contains 100g of butane and you want to make hot chocolate for as many people as possible. If all the heat from your stove goes into heating the water, how much water can you heat from 25oC to 100oC?

And a simplified practice AP question: H+ (aq) + OH‐(aq) → H2O(l) A student is asked to determine ∆H for the net ionic reaction represented above. The student combines equal volumes of 1.0 M HCl and 1.0 M NaOH in an open polystyrene cup calorimeter. The heat released by the reaction is determined by using the equation q=mC∆T.

Assume the Following: ‐ Both solutions are the same temperature before they are combined ‐ The densities of all solutions are the same as that of water ‐ Any heat lost to the calorimeter or to the air is negligible

a) Give appropriate units for each of the terms for the equation q=mC∆T. (example: in PV=nRT, pressure is

measured in atm)

b) List the measurements that must be made during the experiment in order to obtain the value of q.

c) Heat is released by the reaction. Is the ∆H of this reaction positive or negative? Justify your answer.

d) Suppose that a significant amount of heat were lost to the air during the experiment. What effect would this have on the calculated value of ∆H? Justify your answer.

Chapter 6 Pre‐assignment Name___________________________________

6.1: Waves: Define the following terms: Amplitude Wavelength Node Antinode Frequency Label :Amplitude, Wavelength, nodes and antinodes. Write the equation used to express frequency in terms of wavelength and speed: Wavelength/spectrum relationships:


A. B.

If you are told that one of the above waves (A or B) is blue and one is red, which is which? If one of the waves is an X ray and the other is a microwave, which one is which? Name a wavelength that would fit gamma rays? Units of measure for length:

Unit Symbol Length (m) Type Angstrom Å 10-10 X ray Nanometer nm 10-9 Ultraviolet/visible Micrometer µm 10-6 Infrared Millimeter mm 10-3 Infrared Centimeter cm 10-2 Microwave Meter m 1 TV, radio Use dimensional analysis to convert the following: 1 meter angstroms 47 nanometer millimeters 59.6 centimeters micrometers 79 Nanometers meters 0.00537 centimeters Nanometers 0.00000047981 meters Angstroms 6.2 : Origins of Modern Quantum Theory: For the Bohr description Provide the modern replacement orbits quantum number electron as particle with definite location


Formulas to know about waves:

c = speed of light = 3 x 108ms‐1

ν = frequency of wave m = mass

Practice problem for wave equations: Find the energy of a photon with a wavelength of 500 x 10‐9 m 6.3 Energy States of the Hydrogen Atom: Is energy given off or taken in as an electron moves from n=5 to n=2? Using the equation on p. 208, what is the energy difference between an electron in an n=5 state and an n=2 state? What is the wavelength of light involved in an electron moving from n=5 to n=2?

Using the figure to the right or the equation on p. 208, predict which transition produces the longest wavelength spectral line: n=3 to n=2, n=4 to n=3, or n=5 to n=4.


6.5 Modern Quantum Theory: Quantum Numbers: Each electron has a set of four numbers, called quantum numbers, that specify it completely; no two electrons in the same atom can have the same four. The four quantum numbers are:

• n describes the electron shell, or energy level. o The value of ranges from 1 to "n", where "n" is the shell containing the

outermost electron of that atom • describes the angular momentum of the electron, so it determines the shape of the

orbital. can be any integer from 0 to n‐1. = 0 means it’s an s orbital, = 1 describes a p orbital, etc.

• ml is the magnetic quantum number, and designates the orientation of the orbital (e.g., whether it’s a px, py, or pz orbital oriented on the x, y, or z axis, respectively)

• ms describes the spin of the electron and it can be either ‐1/2 or +1/2

name symbol orbital meaning range of values value example principal quantum number shell

azimuthal quantum number (angular momentum)

subshell (s orbital is listed as 0, p orbital as 1 etc.)

for :

magnetic quantum number, (projection of angular momentum)

(orientation of the subshell's shape)

for :

spin projection quantum number

spin of the electron (-1/2 = counter-clockwise, 1/2 = clockwise)

for an electron, either:

6.6 The orbitals themselves: Orbitals are simply are mathematical predictions of where the electron density is located. Each Orbital can hold two electrons and there are 4 different types of orbitals we will use: s, p, d, and f. There is at most one s orbital per energy level, three p orbitals per energy level, and five d orbitals per energy level. The first energy level to contain s, p, and d orbitals is the one where n= _________? There are _______ f orbitals per energy level, starting with the one where n = ________?


http://en.wikipedia.org/wiki/Angular_momentum




s and p Orbitals


tals d Orbi f orbitals

6.8 Specifying the electron configuration for multi‐electron atoms: Rule 1 – Aufbau Principle: Lowest energy orbitals fill first. Thus, the filling pattern is 1s, 2s, 2p, 3s, 3p, 4s, 3d, etc. Since the orbitals within a subshell are degenerate (of equal energy), the entire subshell of a particular orbital type is filled before moving to the next subshell of higher energy. Rule 2 - Pauli Exclusion Principle - Only two electrons are permitted per orbital and they must be of opposite spin. If one electron within an orbital possesses a clockwise spin, then the second electron within that orbital will possess a counterclockwise spin. Two electrons with opposite spins found in the same orbital are referred to as being paired. Rule 3- Hund's Rule - The most stable arrangement of electrons in a subshell occurs when the maximum number of unpaired electrons exist, all possessing the same spin direction. This occurs due to the degeneracy of the orbitals, all orbitals within a subshell are of equal energy. Electrons are repulsive to one another and only pair after all of the orbitals have been singly filled.

Fill in the energy orbital diagram (the drawing to the right) for oxygen, following the above rules. Now write the electron configuration for oxygen: Ex. Electron Configuration of Br: 1s2 2s2p6 3s2p6d10 4s23p5

You can also abbreviate like this: [Ar] 4s23d104p5

Having trouble remembering the ordering of the orbitals? Make sense of the periodic table, above, or use the mnemonic device to theright.

Write the electron configuration in both full and abbreviated format for: Nitrogen Phosphorous Germanium Neon Sodium


Iodine

Welcome to AP Chemistry!€¦ · AP Chemistry Summer Assignment 2011‐2012 Page 1 ... so that...

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