2015 Summer Assignment...Tips for Learning the Ions Polyatomic Anions Most of the work on...

AP Chemistry Summer Enrichment Assignment Dear AP Chemistry Students and Parents, Welcome to AP Chemistry! I am so excited to start the new school year and to meet all of my new students! I hope that you are having a great end to this school year and that you have a wonderful summer planned. AP Chemistry is a highly rigorous course that requires lots of effort and motivation. If you are willing to put in the time, the rewards are huge! ☺ For the summer, I believe that you should have a break, relax, and enjoy yourselves…but I don’t want you forgetting what you had previously learned in honors chemistry. If you find yourself struggling to answer any of the questions or problems included in this assignment, I expect that you will first reference your textbook for clarification. If you are still experiencing difficulty, you should contact me in order to get help. Due dates for the assignment are listed below. There are two parts to the assignment: Part 1 is the memorization portion -‐ memorizing the common polyatomic ions & solubility rules. Part 2 is our first unit, covering chapters 1-‐3 of your text. These chapters are a review of first year chemistry. The assignment may be done in any order you choose, however I strongly suggest you work through your text, the notes, and then take the multiple choice quiz, as the notes are designed to prepare for you this assessment. You do not need to print out this entire packet! You may choose to print only the quizzes and honor code sheet to turn in at the beginning of school. I will also make the quizzes available as Google forms on my website. Please note, many multiple-‐choice questions on the AP Exam involve math but you will not have a calculator to complete them. You should estimate and round numbers to make them easier to work with, in your efforts to solve the problem. For this reason, you are not allowed the use of a calculator on the multiple-‐choice quizzes on the summer assignment. It is my suggestion that you purchase a review book for AP Chemistry. These are great resources for unit test review and the AP Exam! If you have any questions, please feel free to see me or e-‐mail me. Take care and have a great summer!!! Thanks, Mrs. Elyse Vaughan

Part 1: Common ions & solubility rules This part of the summer assignment for AP Chemistry is quite simple (but not easy). You need to master the formulas, charges, and names of the common ions. In the first week of the school year, you will be given a quiz on the ions & on solubility rules. You will be asked to:

• write the names of these ions when given the formula and charge • write the formula and charge when given the names • write net ionic equations predicting products based on solubility rules

I have included several resources in this packet. First, there is a list of the ions that you must know on the first day. This list also has some suggestions for making the process of memorization easier. There are naming patterns that greatly simplify the learning of the polyatomic ions as well. Second, there is a list of the solubility rules. Remember that the rules are applied in order, so any earlier rule overrules the later rules. Also included is a copy of the periodic table used in AP Chemistry. Notice that this is not the table used in first year chemistry. The AP table is the same that the College Board allows you to use on the AP Chemistry test. Notice that it has the symbols of the elements but not the written names. You need to take that fact into consideration when studying for the afore-‐mentioned quiz! Doubtless, there will be some students who will procrastinate and try to do all of this studying just before the start of school. Those students may even cram well enough to do well on the initial quiz. However, they will quickly forget the ions, and struggle every time that these formulas are used in lecture, homework, quizzes, tests and labs. All research on human memory shows us that frequent, short periods of study, spread over long periods of time will produce much greater retention than long periods of study of a short period of time. I could wait and throw these at you on the first day of school, but I don’t think that would be fair to you. Use every modality possible as you try to learn these – speak them, write them, visualize them.

Part 2: Unit 1 The first test will cover Unit 1 (parts 1, 2, and 3), which is a review of first year chemistry. This assignment is designed to prepare you for that test. All documents referenced can be found on my website: https://sites.google.com/a/nsacademy.org/ap-‐chemistry-‐vaughan/?pli=1 All quizzes will be posted on my web site by May 22nd, 2015, and should be taken and submitted via e-‐mail by Monday, August 24th. They are all timed – you should take no more than 24 minutes. No calculators or notes may be used in completing the quizzes. Signing the honor pledge indicates that you have not used your notes or a calculator to complete the quiz, and that you took only 24 minutes. Quizzes may be taken only once and will count as 1 quiz grade. Unit 1 (Part 1): Chemical Foundations

• For Review Purposes: o This section of material is covered in Chapter 1 (pages 1-‐41) in your textbook. o Notes (AP Unit 1 Part 1 Notes.doc ) are provided. I suggest completing all boxed problems (exercise 1.1

– 1.13 in the notes) using your text and the notes provided. Answers are provided so that you can make sure you did the problem correctly.

o A video review of these notes is available at http://vimeo.com/14216778 that you can use as an alternative source of review if you prefer.

• Assignment: o Quiz -‐ I strongly suggest completing the review listed above before taking the quiz!

Unit 1 (Part 2): Atoms, Molecules, and Ions

• For Review Purposes: o This section of material is covered in Chapter 2 (pages 42-‐80) in your text. o Notes (AP Unit 1 Part 2 Notes.doc) are provided. I suggest completing all boxed problems (exercise 1.14


o A video review of these notes is available at http://vimeo.com/14217141 that you can use as an alternative source of review if you prefer.

• Assignment: o Quiz -‐I strongly suggest completing the review listed above before taking the quiz!

Unit 1 (Part 3): Stoichiometry

• For Review Purposes: o This section of material is covered in Chapter 3 (pages 81-‐137) in your text. o Notes (AP Unit 1 Part 3 Notes.doc) are provided. I suggest completing all boxed problems (exercise 1.24


o A video review of these notes is available at http://vimeo.com/13588248 and http://www.vimeo.com/13217975 that you can use as an alternative source of review if you prefer.

• Assignment: o Quiz -‐I strongly suggest completing the review listed above before taking the quiz! o Complete the two free response questions in the document “AP Chemistry Summer Assignment Stoich

FR”. This will be collected on the first day of class.

AP Chemistry Summer Assignment Honor Pledge

I have not: • Received help from another individual in completing these quizzes. • Given help to another individual in completing these quizzes. • Used a calculator or my notes to complete these quizzes. • Spent more than 24 minutes on any one quiz.

All quizzes will be available for you to take on Friday, May 22nd. Quizzes may only be taken one time and must be completed within 24 minutes. All quizzes must be complete by Monday, August 24th. You must sign the pledge below and have a parent sign in order to earn credit.

Unit 1 Part 1 MC Quiz Taken on ____________________________ (date) Student Signature _________________________________ Parent Signature ___________________________________ Unit 1 Part 2 MC Quiz Taken on ____________________________ (date) Student Signature _________________________________ Parent Signature ___________________________________ Unit 1 Part 3 MC Quiz Taken on ____________________________ (date) Student Signature _________________________________ Parent Signature ___________________________________

Please print this document and give to Mrs. Vaughan on the first day of class.

Tips for Learning the Ions 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 SO42-; sulfite has the same charge but one less oxygen (SO3

2-) b. nitrate is NO3

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

2. 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”) Hypochlorite is “under” chlorite, meaning it has one less oxygen

c. The prefix “hyper” means “above” or “too much” (think “hyperkinetic”) The prefix “per” is derived from “hyper” so perchlorate (hyperchlorate) has one more oxygen than chlorate.

d. Notice how this sequence increases in oxygen while retaining the same charge: ClO- ClO2

- ClO3- ClO4

- hypochlorite chlorite chlorate perchlorate

3. If you know that a sulfate ion is SO42- 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.

Example: PO4

3- HPO42- H2PO4

- phosphate hydrogen phosphate dihydrogen phosphate

Ions to Memorize -1 Charge -2 charge -3 charge

Amide NH2 Carbonate CO3 Phosphate PO4 Acetate C2H3O2 Chromate CrO4 +1 charge Bromate BrO3 Dichromate Cr2O7 Ammonium NH4

Chlorate ClO3 Oxalate C2O4 Hydronium H3O+ Cyanide CN Peroxide O2 +2 charge Hydroxide OH Sulfate SO4 Zinc Zn Iodate IO3 Silicate SiO3 Mercury (I) Hg2 Nitrate NO3 Thiosulfate S2O3 +3 charge Permanganate MnO4 Aluminum Al Thiocyanate SCN

Solubility Rules Apply these rules in order. 1. All group 1 and ammonium (NH4

+) salts are soluble. 2. All NO3

-, ClO3-, ClO4

-, and C2H3O2- salts are soluble.

3. All Ag+, Pb22+, and Hg2+ salts are insoluble.

4. All Cl-, Br-, and I- salts are soluble. 5. All SO4

2- compounds are soluble except those of the heavy group 2 metals (Ca2+, Ba2+, Sr2+) 6. All OH- salts are insoluble except those of group 1 and heavy group 2 metals (Ca2+, Ba2+, Sr2+) 7. All sulfide (S2-) compounds are insoluble except those of group 1 & 2. 8. All SO3

2-, CO3-, CrO4

2-, and PO43- compounds are insoluble.

Unit 1: Chemical Foundations (Review) Adapted from Chemical Foundations (NMSI, Rene McCormick) and Chemistry (Brown/LeMay)

1.1 An Overview of Chemistry # Chemistry – the study of MATTER and ENERGY! # Matter – the physical material of the universe; anything

with mass and volume (takes up space), exhibits intertia $ Composed of aprox. 100 different types of atoms

$ Atoms can be broken down or combined to form new substances (ex: water) (this means reactions can be reversible!)

# Property – any characteristic that makes a particular type of matter distinguishable from other types

# Element – a group of 1 type of atom; elements combine to make compounds # Molecules – 2+ atoms joined together by covalent bonds

1.2 The Scientific Method (p. 13)

# Systematic way of finding a possible answer to a problem # Steps in the Scientific Method:

$ Make observations -‐ qualitative or quantitative (quant. obs. = a measurement) $ Formulate a hypothesis – a possible explanation for the observation $ Perform experiments – design experiment to test the hypothesis (try to control as

many variables as you can, REPEAT experiments to improve validity) # Experiments may lead to new observations, therefore new experiments

$ If experiments are repeated and similar results are found by many different researchers, the hypotheses may become a theory (an attempt to explain WHY the observation happened)

# Theory can be similar to a model (an example we use to explain a natural phenomenon – if new evidence is found, the model/theory may change!)

# Example: geocentric model of the universe, atomic theory # Scientific law – summary of observations (theories attempt to explain these observations)

A LAW describes WHAT happens while a THEORY describes WHY it happens! Examples: Law of Conservation of Mass – mass before and after a reaction must be the same (massreactants = massproducts) Law of Conservation of Energy (1st Law of Thermodynamics) – Energy can’t be created nor destroyed, it can only change forms or be transferred

# Analysis of results of experiments are subject to human interpretation (data misinterpretation, emotional attachments, politics, ego, MONEY, religious beliefs)

$ Galileo – due to religious influence, forced to “revise” his astronomical observations $ Lavoisier – “father of modern chemistry” was beheaded due to his political affiliations $ Nuclear devices, fertilizers, explosives

(rapid change of a solid/liquid into a gas where the molecules quickly become very far apart exerting large amounts of energy)

1.3 Units of Measurement (pp. 14 – 15) # Measurement – a quantitative observation; always

contains a number AND a unit! # Measurement systems – English (U.S. and a few

countries in Africa) and Metric (everybody else!) # SI System – (Le Systeme International) in 1960 an

international agreement was reached to set up a system of units so that scientists all over the world could communicate in the same “language”; all units based on the Metric system

# Volume – derived from length (think about a cube…1.0 cm on a side, volume = 1.03 or 1.0 cm3)

# Mass – measure of the resistance of an object to a change in its state of motion (intertia); the amount of matter present (NOT the same as weight! Weight refers to the response of mass to gravity)

# Temperature – the average kinetic energy of a substance (measured in Kelvin, °C, or °F)

# Derived unit – based on the combination of units (ex: measure length on 3 sides of cube to get

volume)

# Scientific Notation – exponential form to represent very large or very small numbers $ Negative exponents used w/ very small numbers (Ex: 0.001 " 1x10-‐3) $ Positive exponents used w/ very large numbers (Ex: 1000 " 1x103) $ The 10x is used to represent how many decimal places are in the number therefore are not used to determine

the significant figures in the number

Ex: 101.09 would be written 1.0109x102 (5 significant figures)

1.4 Uncertainty in Measurement (pp. 20 – 25) # Exact numbers (ex: 12 eggs in 1 dozen, 60 s in 1 min, anything that can be counted) do not contain error # Inexact numbers: any number that is measured, numbers with some error (there is always some estimation when taking a

measurement) # When taking a measurement, always record one digit more than the last marking on the instrument (that is the estimated

digit) # Precision: how close individual measurements are to one another (consistent, repeatability) # Accuracy: how close an individual measurement is to the “right answer”

The results of several dart throws show the difference between precise and accurate. (a) Neither accurate nor precise (large random errors). (b) Precise but not accurate (small random errors, large systematic error). (c) Bull’s-‐eye! Both precise and accurate (small random errors, no systematic error).

# Types of Error

$ Random error (indeterminate) – equal probability of a measurement being high or low $ Systematic error (determinate) – occurs in the same direction each time

# Significant figures: all digits of a measured quantity (including the estimated digit) $ The Rules!

# Non-‐zero digits are significant # A zero is significant if it is terminating AND right of the decimal OR if it is

sandwiched between significant figures

Exercise 1.1 Precision and Accurcy To check the accuracy of a graduated cylinder, a student filled the cylinder to the 25-‐mL mark using water delivered from a buret and then read the volume delivered. Following are the results of five trials: Trial Volume Shown by Volume Shown Graduated Cylinder by the Buret 1 25 mL 26.54 mL 2 25 mL 26.51 mL 3 25 mL 26.60 mL 4 25 mL 26.49 mL 5 25 mL 26.57 mL Average 25 mL 26.54 mL Is the graduated cylinder accurate? No b/c volume shown by the buret is significantly different from the grad cyl. This is a systematic error b/c grad is low each time.

EXACT AND COUNTING NUMBERS HAVE AN INFINITE NUMBER OF SIG FIGS!

# Significant Figures in calculations: $ Multiplication/division – answer has the same number of sig figs as the measurement with the FEWEST number of

sig figs

4.56 × 1.4 = 6.38 6.4

$ Addition/subtraction – answer has the same number of places after the decimal as the measurement with the FEWEST number of places after the decimal point

12.11 18.0 ← limiting term 1.013 31.123 31.1

$ Log functions – the significant figures in log functions are only counted after the decimal

pH of solution with [H+] = 1.50x10-‐3 M pH = 2.824 (3 sig figs)

$ When multiple operations are present, sig figs should be determined by order of operations (this is the CORRECT

thing to do – although most of the time you can look at the problem and go with the fewest # of sig figs from the given quantities)

# Rounding Rules: $ Round at the end of all calculations to get the appropriate number of significant figures $ Look at the significant figure one place beyond your desired number of significant figures. If greater than 5, round

up, if less than 5 drop the digit. $ Don’t double round

Ex: Round 4.348 to 2 s.f " 4.3

1.5 Dimensional Analysis (pp. 25 – 29) # Problem solving strategy # Conversion factor – fraction whose numerator and denominator are the same quantity

expressed in different units $ A pin measures 2.85 cm in length. What is its length in inches?

2.54 cm = 1 in Conversion factors could be

2.54 𝑐𝑚 or __1 𝑖𝑛__

1 𝑖𝑛 2.54 𝑐𝑚

Pick the one that lets you cancel out units!

2.85 𝑐𝑚 × 1 𝑖𝑛/2.54 𝑐𝑚 = 1.12 in

Exercise 1.2 Significant Figures Give the number of significant figures for each of the following results. a. A student’s extraction procedure on tea yields 0.0105 g of caffeine. 3 b. A chemist records a mass of 0.050080 g in an analysis. 5 c. In an experiment, a span of time is determined to be 8.050 × 10-‐3 s. 4

Exercise 1.3 Unit Conversions A pencil is 7.00 in. long. What is its length in centimeters?

17.8 cm

1.6 Temperature (pp. 15 – 16)

# Measure of the average kinetic energy of a substance # Units are Kelvin, °C, °F

$ 1 Kelvin degree = 1 Celsius degree $ 1 Celsius degree > 1 Fahrenheit degree

# Conversions

TF = TC + 32ºF

TK = TC + 273 K TC = TK -‐ 273ºC

Exercise 1.8 Unit Conversions The concentration of carbon monoxide in an urban apartment is 48 µg/m3. What mass of carbon monoxide (in grams) is present in a room measuring 9.0 ft x 14.5 ft x 18.8 ft?

3.3x10-6 g

Exercise 1.4 Unit Conversions You want to order a bicycle with a 25.5-‐in. frame, but the sizes in the catalog are given only in centimeters. What size should you order?

64.8 cm

Exercise 1.5 Unit Conversions A student has entered a 10.0-‐km run. How long is the run in miles?

6.21 mi

Exercise 1.6 Unit Conversions The speed limit on many highways in the United States is 55 mi/h. What number would be posted in kilometers per hour?

89 km/hr Exercise 1.7 Unit Conversions A student calculates the volume of a cube as 2.35 cm3. What is the volume of the cube in mm3?

2350 mm3

Exercise 1.9 Temperature Conversions Normal body temperature is 98.6°F. Convert this temperature to the Celsius and Kelvin scales.

37 C & 310 K

1.7 Density (pp. 17 – 20) # Density: amount of mass in a certain volume of a substance

# 𝐷=𝑚𝑎𝑠𝑠/𝑣𝑜𝑙𝑢𝑚𝑒 # Units for density of a solid (g/cm3) of a liquid (g/mL) # 1 mL = 1 cm3

Exercise 1.12 Density A cylindrical rod formed from silicon is 16.8 cm long and has a mass of 2.17 kg. The density of silicon is 2.33 g/cm3. What is the diameter of the cylinder?

8.40 cm

Exercise 1.10 Temperature Conversions One interesting feature of the Celsius and Fahrenheit scales is that -‐40°C and -‐40°F represent the same temperature. Verify that this is true.

Plug in -40 into one of the equations and solve for the unknown. You should get -40 also.

Exercise 1.11 Temperature Conversions Liquid nitrogen, which is often used as a coolant for low-‐temperature experiments, has a boiling point of 77 K. What is this temperature on the Fahrenheit scale?

-320.8 F

Exercise 1.11 Density A chemist, trying to identify the main component of a compact disc cleaning fluid, finds that 25.00 cm3 of the substance has a mass of 19.625 g at 20°C. The following are the names and densities of the compounds that might be the main component. Compound Density in g / cm3 at 20°C Chloroform 1.492 Diethyl ether 0.714 Ethanol 0.789 Isopropyl alcohol 0.785 Toluene 0.867 Which of these compounds is the most likely to be the main component of the compact disc cleaner?

Density of the liquid is 0.7850 g/mL… Ethanol

1.8 Classification and Properties of Matter (pp. 4 – 13) # Describe matter by physical state or what it’s made of

$ States of matter: # solid (def. shape/def. volume, almost impossible to compress, molecules tightly packed in specific

arrangements, molecules “wiggle” around a fixed point) # liquid (def. volume/indef. shape, can’t be easily compressed, molecules packed closer together, molecules

still vibrate but also have rotational and translational motion and can slide past one another but are still close together

# gas (indef. Shape/indef. volume, particles, molecules are in constant random motion, molecules vibrate, rotate, translate and are independent of each other…so they are VERY far apart and easily compressed, molecules collide with each other and walls of container)

# vapor – the gas phase of a substance that is normally a solid or liquid at room temperature (ex: water vapor)

# fluid – that which can flow; gases and liquids $ Composition

# Pure substance: (compounds AND elements) has specific properties, composition is always the same, atoms are CHEMICALLY combined, therefore must be chemically separated (common method is electrolysis)

$ Elements – can’t be broken down into anything except atoms of that element (Robert Boyle)

$ Atoms can be broken down into: # Nuclei and electrons # p+, n0, and e-‐ # quarks

Electrolysis is an example of a chemical change. In this apparatus, water is decomposed to hydrogen gas (filling the red balloon) and Oxygen gas (filling the blue balloon).

$ Compounds – 2+ elements bonded in the same proportion (ex: water is H2O – always 11% hydrogen and 89% oxygen)

# Compounds have different properties than the elements that make up the compound # Sometimes compounds have different properties depending on the form their in (ex:

solid NaCl vs. aqueous NaCl) # Mixture: PHYSICAL combinations of 2+ pure substances (each part of the mixture keeps its owns

properties); must be PHYSICALLY separated $ Heterogeneous mixture: composition is not uniform throughout mixture $ Homogeneous mixture (solution): composition is uniform

# Separation of Mixtures $ Each part of a mixture keeps its own properties, so you can separate the mixture based on

those properties $ Methods of separation:

# Filtration – separate based on size of particles # Distillation – separate based on boiling point # Chromatograpy – separate based on substances’ ability to

attract to different surfaces

Exercise 1.13 Density (a) Is density an extensive or intensive property? Justify your answer.

(b) Is density temperature dependent? Justify your answer.

Intensive b/c doesn’t matter how much is present; yes (ex: ice floats on water)

Paper chromatograph of ink. (a) A line of the mixture to be separate Is placed at one end of a sheet of porous paper. (b) The paper acts as a wick to draw up the liquid. (c) The component with the weakest attraction for the paper travels faster than those that cling to the paper.

# Properties of Matter $ Physical properties: properties describing appearance, can be measured without changing the identity of the

compound $ Chemical properties: describes how the substance reacts or what its made of $ Intensive properties: properties that DO NOT depend on how much of the substance is present (ex: melting point,

boiling point, density, specific heat capacity, heat of fusion, temperature) $ Extensive properties: properties that DO depend on how

much is there (ex: mass, volume) $ Physical change: changes appearance but not what its made

of (ex: changes in state of matter) $ Chemical change (reaction): changes into new substance,

atoms are rearranged

AP Chemistry Summer Assignment: Quiz 1

1. Assuming identical conditions, in which state of matter do the particles present have the greatest amount of energy?

a. Solids b. Liquids c. Gases d. Both solid and liquids e. Both liquids and gases

2. If matter is uniform throughout, cannot be separated into other substances by physical processes, and

cannot be decomposed into other substances by chemical processes, it is a. An element b. A compound c. A homogenous mixture d. A heterogeneous mixture e. A mixture of compounds

3. Which of the following is an illustration of the law of constant proportions?

a. Water boils at 100oC at 1 atm pressure. b. Water is 11% hydrogen and 89% oxygen by mass. c. Water is a compound. d. Water and salt have different boiling points. e. Water reacts with nonmetal oxides making acids.

4. Which of the following is a mixture?

a. Pure water b. Sea water c. Sodium chloride d. Sodium e. Carbon dioxide

5. Which of the following is often easily separated into its components by simple techniques such as

filtering or decanting? a. Heterogeneous mixture b. Compounds c. Homogenous mixture d. Elements e. Solutions

6. Considering a mixture consisting of sand in salt water; this mixture could be separated into its three components (sand, salt, and water) by first __________________the mixture and then ______________the remaining mixture.

a. Distilling, distilling b. Distilling, filtering c. Filtering, distilling d. Filtering, filtering e. Evaporating, filtering

7. Which of the following are chemical processes?

I. Rusting of a nail II. Freezing of water

III. Decomposition of water into hydrogen and oxygen gas IV. Dissolving of oxygen gas in water

a. II, III, and IV b. I, III, and IV c. I and III d. I and II e. I and IV

8. Which of the following is not an intensive property?

a. Density b. Temperature c. Melting point d. Volume e. Pressure

9. The correct value for the density of water is 1.0 g/ml. Which of the following sets of measurements is

precise but not accurate? a. 1.1 g/mL, 1.0 g/mL, 0.9 g/mL b. 1.0g/mL, 1.0g/mL, 0.8g/mL c. 0.6g/mL, 0.7g/mL, 0.6g/mL d. 1.1 g/mL, 0.1g/mL, 2.1 g/mL e. 1.1 g/mL, 1.2 g/mL, 0.8 g/mL

10. Which of the following base quantities has the wrong unit?

a. Mass in kg b. Amount of substance in g c. Length in m d. Temperature in K e. Time in s

11. Convert 30 degrees Celsius to Kelvin. a. 30 b. 330 c. 303 d. 243 e. 273

12. Convert 80 degrees Fahrenheit to Celsius.

a. 27 b. 353 c. 176 d. -193 e. 80

13. Which of the following shows the relative temperature correctly?

a. 12oC > 310 K b. 43oC < 300 K c. 25oC > 250 K d. 158oC > 450 K e. All of the above show the relative temperatures correctly.

14. Of the following, which is the smallest mass?

a. 25 kg b. 2.5 x 10-2 mg c. 2.5 x 1015 pg d. 2.5 x 109 fg e. 2.5 x 1010ng

15. One angstrom, symbolized Å, is 10-10 m. If you have 1 cm3, how many Å3 that is?

a. 1024 b. 10-24 c. 1030 d. 10-30 e. 10-9

16. 45 m/s = _________km/h

a. 2.7 b. 0.045 c. 1.6 x 102 d. 2.7 x 103 e. 1.6 x 105

17. The density of mercury is 13.6 g/cm3. The density of mercury is a. 1.36 x 10-2 kg/m3 b. 1.36 x 104 kg/m3 c. 1.36 x 108 kg/m3 d. 1.36 x 10-5 kg/m3 e. 1.36 x 10-4 kg/m3

18. In which one of the following numbers are all of the zeros significant?

a. 100.090090 b. 143.29 c. 0.05843 d. 0.1000 e. 0.0030020

19. How many significant figures should there be in the answer to the following computation?

(10.07 - 7.395)/2598.08 a. 1 b. 2 c. 3 d. 4 e. 5

20. A wooden object has a mass of 10.782 g and occupies a volume of 13.73 mL. What is the density of the

object determined to an appropriate number of significant figures? a. 8 x 10-1 g/mL b. 7.9 x 10-1 g/mL c. 7.86 x 10-1 g/mL d. 7.859 x 10-1 g/mL e. 7.8586 x 10-1 g/mL

Answer Sheet Quiz 1 Question Answer

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Unit 1: Atoms, Molecules, and Ions (Review) Adapted from Atoms, Molecules, Ions (NMSI, Rene McCormick) and Chemistry (Brown/LeMay)

1.9 Early History of Chemistry (pp. 38-‐39) # 1,000 B.C.—processing of ores to produce metals for weapons and ornaments; use of embalming fluids # 400 B.C.—Greeks—proposed all matter was make up of 4 “elements” : fire, earth, water and air # Democritus (460 – 370 BC): hypothesized that matter is made of tiny indivisible particles which he named ‘atomos’ (atoms) # Plato & Aristotle: hypothesized that there’s no such thing as an indivisible particle (meaning that every particle can be

divided into smaller pieces that make it up) # Next 2,000 years—alchemy—a pseudoscience where people thought they could turn metals into gold. Some good

chemistry came from their efforts—lots of mistakes were made! # 16th century—Georg Bauer, German , refined the process of extracting metals from ores & Paracelsus, Swiss, used minerals

for medicinal applications # Robert Boyle, English—first “chemist” to perform quantitative experiments of pressure versus volume. Developed a

working definition for “elements”. # Isaac Newton (1642-‐1727): studied properties of air that led him to believe atoms (indivisible particles) existed # 17th & 18th Centuries—Georg Stahl, German—suggested “phlogiston” flowed OUT of burning material. An object stopped

burning in a closed container since the air was “saturated with phlogiston” # Joseph Priestley, English—discovered oxygen which was originally called “dephlogisticated air” # John Dalton (1800s): Dalton’s Atomic Theory described the relationship between atoms and elements

1.10 Fundamental Chemical Laws (pp. 38 – 39) # Late 1700’s – combustion was studied extensively which led to the discoveries of CO2, N2, H2, & O2 # More experimentation led to discovery of more elements # Antoine Lavoisier:

$ Explained the true nature of combustion $ Published first modern chemistry textbook $ Stated Law of Conservation of Mass: Mass can neither be created nor destroyed. $ Pushed for quantitative experimentation $ Beheaded for ties to government during French Revolution

# Joseph Proust: $ Law of Definite Proportions: A given compound always contains the exact same proportions of elements by mass.

# Atom: smallest particle of an atom that still has the properties of that element $ An element contains only 1 type of atom $ A compound contains atoms of 2 or more elements

# John Dalton: $ Law of Multiple Proportions (1808): atoms can combine in different ratios to produce different

compounds (carbon and oxygen can combine to form carbon monoxide and carbon dioxide – the carbon monoxide is formed in a 1:1 ratio – 1 carbon for every 1 oxygen while the carbon dioxide is formed in a 1:2 ratio)

Dalton considered compounds of carbon and oxygen and found:

Mass of Oxygen that combines with 1 gram of C

Compound I 1.33 g Compound II 2.66 g

Therefore Compound I may be CO while Compound II may be CO2.

Exercise 1.14 Law of Multiple Proportions The following data were collected for several compounds of nitrogen and oxygen: Mass of Nitrogen That Combines With 1 g of Oxygen Compound A 1.750 g Compound B 0.8750 g Compound C 0.4375 g

Ratio for A to C is different even though all contain same elements.

Show how these data illustrate the law of multiple proportions. 1.11 Dalton, Gay-‐Lussac, Avogadro (pp. 38 – 39) Dalton’s ATOMIC THEORY OF MATTER: (based on knowledge at that time): 1. All matter is made of atoms. These indivisible and indestructible objects are the ultimate chemical particles. 2. All the atoms of a given element are identical, in both weight and chemical properties. However, atoms of different

elements have different weights and different chemical properties. 3. Compounds are formed by the combination of different atoms in the ratio of small whole numbers. 4. A chemical reaction involves only the combination, separation, or rearrangement of atoms; atoms are neither created nor

destroyed in the course of ordinary chemical reactions. **TWO MODIFICATIONS HAVE BEEN MADE TO DALTON’S THEORY: 1. Subatomic particles were discovered. 2. Isotopes were discovered.

# 1809 Joseph Gay-‐Lussac, French—performed experiments [at constant temperature and pressure] and measured volumes of gases that reacted with each other.

# 1811 Avogardro, Italian—proposed his hypothesis regarding Gay-‐Lussac’s work [and you thought he was just famous for 6.02 x 1023] He was basically ignored, so 50 years of confusion followed.

AVOGADRO’S HYPOTHESIS: At the same temperature and pressure, equal volumes of different gases contain the same number of particles.

1.12 Early Experiments to Characterize the Atom (pp. 39 – 42)

# Dalton, Gay-‐Lussac, & Avogadro " foundation for atom being the basis of chemistry # Through research, found that atoms are made of subatomic particles # Remember that opposite charges attract and like charges repel! # Research about ELECTRONS:

$ Cathode rays (1800s): scientists studied how charges traveled through vacuums

$ J.J. Thomson (1898 – 1903):

# Found that when high voltage was applied to an evacuated tube a “ray” (called a cathode ray b/c it came from the (-‐) electrode (cathode) when YOU apply a voltage across it) was produced

$ The ray was produced at the negative electrode $ Repelled by the (-‐) pole of an applied electric field, E

# Hypothesized that cathode rays were made of (-‐) charged particles (later called electrons) # Measured the deflection of beams of e-‐ to determine the mass to charge ratio of the e-‐

𝑒𝑚=(1.76x108 coulombs/g)

(e is the charge on the electron in Coulombs (C) and m is the mass) # Hypothesized that electrons made up only a little of the mass of an atom # Found that the cathode rays behaved the same way no matter what material was used

# Thomson discovered that he could repeat this deflection & calculation using electrodes of different metals, so all metals contained electrons and therefore all atoms contained electrons

# He knew all atoms were neutral, so he figured there must be some (+) charge in the atom too…led to the plum pudding model

$ Robert Milikan (1868-‐1953): # 1909 – Milikan sprayed charged oil drops into a chamber.

He stopped the effects of gravity by adjusting the voltage across 2 charged plates. The voltage needed to keep the oil drops from falling and the mass of the drops were used to calculate the charge on the oil drop which was a whole number multiple of the charge on an electron.

# Measured the charge of an electron # From the charge he calculated the mass (from Thomson’s

ratio) # Mass of electron: 9.10x10-‐28 g (really small compared to

protons and neutrons) # Research about the NUCLEUS:

$ Henri Becquerel (1852-‐1908): found that a piece of mineral containing uranium could produce its image on a photographic plate in the absence of light. He called this radioactivity and attributed it to a spontaneous emission of radiation by the uranium in the sample.

$ Pierre and Marie Curie (1867-‐1934) : further research on radioactive properties of compounds $ Three types of radiation:

# alpha, α-‐-‐equivalent to a helium nucleus; the largest particle radioactive particle emitted; 7300 times the

mass of an electron. Since these are larger that the rest, early atomic studies often involved them.

# beta, β-‐-‐a high speed electron. OR

# gamma, γ-‐-‐pure energy, no particles at all! Most penetrating, therefore, most dangerous.

$ Ernest Rutherford (1871 – 1937): # Carried out experiments to test

Thomson’s model of the atom # Gold Foil Experiment: directed α particles

at a thin sheet of gold foil. He thought that if Thomson’s model was correct then the α particles would go through the foil easily b/c they were so massive.

Results: # Most α particles did go straight through the foil, but many more than expected were deflected and some were reflected # Knew the plum pudding model was incorrect # particles that were deflected must have come close to a dense positive center

in the atom # The particles that were reflected had a “direct” hit with the positive center # Conceived the idea of a nuclear atom, one that had a dense positive core

Particle Mass (kg) Charge

e-‐ 9.11 × 10-‐31 1-‐

p+ 1.67 × 10-‐27 1+

n0 1.67 × 10-‐27 None

# The nucleus contains most of the mass of the atom, however very little volume (making it very dense) therefore most of the atom is empty space

$ James Chadwick (1891-‐1972): discovered neutrons

1.13 Modern View of Atomic Structure (Introduction) (pp. 43 – 46) # Elements -‐ All matter composed of only one type of atom is an element. There are 92 naturally

occurring, all others are manmade. # atom-‐-‐the smallest particle of an element that retains the chemical properties of that element.

(Entire atom has a diameter of about 100-‐500 pm or 1-‐ 5 Angstrom (Å) [ 1 Å = 10-‐10 m]) $ nucleus-‐-‐contains the protons and the neutrons; the electrons are located outside the

nucleus. Diameter = 10-‐13 cm. The electrons are located 10-‐8cm from the nucleus. A mass of nuclear material the size of a pea would weigh 250 million tons! Very dense!

$ proton-‐-‐positive charge, responsible for the identity of the element, defines atomic number

$ neutron-‐-‐no charge, same size & mass as a proton, responsible for isotopes, alters atomic mass number

$ electron-‐-‐negative charge, same size as a proton or neutron, BUT 1/2,000 the mass of a proton or neutron, responsible for bonding, hence reactions and ionizations, easily added or removed.

# atomic number(Z)-‐-‐The number of p+ in an atom. All atoms of the same element have the same number of p+. # mass number(A)-‐-‐The sum of the number of neutrons and p+ for an atom. A different mass number does not mean a

different element-‐-‐just an isotope. $ Atomic Mass Units (amu): used to measure mass of atoms b/c if measured in g, the masses would be really small –

too small to really easily use the numbers in calculations $ 1 amu = 1.66054x10-‐24 g

mass number → ←element symbol atomic number→

Exercise 1.15 Atomic Symbols Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and how many neutrons does this atom have? F, 9 e and 10 n

# Isotope – atoms having the same atomic number (# of p+) but a

different # of neutrons $ Most elements have at least 2 stable isotopes (Al, F, & P only

have 1) $ H isotopes have special names:

# 0 neutrons = hydrogen # 1 neutron = deuterium # 2 neutrons = tritium

1.14 Intro to the Periodic Table (pp. 49 – 52)

# Once many elements were known, scientists needed an organizational system # Periodic table grouped elements with similar properties # Grouped according to increasing atomic number # Rows = periods (don’t necessarily share properties) # Columns = groups or families (share similar chemical and physical properties due to electron configurations)

Current Name Original Name Symbol

Antimony Stibium Sb

Copper Cuprum Cu

Iron Ferrum Fe

Lead Plumbum Pb

Mercury Hydrargyrum Hg

Potassium Kalium K

Silver Argentum Ag

Sodium Natrium Na Tin Stannum Sn

Tungsten Wolfram W

$ Group 1 – Alkali Metals $ Group 2 – Alkaline Earth Metals $ Groups 3 – 12 – Transition Metals (bottom 2 rows are inner transition metals) $ Group 6 (oxygen group) – Chalcogens $ Group 7 – Halogens $ Group 8 – Noble Gases (also called rare gases or inert gases)

# Elements to the left of the staircase are metals (except H) $ metals—malleable, ductile & have luster; most of the

elements are metals—exist as cations in a “sea of electrons” which accounts for their excellent conductive properties; form oxides [tarnish] readily and form POSITIVE ions [cations]. Why must some have such goofy symbols?

# Elements to the right of the staircase are nonmetals # Elements touching the staircase are metalloids or semimetals

1.15 Molecules and Molecular Compounds (pp. 52 – 55) # Atoms are the smallest sample of an element, but most do not exist as single atoms by themselves – most of the time you

find compounds, not pure elements $ Chemical bond – force that hold atoms together $ Electrons are responsible for bonding and chemical reactivity

# Covalent bonds – atoms share electrons and make molecules # Molecule – 2 or more atoms bonded together (a package of atoms that act as 1 group) Ex: H2, CO2, H2O, NH3, O2, CH4

$ Smallest unit of a compound that retains the chemical properties of the compound (the characteristics of the constituent elements are lost)

$ Diatomic molecule – molecule made up of 2 atoms of the same element (H2, N2, O2, F2, Cl2, I2, Br2) # Chemical formula – tells what and how many atoms of each element are present in a compound # Molecular compounds are made of molecules (mostly nonmetal atoms)

# Molecular formulas – formula showing the actual number of atoms in a compound # Empirical formula – formula showing the ratio of atoms in a compound (only gives relative amounts of atoms, not actual);

the “reduced” formula $ Ex: Molecular Formula: C3H9 Empirical Formula: CH3

# Empirical formulas are mainly used when trying to identify unknown compounds # Structural formulas – formula showing how atoms are bonded; a ‘picture’ of a molecule showing what atoms are

connected, bonds are shown by lines [representing shared e-‐ pairs]; may NOT indicate shape

H O H O H H

1.16 Ions and Ionic Compounds (pp. 55 – 59) # ions-‐-‐formed when electrons are lost or gained in ordinary chem. reactions; affect size of atom dramatically

Generally metals lose electrons to form positive ions called cations Generally nonmetals gain electrons to form negative ions called anions

# Most atoms lose or gain electrons in order to get the same number of electrons as one of the noble gases; charge can be predicted by an atom’s position on the periodic table

$ cations-‐-‐(+) ions; often metals since metals lose electrons to become + charged $ anions-‐-‐(-‐) ions; often nonmetals since nonmetals gain electrons to become – charged $ polyatomic ions-‐-‐ contain atoms joined as in a molecule, only the group of atoms has a charge; properties of atoms

and the ions they form are usually very different

# ionic solids—Electrostatic forces hold ions together. Strong \ ions held close together \ solids. # Positive ions and negative ions are attracted to one another (b/c opposite charges), making an ionic

compound # Ionic compounds usually contain a metal ion and a nonmetal ion # Ions in ionic compounds form in a rigid lattice structure. It takes many ions to make a unit

of an ionic compound. So, we pretty much have to write empirical formulas for ionic compounds…we can only give the ratio of positive ions to negative ions since there are many options for how many ions can be present to make a unit of an ionic compound.

1.17 Naming Simple Compounds (pp. 60 – 69) # Organic compounds – compounds containing carbon (usually with H, N, O, S, or halogens) # Inorganic compounds – everything else! # Names and Formulas of Ionic Compounds

$ Naming Cations: # Cations formed from metal atoms have the same name as the metal Ex: Na+ sodium ion; Al3+ aluminum ion # If a metal can form a cation with more than 1 charge (mostly these are transition metals), the charge is shown with

a roman numeral (or the latin name of the ion) Ex: Fe2+ iron(II) or ferrous ion Fe3+ iron(III) or ferric ion Cu+ copper(I) or cuprous ion Cu2+ copper(II) or cupric ion

**for Latin names, the higher charge gets the “-‐ic” ending where the lower charge is “-‐ous” # Cations formed from nonmetal atoms (must be polyatomic ions) end in –ium Ex: NH4

+ ammonium ion

Common Cations 1+ 2+ 3+

Al3+ aluminum ion Cr3+ cobalt(III) or cobaltic ion Fe3+ iron(III) or ferric ion

H+ hydrogen ion Li+ lithium ion Na+ sodium ion K+ potassium ion Cs+ cesium ion Ag+ silver ion**

Mg2+ magnesium ion Ca2+ calcium ion Sr2+ strontium ion Ba2+ barium ion Zn2+ zinc ion** Cd2+ cadmium ion** 4+

Al3+ aluminum ion Cr3+ cobalt(III) or cobaltic ion Fe3+ iron(III) or ferric ion

4+

NH4+ ammonium ion

Cu+ copper(I) or cuprous ion Co2+ cobalt(II) or cobaltous ion Cu2+ copper(II) or cupric ion Fe2+ iron(II) or ferrous ion Mn2+ manganese(II) or manganous ion Hg2

2+ mercury(I) or mercurous ion Hg2+ mercury(II) or mercuric ion Ni2+ nickel(II) or nickelous ion Pb2+ lead(II) or plumbous ion** Sn2+ tin(II) or stannous ion**

Sn4+ tin(IV) or stannic ion** Pb4+ lead(IV) or plumbic ion**

$ Naming Anions: # Most anions are named by taking the name of the nonmetal atom they came from and changing the ending to

–ide Ex: Cl-‐ chloride ion O2-‐ oxide ion N3-‐ nitride ion

# Some polyatomic ions end in –ide as well (most end in –ate or –ite) Ex: OH-‐ hydroxide ion CN-‐ cyanide ion O2

2-‐ peroxide ion # Polyatomic ions containing oxygen (oxyanions):

• Most common ion ends in –ate • Ion with same charge but 1 less oxygen ends in –ite • Per-‐ -‐ate and hypo-‐ -‐ite are used when there are more than 2 polyatomic ions with the same

elements with the same charge • Very helpful to memorize patterns than to try and memorize each ion individually! • Anions with H in front are named by adding hydrogen (or bi-‐) to the front of the ion name. The H+

reduces the charge of the ion by 1

Common Anions 1-‐ 2-‐ 3-‐

H-‐ hydride ion F-‐ fluoride ion Cl -‐ chloride ion Br-‐ bromide ion I-‐ Iodide ion CN-‐ cyanide ion OH-‐ hydroxide ion C2H3O2

-‐ acetate ion NO3

-‐ nitrate ion NO2

-‐ nitrite ion MnO4

-‐ permanganate ion ClO-‐ hypochlorite ion ClO2

-‐ chlorite ion ClO3

-‐ chlorate ion

O2-‐ oxide ion S2-‐ sulfide ion CO3

2-‐ carbonate ion CrO4

2-‐ chromate ion Cr2O7

2-‐ dichromate ion C2O4

2-‐ oxalate ion SiO3

2-‐ silicate ion SO4

2-‐ sulfate ion SO3

2-‐ sulfite ion S2O3

2-‐ thiosulfate ion O2

2-‐ peroxide ion HPO4

2-‐ biphosphate (hydrogenphosphate)ion HPO3

2-‐ biphosphite (hydrogen phosphite) ion

N3-‐ nitride ion P3-‐ phosphide ion PO4

3-‐ phosphate ion PO3

3-‐ phosphite ion

ClO4-‐ perchlorate ion

BrO-‐ hypobromite ion BrO2

-‐ bromite ion BrO3

-‐ bromated ion BrO4

-‐ perbromate ion IO3

-‐ iodate ion HCO3

-‐ bicarbonate (hydrogen carbonate) ion HSO4

-‐ bisulfate (hydrogen sulfate) ion HSO3

-‐ bisulfite (hydrogen sulfite) ion H2PO4

-‐ dihydrogen phosphate ion H2PO3

-‐ dihydrogen phosphate ion

$ Naming Ionic Compounds Name the cation then name the anion! Ex: NaCl sodium chloride FeCl3 iron(III) chloride or ferric chloride NaClO3 sodium chlorate

$ Writing formulas for Ionic Compounds Write the formula for the cation. Write the formula for the anion. Cross over & Reduce. Ex: aluminum phosphide Al3+ P3-‐ " AlP

Exercise 1.16 Nomenclature Name each binary compound. a. CsF b. AlC13 c. LiH

cesium fluoride; aluminum chloride; lithium hydride

Exercise 1.17 Nomenclature Give the systematic name of each of the following compounds. a. CuCl b. HgO c. Fe2O3 d. MnO2 e. PbC12

copper(I) chloride; mercury(II) oxide; iron(III) oxide; manganese(IV) oxide; lead(II) chloride

TYPE II: Involve a transition metal that needs a roman numeral Mercury (I) is Hg2

+2 Exceptions: these never need a roman numeral even though transition metals. MEMORIZE Ag+, Cd2+, Zn2+

Exercise 1.18 Nomenclature Give the systematic name of each of the following compounds. a. CoBr2 b. CaCl2 c. Al2O3 d. CrCl3

cobalt(II) bromide; calcium chloride; aluminum oxide; chromium(III) oxide

Exercise 1.19 Nomenclature Give the systematic name of each of the following compounds. a. Na2SO4 b. KH2PO4 c. Fe(NO3)3 d. Mn(OH)2 e. Na2SO3 f. Na2CO3 g. NaHCO3 h. CsC1O4 i. NaOC1 j. Na2SeO4 k. KBrO3

sodium sulfate; potassium dihydrogen phosphate; iron(III) nitrate; manganese(II) hydroxide; sodium sulfite; sodium carbonate; sodium bicarbonate; cesium perchlorate; sodium hypochlorite; sodium selenate; potassium bromate # Names and Formulas of Acids

$ Acid – compound that produces H+ in solution (when dissolved in water) $ Usually has H as the first element in the formula! $ Acids are named based on the type of anion present:

# Monotomic anion (nonmetal ion) – named hydro___ic acid (the stem of the element name goes in the ___) # Polyatomic anion – named using convention –ate " -‐ic and –ite " -‐ous

Ex: HCl – hydrochloric acid HF – hydrofluoric acid HNO3 – nitric acid HNO2 – nitrous acid

H2SO4 – sulfuric acid H2SO3 – sulfurous acid

**If polyatomic ion ends in –ate, the acid will end in –ic. If the polyatomic ion ends in –ite, the acid will end in –ous.

o Writing Formulas for Acids H+ is always the cation Write the formula for the anion Cross over Ex: phosphoric acid H+ PO4

3-‐ " H3PO4

# Names and Formulas of Molecular Compounds $ Element further to the left on the periodic table is usually named first (except O – usually last except with F) $ If both elements are in the same group, the one with the larger atomic number is named first $ 2nd element ends in –ide $ Greek prefixes are used to indicate how many atoms of each element are present (except ‘mono’ – it’s only used on the

2nd element if needed) $ Prefixes: 1-‐ mono, 2 – di, 3 – tri, 4 – tetra, 5 – penta, 6 – hexa, 7 – hepta, 8 – octa, 9 – nona, 10 -‐ deca

Ex: Cl2O dichlorine monoxide CO carbon monoxide N2O4 dinitrogen tetroxide

Exercise 1.20 Nomenclature Name each of the following compounds. a. PC15 b. PC13 c. SF6 d. SO3 e. SO2 f. CO2

phosphorus pentachloride; phosphorus trichloride; sulfur hexafluoride; sulfur trioxide; sulfur dioxide; carbon dioxide Exercise 1.21 Nomenclature Give the systematic name for each of the following compounds. a. P4O10 b. Nb2O5 c. Li2O2 d. Ti(NO3)4 e. H3PO4 f. H3N tetraphosphorus decaoxide; niobium(V) oxide; lithium peroxide; titanium(IV) oxide; phosphoric acid; hydronitric acid

Exercise 1.22 Nomenclature Given the following systematic names, write the formula for each compound. a. Vanadium(V) fluoride b. Dioxygen difluoride c. Rubidium peroxide d. Gallium oxide VnF5; O2F2; Rb2O2; Ga2O3 # Naming Simple Organic Compounds

$ Hydrocarbon – compounds that contain only carbon and hydrogen $ Carbon can bond 4 times – no more, no less $ Alkanes -‐ compounds with only C-‐C single bonds

# All alkanes end in –ane # General formula: CnH2n+2

$ Alcohols – compounds with –OH bonded to a carbon atom # End in –ol

$ Prefix indicates how many carbons are present – only first 3 are mentioned here Meth = 1 carbon Eth = 2 carbons Prop = 3 carbons

Ex: methane (CH4) Methanol (CH3OH)

Exercise 1.23 Nomenclature Given the following systematic names, write the formula for each compound. a. ethene b. propyne C2H2; C3H4; C3H7OH; C2H5OH c. propanol d. ethanol

# Exceptions: these lovely creatures have been around longer than the naming system and no one wanted to adapt!! $ Water (H2O) $ Ammonia (NH3) $ Hydrazine (N2H4) $ Phosphine (PH3) $ nitric oxide (NO) $ nitrous oxide (“laughing gas”) (N2O)


1. What is the correct name for KClO3? a. Potassium Chloride b. Potassium Chlorate c. Potassium Chlorite d. Potassium Hypochlorite e. Potassium Perchlorate

2. What is the correct name for H2SO3?

a. Sulfuric acid b. Sulfurous acid c. Persulfic acid d. Hydrosulfuric acid e. Hyposulfuric acid

3. What is the correct name for HF?

a. Hydrofluoric acid b. Hydrogen fluoride c. Fluoric acid d. Fluorous acid e. Hydrogen (I) fluoride

4. What is the correct name for Cu2CO3?

a. Copper (II) carbonate b. Copper (I) carbonate c. Copper carbonate d. Copper carbon trioxide e. Copper (III) carbon oxide

5. What is the name of the following substance: P2Cl6 ? a. phosphorus (II) chloride b. diphosphorus hexachloride c. phosphorus (IV) chloride d. phosphorus chloride e. diphosphorus heptachloride

6. What is the correct name for OF2?

a. Oxygen difluoride b. Oxygen fluoride c. Monoxygen difluoride d. Oxyfluoric acid e. Oxygen (II) fluoride

7. The formula of ammonium nitrite is

a. (NH4)3N b. NH3N c. NH4NO2 d. NH3NO3 e. NH4NO3

8. Which of the following polyatomic ions has the same charge as the hydroxide ion?

a. Ammonium b. Permanganate c. Oxide d. Nitride e. Chromate

9. Which of the following is paired incorrectly? a. Ethane; C2H6 b. Butane; C4H8 c. Methanol; CH2O d. Ethyne; C2H2 e. Propenal; C3H5OH

10. Element M reacts with fluorine to form an ionic compound with the formula MF3. The M-ion has 18

electrons. What is the element M? a. P b. Sc c. Ar d. Ca e. Cr

11. Which pair of elements would you expect to exhibit the greatest similarity in their physical and chemical

properties? a. H and Li b. Cs and Ba c. Ga and Ge d. Ca and Sr e. C and O

12. When forming a positive ion, which of the following describes the change that an atom undergoes?

a. It loses protons b. It gains electrons c. It gains protons d. It loses electrons e. It loses neutrons

13. What is the charge on the most common ion of sulfur? a. +2 b. +1 c. -1 d. -2 e. -3

14. Which of the following species has 13 neutrons, 10 electrons and a +1 charge?

a.

b. c.

d.

e.

15. Consider the following selected postulates of Dalton’s atomic theory: I. Each element is composed of extremely small particles called atoms.

II. Atoms are indivisible. III. Atoms of a given element are identical. IV. Atoms of different elements are different and have different properties.

Which of the postulates is (are) no longer valid? a. I and II b. II only c. II and III d. III only e. III and IV

16. A chemical reaction is observed in a sealed container where new products are formed, a gas is released

and there is a change of color. There is no change in mass. Which part of Dalton’s theory does this illustrate?

a. Law of conservation of mass b. Law of constant composition c. The existence of isotopes d. The existence of the small particles called atoms e. The fact that all atoms of the same element are identical

17. Which of the following statements is correct?

a. An electron has approximately 1/2000 mass of a proton b. A proton has approximately 1/10 mass of a neutron c. Neutrons have no mass and no charge d. All atoms of a particular element are identical e. Different ratios of atoms produce the same compounds

18. Who is credited with the discovery of the neutron? a. Millikan b. Rutherford c. Chadwick d. Bohr e. Thompson

19. What do atoms of the isotopes 37Cl and 35Cl have in common?

a. They have same number of protons b. They have same number of neutrons c. They have same half-life d. They have same molar mass e. They have same diffusion constants

20. In Rutherford’s nuclear-atom model

a. The light subatomic particles, protons and neutrons reside in the nucleus b. Mass is spread essentially uniformly through the atom c. Protons are negatively charged and have much bigger mass than atom d. Nucleus is positive and essentially smallest part of the atom e. Electrons are negatively charged and take less space of the atom than the protons.


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Unit 1: Stoichiometry (Review) Adapted from Stoichiometry (NMSI, Rene McCormick) and Chemistry (Brown/LeMay)

1.18 Atomic Masses (pp. 46 – 48) # 12C—Carbon 12—In 1961 it was agreed that this would serve as the standard and would be defined to have a mass of

EXACTLY 12 atomic mass units (amu). All other atomic masses are measured relative to this. # mass spectrometer—a device for measuring the mass of atoms or molecules

$ atoms or molecules are passed into a beam of high-‐speed electrons $ this knocks electrons OFF the atoms or molecules

transforming them into cations $ apply an electric field $ this accelerates the cations since they are repelled

from the (+) pole and attracted toward the (−) pole $ send the accelerated cations into a magnetic field $ an accelerated cation creates it’s OWN magnetic field

which perturbs the original magnetic field $ this perturbation changes the path of the cation $ the amount of deflection is proportional to the mass;

heavy cations deflect little $ ions hit a detector plate where measurements can be

obtained.

$

Exact by definition # average atomic masses—atoms have masses of whole numbers, HOWEVER samples of quadrillions of atoms have a few

that are heavier or lighter [isotopes] due to different numbers of neutrons present # percent abundance-‐-‐percentage of atoms in a natural sample of the pure element represented by a particular isotope # percent abundance = number of atoms of a given isotope × 100

Total number of atoms of all isotopes of that element # counting by mass—when particles are small this is a matter of convenience. Just as you buy 5 lbs of sugar rather than a

number of sugar crystals, or a pound of peanuts rather than counting the individual peanuts….this concept works very well if your know an average mass.

# mass spectrometer to determine isotopic composition—load in a pure sample of natural neon or other substance. The

areas of the “peaks” or heights of the bars indicate the relative abundances of , , and

Exercise 1.24 Average Atomic Mass When a sample of natural copper is vaporized and injected into a mass spectrometer, the results shown in the figure are obtained. Use these data to compute the average mass of natural copper. (The mass values for 63Cu and 65Cu are 62.93 amu and 64.93 amu, respectively.) 63.55 amu

1.19 The Mole (pp. 90 – 96) # mole—the number of C atoms in exactly 12.0 grams of 12C; also a number, 6.02 × 1023 just as the word “dozen” means 12 and

“couple” means 2. # Avogadro’s number—6.02 × 1023, the number of particles in a mole of anything # Calculated to help chemists estimate very large numbers of very small particles

Exercise 1.25 Mole Conversions – Mass/Particles Americium is an element that does not occur naturally. It can be made in very small amounts in a device known as a particle accelerator. Compute the mass in grams of a sample of americium containing six atoms.

2x10-21 g Exercise 1.26 Mole Conversions – Mass/Particles Aluminum (A1) is a metal with a high strength-‐to-‐mass ratio and a high resistance to corrosion; thus it is often used for structural purposes. Compute both the number of moles of atoms and the number of atoms in a 10.0-‐g sample of aluminum.

2.23x1023 atoms Exercise 1.27 Mole Conversions – Moles/Mass Cobalt (Co) is a metal that is added to steel to improve its resistance to corrosion. Calculate both the number of moles in a sample of cobalt containing 5.00 × 1020 atoms and the mass of the sample.

8.31x10-4 mol; 0.0489 g Co 1.20 Molar mass, Molecular weight, and Formula weight (pp. 88 – 89, 92 -‐ 93)

# molar mass, MM-‐-‐the mass in grams of Avogadro’s number of molecules; i.e. the mass of a mole! # molecular weight, MW-‐-‐sum of all the atomic weights of all the atoms in the formula (must have a correct formula!) # empirical formula-‐-‐that ratio in the network for an ionic substance. # formula weight-‐-‐same as molecular weight, just a language problem [“molecular” implies covalent bonding while A

formula implies ionic bonding {just consider this to be a giant conspiracy designed to keep the uneducated from ever

Moles

Mass (g)

# of Particles

Molarity (mol/L)

Volume (L)

Molar Mass g/mol

6.02x1023 part = 1 mol Mol/L

22.4 L = 1 mol (gas at STP)

understanding chemistry—kind of like the scoring scheme in tennis}. We’ll use MM for all formula masses. # Units: grams per mole (g/mol) # The mass of 1 atom in amu = the mass of 1 mole of atoms of that element in g

Ex: 1 formula unit of NaCl has a mass of 58.5 amu but 1 mole of NaCl (6.02x1023 units of NaCl) has a mass of 58.5 g # Significant Figures: It is correct to use as many significant figures for the molar mass as you are given in your problem.

However, most of the time you’ll be fine if you round the molar masses to 2 places after the decimal.

Exercise 1.28 Calculating Molar Mass Juglone, a dye known for centuries, is produced from the husks of black walnuts. It is also a natural herbicide (weed killer) that kills off competitive plants around the black walnut tree but does not affect grass and other noncompetitive plants [a concept called allelopathy]. The formula for juglone is C10H6O3. a. Calculate the molar mass of juglone.

174.16 g/mol b. A sample of 1.56 x 10-‐2 g of pure juglone was extracted from black walnut husks. How many moles of juglone does this sample represent?

8.96x10-5 mol

Exercise 1.29 Calculating Molar Mass Calcium carbonate (CaCO3), also called calcite, is the principal mineral found in limestone, marble, chalk, pearls, and the shells of marine animals such as clams. a. Calculate the molar mass of calcium carbonate.

100.09 g/mol b. A certain sample of calcium carbonate contains 4.86 moles. What is the mass in grams of this sample? What is the mass of the CO3

2-‐ ions present?

486 g; 292 g Exercise 1.30 Molar Mass and Molecules Isopentyl acetate (C7H14O2), the compound responsible for the scent of bananas, can be produced commercially. Interestingly, bees release about 1µg (1 × 10-‐6 g) of this compound when they sting. The resulting scent attracts other bees to join the attack. How many molecules of isopentyl acetate are released in a typical bee sting?

5x1015 molecules How many atoms of carbon are present?

3x1016 atoms

# ELEMENTS THAT EXIST AS MOLECULES Pure hydrogen, nitrogen, oxygen and the halogens exist as DIATOMIC molecules under normal conditions. MEMORIZE!!! Be sure you compute their molar masses as diatomics. Others to be aware of, but not memorize:

• P4-‐-‐tetratomic form of elemental phosphorous; an allotrope • S8—sulfur’s elemental form; also an allotrope

• Carbon-‐-‐diamond and graphite "covalent networks of atoms 1.21 Percent Composition (pp. 89 – 90)

# Percent Composition – the percent by mass of each element in the compound # Two common ways of describing the composition of a compound: in terms of the number of its constituent atoms and

in terms of the percentages (by mass) of its elements.

Percent (by mass) Composition: law of constant composition states that any sample of a pure compound always consists of the same elements combined in the same proportions by mass.

% comp = mass of desired element × 100 Total mass of compound

Consider ethanol, C2H5OH

Mass % of C = 2 mol × = 24.02 g

Mass % of H = 6 mol × = 6.06 g

Mass % of O = 1 mol × = 16.00g

Mass of 1 mol of C2H5OH = 46.08 g NEXT THE MASS PERCENT CAN BE CALCULATED: Mass percent of C = 24.02 g C × 100% = 52.14% 46.08 g Repeat for the H and O present.

Exercise 1.31 Calculating Mass Percent Carvone is a substance that occurs in two forms having different arrangements of the atoms but the same molecular formula (C10H14O) and mass. One type of carvone gives caraway seeds their characteristic smell, and the other type is responsible for the smell of spearmint oil. Compute the mass percent of each element in carvone.

79.94% C; 9.41% N; 10.65% O Exercise 1.32 Calculating Mass Percent Penicillin, the first of a now large number of antibiotics (antibacterial agents), was discovered accidentally by the Scottish bacteriologist Alexander Fleming in 1928, but he was never able to isolate it as a pure compound. This and similar antibiotics have saved millions of lives that might have been lost to infections. Penicillin F has the formula C14H20N2SO4. Compute the mass percent of each element.

53.82% C; 6.47% H; 8.97% N; 10.26% S; 20.48% O

1.22 Determining the Formula of a Compound (pp. 96 – 100)

# Empirical formula – gives the relative number of atoms of each element present in a compound $ hydrates—“dot waters” used to cement crystal structures. $ anhydrous-‐-‐without water

# To calculate E.F. from % composition data: 1. Assume 100 g of the sample (so the % = mass) (if given grams go straight to step 2) 2. Convert grams of each element in the sample to moles 3. Calculate the mole ratio (divide all moles by the lowest number of moles) 4. Ratio = subscripts in the E.F.

**If in #4 you don’t get whole numbers, multiply all by a number to get a whole number **E.F. must have the lowest whole number subscripts possible (lowest ratio possible)

# Calculating molecular formula from empirical formula $ The subscripts in the M.F. are always a whole number multiple of the subscripts in the E.F. [(empirical formula)n, where

n is an integer] $ To find M.F., you have to know the molar mass of the molecular formula

Whole Number Multiple = molar mass/empirical mass # Combustion Analysis

$ Way of analyzing unknown compounds to find composition $ When faced with a compound of “unknown” formula, one of the most common techniques is to combust it with

oxygen to produce CO2, H2O, and N2 which are then collected and weighed. $ When an organic compound (a hydrocarbon) undergoes combustion, all of the C from the compound is converted into

CO2 and all of the hydrogen is converted into H2O. You can find out how much of each was in the original sample using stoichiometry.

Example: A compound is composed of carbon, nitrogen and hydrogen. When 0.1156 g of this compound is reacted with oxygen [burned, combusted], 0.1638 g or carbon dioxide and 0.1676 g of water are collected. What is the empirical formula of the compound? Compound + O2 " CO2 + H2O + N2 but NOT balanced!!

(You can see that all of the carbon ended up in CO2 so…when in doubt, FIND THE NUMBER OF MOLES!!)

0.1638 g CO2 ÷ 44.01 g/mol = 0.003781 moles of CO2 x 1 mol C/1 mol CO2 = 0.003781 moles of C

(Next, you can see that all of the hydrogen ended up in H2O, so….FIND THE NUMBER OF MOLES!!)

0.1676 g H2O ÷ 18.02 g/mol = 0 .009301 moles of H2O (BUT there are 2 moles of H for each mole of water [ think “organ bank” one heart per body, one C per molecule of carbon dioxide—2 lungs per body, 2 atoms H in water and so on…] so DOUBLE THE NUMBER OF MOLES TO GET THE NUMBER OF MOLES OF HYDROGEN!!)

0.009301 mol H2O = 0.01860 moles of H The rest must be nitrogen, BUT we only have mass data for the sample so convert your moles of C and H to grams:

g C = 0.003781 moles C × 12.01 = 0.04540 grams C +

g H = 0.01860 moles H × 1.01 = 0.01879 grams H 0.06419 grams C + H

SUBTRACT! 0.1156 g sample – 0.06419 g thus far = grams N left = 0.05141 g N so….

0.05141 g N ÷ 14.01 = 0.003670 moles N

Chemical formulas represent mole to mole ratios, so…divide the number of moles of each by the smallest # of moles of any one of them to get a guaranteed ONE in your ratios…multiply by 2, then 3, etc to get to a ratio of small whole numbers!!

Element # moles ALL Divided by 0.003670 C 0.003781 1 H 0.01860 5 N 0.003670 1

Therefore the correct EMPIRICAL formula is CH5N. Next, if we are told that the MM is 31.06 g/mol, then simply use this relationship: (Empirical mass) × n = MM (12.01 + 5.05 + 14.01) × n = 31.06

Solve for n

n = 0.999678… or essentially one, so the empirical formula and the molecular formula are the same. Exercise 1.33 Determining Empirical & Molecular Formulas Determine the empirical and molecular formulas for a compound that gives the following analysis (in mass percents):

71.65% C1 24.27% C 4.07% H The molar mass is known to be 98.96 g/mol.

CH2Cl; C2H4Cl2 Exercise 1.34 Determining Empirical & Molecular Formulas A white powder is analyzed and found to contain 43.64% phosphorus and 56.36% oxygen by mass. The compound has a molar mass of 283.88 g/mol. What are the compound’s empirical and molecular formulas?

P2O5; P4O10 Exercise 1.35 Determining a Molecular Formula Caffeine, a stimulant found in coffee, tea, and chocolate, contains 49.48% carbon, 5.15% hydrogen, 28.87% nitrogen, and 16.49% oxygen by mass and has a molar mass of 194.2 g/mol. Determine the molecular formula of caffeine.

C8H10N4O2 1.23 Chemical Reactions and Equations (pp. 80 – 88)

# Chemical reactions are the result of a chemical change where atoms are reorganized into one or more new arrangements. Bonds are broken [requires energy] and new ones are formed [releases energy].

# Chemical equation – a representation of a chemical reaction $ Reactants are listed on the left side of the arrow, product on the right $ Coefficients represent the relative number of molecules and/or moles required for the reaction $ Energy is included SOMETIMES (called a thermochemical equation) to show that the rxn is either endothermic or

exothermic (this isn’t info you have to figure out, you’ll be told if you’re supposed to include it in the equation) $ The time required for a rxn to occur is not included

# Law of Conservation of Mass (Lavoisier) – atoms (matter) can’t be created or destroyed, so the number of atoms on the left of the arrow has to be the same as the atoms on the right

$ To satisfy the law of conservation of mass, balance the equation # Write correct formulas for all compounds in the reaction # Never change a subscript in an equation # Use coefficients to make the number of atoms equal on each side of the arrow

$ States of matter for each compound in the reaction are represented by: # (s) = solid # (l) = liquid # (g) = gas # (aq) = aqueous

# Information given by a chemical equation

# Writing & Balancing Equations

$ Begin with the most complicated-‐looking thing (save the elemental thing for last). $ If you get stuck, double the most complicated-‐looking thing. $ MEMORIZE THE FOLLOWING:

# metals + halogens " MaXb # CH (and/or O) + O2 " CO2(g) + H2O(g) # H2CO3 [any time formed!] " CO2 + H2O; in other words, never write carbonic acid as a product, it

spontaneously decomposes [in an open container] to become carbon dioxide and water. # metal carbonates " metal OXIDES + CO2

# Types of Chemical Reactions (general) $ Combination (Synthesis) Reactions (A + B " AB)

# Reactants can be single elements or compounds or both! # Product must be one compound

$ Decomposition Reactions (AB " A + B) # Generally require energy (endothermic) # Reactant must be one compound # Products can be single elements or compounds, or both!

$ Combustion in Air # Rapid reactions in which a flame is produced # Most use O2 in air as a reactant # Organic Combustion

$ CxHy + O2 " CO2 + H2O (complete combustion) $ Sometimes incomplete combustion occurs…CxHy + O2 " CO + H2O (it depends on how much

oxygen is there as to whether combustion is complete or incomplete) $ Unless you are told otherwise, assume complete combustion! $ The water product can be liquid or gaseous depending on the reaction conditions

Exercise 1.36 Balancing Equations Chromium compounds exhibit a variety of bright colors. When solid ammonium dichromate, (NH4)2Cr2O7, a vivid orange compound, is ignited, a spectacular reaction occurs, as shown in the two photographs on page 105. Although the reaction is actually somewhat more complex, let’s assume here that the products are solid chromium(III) oxide, nitrogen gas (consisting of N2 molecules), and water vapor. Balance the equation for this reaction.

(NH4)2Cr2O7 " Cr2O3 + N2 + 4H2O

Exercise 1.37 Balancing Equations At 1000ºC, ammonia gas, NH3(g), reacts with oxygen gas to form gaseous nitric oxide, NO(g), and water vapor. This reaction is the first step in the commercial production of nitric acid by the Ostwald process. Balance the equation for this reaction.

4NH3 + 5O2 " 4NO + 6H2O 1.24 Stoichiometric Calculations: Amounts of Reactants and Products (pp. 100 – 104)

# Coefficients in a balanced equation represent the relative numbers of molecules (or moles) in a reaction # Mole ratio allows me to convert between compounds present in a reaction # SOOOOOO important to be good at this…it’s not going away! If you see a reaction, think stoich! # Can solve using dimensional analysis or a new way – either way is fine, whatever is easiest for you!

$ you have to be proficient at the following no matter which method you choose!: # Writing CORRECT formulas—this requires knowledge of your polyatomic ions and being able to use

the periodic table to deduce what you have not had to memorize. Review section 2.8 in your Chapter 2 notes or your text.

# Calculate CORRECT molar masses from a correctly written formula # Balance a chemical equation # Use the mole map to calculate the number of moles or anything else!

Example: What mass of oxygen will react with 96.1 grams of propane?

Option 1: Solve using dimensional analysis: 1. Write the balanced equation: C3H8 + 5O2 " 3CO2 + 4H2O 2. Use D.A. to go from grams of propane to grams of oxygen:

96.1 𝑔 𝐶3𝐻8 ×1 𝑚𝑜𝑙 𝐶3𝐻844.11 𝑔 𝐶3𝐻8×5 𝑚𝑜𝑙 𝑂21 𝑚𝑜𝑙 𝐶3𝐻8×32 𝑔 𝑂21 𝑚𝑜𝑙 𝑂2=349 𝑔 𝑂2 Option 2: Solve using the table method: 1. Make a table

Molar Mass:

Balanced Eq’n

mole:mole # moles amount

2. Write a chemical equation paying special attention to writing correct chemical formulas!

Molar Mass:

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O


3. Calculate the molar masses and put in parentheses above the formulas—soon you’ll figure out you don’t have to do this for

every reactant and product, just those you’re interested in.

Molar (44.11) (32.00) (44.01) (18.02)

Mass: Balanced

Eq’n C3H8

+

5 O2 "

3 CO2 +

4 H2O


4. Look at the coefficients on the balanced equation, they ARE the mole:mole ratios!

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles amount

5. Next, re-‐read the problem and put in an amount—in this example it’s 96.1 g of propane.

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles amount 96.1 g

6. Find the number of moles of something…anything! Use the mole map – start at 96.1 g, divide by molar mass to get the # of

moles of propane.

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles 2.18 amount 96.1 g

7. Use the mole:mole ratio to find moles of EVERYTHING! If 1 = 2.18, then oxygen is 5(2.18) etc…(if the first you find is not a “1” just devide to make it “1”) Leave all digits in your calculator, I only rounded to save space!

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced C3H8 5 O2 3 CO2 4 H2O

Eq’n + " + mole:mole 1 5 3 4

# moles 2.18 10.9 6.53 8.71 amount 96.1 g

8. Re-‐read the problem to determine which amount was asked for…here’s the payoff…AP problems ask for several amounts!

First we’ll find mass of oxygen required since that’s the problem asked. 10.9 mol x 32 g/mol = 349 g O2

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles 2.18 10.9 6.53 8.71 amount 96.1 g 349 g

9. What if another part of the question asked for liters of CO2 at STP (1 atm, 273 K)? Use the mole map. Start in the middle

with 6.53 moles x 22.4 L/mol = 146 L

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles 2.18 10.9 6.53 8.71 amount 96.1 g 349 g 146 L

10. What if another part asked how many water molecules are produced? Use the mole map. Start in the middle with 8.71

moles x 6.02x1023 molecules/1 mol = 5.24x1024 molecules water

Molar Mass:

(44.11) (32.00) (44.01) (18.02)

Balanced Eq’n

C3H8

+

5 O2 "

3 CO2 +

4 H2O

mole:mole 1 5 3 4 # moles 2.18 10.9 6.53 8.71 amount 96.1 g 349 g 146 L 5.24x1024

molec. Either way – table method or D.A – whatever is easier and makes more sense to you!

Exercise 1.38 Stoichiometry Solid lithium hydroxide is used in space vehicles to remove exhaled carbon dioxide from the living environment by forming solid lithium carbonate and liquid water. What mass of gaseous carbon dioxide can be absorbed by 1.00 kg of lithium hydroxide? 919 g

Exercise 1.39 Stoichiometry Baking soda (NaHCO3) is often used as an antacid. It neutralizes excess hydrochloric acid secreted by the stomach:

NaHCO3(s) + HCl(aq) → NaCl(aq) + H2O(l) + CO2(aq)

Milk of magnesia, which is an aqueous suspension of magnesium hydroxide, is also used as an antacid: Mg(OH)2(s) + 2HCl(aq) → 2H2O(l) + MgCl2(aq)

Which is the more effective antacid per gram, NaHCO3 or Mg(OH)2 ?

Mg(OH)2

1.25 Calculations involving a Limiting Reactant (pp. 104 – 108) # How to Recognize a L.R. problem: you’ll be given 2 amounts for the reactants in the problem! One of the reactants will be

the limiting reactant and the other will be in excess $ Limiting reactant-‐ reactant that is completely used up in a reaction (also called the limiting reagent) $ Excess reactant – reactant that is not used up in the reaction (you’ll have some left over when the reaction is done)

Example: Suppose 25.0 kg of nitrogen reacts with 5.00 kg of hydrogen to form ammonia. What mass of ammonia can be produced? Which reactant is the limiting reactant? What is the mass of the reactant that is in excess?

# To Solve: (use D.A. or the table method – I’ll show table method since that’s new!) 1. Set up your table like before only now you’ll have 2 amounts to start with:

Molar Mass: (28.04) (2.02) (17.04) Balanced Eq’n N2 + 3 H2 " 2 NH3

mole:mole 1 3 2 # moles amount 25,000 g 5,000 g

2. Find the # of moles of both reactants you’re given info about

Molar Mass:

(28.02) (2.02) (17.04)

Balanced Eq’n

N2 + 3 H2 "

2 NH3

mole:mole 1 3 2 # moles 892 moles 2,475 moles amount 25,000 g 5,000 g

3. To find which reactant is limiting, pick one – either N2 or H2 – it doesn’t matter. Let’s pick H2. Calculate how much N2 is used if all of the H2 is used up.

WHAT IF I used up all the moles of hydrogen? I’d need 1/3 × 2,475 moles = 825 moles of nitrogen. Clearly I have EXCESS moles of nitrogen!! Therefore, hydrogen limits me.

OR WHAT IF I used up all the moles of nitrogen? I’d need 3 × 892 moles = 2,676 moles of hydrogen. Clearly I don’t have enough hydrogen, so it limits me!! Therefore nitrogen is in excess.

Either way, I’ve established that hydrogen is the limiting reactant so I modify the table:

Molar Mass: (28.02) (2.02) (17.04) Balanced Eq’n N2 + 3 H2

" 2 NH3

mole:mole 1 3 2 # moles 825 mol used

892 moles

2,475 moles

1650 mol produced

amount 825 mol (28.02) =

23,116 g used

25,000 g

5,000 g

1650 mol (17.04)

= 28,116 g produced

1,884 g excess!! Here’s the question again, let’s clean up any sig.fig issues: Suppose 25.0 kg of nitrogen reacts with 5.00 kg of hydrogen to form ammonia. (3 sig. fig. limit)

What mass of ammonia can be produced? 23,100 g produced = 23.1 kg (always polite to respond in the unit given. Which reactant is the limiting reactant? hydrogen—once that’s established, N2 doesn’t matter anymore!

What is the mass of the reactant that is in excess? 1,884 g = 1.88 kg excess nitrogen!!

Exercise 1.40 Stoichiometry: Limiting Reactant Nitrogen gas can be prepared by passing gaseous ammonia over solid copper(II) oxide at high temperatures. The other products of the reaction are solid copper and water vapor. If a sample containing 18.1 g of NH3 is reacted with 90.4 g of CuO, which is the limiting reactant? How many grams of N2 will be formed?

10.7 g

# Theoretical yield – how much product you’ll make when all of the limiting reactant has been used up (this assumes perfect conditions and gives a maximum amount – not likely!)

# Actual yield – amount of product actually made 𝐴𝑐𝑡𝑢𝑎𝑙 𝑌𝑖𝑒𝑙𝑑/𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑×100=% 𝑌𝑖𝑒𝑙𝑑

Exercise 1.41 Calculating Percent Yield Methanol (CH3OH), also called methyl alcohol, is the simplest alcohol. It is used as a fuel in race cars and is a potential replacement for gasoline. Methanol can be manufactured by combination of gaseous carbon monoxide and hydrogen. Suppose 68.5 kg CO(g) is reacted with 8.60 kg H2(g). Calculate the theoretical yield of methanol. If 3.57 × 104 g CH3OH is actually produced, what is the percent yield of methanol?

52.3 %


$ The atomic masses in periodic table are not integral numbers. For example, carbon is listed as 12.01115 instead of 12.0000. Why? f. Our technology does not allow for exact measurement of such a small quantity. g. Atoms gain or lose electrons easily and that changes their mass significantly. h. Atomic masses listed in the periodic table are weighted averages of naturally occurring isotopes. i. Atomic masses are measured in real samples that are always contaminated with other elements. j. There is a theoretical uncertainty in the masses of atoms.

21. What is the average atomic mass of a sample of naturally occurring X, if it is comprised of the following

isotopes: 47% 51X, 36% 52X, and 17% 56X? a. 56.00 b. 51.00 c. 53.00 d. 54.34 e. 52.21

22. What is the mass of one atom of sulfur?

a. 6.02 x 1023 g b. 32.0 g c. 5.32 x 10-23 g d. 1.93 x 1025 g e. 6.02 x 10-23 g

23. What is molar mass of urea (NH2)2CO, a compound used as a nitrogen fertilizer?

a. 44.0 b. 43.0 c. 60.1 d. 8.0 e. 32.0

24. Two moles of a particular group I bromide have mass of 206 g. Identify the group I metal.

a. Li b. Na c. K d. Rb e. Cs

25. How many individual ions does one mole of barium phosphate contain?

a. 6.02 x 10 23 b. 3.01 x 10 24 c. 5 d. 2 e. 1.2 x 10 24

26. What is the % by mass of Beryllium chloride? a. 50% Be, 50% Cl b. 33% Be, 67% Cl c. 25% Be, 75% Cl d. 11.25% Be, 88.75% Cl e. 10.1% Be, 89.9% Cl

27. Which hydrocarbon pair below have identical mass percentage of C?

a. C3H4 and C3H6 b. C2H4 and C3H4 c. C2H4 and C4H2 d. C2H4 and C3H6 e. C3H9 and C6H6

28. A 16.0 g sample of a hydrocarbon undergoes combustion to produce 36.0 g of water. What is the

percentage of hydrogen in the hydrocarbon? a. 2.11% b. 9.23% c. 25.0% d. 54.9% e. 73.1%

29. An empirical formula always indicates

a. Which atoms are attached to which in a molecule b. How many of each atom are in the molecule c. The simplest whole-number ratio of different atoms in a compound d. The isotopes of each atom in a compound e. The geometry of a molecule

30. The empirical formula of a compound with molecules containing 12 carbon atoms, 14 hydrogen atoms,

and 6 oxygen atoms is a. C12H14O6 b. CHO c. CH2O d. C6H7O3 e. C2H4O

31. A sulfur oxide is 50% by mass sulfur. What is its empirical formula?

a. SO b. SO2 c. S2O d. S2O4 e. SO3

32. What is the molecular formula of a compound that is 5.88% hydrogen, the remainder being oxygen, and that has a molar mass of 34 g/mol?

a. H5O b. H5O2 c. H2O2 d. H2O e. HO

33. When the chemical reaction is written and balanced with the lowest possible integers, what coefficient

appears in front of the methane? Reaction: Methane combusts in oxygen. a. 1 b. 2 c. 3 d. 4 e. 6

34. When following chemical reaction is balanced using lowest possible integers, what is the total sum of

the coefficients? H3PO4 + Ba(OH)2 ! Ba3(PO4)2 + H2O a. 2 b. 3 c. 4 d. 12 e. 15

35. Consider reaction: 4FeS2 + 11O2 !2Fe2O3 + 8SO2. How many moles of FeS2 will react with 6.00 mol of

oxygen? a. 6 b. 11 c. 2.18 d. 2.75 e. .545

36. Consider reaction: CaC2 + 2H2O! Ca(OH)2 + C2H2. If 16.0 g of calcium crabide is consumed in this

reaction how many liters of acetylene has been produced under STP? a. 5.58 b. 11.6 c. 22.4 d. 44.8 e. 34.0

37. Solid aluminum and gaseous oxygen react in a combination reaction to produce aluminum oxide. What

is the maximum amount of aluminum oxide produced in reaction of 2.7 g of Al and 2.7 g of O2? a. 0.023 mol b. 0.050 mol c. 0.056 mol d. 0.100 mol

e. 0.125 mol

38. If 0.00250 mol of copper(II) carbonate are reacted with .00400 mol of HCl, to produce 0.06600 g of carbon dioxide, 0.02700 g of water and some other mass of copper(II) chloride, which is the limiting reactant?

a. Hydrochloric acid b. Copper(II) chloride c. Copper(II) carbonate d. Water e. Carbon dioxide

39. In the reaction: 2C4H10 + 13O2 ! 8CO2 + 10 H2O, 10.00 mol of carbon dioxide are formed from 10.0

mol of hydrocarbon. What is percentage yield for this reaction? a. 10% b. 25% c. 50% d. 80% e. 90%


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

AP Chemistry Summer Assignment Stoichiometry Free Response Questions CLEARLY SHOW THE METHOD YOU USED AND STEPS INVOLVED IN ARRIVING AT YOUR ANSWERS. It is to your advantage to do this, because you may earn partial credit if you do and you will receive little or no credit if you do not. Attention should be paid to significant figures. Be sure to write all your answers to the questions on your own paper, labeling each question part (ex: 1a.) 1. Answer the following questions that relate to chemical reactions.

(a) Iron(III) oxide can be reduced with carbon monoxide according to the following equation.

Fe2O3(s) + 3 CO(g) → 2 Fe(s) + 3 CO

2(g)

A 16.2 L sample of CO(g) at 1.50 atm and 200.°C is combined with 15.39 g of Fe2O3(s).

(i) How many moles of CO(g) are available for the reaction? (ii) What is the limiting reactant for the reaction? Justify your answer with calculations. (iii) How many moles of Fe(s) are formed in the reaction?

(b) In a reaction vessel, 0.600 mol of Ba(NO

3)2(s) and 0.300 mol of H

3PO

4(aq) are combined with distilled

water to a final volume of 2.00 L. The reaction represented below occurs.

3 Ba(NO3)2(aq) + 2 H

3PO

4(aq) → Ba

3(PO

4)2(s) + 6 HNO

3(aq)

(i) Calculate the mass of Ba

3(PO

4)2(s) formed.

(ii) Calculate the pH of the resulting solution. (iii) What is the concentration, in mol L

–1, of the nitrate ion, NO

3

– (aq), after the reaction reaches

completion? 2. The molecular formula of a hydrocarbon is to be determined by analyzing its combustion products and

investigating its colligative properties. (a) The hydrocarbon burns completely, producing 7.2 grams of water and 7.2 liters of CO2 at standard

conditions. What is the empirical formula of the hydrocarbon?

(b) Calculate the mass in grams of O2 required for the complete combustion of the sample of the hydrocarbon described in (a).

(c) The hydrocarbon dissolves readily in CHCl3. The freezing point of a solution prepared by mixing

100. grams of CHCl3 and 0.600 gram of the hydrocarbon is −64.0oC. The molal freezing-‐point depression constant of CHCl3 is 4.68°C/molal and its normal freezing point is −63.5°C. Calculate the molecular weight of the hydrocarbon.

(d) What is the molecular formula of the hydrocarbon?

2015 Summer Assignment...Tips for Learning the Ions Polyatomic Anions Most of the work on...

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