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Unit 1

Lesson Description Homework

1.1 Brief Course Introduction, Syllabus, Measurement Activity

1.2 Lab Safety Rules, Safety Worksheet, Equipment Identification 1.2

1.3 Accuracy vs. Precision, Intensive vs. Extensive Properties Error Types,

Desktop Activity: Mini Lab Accuracy vs. Precision

1.3

1.4 Exponential Numbers/Scientific Notation , Powers of 10 video, SI units,

Dimensional Analysis

1.4

1.5 Significant Figures 1.5

1.6 & 1.7 Classification of Matter, Phases of State, and Physical vs. Chemical

Changes, and Begin Foul Water Lab

1.6 & 1.7

1.8 Finish Foul Water Lab

1.9 Introduction to the mole and mole conversion 1.9

1.10 Introduction to Stoichiometry and Percent Yield 1.10

1.11 Stoichiometry Lab Review for Exam

1.12 Exam 1 Review

Ions to memorize for this unit:

NH4+ Ammonium

NO2- Nitrite

NO3- Nitrate

ClO2- Chlorite

ClO3- Chlorate

ClO4- Perchlorate

IO3- Iodate

Learning Objectives:

Lab Safety

SI units

Differentiate between Accuracy vs. Precision

Differentiate between Extensive vs. Intensive

Properties

Differentiate between determinate and

indeterminate error

Scientific Notation

Dimensional Analysis

Significant Figures

Mole

Stoichiometry

Unit 1 Vocabulary

scientific notation

SI units

significant figures

accuracy vs. precision

intensive property

extensive property

determinate error

indeterminate error

physical change

chemical change

molecular weight

element

compound

chemical symbol

ions

isotopes

atomic mass unit

atomic number (Z)

mass number (A)

Mole

Stoichiometry

Molar mass

Avogadro’s number

Stoichiometry

Molecule

Subscript

Coefficient

Chemical reaction

Yield

Percent yield

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1.1 Measurement Activity

Materials: shoe, meter stick, calculator

Procedure:

1. Break up into lab groups

2. Estimate the length of the classroom in: Estimated mm _____________________ Estimated cm _____________________ Estimated m ______________________ Estimated km _____________________

3. Pick a shoe from one of your group members. Use the length of that shoe to estimate the length of the room in shoes

Estimated shoe lengths __________________

4. Measure the length of the room in shoe lengths Actual shoe lengths __________________

5. Use the meter stick and the actual shoe lengths to come up with a way to determine the actual length of the room for the following:

Actual mm _____________________ Actual cm _____________________ Actual m ______________________ Actual km _____________________

Questions to discuss as a class:

A. How did you convert between shoe lengths and other units? B. What is the actual length of the room? Did all groups come up with the same length? Why or why not? C. Which group’s estimate is the more accurate? Why? D. Could an object other than a shoe be used to do this? E. If the standard distance around a running track is 440 yards, how many shoes lengths is this? How many

miles is one lap? How many laps would you need to do to run a mile? How about 1 km? Note:

1 yard = 0.9144 meters

1 yard = 3 feet

1 mile = 5280 feet

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1.2 Safety in the Science Laboratory

In the laboratory, you will be working with equipment and materials that can cause injury if they are not handled

properly. However, the laboratory can be a safe place to work if you are careful. Accidents do not just happen;

they are caused – by carelessness, haste, and disregard of safety rules and practices.

Safety rules to be followed in the laboratory are listed below. Before beginning any lab work, read these rules,

learn them, and follow them carefully. If you have any questions about these rules, ask your teacher before stating

lab work.

General Precautions:

1. Be prepared to work when you arrive in the laboratory. Familiarize yourself with the lab procedures before beginning the lab.

2. Carefully follow all oral and written instructions. Never do any unauthorized experiments. 3. Notify your teacher of any medical problems, like allergies or asthma, before beginning the lab. 4. Never eat or drink during lab. 5. Keep your work area tidy. Only items necessary for the experiment should be on the lab counter. Place

purses & backpacks on your seat out of the aisles. 6. Dress appropriately for lab.

a. Always wear goggles and a lab apron. b. Tie back long hair. c. Wear closed-toed shoes. d. Do not wear loose clothing, like ties or shirts with baggy sleeves. e. Do not wear dangling jewelry, including lanyards. f. Avoid wearing contacts on lab days.

7. Do not engage in horseplay. 8. Set up the apparatus as describe by your teacher. 9. Use the prescribed instrument for handling the apparatus. Ex) beaker tongs or test tube clamp. 10. Keep all combustible materials away from the flame. 11. Never smell any chemical directly. When testing for odors, use a wafting motion to direct the odors to

your nose. 12. Conduct any experiment involving poisonous vapors in the fume hood. 13. Dispose of waste materials as instructed by your teacher. 14. Clean and wipe your work area and wash your hands thoroughly when done.

Handling Chemicals:

15. Read and double-check labels on chemical bottles before removing any chemical. 16. Take only as much chemical as you need. 17. To avoid contamination, do not return any unused chemicals to stock bottles. 18. Avoid touching chemicals with your hands. If you do, wash them immediately. 19. When mixing an acid and water, always add the Acid to the water, never the reverse.

Handling Glassware:

20. Always carry glassware away from your body. 21. When inserting a thermometer or piece of glassware into a rubber stopper, lubricate the glass to avoid

breakage. 22. Do not place hot glassware directly on the table. 23. Allow hot glassware time to cool before touching it. Remember hot glass looks like cool glass! 24. Never handle broken glass with your hands. Use the dust pan and brush. Dispose in the cardboard

container.

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Heating Substances:

25. Use extreme caution with the gas burners. Keep your head and clothing away from the flame. 26. Never leave a burner unattended. 27. Never heat anything unless instructed to do so. 28. When heating a substance in a test tube, point the mouth of the test tube away from yourself or anyone

else. Never look into a container that is being heated. 29. Never heat a closed container.

If an Emergency Occurs:

30. If an injury should occur, it is important to remain calm. 31. Notify your instructor immediately if there is an injury or spill. 32. Be familiar with first-aid practices. 33. Know the location and proper use of emergency equipment such as the fire extinguisher, shower, eye

wash and fire blanket. 34. Know how to summon assistance. 35. Know how to interpret NFPA diamonds on doors and chemicals

Draw a sketch of your chemistry laboratory below and mark the locations of these items:

fire extinguisher

emergency exits

safety shower

fire blanket

eyewash

aprons

goggles

lab stations

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Be Able to Identify the Following Pieces of Lab Equipment

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1.3 Accuracy vs. Precision

In the field of chemistry, the accuracy of a measurement system is the degree of closeness of measurements of a

quantity to that quantity's actual (true) value. The precision of a measurement system, also called reproducibility

or repeatability, is the degree to which repeated measurements under unchanged conditions show the same

results.

Intensive vs. Extensive Properties

Intensive - Properties that do not depend on the amount of the matter present. Color Odor Luster - How shiny a substance is. Malleability - The ability of a substance to be beaten into thin sheets. Ductility - The ability of a substance to be drawn into thin wires. Conductivity - The ability of a substance to allow the flow of energy or electricity. Hardness - How easily a substance can be scratched. Melting/Freezing Point - The temperature at which the solid and liquid phases of a substance are in

equilibrium at atmospheric pressure. Boiling Point - The temperature at which the vapor pressure of a liquid is equal to the pressure on

the liquid (generally atmospheric pressure). Density - The mass of a substance divided by its volume

Extensive - Properties that do depend on the amount of matter present. Mass - A measurement of the amount of matter in an object (grams). Weight - A measurement of the gravitational force of attraction of the earth acting on an object. Volume - A measurement of the amount of space a substance occupies. Length

Practice Questions on Intensive vs. Extensive

1. Which of the following is an intensive chemical property of a box of raisins?

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a. calories per serving b. total grams c. total number of raisins d. total calories

2. Which of the following are extensive properties of a sample of gold?

a. density of 19.3 g/cm3 b. melts at 1064o C c. weighs 30.0 g d. yellow

3. Which of the following are extensive properties of coffee in a mug?

a. total mg of caffeine b. cream added per mL of coffee

c. temperature d. percent sugar

Types of Error

Determinate Error: have a definite direction and magnitude and have an assignable cause (their cause can be determined). Determinate error is also called systematic error. Determinate error can (theoretically) be eliminated.

Indeterminate Error: arise from uncertainties in a measurement as discussed above. Indeterminate error is also called random error, or noise. Indeterminate error can be minimized but cannot be eliminated.

Percent Error: expresses as a percentage the difference between an approximate or measured value and an exact or known value.

PE = accepted – experimental/accepted x 100%

Example: If a student measures the length of the desk to be 1.23 m and the actual length of the desk is 1.25 m,

what is the student’s PE?

Accuracy and Precision Mini Lab at Desk

Materials: ruler, sheet of paper

Accuracy = how close a measurement is to some accepted, true value

Precision = a term used to describe how close repeated measurements are to each other

Lab Objective: The student will be able to distinguish between the accuracy and precision of estimates made of

the measure of several different quantities.

Procedure 1: Estimating Lengths

1. Tear a sheet of paper into 8 equal size rectangles 2. Without a ruler, draw a free hand line that you estimate to be 5.00 cm long on one of the sheets. Turn

the sheet of paper over. 3. Without a ruler, repeat drawing of what you estimate to be 5.00 cm long lines on 3 more sheets of paper,

each time turning the sheet over so that you are drawing a line without being able to look at your previous estimates.

4. Now measure the length of each line with your metric ruler. Enter these measurements in the data table below. Remember that you will have two digits to the right of the decimal.

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Measured Length: Deviation from Average:

absolute value of:

avg. length – measured length

Percent Error:

accepted – experimental/accepted x 100%

Trial 1

Trial 2

Trial 3

Trial 4

Avg. = Avg. = Avg. =

5. Hide the first four sheets from sight. Repeat the above process with the remaining 4 pieces of paper, each time turning them over so that you cannot see the other estimates.

6. Again measure the actual lengths. Enter the lengths on the data table.

Measured Length: Deviation from Average:

absolute value of:

avg. length – measured length

Percent Error:

accepted – experimental/accepted x 100%

Trial 5

Trial 6

Trial 7

Trial 8

Avg. = Avg. = Avg. =

Questions:

1. Which of your two sets of estimates was most accurate (1-4 or 5-8)? Explain why you think this occurred and

justify your answer. (Use data from your table.)

2. Which of the two sets of estimates was most precise? Again, justify your answer?

4. Is the most accurate set also the most precise? Do they have to be? Explain.

5. Would the following be determinate (D) or indeterminate (I) error in the above activity?

a. Your ruler is missing calibration marks on the portion used to measure the length. b. You always make your measurement starting at 1 cm instead of 0 cm. c. Your ruler has inaccurate calibration marks made by the manufacturer.

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1.4 Math Skills for the Laboratory

1. Exponential Numbers

The numbers that we deal with in the laboratory are often very large or very small. Consequently, these numbers are expressed in scientific notation, using exponential numbers. These rules apply to the use of exponents:

When n is a positive integer, the expression 10n means “multiply 10 by itself n times”. Thus,

101 = 10 102 = 10 X 10 = 100 103 = 10 X 10 X 10 = 1,000 etc.

When n is a negative integer, the expression 10 n means “multiply 1/10 by itself n times”. Thus,

10-1 = 0.1 10-2 = 0.1 X 0.1 = 0.01

10-3 = 0.1 X 0.1 X 0.1 = 0.001 etc.

Examples: 2 x 101 = 2 X 10 = 20

2.62 x 102 = 2.62 X 100 = 262

5.30 x 10-1 = 5.30 X 0.1 = 0.530

8.1 x 10-2 = 8.1 X 0.01 = 0.081

In scientific notation, all numbers are expressed as the product of a number (between 1 and 10) and a whole number power of 10. This is also called exponential notation. To express a number in scientific notation, do the following:

1. First express the numerical quantity between 1 and 10. 2. Count the places that the decimal point was moved to obtain this number. If the decimal point has to be

moved to the left, n is a positive integer; if the decimal point has to be moved to the right, n is a negative integer. Examples: 8162 requires the decimal to be moved 3 places to the left

= 8.162 x 103

0.054 requires the decimal to be moved 2 places to the right

= 5.4 x 10-2

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Practice:

Express the following numbers in scientific notation.

20,205 = 0.000192 =

5,800000,000 = ______________ 0.0000034 =______________

40,230,000 = 543.6 =

34.5 x 103 = 0.004 x 10-3 =

0.72 x 10-6 = 0.029 x 102 =

2. Addition and Subtraction of Exponential Numbers

Before numbers in scientific notation can be added or subtracted, the exponents must be equal.

Example: (5.4 x 103) + (6.0 x 102) =

(5.4 x 103) + (0.60 x 103) =

(5.4 + 0.60) x 103 = 6.0 x 103

Practice:

(5.4 x 10-8) + (6.6 x 10-9) = (4.4 x 105) - (6.0 x 106) =

(3.24 x 104) + (1.1 x 102) = (0.434 x 10-3) - (6.0 x 10-6) =

3. Multiplying and Dividing Exponential Numbers

A major advantage of scientific notation is that it simplifies the process of multiplication and division. When

numbers are multiplied, exponents are added; when numbers are divided, exponents are subtracted.

Examples: (3 x 104)(2 x 102) = (3 X 2)(104+2) = 6 x 106

(3 x 104) (2 x 102) = (3 2)(104-2)= 1.5 x 102

OR (3 x 104) = (3/2)(104-2) = 1.5 x 102

(2 x 102)

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Practice:

All answers should be left in scientific notation.

(3.4 x 103)(2.0 x 107) = ___________ (5.4 x 102) (2.7 x 104) =_______________

(4.6 x 101)(6.7 x 104) = ___________ (8.4 x 10-3) (4.0 x 105) = ______________

Combine everything you have learned and perform the following calculation. Write your answer in scientific

notation.

(3.24 x 108)(14,000)/(3.5 x 10-3) = _________________

4. Prefixes Used for Powers of 10

Table 1: SI Prefixes and Symbols

Factor Decimal Representation Prefix Symbol

1018 1,000,000,000,000,000,000 exa E

1015 1,000,000,000,000,000 peta P

1012 1,000,000,000,000 tera T

109 1,000,000,000 giga G

106 1,000,000 mega M

103 1,000 kilo k

102 100 hecto h

101 10 deka da

100 1

10-1 0.1 deci d

10-2 0.01 centi c

10-3 0.001 milli m

10-6 0.000 001 micro m

10-9 0.000 000 001 nano n

10-12 0.000 000 000 001 pico p

10-15 0.000 000 000 000 001 femto f

10-18 0.000 000 000 000 000 001 atto a

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5. SI Units Units are specific quantities (amounts) used for measuring. You could grab a stick from outside and measure a

table in number of sticks. But if each of us grab a stick from outside each stick is likely to be a different length, so

we use standard units to measure things. A meter stick is a standard unit of about 3 feet and three inches and

every meter stick is the same length. If we all measure a table with meter sticks we should get the same number

of meter sticks in length. Scientists use the International System (SI) of standard measurements. Why do we

abbreviate International System as SI? It comes from the French, who invented the metric system and SI, and the

words are reversed in the French language. The meter and other metric units (based on the number 10) are one

small part of SI. Many Americans remain unfamiliar with the metric system because, unlike most countries, we are

one of only a small handful of countries that still use the British/ Imperial units of measurement on a daily basis.

Some examples of metric units (part of the larger set of SI units!) are liters (l), meters (m), centimeters (cm) and

millimeters (mm). Some examples of British/ Imperial are: gallons (gal), pounds (lbs), inches (in) and feet (ft).

SI Units used in Chemistry

Quantity Unit Symbol Definition

Length meter m The path travelled by light in vacuum during a time interval of

1/299792458 seconds. This fixes the speed of light to exactly 299792458 m/s.

Mass kilogram kg Mass of the platinum-iridium prototype at Bureau international des poids et

mesures in Sevres.

Time second s One second equals 9192631770 periods of the radiation due to the transition

between the two hyperfine levels of the ground state of Cesium 133.

Temperature kelvin K One degree K equals 1/273.16 of the thermodynamic temperature of the triple

point of water.

Quantity of

substance

mole mol The amount of a substance composed of as many specified elementary units

(molecules, atoms) as there are atoms in 0.012 kg of Carbon 12.

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Length SI Unit: meter (m)

1 meter = 1.0936 yards

1 centimeter = 0.39370 inch

1 inch = 2.54

centimeters

1 Kilometer = 0.62137 mile

1 mile = 5280 feet

= 1.6093

kilometers

1 angstrom = 10-10 meter

Volume

SI Unit: cubic meter (m3)

1 liter = 10-3 m3

= 1 dm3

= 1.0567 quarts

1 gallon = 4quarts

= 8pints

= 3.7854 liters

1 quart = 32 fluid ounces

= 0.94633 liter

Mass SI Unit: Kilogram (Kg)

1 Kilogram = 1000 grams

= 2.2046 pounds

1 pound = 453.59 grams

= 0.45359

Kilogram

= 16 ounces

1 ton = 2000 pounds

= 907.185

Kilograms

1 metric ton = 1000 Kilograms

= 2204.6 pounds

1 amu = 1.66056 x 10-27

Kilograms

Temperature (use formulas to convert) SI Unit: kelvin (K)

0 K = -273.15oC

= -459.67oF

K = oC + 273.15 oC = 5/9(oF – 32) oF = 9/5(oC) + 32

Pressure

SI Unit: pascal (Pa)

1 pascal = 1 N

m2

= 1 Kg

m s2

1 atmosphere = 101.325 kPa

= 760 torr

(mmHg)

= 14.70 psi

1 bar = 105 pascals

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6. Conversion Factors and Dimensional Analysis

The use of a conversion factor is often useful in doing more complex conversions. A conversion factor is simply the ratio between the two units of measurement.

Examples: Give conversion factors for the following pairs of units.

Kilograms and grams 1000g = 1 kg so 1000g/kg or 1 kg/1000g

Liters and milliliters 1 L = 1000 mL so 1 L/1000mL or 0.001 L/mL

meters and centimeters 1 m = 100 cm so 100 cm/m or 0.01 m/cm

Often in chemistry, the measurements we need are not in convenient units. Not only are there metric units and prefixes to consider (distances in mm, cm, m, km..), there are American units (distances in feet, inches, miles…). When solving and communicating math problems, unit conversions are expressed in “dimensional analysis” – which we sometimes just call “railroad tracks” or “unit analysis.” The railroad tracks handles the basic algebra for us. Example: Let’s start with something basic. A wire is 1.3 feet long. How many centimeters is that? I know that 1 inch = 2.54 cm… and 12 inches = 1 foot. Setup your unit conversion railroad track so that units in the numerator (on top) will cancel with units in the denominator (on bottom).

1.300ft 12 in 2.54cm = 39.62 cm

1 ft 1 in

Another example: A snail travels 13 feet / hour. How fast is this in m / sec?

13 ft 12 in 2.54cm 1 m 1 hr 1 min 0.0011 m

1 hr 1 ft 1 in 100 cm 60 min 60 s s

Another example: Convert 30.0 in2 to cm2 For a problem like this you must square your conversion factor. So 1 in = 2.54 cm or 1 in2=6.45 cm2

30 in2 6.45 cm2 = 193.5cm2

1 in2

Notice how top & bottom are always equal QUANTITIES, even though they aren’t the same number.

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When solving a problem using dimensional analysis, remember to do the following:

1. Identify the GIVEN and WANTED values. 2. Write down the per expressions (conversion factors) that share the units of measurement of the given and

wanted values, providing a unit pathway. 3. Align the given quantities and the conversion factors so that the given units of measurement cancel and the

wanted units of measurement are left in the numerator. 4. Write the calculation, including units. 5. Calculate the numerical answer and cancel out units of measurement that disappear when divided by

themselves. 6. Check the answer to be sure both the number and units make sense.

Practice:

1. Convert 555,000. square centimeters square miles.

2. Convert 1.00 square yard square centimeters.

3. Convert 30.0 m/s to km/hr

4. Convert 459 ft/sec --> mi/hr

Additional Questions: Harry Potter and the Amazing Unit Analysis

1. As Hagrid says, wizard money is very easy to understand. There are three coins: knuts, sickles and galleons. There are 29 knuts in one sickle and 17 sickles in one galleon. How many knuts in a galleon?

2. Ron has carefully horded every knut he’s found for 10 years and now he has three huge bags. He didn’t want to count every coin so he weighed the bags and found he had 75 pounds of knuts. One knut weighs 2 ounces. How much does he have in galleons?

16 oz. = 1 lb

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3. Harry is practicing flying on his Firebolt. He does 10 laps around the Quidditch field in 18 minutes. One lap of the field is 700 meters (m). How fast is he going in kilometers (km) per hour?

1000 m = 1 km

4. One of the most important ingredients in Polyjuice Potion (used to make you look like someone else) is dried boomslang skin. As you know, boomslangs are very small which is why boomslang skin is so expensive. It takes 32 boomslangs to make 1 teaspoon (tsp.) of dried boomslang skin. The potion calls for ½ cup (c.) of skin. How many boomslangs have to give their lives for the recipe?

3 tsp. = 1 tablespoon (tbsp.) 16 tbsp. = 1 cup

1.5 Determining Significant Figures

It is important to make accurate measurements and to record them correctly so that the accuracy of the

measurement is reflected in the number recorded. No physical measurement is exact; every measurement has

some uncertainty. The recorded measurement should reflect that uncertainty. One way to do that is to attach an

uncertainty to the recorded number. For example, if a bathroom scale weighs correctly to within one pound, and a

person weighs 145 lbs, then the recorded weight should be 145 + 1 lbs. The last digit, 5, is the uncertain digit, and

is named the doubtful digit.

Another way to indicate uncertainty is the use of significant figures. The number of significant figures in a quantity

is the number of digits that are known accurately plus the doubtful digit. The doubtful digit is always the last digit

in the number. Significant figures in a measurement

apply to measurements or calculations from measurements and do not apply to exact numbers.

are independent of the location of the decimal point

are determined by the measurement process and not the units

For example, a balance can weigh to + 0.01 g. A sample weighs 54.69 g. The doubtful digit is 9.

When an answer given has more numbers than significant, then the last number must be rounded off. If the first

digit to be dropped is <5, leave the doubtful digit before it unchanged. If the first digit to be dropped is >5, then

you round upward by adding a unit to the doubtful digit left behind. For example, a student using the balance

above measures 4.688 g. The correct number will be 4.69 g.

If there is only one digit beyond the doubtful digit in your number, and that digit is exactly 5, the rule is to round it

down half the time and to round it up half the time so that you don’t add a systematic error to your data. To keep

track when to round up and when to round down, the rule of thumb is to always round to an even number in the

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remaining doubtful digit. For example, if a measurement on a balance with a + 0.01 g accuracy is used to measure

4.895 g, you should record 4.90 g. If it reads 4.885 g, you should record 4.88 g as your data.

More explanation:

There is always an uncertainty associated with physical measurements, arising not only from the care with which you take the measurement, but also from the care with which the measuring device is calibrated. If you have done all that you can to minimize error in taking measurement, your recorded values should then reflect the uncertainty (precision) of the measuring tool. This is usually the smallest numerical value that can be estimated with the measuring device. For example, imagine trying to measure the length of the following line segment using a metric ruler:

Is the length of the line between 4 and 5 cm? Yes, definitely. Is the length between 4.0 and 4.5 cm? Yes, it looks that way. But is the length 4.3 cm? Is it 4.4 cm?

Given the precision of the ruler and our ability to estimate where between a set of marked graduations (the tick marks on the ruler) a measurement falls, we are somewhat uncertain about what number to record after the decimal. So, what we can say is that the actual length is around 4.4 cm, but it might be closer to 4.3 cm, or it might be closer to 4.5 cm. In other words, we think the length is 4.4 cm but we might be off by 0.1 cm in either direction. We would record this measurement in this way:

4.4 ± 0.1 cm

Keeping track of the uncertainty would be cumbersome if the uncertainty had to be reported this way each time the measurement itself were reported or used in a calculation. Therefore, we use significant figures to imply the precision of a measurement without having to state the uncertainty explicitly. In this course, we will assume an uncertainty of ±1 in the last recorded digit unless stated otherwise. the measurement recorded above could then be recorded just as:

4.4 cm

and the uncertainty of ±0.1 cm would be implied (but you still, always, have to include the units). Note that when using this method, it is very important that you record all significant digits. If you measured a mass and found it to be 2.0000 ± 0.0001g, it would be wrong to record the mass as:

2 g (Wrong!)

Instead you must include all significant figures, even if they happen to be trailing zeros:

2.0000 g (Right!)

The uncertainty that we have been discussing so far is always associated with individual physical measurements. The value that you are trying to find when you make such a measurement has a true value which is unknown and is fundamentally unknowable. Because there is (unavoidable) uncertainty in your measurements, the values you get when taking a series of measurements will tend to scatter around the true value. For example, if the above line

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were measured with several different rulers, a series of measurements would be obtained, each of which might be slightly different.

The difference between the true value and any given measured value is called the error in the measurement.

Practice:

The uncertainty of a balance measurement is + 0.01 g. Write the numbers that should be record as data with the

correct number of significant figures for the following. Some answers may already be correct.

445.81 g _______________ 6.731 g _______________

5872.30 g ______________ 5.556 g _______________

5.555 g 5.565 g

It is sometimes confusing to determine whether a zero in a number is a significant figure or not. Significant figures

are critical when reporting scientific data because they give the reader an idea of how well you could actually

measure/report your data. Before looking at a few examples, let's summarize the rules for significant figures.

1) ALL non-zero numbers (1,2,3,4,5,6,7,8,9) are ALWAYS significant.

2) ALL zeroes between non-zero numbers are ALWAYS significant.

3) ALL zeroes which are SIMULTANEOUSLY to the right of the decimal point AND at the end of the number

are ALWAYS significant.

4) ALL zeroes which are to the left of a written decimal point and are in a number >= 1 are ALWAYS

significant.

A helpful way to check rules 3 and 4 is to write the number in scientific notation. If you can/must get rid of the

zeroes, then they are NOT significant.

Examples: How many significant figures are present in the following numbers?

Number # Significant Figures Rule(s)

48,923 5 1

3.967 4 1

900.06 5 1,2,4

0.0004 (= 4 E-4) 1 1,4

8.1000 5 1,3

501.040 6 1,2,3,4

3,000,000 (= 3 E+6) 1 1

10.0 (= 1.00 E+1) 3 1,3,4

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Practice:

Determine the correct number of significant figures in the following numbers.

10.01 g 140 g

0.0010 g 140.0 g _______

1.100 g 1100 g

Calculations Using Significant Figures

In adding or subtracting numbers, the answer should contain only as many decimal places as the measurement

having the least number of decimal places. In other words, you answer should reflect the accuracy of the

measurement by correctly placing the doubtful digit. This is best done by lining up the numbers to be added or

subtracted, performing the addition or subtraction, and discarding any digits to the right of the doubtful digit from

the answer.

Example: For a balance that measures to + 0.01 g, the sum of the following measurements yields:

34.60 + 24.555 g = 34.60

+ 24.555

59.155 g = 59.16 g

Practice:

Solve the following and report your answer with the correct number of significant figures and units.

16.0 g + 3.106 g + 0.8 g (from a balance that weight to + 0.1 g)

9.002 m - 3.10 m (from a meter stick that measures to the nearest cm)

When multiplying or dividing, the answer may have only as many significant figures as the measurement with the

least number of significant figures. This is especially important to remember when using a calculator, since your

calculator may give you an answer with 11 digits!

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Examples: (1.13 m)(5.1261 m) = 5.79251786 m2 = 5.79 m2

Significant figures: 3 5 = 3

4.96001 g 4.740 cm3 = 1.0464135 g/cm3 = 1.046 g/cm3

Significant figures: 6 4 = 4

Dimensional Analysis and converting between two sets of units never changes the number of significant figures

in a measurement. Remember, data are only as good as the original measurement, and no later manipulations

can clean them up.

For example: Convert and expression conversion with correct number of significant figures: 30.0 cm/s in/min Significant Figures Practice

1. For the following measurements, indicate how many significant figures (sf's) there are:

a) 34 g ___ b) 56.4 L ___ b) 19.30 mm ___ d) 0.0001 mg ___ e) 101 km ___ f) 0.010100 __

g) 23100 ___ h) 23100. ___ i) 23100.00 __

2. Round off the following numbers to three significant digits:

a) 120000 _______________ b) 4.53619 _______________

c) 0.0008769 _______________ d) 876493 _______________

3. Perform all calculations and express your answer with the appropriate sig figs & units:

a) 67 cm x 55 cm = ______ b) 4.29 m x 9.83 m = _____ c) 870 mm x 430 mm = _____

d) 0.034 g/L x 8.8 L = _____ e) 5.79 m/hr x 2.34 hr = _____ f) 1.405 m x 6.39 m = ______

4. Perform all calculations & express your answers with the appropriate sig figs & units:

a) 67 cm + 45 cm = ______ b) 4.29 m + 9.83 m = ______ c) 170 mm + 250 mm = ______

d) 69.2 cm x 45 cm = ______ b) 5.29 m x 10.83 m = _____ c) 170 mm x 250 mm = ______

d) 90.2 cm / 45 cm = ______ b) 12.29 m / 0.83 m = _____ c) 170. mm / 250. mm = ______

21

1.6 Classification of Matter

Matter:

Can substance be separated by

physical means?

Is the composition variable? Can substance be separated by chemical means? Does it contain more

than one type of atom?

22

Material

1st level of classification 2nd level of classification

Iron filings (Fe) pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Limestone (CaCO3) pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Pure water (H2O) pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Tap water pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Acetylene (C2H2) pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Coca-Cola pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Orange juice pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Air inside a balloon pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Brass (an alloy of copper and zinc)

pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

White bread pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

Gasoline with 10% ethanol pure substance

element or compound

mixture homogeneous (solution)

or heterogeneous (suspension)

23

In the boxes below, draw models of the type of matter indicated.

A mixture of 2 different elements, one of which is diatomic.

A sample of a 3 atom compound made of 2 elements (like water).

A mixture of a monatomic element and a compound made of 2 elements.

Phases of State

Matter can come in one of four phases or a combination of phases.

24

Modeling Matter

Each diagram (A - P) show a sample of substances as viewed at the atomic level. Characterize the

contents of the container in terms of each of the following categories:

Category I. Homogeneous mixture, heterogeneous mixture or pure substance

Category II. Element(s), compound(s) or both

Category III. Solid, liquid, gas or combination of phases

25

1.9 Introduction to Mole

The SI unit used to measure the amount of substance is the mole. Moles are actually quite versatile and are

widely used in Chemistry. The term is a bit generic as is the term dozen. When we think of a dozen bananas we

instantly envision 12 bananas. The term mole is used to “count” the number of particles of a substance, represent

the mass of substances or even volume of gases.

One way moles are used is to count the number of particles of a substance. A mole of a substance is equal to 6.02 x

1023 representative particles of a substance. This is just as easy to remember as remembering a dozen is equal to

12. So, when you think of a mole of bananas you simply envision 6.02 x 1023 bananas. Piece of cake! You may be

wondering at this point what a representative particle is. A representative particle just refers to how a substance is

represented. For example,

Substance Example Representative Particle

Element Ne atoms

Compounds (covalent) H2O molecules

Compounds (ionic) NaCl formula units

Ions Al3+ ions

The number, 6.02 x 1023, is called Avogadro’s number, after Lorenzo Romano Amedeo Carlo Avogadro di Quareqa

e di Carreto - Avogadro for short. Amadeo Avogadro was the pioneer who discovered and quantified amounts of

particles in a given volume. The following link explains the history behind this event.

http://www.carlton.paschools.pa.sk.ca/chemical/molemass/avogadro.htm

What is neat is that 1 mole of Ne = 6.02 x 1023 atoms; 1 mole of water = 6.02 x 1023 molecules; 1 mole of NaCl =

6.02 x 1023 formula units and so forth. Solving mole problems is easy with dimensional analysis (factor-label

method).

EXAMPLES:

1- How many atoms are in 3.2 moles of zinc?

3.2 mol 6.02 x 1023 atoms = 1.9 x 1024 atoms

1 mol

2- How many moles are in 4.50 x 1028 molecules of water?

4.50 x 1028 molecules 1 mol = 7.48 x 104 molecules

6.02 x 1023 molecules

Avogadro’s number is huge. Some cool facts follow.

26

Mole Facts

6.02 X 1023 Watermelon Seeds: Would be found inside a melon slightly larger than the moon.

6.02 X 1023 Donut Holes: Would cover the earth and be 5 miles (8 km) deep.

6.02 X 1023 Pennies: Would make at least 7 stacks that would reach the moon.

Additional Explanation of the Mole

The Mole:

Is the SI unit for the amount of substance and is abbreviated “mol”.

One mole contains 6.02 x 1023 representative particles.

For instance, 1 mole = 6.022 x 1023 atoms

1 mole = 6.02 x 1023 ions

1 mole = 6.02 x 1023 molecules

1 mole = 6.02 x 1023 ion pairs

Avogadro’s Number (or constant) = 6.02 x 1023 atoms.

Compare:

One dozen contains 12 objects.

1 dozen eggs = 12 eggs

1 mol eggs = 6.02 x 1023 eggs

Atomic Mass: the weighted average of the masses of the isotopes of an element. Ex: 1 mole of Na = 22.99 g

(covered in unit 1)

How many atoms are there in 0.500 moles of Na?

0.500 mol Na x 6.02 x 1023 atoms Na = 3.01 x 1023 atoms Na

1 mole of Na

0.250 mol Na x 6.02 x 1023 atoms Na = 1.51 x 1023 atoms Na

1 mole of Na

How many moles are there in 2.50 x 1023 atoms of Na?

2.50 x 1023 atoms x 1 mole = 0.415 moles of Na

6.02 x 1023 atoms Na

Molar Mass

Molecular Mass of a covalent compound is the (atomic) mass of one molecule.

H2O = H 2 x 1.01 = 2.02 g

O + 1 x 15.99 = 15.99 g

18.01 g/mol

27

Calculate the molar (molecular) mass of CH4 (methane)

CH4 =

Formula Mass of an ionic compound is the (atomic) mass of one formula unit.

NaCl = 22.99 + 35.45 = 58.44 g/mol

MnCl2 = 54.94 + 70.90 (2 mol Cl) = 125.84 g/mol

How do you determine the amount of moles in a compound?

1. If you have 9.0 grams of water, you have ____moles of water.

Water - H2O = 18.01 g/mol

9.0 g H2O x 1 mole of H2O = 0.50 mol of H2O

18.01 g H2O

2. If you have 13.0 grams of water, you have ____moles of water.

3. If you weigh out 5.00g of Formaldehyde, how many moles is that? 60.0 g of formaldehyde?

CH2O – Formaldehyde 30.02 g/mol

Counting Moles

2 mol C6H12O6 contains:

2 mole of C6H12O6 molecules

2 x 6 = 12 moles of C atoms

2 x 12 = 24 moles of H atoms

2 x 6 = 12 moles of O atoms

48 moles of atoms in total

6 Ne atoms contain

6 x 1 = 6 moles of the Neon atoms

3 mol BF3 contains:

________ mole of BF3 molecules

________ moles of B atoms

________ moles of F atoms

________ moles of atoms in total

7 mol Ca(NO3)2 contains:

________ mole of the compound Ca(NO3 )2

________ moles of N atoms

________ moles of O atoms

________ moles of atoms in total

Independent Practice:

I. Calculate the molar masses of the following compounds.

1. HCl

2. N2O5

3. MgSO4

II. Determine the amount of moles and particles from the given mass amounts of the following compounds.

1. 38.2 g of HCl

2. 97.1 g of N2O5

3. 85. g of MgSO4

Now you can make calculations from one quantity of a substance to another quantity. Some examples follow.

EXAMPLES: 1- What is the mass in grams of 3.2 x 1024 molecules of CH4?

3.2 x 1024 molcules CH4 1 mol CH4 16.0g CH4 = 85.0 g CH4

6.02 x 1023 molecules 1 mol CH4

2- How many molecules are in 7.50 Liters of O2?

7.50 L O2 1 mol O2 6.02 x 1023 molecules O2 = 2.02 x 10 23

22.4 L O2 1 mol O2

Remember the relationships that we just learned:

A) 1 mole of any substance = 6.02 x 1023 representative particles (atoms, ions, molecules, formula units)

B) 1 mole of any gas at STP = 22.4 liters C) 1 mole of any substance = gram formula mass (molar mass) of that substance (using masses in

periodic table)

MORE MOLE PRACTICE

1. How many atoms of potassium are in 2 moles of K2O? _______________________

2. How many molecules of water make up 5 MOLES? ________________________

3. How many moles are 6.022 x 1023 atoms of sodium? ________________________

4. How many grams are in 2.3 x 10-4 moles of calcium phosphate, Ca3(PO3)2?

5. How many oxygen atoms are in 3.4 x 10-7 grams of silicon dioxide, SiO2?

1.10 Mole ratios in chemical equations introduction to stoichiometry

We have learned that chemical equations show how elements and compounds interact with each other. For example,

consider the equation below.

Methane + oxygen carbon dioxide + water

|--------------reactants-------------| |-------------------------products--------------------|

CH4 (l) + 2O2 (g) CO2 (g) + 2H2O (g)

The word equation and chemical equation using formulae shows how chemists represent a chemical reaction using words. It

reads:

‘One mole of methane reacts with two moles of oxygen gas to produce [or yield] one mole of carbon dioxide gas and

two moles of water. Heat energy is released.’

The arrow represents produces or yield. Scientists never put an equal (=) sign instead of the arrow. The substances reacting

are called reactants and the substances formed are called products. The (l) and (g) represent the state of the substance. (l) is

for a liquid, (g) a gas, (s) a solid and (aq) for an aqueous solution. An aqueous solution is formed when a solid compound is

dissolved in water. For example sodium chloride solution is formed by taking a known amount of solid sodium chloride and

dissolving it in a known volume of water. More about this later.

In a chemical reactions, substances need to combine in the correct mole ratio in order for the reaction to occur. Even if the

number of moles changes the ratio of moles of reactants and products does not change. In the reaction below two moles of

hydrochloric acid reacts with one mole of zinc.

Hydrochloric acid + zinc zinc chloride + hydrogen

2HCl(aq) + Zn(s) ZnCl2(aq) + H2(g)

Applying the mole ratio:

2 moles of hydrochloric acid react with _____ mole of zinc to produce 1 mole of zinc chloride and _____ mole of hydrogen.

1 mole of zinc chloride contains 1 mole of zinc ions and 2 moles of chloride ions

2 moles of hydrochloric acid contains ___ mole of hydrogen atoms and ___ of chlorine atoms.

All chemical reactions MUST be balanced. The numbers and types of atoms before (the reactants) and after (the products)

MUST be the same! When this happens the total mass of the atoms before and after the reaction will be the same. This is

called Law of conservation of mass.

Stoichiometry

WHY CHEMISTS USE IT?

Chemists use stoichiometry to determine quantities of chemicals _________________, and to predict the

quantities __________________ for any chemical reaction. Without stoichiometry, we would not know the

_______________ in which chemicals react or how much they produce.

What does it mean?

The word stoichiometry derives from two Greek words:

stoicheion (meaning "_______________") and metron

(meaning "____________________").

Stoichiometry deals with calculations about the masses

(sometimes volumes) of reactants and products involved

in a chemical reaction. It is a very mathematical part of

chemistry, so be prepared for lots of calculator use.

What You Should Expect…

The most common stoichiometric problem will

present you with a certain amount of a

__________________ and then ask how much of a

____________________ can be formed.

Here is a generic chemical equation:

2A + B2 ---> 2AB

Here is a typically-worded problem:

Given 20.0 grams of A and sufficient B, how

many grams of AB can be produced?

First things first… Mole Ratios

You will always use a mole ratio when working stoichiometry problems. A mole ratio is a

_________________________ that relates the amounts of moles of any two substances involved in a chemical

reaction.

Molar ratios are obtained from ___________________________ chemical equations.

For example: 2H2+ O2 → 2H2O What are the mole ratios in this reaction?

1) 2 mol H2 1 mol O2 (Both of these ratios are the same;

1 mol O2 2 mol H2 the only difference is which one is on top.)

2) 2 mol H2 2 mol H2O

2 mol H2O 2 mol H2

3.) 1 mol O2 2 mol H2O

2 mol H2O 1 mol O2

Example 1: What are some (give 3 )mole ratios in this balanced reaction?

3Ca + 2 H3PO4 Ca3(PO4)2 + 3 H2

Stoichiometry Practice Problems

Given chemical equation: Br2 + NaI NaBr + I2

1. Balance the above chemical equation.

2. How many moles of I2 are produced when 3.00 mol of Br2 completely reacts?

3. How many grams of sodium iodide (NaI) will completely react with 10.0 mol of bromine (Br2)?

4. How many grams of iodine (I2) could be produced when 50.0 g of bromine (Br2) completely reacts?

5. How many grams of sodium bromide (NaBr) could be produced from 0.172 mol of bromine (Br2)?

Given chemical equation: H2 + O2 H2O

6. How many moles of H2O are produced when 2.50 moles of oxygen are used?

7. If 3.6 moles of H2O are produced, how many grams of oxygen must be consumed?

At the heart of any stoichiometry problem…

You are given the amount of one substance and asked to find the number of moles of another

substance. 1. Write and balance the ______________ 2. Convert everything to _____________ 3. Use _____ _______to solve for what you are trying to find Convert everything into the required unit if needed.

Ex: How many moles of calcium chloride will be produced if you start out with 35 .0 g of HCl?

Ca(OH)2 + HCl → CaCl2 + H2O

8. How many grams of O2 can be produced by letting 12.00 moles of KClO3 react if given the following equation: KClO3 KCl + O2

9. Camels store the fat tristearin (C57H110O6) in the hump. As well as being a source of energy, the fat is a source of water.

How many moles of carbon dioxide are produced from 8.73 moles of tristearin?

2 C57H110O6(s) + 163 O2(g) 114 CO2(g) + 110 H2O(l)

1.8 Investigating Matter: Foul Water

Introduction

Your objective is to clean up a sample of foul water, producing as much “clean water” as possible, to a point where it could be

used for hand-washing. (Caution: Do not test any water samples by drinking or tasting them.) You will use several different

water-purification procedures: oil–water separation, sand filtration, and charcoal adsorption and filtration. Before starting,

read the procedure to learn what you will need to do, note safety precautions, and plan necessary data collecting and

observations.

Before starting, read the procedure to learn what you will need to do, note safety precautions, and plan necessary data

collecting and observations.

Procedure

1. Use the data table provided to record your observations throughout the investigation 2. Using a clean beaker, obtain approximately 100 mL (milliliters) of foul water from your teacher. Measure its volume

accurately with a graduated cylinder. Record the actual volume of the water sample in your data table. Leave your sample in the graduated cylinder.

3. Describe in detail the appearance, color, clarity, and odor of your original sample. Record your observations in the “Before treatment” row of your data table.

Oil–Water Separation

As you probably know, if oil and water are mixed and left undisturbed, the oil and water do not noticeably dissolve in each

other. Instead, two layers form. Which layer do you think will float on top of the other? Make careful observations in the

following procedure to check your answer.

4. Allow your sample to sit in the graduated cylinder for at least one minute. 5. Using a clean, dry Beral pipet, carefully remove as much of the upper liquid layer as possible and place it in a clean, dry

test tube. 6. Add several drops of distilled water to the liquid you placed in the test tube. Does the water float on top or sink to the

bottom? Is the liquid you removed in Step 5 water? Explain your reasoning, using evidence from your observations to support your answer.

7. Read and record the volume of the liquid sample remaining in the gradated cylinder. 8. Dispose of the liquid in the test tube as directed by your teacher.

Sand Filtration

In filtration, solid particles are separated from a liquid by passing the mixture through a material that retains the solid

particles and allows the liquid to pass through. The liquid collected after it has been filtered is called the filtrate. A sand filter

traps and removes solid impurities— at least those particles too large to fit between sand grains—from a liquid.

9. Using a straightened paper clip, poke small holes in the bottom of a disposable cup. See Figure 1.5 in your textbook. 10. Add pre-moistened gravel and sand layers to the cup (The bottom gravel layer prevents the sand from washing through

the holes. The top layer of gravel keeps the sand from churning up when the water sample is poured into the cup.) 11. Gently pour the sample to be filtered into the cup. Catch the filtrate in a beaker as it drains through. 12. Dispose of the used sand and gravel according to your teacher’s instructions. (Caution: Do not pour any sand or gravel

into the sink!) 13. Observe the properties of the filtered water sample and measure its volume. Record your results. Save the filtered water

sample for the next procedure.

Charcoal Adsorption and Filtration

Charcoal adsorbs, which means attracts and holds on its surface, many substances that could give water a bad taste, a cloudy

appearance, or an odor.

14. Fold a piece of filter paper, as shown in Figure 1.7 in your textbook. 15. Place the folded filter paper in a funnel. Hold the filter paper in position and moisten it slightly so that it rests firmly

against the base and sides of the funnel cone. 16. Place the funnel in a clay triangle supported by a ring, as shown in Figure 1.8 in your textbook. Lower the ring so that the

funnel stem extends 2 to 3 cm (centimeters) inside a 150-mL beaker. 17. Place no more than one level teaspoon of charcoal in a 125-mL or 250-mL Erlenmeyer flask. 18. Pour the water sample into the flask. Swirl the flask vigorously for several seconds. Then gently pour the liquid through

the filter paper. Keep the liquid level below the top of the filter paper; liquid should not flow between the filter paper and the funnel because that might permit unwanted charcoal and other solid matter to seep into the filtrate.

19. If the filtrate is darkened by small charcoal particles, once again filter the liquid through a clean piece of moistened filter paper.

20. When you are satisfied with the appearance and odor of your charcoal-filtered water sample, pour the filtered water sample into a graduated cylinder. Record the final volume and properties of your purified sample.

21. Follow your teacher’s suggestions about saving or disposing of your purified sample. Place the used charcoal in the container that is provided for that purpose.

22. Wash your hands thoroughly before leaving the laboratory.

Data Analysis

Record all calculations and answers in the space provided.

1. What percentage of your original foul water sample did you recover as purified water? This value is called the percent recovery.

Percent recovery __________

2. What volume of liquid (in milliliters) did you lose during the entire purification process?

Volume lost __________

3. What percent of your original foul-water sample was lost during purification?

Percent lost __________

To answer the following questions, first collect a list of percent recovery values for water samples from each laboratory

group.

4. Construct a histogram showing the percent recovery obtained by all laboratory groups in your class. To do so, organize the data into equal subdivisions, such as 90.0–99.9%, 80.0–89.9%, and so forth. Count the number of data points in each subdivision. Then use this number to represent the height of the appropriate bar on your histogram, as illustrated in Figure 1.9 of your textbook.

5. What was the largest percent recovery obtained by a laboratory group in your class? What was the smallest? The difference between the largest and smallest values in a data set is the range of those data points. What was the range of percent recovery data in your class?

Largest percent recovery __________________________

Smallest percent recovery __________________________

Range of percent recovery __________________________

6. What was the average percent recovery for your class? Compute an average value by adding all values together and dividing the sum by the total number of values. The result is also called the mean value.

Average percent recovery __________

7. The mean is a mathematical expression for the most “typical” or “representative” value for a data set. Another useful expression is the median value, or middle value. To find the median for percent-recovery data, list all values in either ascending or descending order. Then find the value in the middle of the list— the point where there are as many data points above as below. If you have an even number of data points, take the average of the two values nearest the middle. What is the median percent recovery of all of your class laboratory results?

Median percent recovery __________

Investigating Matter: Foul Water

Data Table

Data Table

Volume

(mL)

Color

Clarity

Odor

Presence

of Oil

Presence

of Solids

Before

Treatment

After

oil-water

separation

After sand

filtration

After

charcoal

adsorption

and filtration

1.11 Baking Soda Stoichiometry/Percent Yield Lab

In this lab, you will combine your powers of observation, reasoning, equation balancing, and knowledge of stoichiometric calculations to earn a perfect 10 / 10 (hopefully). Procedure: 1. Obtain a large Pyrex test tube & weigh it on one of the scales in the front of the room. Record this mass in the table at right. Pyrex is a kind of glass that can be subjected to very high (and low) temperatures without shattering. 2. Go back to your lab station & place one large scoop of baking soda (NaHCO3) into the test tube. Then, using the same scale as before, weigh the test tube with the baking soda. Record this mass in the data table. (You should be able to figure out the mass of the baking soda in the test tube.) 3. Holding the test tube nearly horizontal, shake the baking soda gently so that it spreads out a bit as shown: 4. Then tighten the test tube clamp gently around the test tube, just below the lip so that it is positioned nearly horizontally, about 20 cm above the lab desk as shown: 5. Light a burner and adjust it to a cool flame (vent closed) hitting the bottom half of the test tube as shown: Record the time you started heating: _______________ This will initiate a chemical change (a sort of decomposition reaction) that breaks the NaHCO3 down... not into its elements, but into three separate compounds. 6. Put one drop of the green indicator solution on the end of the small swab. Then carefully insert this end into the mouth of the test tube as shown: See if you can observe a distinct color change. If a metal oxide like K2O, Na2O, or MgO is being produced, it will create a basic (alkaline) solution and turn the indicator blue. If a nonmetal oxide like NO2, SO3, or CO2 is being produced, it will create an acidic solution & turn the indicator yellowish. What color does it become? _______ So, is the reaction producing a metal or nonmetal oxide? __________ Look at the chemical formula of the substance you are heating: NaHCO3. So, what common oxide is being produced in the test tube? 7. What do you observe happening in the upper half of the test tube? ___________________________________ What common substance appears to be a second product of this reaction? 8. Move the burner occasionally to a different spot to ensure a thorough heating of the entire bottom half of the test tube. Consider the substance that is left in the test tube... it may look just like the baking soda you started with, but it actually has been converted into something else: sodium carbonate, Na2CO3. This is the third product. Now go down and answer questions 1-4 below, but keep an eye on the time. After you have been heating the test tube for 8~10 minutes, turn off the burner and let the test tube cool for 5~6 minutes, unless the inside walls still have moisture. If so, continue heating for a few more minutes to drive out the moisture before cooling. Questions: 1. You should have figured out from steps #6, 7, 8 above what the three products are. Write the unbalanced chemical equation for the reaction that just took place: Check it with the teacher to make sure you have it right, then go back & balance it. (Hint: it's very easy w/ small #'s)

Data Table

2. Look back at your data table above. What mass of the NaHCO3 did you start with in the test tube? ____________ 3. Starting with that mass of NaHCO3, use stoichiometry & the balanced equation to figure out what mass of sodium carbonate you should have ended up with in the test tube. Show below: 4. So... assuming all the baking soda you started with got converted into sodium carbonate, what should the test tube & its contents weigh now (official predicition)? If your test tube has been cooling for 5~6 minutes, it is ready for the official weigh-in! Bring the test tube, along with this sheet containing your prediction above, up to the balances. Your teacher will weigh it on the same scale you used, but not show you the weight. They will tell you your grade based on how close your prediction was to the actual weight (see table at right). If you are satisfied with your grade, congratulations! You are done. If you are not satisfied, you can go back, correct your mistake and change your prediction for a 2nd attempt. The 2nd attempt will cost you 1 point, and you may end up with a lower score. So, only try the 2nd attempt if you are fairly sure you can correct any mistake you may have made the first time. 5. Observe the substance that remains left behind in the test tube; compare it to the sealed tube of NaHCO3 at your lab station. Do you notice any slight difference between the two? After you have finished all of the above, rinse out the test tube into the sink, and leave it on the table to dry. Take a fresh (dry) test tube and place it in the clamp for the next group. Follow-Up Questions: 6. If you hadn't heated up the test tube long enough, would that make your prediction too high, too low, or no effect? Explain: 7. CO2 is more dense than air. So why did the CO2 you produced from the reaction rise upwards out of the opening of the test tube? 8. Why did the water only condense on the upper half of the test tube?

Official Prediction:

9. Using your original mass of baking soda (NaHCO3) from question 2, calculate the mass of H2O that was produced: (show work): 10. Using your original mass of baking soda (NaHCO3) from question 2, calculate the mass of CO2 that was produced: (show work): 11. Add the two masses from #9 and #10 above along with the calculated mass of sodium carbonate produced from question #3: #9 ___________ + #10 ___________ + #3 ___________ = What total mass of products does this give? 12. How does this mass (#11) compare with the initial mass of NaHCO3 you put in the test tube?? If it is different give some reasons why. 13. Calculate the percent yield of sodium carbonate. 13. If a person accidentally leaves a pan of oil on the stove, it might get so hot it will ignite. This is known as a grease fire. Pouring water on a grease fire is a bad idea, because the water (being more dense than the grease) will sink in the oil, expand rapidly in the heat, and splatter the grease thus spreading the fire. Pouring baking soda on a grease fire is a much better idea. Why? 14. Chemical reactions can be categorized as either exothermic (heat is given off by the reaction) or endothermic (heat is taken in by the reaction). What type of reaction is the decomposition of NaHCO3? ____________________ How do you know this?

Homework: 1.2

1. When diluting acid with water, which is added to which, and why?

2. What is the proper way to heat a test tube?

3. Is food and drink permitted in the lab? Why or Why not.

1.3

1. Classify the following as intensive or extensive properties.

Temperature, mass, color, hardness, volume, length, boiling point, density

2. A handbook lists the density of lead as 11.3 g/mL. Several groups of students are attempting to determine the density of a lead weight by various methods. Complete this table by calculating the average density measured by each group in g/mL.

Group 1 Group 2 Group 3

Trial 1 12.7 11.5 10.9

Trial 2 11.2 11.4 11.3

Trial 3 10.3 11.4 11.1

Average Density (g/mL)

Which group was most accurate?

Which group was most precise?

3. Record interval of each graduated cylinder in the box and give the volume of each graduated cylinder on the line.

1.4

1. On planet Zizzag, city Astric is 35.0 digs from city Betrek. The latest in teenage transportation is a Zeka which can travel a maximum of 115 dillidigs/zip. On Zizzag the planet turns once on its axis each dyne. Their time system divides each dyne into 25 zips. How many dynes will it take Pezzi to get from Astric to Betrek to see his girlfriend?

2. The density of crude oil is 15.0 stones/barrel. What is the density in kg/L?

3. Convert the following from scientific notation to normal notation or vice versa: a. 3.7 x 10-4

c. 0.00000000000008.45 b. 4.28 x 104 d. 82480000000000000

1.5

1. Record each number to 2 sig figs. 233.356

0.002353

1.005

2. To how many sig figs should each answer be rounded?

(6.626 𝑥 10−34)(2.9979 𝑥 108)

4.310 𝑥 10−7= 460836519722 𝑥 10−19

(6.022 𝑥 1023)(0.513)

20.18= 1.531 𝑥 1022

3. Record your answer to the correct number of sig figs.

5.9851 x 104 + 9.967 x 106

1.6

1. Diagram a substance that can be described by the following a. Solid homogeneous mixture containing

a compound and a diatomic element.

b. Gaseous pure substance comprised of a compound

c. Heterogeneous liquid containing two different atoms.

d. Describe the following:

1.7

1. Classify each as a <Phy>sical or <Chem>ical property:

_____ a. odor due to rotting _____ b. red color of apple _____ c. flammability

_____ d. density _____ e. reaction to acids _____ f. melting point

2. Classify each as a <Phy>sical or <Chem>ical change:

_____ a. rusting of iron

_____ b. boiling of water

_____ c. burning of sulfur

_____ d. cooking an egg

_____ e. digestion of food

_____ f. sawing of wood

_____ g. melting of wax

_____ h. dissolving salt in water

1.9

1. Calculate the number of molecules of the compound and the number of atoms of each of the elements in 1.25 moles of C2H6.

2. What is the mass in grams of 8.39 1018 atoms Br.

3. How many atoms is equivalent to 7.632 x 10-21 g W?

‘cules

Atoms H

Atoms C

1.10 1. Al(OH)3(s) + 3HCl(aq) AlCl3(aq) + 3H2O(l)

a. How many moles of HCl are required to react with 0.935 g of Al(OH)3 in the above reaction?

b. How many moles of H2O are produced?

2. 3 H2SO4 + 2 Al(OH)3 Al2(SO4)3 + 6 H2O a. How many grams of H2SO4 are required to react with 79.0 grams of Al(OH)3 in the above

reaction?

b. A student performed this reaction in the lab and produced and calculated a percent yield of 94.7%. How much Al2(SO4)3 did the student produce?