Dear Students,

Welcome to AP Chemistry, a little early. We will have a fabulous year together in which you will be taught everything you need for success in the course and on the AP Chemistry Exam scheduled for Friday, May 07, 2021. I apologize for the summer assignment, but I want to have at least a month to review with you prior to the AP Chemistry Exam. The assignment is mostly review and NOT comprehensive of all material covered in honors chemistry. I would recommend that you start the assignment one week before school starts, as summer should be reserved for fun.

Summer Assignment:

1. Complete the Math Mania sheet without a calculator. Part 1 of the APC exam is without a calculator. We will work on non-calculator math skills throughout the course.

2. Read the Chapter 1 notes provided. These notes, as well as the two other packets, correspond with an older textbook and do not match the chapters in your book perfectly. Complete problems from Hwk 1.1 (the homework for the three chapters has been packaged as a unit) on a separate sheet of paper that has your name in the top right hand corner and Clwk 1.1.

3. Read the Chapter 2 notes provided. Complete problems from Hwk 1.2 on a separate sheet of paper that has your name in the top right hand corner, and Clwk 1.2A and B (1.2B is NOT available online). Note: Organic chemistry is no longer assessed on the APC Exam but is assessed on the SAT 2 Chemistry Subject Test so I have included some information on this concept. Just hold onto it for now.

4. Read the Chapter 3 notes provided. Complete problems from Hwk 1.3 on a separate sheet of paper that has your name in the top right hand corner. We will review this section extensively, as stoichiometry is a critical component of APC! You already know how to determine LRs and theoretical yields, but there are tricks that will make the process less labor intensive and doable without a calculator. The advanced stoichiometry was not covered at the honors or pre-APC level. We will cover it together in class so please do not fret over this concept.

5. Memorize the chemical symbols for elements on the periodic table (PT). For example, Mg represents . You can use a PT on every assessment, but the name of the element is NOT listed. The periodic table that you will use all year and on the AP Chemistry Exam is provided in this packet. Also included are formula sheets (provided on the APC Exam), polyatomic ions (not provided on the APC Exam), and other pertinent information. Please do not lose these reference sheets. You must memorize all of the polyatomic ions that have been provided.

6. Assemble all of your completed work into one packet, in the order listed above, which will be turned in the second day of school.

You will take mini-quizzes through Google Forms the first week of school and Quiz 1.1 at the beginning of the second week of school. These quizzes will assess material from the summer assignment (i.e. compound naming, formula writing, stoichiometry, empirical formula, etc.).

If you have any questions or concerns, please email me at [email protected]. Have a wonderful and safe summer. I look forward to working with all of you next year!

Sincerely,

Dr. Kalish ☺

Goggles, anyone?! Page 1 Page 2 Page 3 Page 4

Activity Series Polyatomic Ions Metal Ion Name Formula Ion Name Formula Ion Name Formula - - 2- Li Acetate CH3COO Dihydrogen phosphate H2PO4 Oxalate C2O4 + - - Rb Ammonium NH4 carbonate or HCO3 Perchlorate ClO4 K bicarbonate 3- - - Ba Arsenate AsO4 Hydrogen sulfate HSO4 Permanganate MnO4 - - 2- Sr Azide N3 Hydroxide OH Peroxide O2 - - 3- Ca Bromate BrO3 Hypochlorite ClO Phosphate PO4 2- - 3- Na Carbonate CO3 Iodate IO3 Phosphite PO3 - - 2- Mg Chlorate ClO3 Iodite IO2 Sulfate SO4 - 2+ 2- Al Chlorite ClO2 (I) Hg2 Sulfite SO3 2- + - Mn Chromate CrO4 Methylammonium CH3NH3 Thiocyanate SCN - 2- 2- Zn Cyanide CN Monohydrogen HPO4 Thiosulfate S2O3 Cr phosphate 2- - 2+ Fe Dichromate Cr2O7 Nitrate NO3 Uranyl UO2 - Cd Nitrite NO2 Co Ni Important Constants Formulas Sn Avogadro’s constant: 1 mole = 6.022 x 1023 atoms, particles, Density = mass Pb molecules Volume 8 H2 Speed of light (c): c = 2.998 x 10 m/s c =    = frequency Sb Planck’s constant: h = 6.6262 x 10-34 J-sec E = h  E = Energy Bi Universal Constant R = 0.0821 L-atm/mole-K or Cu R = 8.314 L-kPa/mole-K or R = 62.4 L-Torr/mole-K Molar Volume (STP) Vm = 22.414 L/mole Hg

Ag Subshell Tree Diagram Prefix Number Pt Mono- 1 Au 7s 7p Di- 2 Tri- 3 6s 6p 6d 5s 5p 5d 5f Tetra- 4 4s 4p 4d 4f Penta- 5 3s 3p 3d Hexa- 6 2s 2p Hepta- 7 1s Octa- 8 Start Nona- 9

Deca- 10 Electronegativities Atomic Number Element Electronegativity Atomic Number Element Electronegativity 1 H 2.2 28 Ni 1.9 3 Li 1.0 29 Cu 1.9 4 Be 1.6 30 Zn 1.6 5 B 2.0 34 Se 2.6 6 C 2.5 35 Br 3.0 7 N 3.0 37 Rb 0.8 8 O 3.4 47 Ag 1.9 9 F 4.0 48 Cd 1.7 11 Na 0.9 50 Sn 2.0 12 Mg 1.3 51 Sb 2.0 13 Al 1.6 53 I 2.7 14 Si 1.9 55 Cs 0.8 15 P 2.2 56 Ba 0.9 16 S 2.6 78 Pt 2.2 17 Cl 3.2 79 Au 2.4 19 K 0.8 80 Hg 1.9 20 Ca 1.0 82 Pb 1.8 25 Mn 1.6 83 Bi 1.9 26 Fe 1.8 84 Po 2.0 27 Co 1.9

Page 5 Page 6 Name: ______Date: ______Period #: _X_

Math Mania 1

Directions: The AP Chemistry Exam requires superior pencil and paper math skills. To better hone your skills, please complete each of the problems below WITHOUT a CALCULATOR; show your work.

1. (3.0 x 10-4) (2.0 x 10-7)

2. (8.0 x 10-5) (7.0 x 10-8)

3. (2.0 x 10-3) (7.0 x 104) (3.0 x 10-4)

4. (8.0 x 10-12) / (4.0 x 10-5)

5. (2.0 x 10-12) / (6.0 x 10-4)

6. (1.0 x 10-2) / (4.0 x 10-3)

7. 4.0 x 10-4 = x2/0.0200

8. 6.0 x 10-6 = x2/0.0200

9. 8.1 x 10-8 = x3

10. 1.25 x 10-13 = x3

11. 0.3 x (5/6) x 8

12. 0.80 x (1/2) x 90

13. 0.25 x (3/8) x 64

14. 0.40 x (1/5) x 25

15. 18.0 grams Al x 1 mol Al/27.0 grams Al x 4 mole Fe/3 mole Al x ______= mass of Fe

16. 6.539 grams Zn x 1 mol Zn/65.39 grams Zn x 2 mole HCl/1 mole Zn x ______= mass of HCl

17. 20.0 grams Ca x 1 mol Ca/40.0 grams Ca x 2 mole Al/3 mole Ca x ______= mass of Al

Page 7 Page 8 Chemistry: Matter and Measurement

I. Introduction:

Chemistry enables us to design all kinds of materials: drugs (disease); pesticides; fertilizers; fuels; fibers (clothing); building materials; plastics; etc.

A. Key Terms:

1. Chemistry: study of the composition, structure, properties, and reactions of matter a. Matter: anything that has mass and occupies space Examples: wood, sand, water, air, , etc.

1) Mass: quantity of matter in an object a) use a balance to measure mass; the unit on a balance is the gram, but the fundamental unit is the kilogram (kg) 2) Volume: space that object occupies a) use a ruler to measure a regular solid, and a graduated cylinder to measure an irregular solid (water displacement) or a liquid b) the units vary: cm3, ml, L [1 ml = 1 cm3]

2. Atoms: smallest distinctive units in a sample of matter (building blocks)

3. Molecules: larger units in which two or more atoms are joined together a. The way in which matter behaves depends on the atoms present and the manner in which they are combined

4. Composition: types of atoms and their relative proportions in a sample of matter

B. Properties:

1. Physical property: characteristic displayed by a sample of matter without it undergoing any change in its composition; what you can see or measure without altering the chemical nature of the material Examples: color, mass, density, state, Tm, Tf, Tb

2. Chemical property: characteristic displayed by a sample of matter as it undergoes a change in composition Examples: flammability, ability to react with acids

C. Types of changes:

1. Physical change: change at the macroscopic level but no change in composition; the same substance must remain after the change Examples: Phase changes, dissolving

Page 9 2. Chemical change or chemical reaction: change in composition and/or the structure of its molecules; one or more substances is altered--new substances are formed Examples: cooking and spoiling of foods; burning, digestion, fermentation a. Reactants → Products b. Evidence of a Chemical Change 1) Evolution of a gas 2) Formation of a ppt. 3) Evolution or absorption of heat (exo- vs. endothermic reactions) 4) Emission of light 5) Color change

D. Classifying Matter

Material s

Homogeneous Heterogeneous Materials Materials

Substance s Homogeneous Heterogeneous Mixtures Mixtures (Solutions) Elements Compounds

1. Material: any specific type of matter a. Homogeneous Materials: uniform matter b. Heterogeneous Materials: nonuniform matter

2. Mixture: consists of two or more different atoms or compounds with no fixed composition; the atoms or compounds are mixed together physically a. Heterogeneous Mixture: variable composition and/or properties throughout; two or more distinct phases; interface Examples: blood, granite; Oil and water b. Homogenous Mixture or Solution: has the same composition and properties throughout Examples: salt water; sugar water

3. Substance: type of matter that has a definite or fixed composition that does not vary from one sample to another Examples: Elements or compounds

Page 10 a. Element: substance that cannot be broken down into other simpler substances by chemical reactions; substances composed of one type of atom; represented by a chemical symbol

b. Compound: substance made up of atoms of two or more elements that combine in fixed proportions or definite ratios; represented by a Example: H2O with 2 H and 1 O

4. Key differences between mixtures and compounds a. The properties of a mixture reflect the properties of the substances it contains; the properties of a compound bear no resemblance to the properties of the elements that comprise the compound. b. Compounds have a definite composition by mass of their combining elements, while the components of a mixture may be present in varying proportions.

E. Scientific Methods: 1. Observation 2. Hypothesis: tentative prediction or explanation concerning some phenomenon 3. Experiment: procedure used to test a hypothesis a. Data 4. Scientific Laws: summary of patterns in a large collection of data 5. Theory: multi-tested and confirmed hypothesis

II. Scientific Measurement: A. International System of Units 1. Length: the SI base unit is the meter 2. Mass: the quantity of matter in an object a. SI base unit is the kilogram 3. Time: SI base unit is the second 4. Temperature: property that tells us in what direction heat will flow a. SI base unit is Kelvin (K)

Seven Fundamental Units in SI Quantity Unit Abbreviation for unit Symbol for quantity Length l meter m Mass m kilogram kg Time t second s Thermodynamic Temperature T kelvin K Amount of Substance n mole mol Electric Current I ampere A Luminous Intensity Iv candela cd

Page 11 B. Density 1. mass per unit volume of a substance Lowest d=m/V density

Density of water is 1.00 g/ml

Problems…

Common SI prefixes

Prefix Unit abbreviation Exponential Multiplier Meaning

Giga- G 109 1 000 000 000 Mega- M 106 1 000 000 Kilo- k 103 1 000 Greatest Hecto- h 102 100 density Deca- da 101 10 100 1 deci- d 10-1 1/10 centi- c 10-2 1/100 milli- m 10-3 1/1000 micro-  10-6 1/1 000 000 nano-  10-9 1/1 000 000 000 pico- p 10-12 1/1 000 000 000 000 femto- f 10-15 1/1 000 000 000 000 000

C. Conversions within the same quantity vs. those between different quantities 1. Same Quantity Examples: 1) 5.00 m = 500. cm 2) 3.75 mm = 0.00375 m 3) 579000 ug = 0.579000 g 4) 5.65 L = 5650 ml

2. Different Quantities: Factor-Label Method or Dimensional Analysis

Equalities and Conversion Factors 2.54 centimeters = 1 in 2 cups = 1 pint 28.3 grams = 1 ounce 12 inches = 1 foot 2 pints = 1 quart 1 kilogram = 2.2 pounds 3 feet = 1 yard 4 cups = 1 quart 453.6 g = 1 lb 5280 feet = 1 mile 1 liter = 1.06 quarts 28.35 g = 1 ounce 1 meter = 39.37 inches 16 ounces = 1 pound 3.785 L = 1 gal 1 meter = 1.09 yards 1 mole = 6.022 x 1023 atoms 29.57 ml = 1 fl oz 1 km = 0.62 miles 1 ml = 1 cm3 1 mole gas @STP = 22.4 L

Page 12 Page 13 Number Digits to Count Example Number of Significant Digits Nonzero Digits All 3279 4 11.2 3

Leading Zeroes (zeroes None 0.0045 2 before an integer) 0.000005 1

Captive Zeroes (zeroes All 5.007 4 between two integers) 6,000,008 7

Trailing Zeroes (zeroes Counted only if the 100 1 after the last integer) number contains a 100. 3 decimal point 100.0 4 0.0100 3

Scientific Notation All BUT be careful 1.7 x 10-4 2 of “incorrectly” 1.30 x 10-2 3 written scientific 0.005 x 102 1 notation

4. Rules: a. Multiplying and Dividing: Round the calculated result to the same number of significant figures as the measurement having the least number of significant figures. [carry all numbers through and then round off]

Examples: 3.45 cm * 4.5555 cm =15.716475 cm2 → 15.7 cm2 (Answer is expressed in 3 significant figures)

3500 m x 426000 m = 1 491 000 000 m2 → 1 500 000 000 m2 or 1.5 x 109 (2 sig figs)

b. Addition and Subtraction: The answer can have NO more digits to the right of the decimal point than there are in the measurement with the smallest number of digits to the right of the decimal point. Example: 3.45 cm + 100.1 cm = 103.55 cm → 103.6 cm (Express answer to the tenth place)

c. Rounding Fives: If the last significant digit before the five is odd, round up. If the last significant digit before the five is even (and there are not any numbers other than zero after the five), do NOT round up (leave it alone). Example: 3.15 → For two significant digits or to the tenth place, round to 3.2 Example: 3.45 → For two significant digits, round to 3.4 Example: 3.451 → For two significant digits, round to 3.5

Page 14 Page 15 Page 16 Homework Problems:

Hwk 1.1: Chapter 1 1) Which of the following are examples of matter? a) c) The human body e) Gasoline b) Air d) Red light f) An idea 2) Which of the following is NOT a physical property? a) Solid iron melts at a temperature of 1535 oC. c) Natural gas burns. b) Solid as a yellow color. d) Diamond is extremely hard. 3) Which of the following describe a chemical change, and which a physical change? a) Sheep are sheared and the wool is spun into c) Milk sours when left out. yarn. d) Silkworms feed on mulberry leaves and b) A cake is baked from a mixture of flour, produce silk. baking powder, sugar, eggs, shortening, and e) An overgrown lawn is mowed. milk. 4) Which of the following represent elements? Explain. a) C c) Cl e) Na b) CO d) CaCl2 f) KI 5) Which of the following are substances, and which are mixtures? Explain. a) gas used to fill a balloon c) A premium red wine b) Juice squeezed from a lemon d) Salt used to de-ice roads 6) Indicate whether the mixture is homogeneous or heterogeneous. a) Gasoline c) Italian salad dressing b) Raisin pudding d) Coke 7) Convert the following quantities: a) 546 mm to meters c) 181 pm to µm e) 46.3 m3 to L (careful) b) 87.6 mg to kg d) 1.00 h to µs f) 55 mi/h to km/min 8) How many significant figures are there in each of the following quantities? a) 4051 m c) 0.0430 g e) 1.60 x 10-9 s b) 0.0169 s d) 5.00 x 109 m f) 0.0150 oC 9) Perform the indicated operations and provide answers in the indicated unit with the correct number of significant digits. a) 13.25 cm + 26 mm – 7.8 cm + 0.186 m (in cm) b) 48.834 g + 717 mg – 0.166 g + 1.0251 kg (in kg) 10) Calculate the density of a salt solution if 50.0 ml has a mass of 57.0 g. 11) A glass container has a mass of 48.462 g. A sample of 4.00 ml of antifreeze solution is added, and the container with the antifreeze has a mass of 54.513 g. Calculate the density of the antifreeze solution expressed in the correct number of significant figures. 12) A rectangular block of gold-colored material measures 3.00 cm x 1.25 cm x 1.50 cm and has a mass of 28.12 g. Can the material be gold if the density of Au is 19.3 g/cm3? Calculate the percent error.

Hwk 1.2: Chapter 2 1) When 24.3 g of magnesium is burned in 16.0 g of , 40.3 g of magnesium oxide is formed. When 24.3 g of magnesium is burned in 80.0 g of oxygen, (a) What is the total mass of substances present after the reaction? (b) What mass of magnesium oxide is formed? (c) What law(s) is/are illustrated by this reaction? 2) What is the atomic nucleus? Which subatomic particle(s) is/are found in the nucleus? 3) Which of the following pairs of symbols represent ? Which are isobars? 70 70 7 8 a) 33 E and 34 E d) 3 E and 4 E 57 66 22 44 b) 28 E and 28 E e) 11 E and 22 E 186 186 c) 74 E and 74 E 4) What do values represent? 5) What type of information is conveyed by each of the following representations of a molecule? a) Empirical formula b) Molecular formula c) Structural formula

Page 17 6) A substance has the molecular formula C4H8O2. (a) What is the empirical formula of this substance? (b) Can you write a structural formula from an empirical formula? Explain. 7) Are hexane, C6H14, and cyclohexane, C6H12, isomers? Explain. 8) For which of the following is the molecular formula alone enough to identify the type of compound? For which must you have the structural formulas? a) An organic compound c) An alcohol e) A carboxylic acid b) A hydrocarbon d) An alkane 9) Explain the difference in meaning between each pair of terms: a) A and on the periodic table c) An acid and a salt (P.T.) d) An isomer and an b) An ion and ionic substance 10) Indicate the numbers of electrons and neutrons in the following atoms: a) B-11 c) Kr-81 b) Sm-153 d) Te-121 11) in nature consists of two isotopes, Eu-151, with a mass of 150.92 amu and a fractional abundance of 0.478, and Eu-153, with a mass of 152.92 amu and a fractional abundance of 0.522. Calculate the weighted average atomic mass of Europium. 12) The two naturally occurring isotopes of are N-14, with an atomic mass of 14.003074 amu, and N-15, with an atomic mass of 15.000108 amu. What are the percent natural abundances of these isotopes? {Hint: set one at x and the other at 1-x} 13) The two naturally occurring isotopes of are Rb-85, with an atomic mass of 84.91179 amu, and Rb-87, with an atomic mass of 86.90919 amu. What are the percent natural abundances of these isotopes? {Hint: set one at x and the other at 1-x} 14) Identify the elements represented by the following information. Indicate whether the element is a metal or nonmetal. a) Group 3A (13), period 4 d) Group 1A (1), period 2 b) Group 1B (3), period 4 e) Group 4A (14), period 2 c) Group 7A (17), period 5 f) Group 1B (3), period 4 15) Write the chemical symbol or a molecular formula for the following, whichever best represents how the element exists in the natural state. a) c) e) b) Sulfur d) 16) Which of the following are binary molecular compounds? a) iodide c) Chlorofluorocarbons e) Sodium cyanide b) Hydrogen bromide d) Ammonia 17) Write the chemical formula or name the compound: a) PF3 d) Phosphorus f) Dinitrogen pentoxide b) I2O5 pentachloride c) P4S10 e) Sulfur hexafluoride 18) Write the chemical symbol or name for the following monatomic ions: a) ion c) Sulfide ion e) Ba2+ b) (II) ion d) Fe3+ f) Se2- 19) Write the chemical formula or name for the following polyatomic ions: - 2- a) HSO4 d) CrO4 f) Dichromate ion - b) NO3 e) Hydrogen phosphate g) Perchlorate ion - c) MnO4 ion h) Thiosulfate ion 20) Name the following ionic compounds: a) Li2S f) KOH k) K2Cr2O7 b) FeCl3 g) NH4CN l) Ca(ClO2)2 c) CaS h) Cr(NO3)3 9H2O m) CuI d) Cr2O3 i) Mg(HCO3)2 n) Mg(H2PO4)2 e) BaSO3 j) Na2S2O3 5H2O o) CaC2O4 H2O

Page 18 21) Write the chemical formula for the following ionic compounds: a) sulfide f) Magnesium j) sulfate b) Barium carbonate g) Cobalt(II) nitrate heptahydrate c) Aluminum bromide h) Magnesium dihydrogen k) Sodium hydrogen hexahydrate phosphate phosphate d) Potassium sulfite i) Potassium nitrite l) Iron(III) oxide e) (I) sulfide 22) Name the following acids: a) HClO(aq) d) HF(aq) g) H2SO3(aq) b) HCl(aq) e) HNO3(aq) h) H2C2O4(aq) c) HIO4(aq) f) H2SO4(aq) 23) Write the chemical formula for the following acids. a) Hydrobromic acid e) Acetic acid b) Chlorous acid f) Phosphorous acid c) Perchloric acid g) Hypoiodous acid d) Nitrous acid h) Boric acid

Hwk 1.3: Chapter 3 1) What are the empirical formulas of the compounds with the following molecular formulas? a) H2O2 c) C10H8 b) C6H16 d) C6H16O 2) Calculate the molecular or formula mass of the following. a) C2H5NO2 d) K3[Co(NO2)6] b) Na2S2O3 e) Chlorous acid c) Fe(NO3)3 9H2O f) Ammonium hydrogen phosphate 3) Calculate the mass, in g, of the following: a) 4.61 mol AlCl3 c) 0.615 mol (III) oxide b) 0.314 mol HOCH2(CH2)4CH2OH 4) Calculate the mass percent nitrogen in the compound having the condensed structural formula, CH3CH2CH(CH3)CONH2. 5) Calculate the mass percent of in the , Be3Al2Si6O18. Calculate the maximum mass of Be obtainable from 1.00 kg of Be. 6) The empirical formula of apigenin, a yellow dye for wool, is C3H2O. The molecular mass of the compound is 270 amu. What is the molecular formula? 7) Resorcinol, used in manufacturing resins, drugs, and other products, is 65.44 %C, 5.49 %H, and 29.06 %O by mass. Its molecular mass is 110. amu. What is the molecular formula? 8) Sodium tetrathionate, an ionic compound formed when sodium thiosulfate reacts with is 17.01 % Na, 47.46 % S, and 35.52 % O by mass. The formula mass is 270 amu. What is its formula? 9) A 0.0989 g sample of an alcohol is burned in oxygen to yield 0.2160 g CO2 and 0.1194 g H2O. Calculate the mass percent composition and empirical formula of the compound. 10) Balance the following equations: a) TiCl4 + H2O → TiO2 + HCl e) Al2(SO4)3 + NaOH → Al(OH)3 + Na2SO4 b) WO3 + H2 → W + H2O f) Ca3P2 + H2O → Ca(OH)2 + PH3 c) C5H12 + O2 → CO2 + H2O g) Cl2O7 + H2O → HClO4 d) Al4C3 + H2O → Al(OH)3 + CH4 h) MnO2 + HCl → MnCl2 + Cl2 + H2O 11) Write a balanced chemical for each of the following: a) Decomposition of solid potassium chlorate upon heating to generate solid and oxygen gas b) Combustion of liquid 2-butanol c) Reaction of gaseous ammonia (NH3) and oxygen gas to generate nitrogen monoxide gas and water vapor. d) The reaction of chlorine gas, ammonia vapor, and aqueous sodium hydroxide to generate water and an aqueous solution containing sodium chloride and hydrazine (N2H4, a chemical used in the synthesis of pesticides).

You did not cover this concept in Chem I H. There is a sheet on page 47 if you want to try it, but we will cover it in full when we return to school. You may hold off on this one if you wish.

Page 19 12) Toluene and nitric acid are used in the production of trinitrotoluene (TNT), an explosive. ___C7H8 + ___HNO3 → ___C7H5N3O6 + ___H2O a) What mass of nitric acid is required to react with 454 g of C7H8? b) What mass of TNT can be generated when 829 g of C7H8 reacts with excess nitric acid? 13) Acetaldehyde, CH3CHO (D = 0.789 g/ml), a liquid used in the manufacture of perfumes, flavors, dyes, and plastics, can be produced by the reaction of ethanol with oxygen gas. ___CH3CH2OH + ___O2 → ___ CH3CHO + ___H2O a) How many liters of liquid ethanol (D = 0.789 g/ml) must be consumed to generate 25.0 L acetaldehyde? 14) trifluoride reacts with water to produce boric acid and fluoroboric acid. 4BF3 + 3H2O → H3BO3 + 3HBF4 a) If a reaction vessel contains 0.496 mol BF3 and 0.313 mol H2O, identify the limiting reactant. b) How many moles of HBF4 should be generated? 15) A student needs 625 g of zinc sulfide, a white pigment, for an art project. He can synthesize it using the reaction, Na2S(aq) + Zn(NO3)2(aq) → ZnS(s) + 2NaNO3(aq) a) What mass of zinc nitrate will he need if he can make the zinc sulfide at an 85.0 % yield? 16) Calculate the molarity of each of the following aqueous solutions: a) 2.50 mol H2SO4 in 5.00 L solution b) 0.200 mol C2H5OH in 35.0 ml of solution c) 44.35 g KOH in 125.0 ml of solution d) 2.46 g oxalic acid in 750.0 ml of solution e) 22.00 ml triethylene glycol, (CH2OCH2CH2OH)2 (D = 1.127 g/ml) in 2.125 L of solution f) 15.0 ml isopropylamine, CH3CH(NH2)CH3, (D = 0.694 g/ml) in 225 ml of solution 17) A stock bottle of nitric acid indicates that the solution is 67.0 % HNO3 by mass (67.0 g HNO3/100.0 g solution) and has a density of 1.40 g/ml. Calculate the molarity of the solution. 18) A stock bottle of solution is 50.0 % KOH by mass (50.0 g KOH/100.0 g solution) and has a density of 1.52 g/ml. Calculate the molarity of the solution. 19) If 50.00 ml of 19.1 M NaOH is diluted to 2.00 L, calculate the molarity of NaOH in the diluted solution.

Page 20 Name: ______Date: ______Period #: ______

Matter--Substances vs. Mixtures

All matter can be classified as wither a substance (element or compound) or a mixture (heterogeneous or homogeneous).

Directions: Classify each of the following as a substance or a mixture. If it is a substance, write element or a compound in the substance column. If it is a mixture, write heterogeneous or homogeneous in the mixture column.

Type of Matter Substance Mixture 1. Chlorine 2. Water 3. Soil 4. Sugar water 5. Oxygen 6. dioxide 7. Rocky road ice cream 8. Pure Alcohol 9. Air 10. Iron

------

Physical vs. Chemical Changes

In a physical change, the original substance still exists. It has changed in form only. In contrast, a new substance is produced when a chemical change occurs. Energy always accompanies chemical changes.

Directions: Classify each of the following as a chemical (C) or physical (P) change.

1. Sugar dissolves in water. ______

2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water, and heat. ______

3. A pellet of sodium is sliced in two. ______

4. Water is heated and changed to steam. ______

5. Potassium chlorate decomposes to potassium chloride and oxygen gas. ______

6. Iron rusts. ______

7. When placed in water, a sodium pellet catches on fire as hydrogen gas is liberated and sodium hydroxide forms. ______

8. Evaporation ______

9. Ice Melting ______

10. Milk sours. ______

Page 21 11. Sugar dissolved in water. ______

12. Wood rotting ______

13. Pancakes cooking on a griddle ______

14. Grass growing in a lawn ______

15. A tire is inflated with air. ______

16. Food is digested in the stomach. ______

17. Water is absorbed by a paper towel. ______------

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 altered fundamentally when observing a chemical property. For example, iron reacts with oxygen to form rust, which is also known as iron oxide.

Directions: Classify each of the following properties as either chemical or physical by denoting with a check mark.

Physical Property Chemical Property

1. Blue color 2. Density 3. Flammability 4. Solubility 5. Reacts with acid to form H2 6. Supports combustion 7. Sour taste 8. Melting Point 9. Reacts with water to form a gas 10. Reacts with base to form water 11. Hardness 12. Boiling Point 13. Can neutralize a base 14. Luster 15. Odor

Page 22 Atoms, Molecules, and Ions

I. Laws and Theories: A Brief Historical Introduction A. Laws of Chemical Combination 1. Lavosier (1743-1794): The Law of Conservation of Mass a. The total mass remains constant during a chemical reaction Example: HgO → Hg + O2 Mass of reactants = Mass of products

2. Proust (1754-1826): The Law of Constant Composition or Definite Proportions a. All samples of a compound have the same composition or the same proportions by mass of the elements present Example: NaCl is 39.34 % Na and 60.66 % Cl

Example: O:Mg in MgO is 0.6583:1. What mass of MgO will form when 2.000 g Mg is converted to MgO by burning in pure O2?

2.000 g Mg x 0.6583 O = 1.317 g O 1 Mg

2.000 g Mg + 1.317 g O = 3.317 g MgO

B. John (1766-1844) and the Atomic Theory of Matter (1803) 1. Law of Multiple Proportions a. When two or more different compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of a second element are in the ratio of small whole numbers Examples: CO vs. CO2; SO2 vs. SO3

2. Atomic Theory a. All matter is composed of extremely small, indivisible particles called atoms b. All atoms of a given element are alike in mass and other properties, but atoms of one element differ from the atoms of every other element c. Compounds are formed when atoms of different elements unite in fixed proportions d. A chemical reaction involves a rearrangement of atoms. No atoms are created, destroyed or broken apart in a chemical reaction.

3. Dalton used the Atomic Theory to restate the Law of Conservation of Mass: Atoms can neither be created nor destroyed in a chemical reaction, and as a consequence, the total mass remains unchanged.

C. The Divisible Atom 1. Subatomic Particles a. Proton 1) Relative mass = 1

Page 23 2) positive electrical charge = +1 b. Neutron 1) Relative mass = 1 (although slightly greater than a proton) 2) no charge = 0 c. Electron 1) mass = 1/1836 of the mass of a proton 2) negative electrical charge = -1

Particle Symbol Approximate Relative Location in Relative Mass Charge Atom Proton p+ 1 1+ Inside nucleus Neutron n 1 0 Inside nucleus Electron e- 0.000545 1- Outside nucleus

2. An atom is neutral (has no net charge) because p = e-. 3. The number of protons (Z) determines the identity of the element. 4. Mass number (A)= protons + neutrons a. neutrons = A – Z

52 - Example: 24Cr Determine the number of p, e , and n

24 p, 24 e-, and 28 n

5. Isotopes: atoms that have the same number of protons but different numbers of neutrons Examples: 1H, 2H, 3H [or H-1, H-2, H-3] All have 1 p, but n is 0,1,2, respectively 32S , 35S Both have 16 p, but n is 16 and 19, resp. 59Co , 60Co Both have 27 p, but n is 32 and 33, resp.

6. Isobars: atoms with the same mass number but different atomic numbers a. Example: 14C, 14N

D. Atomic Masses 1. Dalton arbitrarily assigned a mass number to one atom (H-1) and determined the masses of other atoms relative to it. 2. Current atomic mass standard is the pure isotope C-12. 3. Atomic mass unit (amu): 1/12 the mass of C-12. 4. Atomic Mass: weighted average of the masses of the naturally occurring isotopes of that element Example: Ne-20: 90.51 %, 19.99244 u Ne-21: 0.27 %, 20.99395 u Ne-22: 9.22 %, 21.99138 u

Average weighted atomic mass = (0.9051 x 19.99244 amu) + (0.0027 x 20.99395 amu) + (0.0922 x 21.99138 amu) = 20.17941 amu

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Advanced Example: Two isotopes of are As-75 and As-77, with atomic masses of 74.503 amu, and 76.752 amu, respectively. What are the percent natural abundances of these isotopes? {Hint: set one at x and the other at 1-x}

(x * 74.503 amu) + ((1-x) (76.752 amu)) = 74.92 amu Solve for x algebraically

x = .81458 1-x = .18542 Percent abundance of As-75 is 81.458 % and As-77 is 18.542 %

E. The Periodic Table 1. Dmitri Mendeleev’s (1869) Periodic Table: a. Arranged elements in order of increasing atomic mass, from left to right in rows, and from top to bottom in groups b. Elements that most closely resemble each other are in the same vertical group (more important than increasing mass). c. The group similarity recurs periodically (once in each row) d. Gaps for missing elements; predict characteristics of yet to be discovered elements based on their placement 2. Modern Periodic Table a. Elements are placed according to increasing atomic number b. Groups or Families: vertical columns c. Periods: horizontal rows d. Two series “pulled out” 1) Lanthanide and Actinide Series e. Classes 1) Most elements are Metals, which are to the left (NOT touching) the stair-step line a) luster, good conductors of heat and electricity b) malleable (hammered into thin sheets or foil), ductile (drawn into wires) c) Solids at room temperature (except mercury) 2) Nonmetals are to the right (NOT touching) of the stair-step line a) poor conductors of heat and electricity b) many are at RT 3) Metalloids: touch the vertical and or horizontal of the stair-step line (except Al and Po II. Introduction to Molecular and Ionic Compounds A. Key Terms 1. Chemical Symbols are used to represent elements 2. Chemical Formulas are used to represent compounds a. Subscripts indicate how many atoms of each element are present or the ratio of ions B. Molecules and Molecular Compounds 1. Molecule: group of two or more atoms held together in a definite spatial arrangement by covalent bonds 2. Molecular Compound: molecules are the smallest entities, and they determine the properties of the substance 3. Empirical Formula: simplest formula for a compound

Page 25 a. indicates the elements present in their smallest integral ratio Example: CH2O = 1 C: 2 H: 1 O 4. Molecular Formula: true formula for a compound {n = MFmass/EFmass} a. indicates the elements present and in their actual numbers Example: C6H12O6 = 6 C: 12 H: 6 O 5. Diatomic Elements: two-atom molecules, which don’t exist as single atoms in nature a. Br2, I2, N2, Cl2, H2, O2, F2 6. Polyatomic Elements: many-atom molecules a. S8, P4 7. Structural Formulas: shows the arrangement of atoms a. lines represent covalent bonds between atoms

C. Writing Formulas and Names of Binary Molecular Compounds 1. Binary Molecular Compounds: comprised of two elements, which are usually nonmetals a. The first element symbol is usually the element that lies farthest to the left of its period and/or lowest in its group (exceptions: H and O) [Figure 2.7] b. Molecular compounds contain prefixes for subscripts (exception: mono is not used for the first element) c. The name consists of two words: Prefix Number (prefix) element prefix –ide form mono- 1 di- 2 tri- 3 * rule with oxide tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 nona- 9 deca- 10

D. Ions and Ionic Compounds 1. Ion: charged particle due to the loss or gain of one or more electrons a. Monatomic Ion: a single atom loses or gains one or more e- 1) use the PT to predict charges 2) more than one ion can form with transition elements b. Cation: positively charged ion [usually a metal] c. Anion: negatively charged ion [usually a nonmetal] d. Polyatomic Ion: a group of covalently bonded atoms loses or gains one or more e- e. Ionic Compounds: comprised of oppositely attracted ions held together by electrostatic attractions; no identifiable small units 2. Formulas and Names for Binary Ionic Compounds a. Cation anion (–ide form) b. Cation (Roman Numeral) anion (–ide form) 3. Polyatomic Ion: charged group of bonded atoms a. suffixes are often –ite (1 less O) and –ate

Page 26 b. prefixes are often hypo- (1 less O than –ite form) and per-(1 more O than –ate form) c. Example: Hypochlorite ClO- - Chlorite ClO2 - Chlorate ClO3 - Perchlorate ClO4 4. Hydrates: ionic compounds in which the formula unit includes a fixed number of water molecules together with cations and anions

a. Example: CaCl2 o 6H2O Calcium chloride hexahydrate b. Anhydrous: without water

E. Acids, Bases, and Salts 1. Basic Characteristics of Acids and Bases when dissolved in water a. Acids: 1) taste sour 2) sting or prick the skin 3) turn litmus paper red 4) react with many metals to produce ionic compounds and H2(g) 5) react with bases b. Bases 1) taste bitter 2) feel slippery or soapy 3) turn litmus paper blue 4) react with acids 2. The Arrhenius Concept (1887) a. Acid: molecular compound that ionizes in water to form a solution containing H+ and anions b. Base: compound that ionizes in water to form a solution containing OH- and cations c. Neutralization: the essential reaction between and acid and a base, called neutralization, is the combination of H+ and OH- ions to form water and a salt 1) Example: HCl + NaOH → NaCl + H2O

3. Formulas and Names of Acids, Bases, and Salts a. Arrhenius Bases: cation hydroxide 1) Examples: NaOH = Sodium hydroxide KOH = Potassium hydroxide Ca(OH)2 = Calcium hydroxide b. Molecular Bases: do not contain OH- but produce them when the base reacts with water 1) Example: NH3 = Ammonia c. Binary Acids: H combines with a nonmetal 1) Examples: HCl(g) = Hydrogen chloride HCl(aq) = Hydrochloric acid HI(g) = Hydrogen iodide HI(aq) = Hydroiodic acid H2S(g) = Hydrogen sulfide H2S (aq) = Hydrosulfuric acid

Page 27 d. Ternary Acids: H combines with two nonmetals 1) oxoacids: H combines with O and another nonmetal a) Examples: Hypochlorous Acid HClO Chlorous Acid HClO2 Chloric Acid HClO3 Perchloric Acid HClO4 Sulfurous Acid H2SO3 Sulfuric Acid H2SO4 b) ate-ic ite-ous III. Introduction to Organic Compounds (Carbon-based Compounds) {NO longer assessed in APC but helpful for SAT II Chemistry} A. Alkanes: Saturated Hydrocarbons (contain H and C) 1. molecules contain a maximum number of H Atoms Stem Number meth- 1 2. Formula: CnH2n+2 eth- 2 a. Methane: CH4 prop 3 b. Ethane: C2H6 but- 4 c. Propane: C3H8 pent- 5 d. Butane: C4H10 hex- 6 1) Two possible structural formulas hept- 7 --except methane, ethane, & propane oct- 8 non- 9 2) Compounds with the same molecular formula dec- 10 but different structural formulas are known as isomers, and they have different properties. B. Cyclic Alkanes 1. Formula: CnH2n 2. prefix = cyclo- C. Alkenes: unsaturated hydrocarbon 1. Formula: CnH2n a. Ethene: C2H4 b. Propene: C3H6 c. Butene: C4H8 D. Alkynes: unsaturated hydrocarbon 1. Formula: CnH2n-2 a. Ethyne: C2H2 b. Propyne: C3H4 c. Butyne: C4H6 E. Homology 1. a series of compounds whose formulas and structures vary in a regular manner also have properties that vary in a predictable manner a. Example: Both the densities and boiling points of the straight-chain alkanes increase in a continuous and regular fashion with increasing numbers of C.

F. Types of Organic Compounds: 1. Functional Group: atom or group of atoms attached to or inserted in a hydrocarbon chain or ring that confers characteristic properties to the molecule a. usually where most of the reactions of the molecule occur

Page 28

2. Alcohols (R-OH) where R represent the hydrocarbon a. Examples: CH3OH = methanol CH3CH2OH = ethanol CH3CH2CH2OH = 1-propanol CH3CH(OH)CH3 = 2-propanol or isopropanol b. Not bases!

3. Ethers (R-O-R’) where R’ can represent a different hydrocarbon than R a. Example: CH3CH2OCH2CH3 = Diethyl ether

4. Carboxylic Acids (R-COOH) a. Examples: HCOOH = methanoic or formic acid CH3COOH = ethanoic or acetic acid

b. the H of the COOH group is ionizable; the acid is classified as a weak acid

5. Esters (R’-COOR) a. Flavors and fragrances b. Examples: CH3COOCH2CH3 = ethyl acetate CH3COOCH2CH2CH2CH2CH3 = pentyl acetate

6. Ketones (R-CO-R’)

7. Aldehydes (R-CO-H)

8. Amines (R-NH2, R-NHR’, R-NR’R”) a. most common organic bases; related to ammonia b. one or more organic groups are substituted for the H in NH3 c. Examples: CH3NH2 = methyl amine CH3CH2NH2 = ethyl amine

Page 29 Page 30 Page 31 Page 32 Name: ______Date: ______Period #: ______Molecular Formula Writing and Naming Name the following compounds:

1. SF4 ______

2. B3Cl6 ______

3. PBr5 ______

4. N2O5 ______

5. S3 I9 ______

6. S2O4 ______------Write the chemical formula for each of the following compounds: 7. carbon dioxide ______8. sulfur hexafluoride ______9. dinitrogen tetroxide ______10. trisulfur heptaiodide ______11. disulfur pentachloride ______12. triphosphorus monoxide ______------Ionic Formula Writing and Naming Directions: Name the following ionic compounds.

13. MgCl2 ______

14. NaF ______

15. Na2O ______

16. Al2O3 ______17. KI ______

18. AlF3 ______

19. Mg3N2 ______

20. FeCl3 ______

21. MnO2 ______

22. CrN2 ______------Compounds that include Polyatomic Ions:

23. Ca(OH)2 ______

24. (NH4)2S ______

25. Al2(SO4)3 ______

26. H3PO4 ______

27. Ca(NO3)2 ______

Page 33 28. CaCO3 ______

29. Na2SO3 ______

30. Co(CH3COO)3 ______

31. Cu2(SO3)3 ______

32. Pb(OH)2 ______------Directions: Write the correct formula for each of the following compounds.

1. Magnesium sulfide ______2. Calcium phosphide ______3. Barium chloride ______4. Potassium nitride ______5. Aluminum sulfide ______6. Magnesium oxide ______7. Calcium fluoride ______8. fluoride ______9. Barium iodide ______10. Aluminum nitride ______11. nitride* ______12. (II) bromide ______13. (IV) phosphide ______14. (II) sulfide ______------Compounds that include Polyatomic Ions:

15. Aluminum phosphate ______16. Sodium bromate ______17. Aluminum sulfite ______18. Ammonium sulfate ______19. Ammonium acetate ______20. Magnesium chromate ______21. Sodium dichromate ______22. Zinc hydroxide* ______23. Copper(II) nitrite ______24. (II) hydroxide ______25. Iron(II) sulfate ______26. Iron(III) oxide ______------

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