AP Chemistry Summer Assignment 2018

This summer assignment is set up in two parts. Use your notes from Honors Chemistry, any

books, or any other resources needed to complete both parts of the assignment.

The first part is a review of the first three chapters (all basic concepts from first year

chemistry) in AP chemistry for you to study and practice. At the beginning of the year the

chapters 1-3 test will be scheduled. Any questions on the assignment will be dealt with in the

first few days of the year.

The second part of the assignment is a more comprehensive review of topics covered

throughout first year chemistry. Many of those topics will be encountered in the first half of AP

chemistry but covered in greater depth and connectivity.

All work and problems will be checked on the first Wednesday of the year in class. Any

questions will also be addressed in those first few days. Solutions to the problems will be

posted on iTunesU at the end of the first week of classes.

After the introduction information for the beginning of the year and review questions

answered the class will begin the next section in preparation of chapter #4 content.

If you have any questions or problems my email is

james.poindexter@bellevernonarea.net and I will check it periodically throughout the summer.

Part One:

Vocabulary, theory and problems to be reviewed (listed by chapter) be able to define,

explain and apply the theory in questions and do the problems for practice for a check on the

first day of school in August:

From AP Chapter #1

Atoms and molecules

Chemistry

Scientific Method

Law of conservation of mass

Law of conservation of energy

Matter

Solid, liquid, gas, and plasma

Crystalline and amorphous

Pure substance (element and compound)

Mixture (heterogeneous and homogeneous)

Physical changes and physical properties

Chemical changes (r’xns) and chemical properties

Energy (work)

Potential energy and kinetic energy

SI system of measurements

Know the seven fundamental Si units and the quantities they measure

Fundamental and derived units

Intensive physical property and extensive physical property

Density

Sig figs

Accuracy and precision

Problems:

Unit conversions (SI prefix and English system)

Celsius to Kelvin conversions

Density

Specific heat

From AP Chapter #2

Law of conservation of mass

Law of definite proportions

Law of multiple proportions

Dalton’s atomic theory

Cathode ray tube experiment (J.J. Thomson)

Electron

Rutherford’s gold foil experiment

Nucleus (protons and neutrons)

Atomic mass unit

Atomic number

Mass number

Isotopes

Ions (and ion formation)

Periodic law

Metals, nonmetals, and metalloids

Periods and groups (families)

Main-group elements, transition elements, noble gases, alkali metals, alkaline-earth metals and

halogens

Atomic mass (average atomic mass)

Mass spectrometry (how does it work?)

Mole (Avogadro’s number)

Problems:

Hyphen notations

Symbol notations

Calculating average atomic mass

Converting between mass-moles-and atoms

From AP Chapter #3

Ionic bond

Cation and Anion

Covalent bond

Polar-covalent bond and nonpolar-covalent bond

Chemical formula

Empirical formula and molecular formula

Ways of representing molecules

Elements can be individual atoms, diatomic elements, and some are groups of atoms bonded

Molecular compounds (molecules)

Ionic compounds (formula units)

Formula mass (molar mass)

Chemical r’xn

Reactants and products

Problems:

Nomenclature of ionic and covalent compounds

Nomenclature of acids and hydrates

*Know what metals need Roman numerals for charges (the only s and p block metals that get

them are Sn and Pb; the only d and f block metals that don’t get them are Sc, Zn, Cd, and Ag)

*Know how to name polyatomic ions

Calculating molar mass and using it in conversions

Finding percent composition and finding formulas from composition data

Formulas and composition as conversion factors

Combustion analysis

Writing and balancing chemical equations

Practice problems from chapters 1-3:

1) Perform the following conversions:

0.0615 cm to m

2250 cg to kg

0.0019 km to m

340000 g to mg

24.3 ft to m

78°C to K

814 miles to km

2) If a 327 g sample of an alloy (specific heat = 0.729 J/gC) absorbs 1050 J of energy and ended

at 41°C, what was the initial temperature?

3) What volume of a solution will have a mass of 1130 g if its density was found to be 1.24

g/mL?

4) What mass of a metal, with a specific heat of 0.428 J/gC°, will release 547 J of energy and

cool from 36°C to 29°C?

5) What is the density of a cork sample if a 48.5 g sample occupied 60.0 mL?

6) Find the molar mass for each of the following compounds:

a) Al2(SO4)3 b) Mg3N2 c) Ni(ClO3)3

7) Perform the following conversions :

a) 4.26 X1022 molecules Mg3N2 to grams

b) 143 g Ni(ClO3)3 to moles

c) 6.27 mol Al2(SO4)3 to grams

8) Find the empirical formula for the compound that is 35.9% Al and 64.1% S by mass.

9) Find the empirical and molecular formulas for a compound that gave the following data:

62.0% C, 10.4% H, and 27.6% O by mass and its formula mass was found to be 290.3 g/mol.

10) How many grams of an ore, which is 64% Cu3(PO4)2, are needed to produce 1.80kg copper

at 82.6% recovery?

11) An ore sample that was 67.3% Ag4SiO3 gave 1.68kg of silver, with a 79.6% recovery. What

mass of the ore was used?

12) The atomic mass for manganese is 54.938 amu. Its two naturally occurring isotopes are

manganese-54, with a mass of 53.876 amu, and manganese-55, with a mass of 55.144 amu.

What is the percent abundance for each isotope?

13) Find the empirical formula for the compound that is 34.0% Mn and 66.0% Cl by mass.

14) Find the empirical and molecular formulas for a compound that gave the following data:

62.0% C, 10.4% H, and 27.6% O by mass and its formula mass was found to be 290.3 g/mol.

15) A 12.00g sample of a mixture of KClO3 and KCl is heated to remove all oxygen. The product,

all KCl, has a mass of 9.00g. What is the percentage KClO3 in the original mixture?

16) When a 3.00g sample of a compound containing only carbon, hydrogen, and oxygen was completely burned, 1.17g of water and 2.87g carbon dioxide were formed. What is the empirical formula for the compound? 17) A 1.5g sample of a compound contains only carbon, hydrogen, and oxygen was burned to produce 1.738g carbon dioxide and 0.711g water. What is its empirical formula?

18) A 5.82g silver coin is dissolved in nitric acid and the silver is precipitated as 7.2g AgCl. Find

the percent Ag in the coin.

19) Find the percent composition of each element in C2H8AsB.

20) How many grams of mercury can be obtained from 150g of an ore that is 85.0% HgS?

Part 2

This is the general review of first year chemistry. (There are some explanations and hints in this

section that may also be useful for part 1 so don’t forget to check there for some helpful hints.)

Resources for further assistance: 1) NMSI- http://bit.ly/1fRGID3 2) Our Videos- http://youtube.com/flippinsciencevideos 3) Khan Academy Videos- http://bit.ly/1lF0uED 4) Bozeman Science Videos- http://www.bozemanscience.com/ap-chemistry/ 5) www.sciencegeek.net -you can use the many review activities in the AP chemistry section to review 6) http://www.chemistrycoach.com/home.htm 7) http://www.collegeboard.com/ap/students/chemistry/index.html 8) http://www.wwnorton.com/college/chemistry/chemistry3/chemtours.aspx The following worksheets are to be completed for the comprehensive review:

AP Chem Summer Assignment Worksheet #1

Atomic Structure 1. a) For the ion 39K+ , state how many electrons, how many protons, and how many 19

neutrons are present?__________________________________________________________ b) Which of these particles has the smallest mass? _____________________________________________________________ 2. An atom has a net charge of -1. It has 18 electrons and 20 neutrons. Give a) its isotopic symbol b) its atomic number c) the charge on its nucleus d) the number of protons. a)_____________________________________________ b)______________________________________________ c)_____________________________________________ d)______________________________________________ 3. What is the number of electrons in Rb+1? ______________________________

4. Determine the number of protons, electrons, and neutrons in 80Br-1

and 79Se-2 ? _________________________________________________________________________________________________________

5. Which of the following statements is wrong for structure of an atom? A) Protons and neutrons are in the center. B) Electrons are moving around the nucleus. C) Electrons are negatively charged particle. D) Neutrons are positively charged particles. E) Mass of one proton is equal to mass of one neutron. 6. Which of the following statements is (are) true for structure of an atom? I. Volume of a nucleus is smaller than volume of its atom. II. The atomic mass number is the sum of proton and neutron numbers. III. The atomic number is the sum of protons and electrons. A) I B) II C) III D) I , II E) I, II , III 7. 1) Proton number 2) Neutron number 3) Chemical properties 4) Physical properties Which of the above is (are) different for isotopes? A) II B) III C) I , IV D) II , III E) II , IV 8. If X+ 2 has 28 electrons and 35 neutrons, what is the atomic mass number of X? A) 68 B) 67 C) 65 D) 63 E) 60

9. If atomic mass number of 24X is 51, what is the number of neutrons of X? A) 27 B) 24 C) 51 D) 75 E) 40 10. The element F has 10 neutrons and a mass number = 19. How many electrons are in its outermost shell? A) 10 B) 3 C) 9 D) 7 E) cannot tell from the information given 11. The most common form of iron has 26 protons and 30 neutrons in its nucleus. State its atomic number, atomic mass, and number of electrons if it's electrically neutral. Atomic number: _______ Atomic mass: ________ # of electrons: __________ 12. Consider the following three atoms: Atom 1 has 7 protons and 8 neutrons; atom 2 has 8 protons and 7 neutrons; atom 3 has 8 protons and 8 neutrons. Which two are

isotopes of the same element? _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 13. Consider fluorine atoms with 9 protons and 10 neutrons. What are the atomic number and atomic mass of this fluorine? Suppose we could add a proton to this fluorine nucleus. Would the result still be fluorine? Explain. What if we added a neutron to the fluorine nucleus? Atomic number: ____________ Atomic mass: _____________ 14. How many neutrons are in the nucleus of an atom of tungsten-184 which has an atomic number of 74? # of neutrons:_________________ 15. Which of the following combinations of particles represents an ion of net charge -1 and of mass number 82? (A) 46 neutrons, 35 protons, 36 electrons (B) 46 neutrons, 36 protons, 35 electrons (C) 46 neutrons, 36 protons, 36 electrons (D) 47 neutrons, 35 protons, 35 electrons (E) 47 neutrons, 35 protons, 36 electrons 16. One species of element M has an atomic number of 10 and a mass number of 20; one species of element N has an atomic number of 11 and a mass number of 20. Which of the following statements about these two species is true? (A) They are isotopes. (B) They are isomers. (C) They are isoelectronic (D) They contain the same number of neutrons in their atoms. (E) They contain the same total number of protons plus neutrons in their atoms. 17. A neutral atom has an atomic number of 30 and a mass number of 62, the atom must contain: (A) 92 neutrons (B) 62 electrons (C) 29 neutrons (D) 30 electrons

18. Atom X has 12 protons, 12 electrons, and 13 neutrons. Atom Y has 10 protons, 10 electrons, and 15 neutrons. It can therefore be concluded that: (A) atoms X and Y are isotopes. (B) atom X is more massive than atom Y. (C) atoms X and Y have the same mass number. (D) atoms X and Y have the same atomic number. 19. A neutral atom which has 42 electrons and a mass number of 93 has (A) an atomic number of 51. (B) a nucleus containing 51 neutrons. (C) a nucleus containing 40 neutrons. (D) a nucleus containing 51 protons. 20. A sodium ion, Na + , contains the same number of electrons as (A) a sodium atom, Na. (B) a magnesium atom, Mg. (C) a potassium ion, K+ . (D) a neon atom, Ne.

AP Chem Summer Assignment Worksheet #2

Electronic Configuration Write the electron configurations using arrows of the following elements: 1) Magnesium _________________________________________________ 2) Cobalt __________________________________________________ 3) Krypton ___________________________________________________ 4) Beryllium ___________________________________________________ 5) Scandium ___________________________________________________

Write the electron configurations of the following elements: 6) Nickel ______________________________________________________ 7) Cadmium______________________________________________________ 8) Selenium ______________________________________________________ 9) Strontium______________________________________________________ 10) Lithium ______________________________________________________ Determine what elements are denoted by the following electron configurations: 11) 1s22s22p63s23p5 ____________________ 12) 1s22s22p63s23p64s23d104p65s2 _____________________ 13) [Kr] 5s24d105p4 ___________________ 14) [Xe] 6s24f145d7 ___________________ 15) [Rn] 7s25f12 ___________________ Determine whether the following electron configurations are or are not valid: 16) 1s22s22p63s23p64s24d104p6 __________________ 17) 1s22s22p63s33d6 _________________ 18) [Rn] 7s25f9 __________________ 19) [Xe] __________________ 20) [Ne] 3p5 3s2 __________________

AP Chem Summer Assignment Worksheet #3

Periodic Table 1. A gaseous ion X2 + has two unpaired electrons in its most stable state. If X is a representative element, to which groups could the element X belong?

a) IA and VIA b) IVA c) VA and VIIA d) IA and IIIA e) VIA and VIIA 2. In the fourth period of the Periodic Table, how many elements have one or more 4p electrons? a) 2 b) 6 c) 10 d) 18 d) 32 3. A horizontal row in the periodic table containing a sequence of elements is a: a) Group b) family c) subgroup d) period 4. The modern periodic table has the elements arranged in order of: a) their date of discovery. b) increasing radii of the atoms. c) increasing number of neutrons. d) increasing number of protons. e) number of isotopes. 5. Which of the following electron configurations is correct for chromium, (atomic number 24)? a) [Ar]4s24p4

b) [Ar]3d54s1

c) [Ar]3d45s2

d) [Kr]4d55s1

e) [Ar]4s14p5

6. The electron configuration of Pt is [Xe]4f145d96s1. How many unpaired electrons are in this atom? a) 1 b) 2 c) 7 d) 8 e) 9 7. Which of the following statements is false? a) The s orbital is spherical. b) There are 10 d orbitals in a d subshell. c) The third energy level has no f orbitals. d) The fifth energy level has a set of f orbitals. 8. For a group in the periodic table, the elements have in common: a) the same number of electrons. b) the same number of electrons in all subshells. c) the same number of filled subshells. d) the same number of electrons in the highest energy subshell containing electrons.

e) the same number of orbitals. 9. Which of the following electron configurations belong to an element that is the most chemically reactive? a) 1s2 b) 1s2 2s2 2p6 c) 1s22s22p5 (d) 1s22s22p63s23p6

10. How many valence electrons does the element with the electron configuration 1s22s22p63s23p64s1 have? a. 0 b) 1 c) 9 d) 19 11. How is the electron configuration similar for each element in a group? ___________________ 12. How is the electron configuration similar for each element in a period? ___________________ 13. Refer to a periodic table. In which period is calcium? a. Period 2 b. Period 4 c. Period 6 d. Period 8 14. Refer to a periodic table. In which group is calcium? a. Group 1 b. Group 2 c. Group 17 d. Group 18 15. An element that has the electron configuration [Ne]3s23p5 is in which period? a. Period 2 b. Period 3 c. Period 5 d. Period 7 16. An element that has the electron configuration [Ne]3s23p5 is in which group? a. Group 2 b. Group 5 c. Group 7 d. Group 17

AP Chem Summer Assignment Worksheet #4

Mole Concept Complete the following problems showing all work and with answers using the correct significant digits.

1) How many grams does 0.500 moles of NaCl weigh? ____________________ 2) How many molecules are there in 0.655 moles of C6H14? ____________________ 3) How many moles are there in 1.204 x 1025

molecules of water? ____________________ 4) How many grams does2.408 x 1023

molecules of SiO2 weigh? ____________________ 5) How many molecules are there in 21.6 grams of CH4? ____________________

6) Calculate the molecular mass for the following: A. KOH ______________ B. N2O2 ___________________ C. Cu2SO3 ______________ D. Sr3 (PO4)2_________________ 7) Convert each of the following from grams to moles. A. 15.O g C2H6 _________________ C. 80 g NaOH__________________ B. 36 g H2O ____________________ D. 50 g CaCO3___________________________

8) Convert moles to grams in each of the following: A. 2 mole NH3_____________________________

B. 0.5 mole H2SO4_____________________________

9) Convert the following to moles: A. 3.01 X 1023

atoms Na ____________________ B. 6.02 X 1024

molecules CO2 _____________________ 10) Calculate the molecular mass of the following compounds: a. acetone, CH3COCH3 _________________________ b. ethyl ether, C4H10 ____________________________ 11) How many moles of hydrogen atoms are there in 1 mole of following molecules? a. C3H8____________________ b. C2H5OH_____________________________ 12) How many moles of atoms are there in 2 moles of following molecules? a. C10H22 ___________________ b. CH3COOH__________________________

AP Chem Summer Assignment Worksheet #5

Chemical Formula /Percent Composition Complete the following problems showing all work and with answers using the correct significant digits.

1. A 0.941 gram piece of magnesium metal is heated and reacts with oxygen. The resulting oxide weighed 1.560 grams. Determine the percent composition of each element in the compound. ___________________ 2. Determine the empirical formula given the following data for each compound: (a) Fe = 63.53%, S = 36.47% (b) Fe = 46.55%, S = 53.45% (a)___________________(b)______________ 3. A compound contains 21.6% sodium, 33.0% chlorine, 45.1% oxygen. Determine the empirical formula of the compound. 4. A 2.500 gram sample of uranium was heated in air. The resulting oxide weighed 2.949 gram. Determine the empirical formula of the oxide. {Hint: Carry out the calculations to four decimal places}. ______________

5. When 1.010 g of zinc vapor is burned in air, 1.257 grams of the oxide is produced. a) What elements are present in the oxide? b) Determine the percent composition of each element in the oxide. c) Determine the empirical formula of the compound. (a)_________________ (b)________________ (c)_______________ 6. A compound has the empirical formula of CH3Br and a vapor density of 6.00 g/L, at 375 K and 0.983 atm. Using these data, determine the following: a) The molar mass of the compound. b) The molecular formula of the compound. (a)________________ (b)_______________ 7. A compound containing the elements C, H, N, and O is analyzed. When a 1.2359 gram is burned in excess oxygen, 2.241 g of CO2 (g) is formed. The combustion analysis showed that the sample contained 0.0648 g of H. a) Determine the mass, in grams, of C in the 1.2359 g sample of the compound. b) When the compound is analyzed for N content only, the mass percent of N is found to be 28.84%. Determine the mass, in grams, of N in the original 1.2359 g sample of compound. c) Determine the mass, in grams, of O in the original 1.2359 g sample of the compound. d) Determine the empirical formula of the compound. (a)_________________ (b)__________________ (c)_________________ (d)______________

AP Chem Summer Assignment Worksheet #6

Stoichiometry Balance all of the equations that need to be balanced. The answers are given in parenthesis at the end of each problem.

1. In the decomposition of sodium hydroxide, how many moles of sodium hydroxide are needed to produce 30.0 moles of water? (60.0 moles NaOH) __________________ 2. In the single replacement reaction of lithium and magnesium nitrate, what mass of lithium combines with 75.0 grams of magnesium nitrate? (7.02 g Li) _________________ 3. How many grams of lead (II) nitrate are needed to produce 60.0 grams of potassium nitrate in the double replacement reaction of potassium iodide and lead (II) nitrate. (98.3 g lead (II) nitrate) _________________ 4. In the synthesis reaction of zinc and sulfur, what mass of zinc sulfide is produced from 100.0 grams of sulfur? (303.9 g ZnS) ________________ 5. A synthesis reaction of calcium and oxygen was completed in a lab and

234.9 grams of calcium oxide were produced from 75.00 grams of oxygen. What is the percent yield? (89.36%) _______________ 6. In the single replacement reaction of magnesium and aluminum phosphate, if 7.00 moles of magnesium react, how many moles of aluminum phosphate would be needed? (4.67 mol AlPO4) _______________ 7. When methane and oxygen react (complete combustion reaction) how many grams of water would be produced from 25.0 grams of methane? (56.1g water) _______________ 8. A 26.3 gram sample of potassium chlorate decomposed and produced 9.45 grams of oxygen. What is the percent yield for oxygen? (91.7%) ______________ 9. If 7.40 grams of calcium hydroxide react with nitric acid to produce 2.01 grams of water, what is the percent yield? (55.8%) ______________ 10. Lime, CaO, reacts with hydrochloric acid to form calcium chloride and water. How many moles of HCl would be required to react with 7.5 moles of lime? How many moles of water would be formed? (15 mol HCl; 7.5 mol water) ______________ For each of the following write balanced chemical equations and then solve the problem.

11. What is the maximum number of grams of PH3 that can be formed when 6.2 g of phosphorus reacts with 6.0 g of hydrogen to form PH3? ______________ 12. Copper is formed when aluminum reacts with cupric sulfate in a single-replacement reaction. How many grams of copper can be obtained when 29.0 g of Al reacts with 156 g or copper (II) sulfate? ______________ 13. If you begin with 1250 g of N2 and 225 of H2 in the reaction that forms ammonia gas (NH3), how much ammonia will be formed? What is the limiting reagent? How much of the reagent is left when the maximum amount of ammonia is formed? ______________

Common Ions and Their Charges Tips for Learning the Ions “From the Table” These ions can be organized into two groups. 1. Their place on the table suggests the charge on the ion, since the neutral atom gains or loses a predictable number of electrons in order to obtain a noble gas configuration. This was a focus in first year chemistry, so if you are unsure what this means, get help BEFORE the start of the year. a. All Group 1 Elements (alkali metals) lose one electron to form an ion with a 1+ charge b. All Group 2 Elements (alkaline earth metals) lose two electrons to form an ion with a 2+ charge c. Group 13 metals like aluminum lose three electrons to form an ion with a 3+ charge d. All Group 17 Elements (halogens) gain one electron to form an ion with a 1- charge e. All Group 16 nonmetals gain two electrons to form an ion with a 2- charge f. All Group 15 nonmetals gain three electrons to form an ion with a 3- charge Notice that cations keep their name (sodium ion, calcium ion) while anions get an “-ide” ending (chloride ion, oxide ion). 2. Metals that can form more than one ion will have their positive charge denoted by a roman numeral in parenthesis immediately next to the name of the metal Polyatomic Anions Most of the work on memorization occurs with these ions, but there are a number of patterns that can greatly reduce the amount of memorizing that one must do. 1. “ate” anions have one more oxygen then the “ite” ion, but the same charge. If you memorize the “ate” ions, then you should be able to derive the formula for the “ite” ion and vice-versa. a. sulfate is SO4

-2, so sulfite has the same charge but one less oxygen (SO3-2)

b. nitrate is NO3-1

so nitrite has the same charge but one less oxygen (NO2-1)

2. If you know that a sufate ion is SO4-2

- then to get the formula for hydrogen sulfate ion, you add a hydrogen ion to the front of the formula. Since a hydrogen ion has a 1+ charge, the net charge on the new ion is less negative by one.

a. Example: phosphate (PO4

-3) hydrogen phosphate (HPO4-2) dihydrogen phosphate (H2PO4

-1) 3. Learn the hypochlorite � chlorite � chlorate � perchlorate series, and you also know the series containing iodite/iodate as well as bromite/bromate. a. The relationship between the “ite” and “ate” ion is predictable, as always. Learn one and you know the other. b. The prefix “hypo” means “under” or “too little” (think “hypodermic”, “hypothermic” or “hypoglycemia”) i. Hypochlorite is “under” chlorite, meaning it has one less oxygen than chlorite c. The prefix “hyper” means “above” or “too much” (think “hyperkinetic”) i. the prefix “per” is derived from “hyper” so perchlorate (hyperchlorate) has one more oxygen than chlorate. The following page has a list of polyatomic ions to know for the beginning of the year. If you know these it will be easier to add any others as we go throughout the year. (Remember that if you know the –ate ion you can figure out the –ite, hypo- -ite, and per- -ate ions)

Sulfate Phosphate Nitrate Ammonium Thiocyanate Carbonate Borate Chromate Dichromate Arsenate Iodate

Permanganate Oxalate Amide Hydroxide Cyanide Acetate Peroxide Chlorate Thiosulfate Bromate