1

Chemistry HP Unit 1 – Scientific Measurement

1. Define matter and distinguish between a pure substance and a mixture; an element and a compound; a homogeneous mixture

and a heterogeneous mixture (by definition and with examples).

Chemistry = Study of matter

Matter = Anything that has mass and takes up space

1. Pure Substance = Every sample in the substance is identical to every other sample. (Examples: Water, H2O, O2, Au)

a. Pure substances are either elements or compounds.

i. Element = Pure substance made of only one type of atom. (Example: Au)

ii. Compound = Pure substance composed of two or more different elements chemically combined.

(Examples: H2O, O2)

2. Mixture = Type of matter composed of two or more substances physically combined. (Examples: A bowl of Raisin Bran cereal, Air (oxygen, nitrogen))

a. Homogeneous Mixture (SOLUTION) = Mixture that is the same throughout; uniform appearance and composition. (Examples: Pitcher of Kool Aid, Air (oxygen, nitrogen), metal alloys)

b. Heterogeneous Mixture = Mixture that is NOT the same throughout. (A bowl of Raisin Bran cereal)

Separation of Mixtures 3. Since mixtures are physically combined, they can be separated by physical methods, which are methods based on physical

properties. a. Filtration = The separation of a mixture based on particle size. b. Distillation = Separation based on boiling point.

i. A mixture of different liquids can be separated as long as they have different boiling points.

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2. Compare physical changes and chemical reactions by listing the key characteristics of each and giving examples.

1. Chemical Change = A change in the chemical composition of a substance. Atoms or molecules are rearranged. Bonds are

broken and formed.

a. Example: the formation of rust:

i. 4Fe + 3O2 2Fe2O3

2. Physical Change = A change in the physical state. The atoms or molecules are NOT rearranged. Bonds are NOT broken and

formed.

a. Example: the freezing of water:

i. H2O(l) H2O(s) + energy

3. Chemical Properties = Characteristics which can be seen as one substance is chemically transformed into another.

a. A chemical’s reactivity with another chemical

b. Examples: iron’s tendency to rust (Fe is becoming Fe2O3), flammability

4. Physical Properties = Anything that can be observed without changing the identity of the substance.

a. Examples: smell, color, solubility, boiling point, volatility

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1-3. Compare the three phases of matter (solid, liquid, and gas) in terms of molecular speed and molecular spacing, and be able to

draw a simple diagram for each state.

Solid Liquid Gas

Volume Definite Definite Takes volume of container

Shape Definite Takes shape of its container

Takes shape of its container

Molecular Speed

Slowest (vibrational)

Medium Fastest

Molecular Spacing

Closest Medium Farthest apart

Draw particle diagrams of Solid, Liquid, Gas

Solid Liquid Gas

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1-4. Convert between scientific notation and standard form of a number.

1. To convert from scientific notation to decimal notation:

a. If the exponent is positive, move the decimal to the right the number of places equal to the value of the exponent.

2 x 103 = 2000

b. If the exponent is negative, move the decimal to the left the number of places equal to the value of the exponent.

2 x 10-3 = 0.002

2. To convert from decimal notation to scientific notation:

a. Multiply the number in decimal notation by 100

2100

2100 x 100

b. Move the decimal of the significand (the number in decimal notation) to just after the first non-zero digit.

i. If you moved the decimal to the right, making the number in decimal notation larger, you want to make

the exponent smaller by the same number of digits.

ii. If you moved the decimal to the left, making the number in decimal notation smaller, you want to make

the exponent larger by the same number of digits.

Example: Convert 2100 to scientific notation:

2100 x 100

2.100 x 100 (moved the decimal 3 places to the left)

2.100 x 103 (since the decimal was moved to the left 3 places, the exponent was raised 3 places, from 0 to 3)

Example: Convert 0.051 to scientific notation:

0.051 x 100

5.1 x 100 (moved the decimal 2 places to the right)

5.1 x 10-2 (since the decimal was moved to the right 2 places, the exponent was lowered 2 places, from 0 to -2)

3. To convert from exponential notation to scientific notation:

a. Move the decimal of the significand to just after the first non-zero digit.

i. If you moved the decimal to the right, making the number in decimal notation larger, you want to make

the exponent smaller by the same number of digits.

ii. If you moved the decimal to the left, making the number in decimal notation smaller, you want to make

the exponent larger by the same number of digits.

Example: Convert 756 x 10-2 to scientific notation.

7.56 x 10-2 (move the decimal to just after the first non-zero digit)

7.56 x 100 (Since the decimal was moved to the left 2 places, the exponent is raised 2 places).

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Example: Convert 0.03450 x 10-4 to scientific notation.

3.450 x 10-4 (move the decimal to just after the first non-zero digit)

3.450 x 10-6 (Since the decimal was moved to the right 2 places, the exponent is lowered 2 places).

WS #2 (Learning Target 1-4. Convert between scientific notation and standard form of a number)

Directions: Write the number(s) given in each problem in decimal form.

1. The age of earth is approximately 4.5 x 109 years. __________________________yr

2. The weight of one atomic mass unit (a.m.u.) is 1.66 x 10-27 kg.

_____________________________________kg

Directions: Write each number in scientific notation.

3) 0.00000216

4) 5400000

5) 60

6) 0.63 × 102

7) 6.7

8) 0.0000002

9) 2000000

10) 71 × 103

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4. When you add or subtract numbers in scientific notation, they must have the same exponentials. If they do not have the

same exponentials, you need to convert one of them so that they do.

a. Convert the exponential of the smaller number to the exponential of the larger number.

b. Since you are raising its exponential, you need to move the decimal of its significand to the left the same number

of places.

Example: 1.70 x 103 + 3.75 x 105 = ?

In this example, we need to raise the exponential of 1.70 x 103 to match the exponential of 3.75 x 105. Since we are raising the exponential two places, we need to move the decimal two places to the left. 1.70 x 103 0.0170 x 105 Since these two number have identical exponentials, now we can add them together. 3.75 x 105

+ 0.0170 x 105

4.7670 05

5. Here are the rules for multiplying and dividing exponents.

a. When you multiply exponentials, you add the exponents.

b. When you divide exponentials, you subtract the exponents.

Example: Multiply 2.00 x 10-2 by 2.00 x 103.

(2.00 x 10-2)(2.00 x 103) = (2.00 x 2.00) (10-2 x 103) = (4.00)(10-2+3)

= 4.00 x 101

Example: Divide 2.00 x 10-2 by 2.00 x 103.

(2.00 x 10-2) ÷ (2.00 x 103) = (2.00 ÷ 2.00) (10-2 ÷ 103) = (1.00)(10-2-3)

= 1.00 x 10-5

6. When raising exponentials to a power, multiply the exponents.

Example: (2.00 x 10-2)3 = (2.00)3 x (10-2)3 = 8.00 x 10-6

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WS #3 (Learning Target 1-4. Convert between scientific notation and standard form of a number)

Directions: Perform the following calculations without a calculator. Write all answers in scientific notation.

11. 7.0 x 103 + (0.0070 x 106)

12. 8.5 x 102 - 2.0 x 10 -1

13. (4 x 10-3) (4 x 10-2)

14. (9.0 x 10-3) ÷ (3.0 x 106)

15. (5.0 x 10 6)2 ÷ (2.5 x 10-9)

16. (4 x 103)(2 x 104)

17. 8.74 x 102 – 2.3 x 103

18. (6.0 x 104)(7.0 x 102)

19. 2.4 x 10-1 – 4.0 x 10-3

20. (2.0 x 102)3 – (8.0 x 105)

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1-5. Convert from one metric unit to another using a metric scale.

Metric Base Units

Measurement Metric Base Unit

Length Meter

Mass Gram

Volume Liter

Basic Metric Prefixes

“times” (x) base unit Exponential Notation

Milli – (m) 1/1000 10-3

Centi – (c) 1/100 10-2

Deci – (d) 1/10 10-1

Deka – (da) 10 101

Hector – (h) 100 102

Kilo – (k) 1000 103

Useful Scientific Metric Prefixes

“times” (x) base unit Exponential Notation

Tera – (T) 1012 1012

Giga – (G) 1 000 000 000 109

Mega – (M) 1 000 000 106

micro – (µ) 0.000001 10-6

nano – (n) 0.000000001 10-9

pico – (p) 10-12 10-12

femto – (f) 10-15 10-15

Other Relationships to Memorize

Is equal to:

1 cm3 1 mL

1 dm3 1 L

1 g H2O at 1ᵒC 1 mL

How do we convert from ANY metric prefix to another?

Example: How many µL is 25 mL?

1. First determine what we have and what we want to get

2. Take the number we have and multiply by its multiple

3. Divide by the multiple of the unit we want to get

4. Write the result with the new unit

1. First determine what we have and what we want to get

25 mL = ______ µL

2. Take the number we have and multiply by its multiple

25 mL = 25 x 10-3 L

3. Divide by the multiple of the unit we want to get

9

25 × 10−3

10−6

4. Write the result with the new unit

25 × 10−3

10−6= 25 × 103 = 2.5 × 104𝜇𝐿

Example: 92 meters is how many centimeters?

Example: How many milliliters is 27.4 L?

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WS #4 (Learning Target 1-5. Convert from one metric unit to another using a metric scale.)

Directions: Perform the following conversions.

1. 42 mg = ______________ cg

2. 1.385 x 102 mm = ______________ μm

3. 87.2 cg = ______________ kg

4. 4.67 x 104 km = ______________ dm

5. 88.5 g = ______________ mg

6. 7.43 x 105 m = _______________ pm

7. 687 500 000 ng = ______________ mg

8. 1.30 x 106 mm = ______________ cm

9. 0.048 5 kg = ______________ g

10. 3.95 x 10-3 km = ______________ m

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1-6. Convert temperature between Fahrenheit, Celsius, and Kelvin.

Conversion Factors

℉ =9

5℃ + 32

℃ =5

9(℉ − 32)

𝐾 = ℃ + 273.15

Sample Problem: Convert 10 ᵒC to ᵒF.

Sample Problem: Convert 32 ᵒF to ᵒC.

Sample Problem: Convert -50o C to Kelvin.

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WS #5 (Learning Target 1-6. Convert temperature between Fahrenheit, Celsius, and Kelvin.)

Convert the following to Fahrenheit

1) 10o C ________

2) 30o C ________

3) 40o C ________

4) 37o C ________

5) 0o C ________

Convert the following to Celsius

6) 32o F ________

7) 45o F ________

8) 70o F ________

9) 80o F ________

10) 90o F ________

11) 212o F ________

Convert the following to Kelvin

12) 0o C ________

13) -50o C ________

14) 90o C ________

15) -20o C ________

Convert the following to Celsius

16) 100o K ________

17) 200o K ________

18) 273o K ________

19) 350o K ________

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1-7. Convert from one metric unit to another using conversion factors.

1. Dimensional Analysis

1. Sample Problem: How many hours are in a year?

2. Sample Problem: You're throwing a pizza party for 15 and figure each person might eat 4 slices. You call up the pizza place

and learn that each pizza will cost you $14.78 and will be cut into 12 slices. How much is the pizza going to cost you?

3. Sample Problem: How many months are in 5 years? (Use dimensional analysis to solve)

Wanted:

Given:

Conversion Factor(s):

? ___________ = ____________ x

4. Sample Problem: How many liters are in 350 mL?

Wanted:

Given:

Conversion Factor(s):

? ___________ = ____________ x

5. Sample Problem: How many hours are equal to 390 min?

Wanted:

Given:

Conversion Factor(s):

? ___________ = ____________ x

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WS #6 (Learning Target 1-7. Convert from one metric unit to another using conversion factors. )

Make the following conversions.

1) 484 days to years

2) 125 mL to liters

3) 5 x 103 kg to grams

4) 0.12 hrs to min.

5) 1.35 nm to meters

6) 25.3 millimol to mol

7) How many yards are there in a mile (1 mile = 5280 feet)?

8) How many hours are there in a fortnight (1 fortnight = 14 days)?

9) 0.025 km to cm

10) 923 cL to GL

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1-8. Convert between imperial and metric units using conversion factors.

Length Mass Volume Area

1 inch = 2.54 centimeters 1 mile = 5280 ft = 1.61 kilometers

1 kilogram = 2.204 pounds 1 ounce = 28.35 grams

1 gallon = 4.55 Liters 1 milliliter = 1 cm3

1 hectare = 10000 m2 = 2.47 acres

Sample Problem: Convert 43 miles into feet

Sample Problem: Convert 22,647 inches into miles

1-9. Convert rate expressions from one set of units to another.

Sample Problem: Convert 60 mph to km/min and meters/second.

Sample Problem: Convert 22 m/s to km/min and mph.

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1-10. Convert squared and cubed terms to different units.

Sample Problem: What is the area in m2 of a 25.0 cm2 box?

Sample Problem: What is the volume in m3 of a 5.93 dm3 cube?

Sample Problem: Convert the volume of this rectangular pyramid to meters3.

(Volumerectangular pyramid = base x width x height x ⅓)

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WS #7 (Learning Targets 1-8, 1-9, 1-10. Conversions.)

1. The distance from Los Angeles to San Diego is approximately 120 miles. What is the distance in kilometers?

______ km

2. A tropical rain forest has an annual rainfall of 635 cm. What is the rainfall in inches? In meters?

______ in

______ m

3. A bag holds 2.5 kilograms of sugar. What is the mass in pounds? In grams? In ounces? (round to the nearest whole number).

______ lb

______ g

______ oz

4. A bottle holds 910 mL of liquid. What is the volume in centimeters cubed? In Liters? In gallons?

______ cm3

______ L

______ gallons

5. The distance from the earth to the sun (one astronomical unit) is 1.50 x108 km. What is this distance in miles? (give the answer in scientific notation with two decimal places)

______ miles

6. A drink container holds 1.82 L. What is the volume in gallons? In milliliters? In centimeters cubed? In meters cubed? (give the answer in scientific notation)

______ gallons

______ mL

______ cm3

______ m3

7. A corn field measures 2000 m by 1500 m; what is the area in meters squared? (give the answer in scientific notation) In kilometers squared? In hectares? In acres?

______ m2

______ km2

______ hectares

______ acres

8. A sheet of paper measures 8.5 inches by 11 inches; what is the area in inches squared? In centimeters squared? (give the answer to one decimal place)

______ in2 ______ cm2

9. A car is travelling at 80 miles per hour (miles/h). What is the speed in kilometers per hour (km/h)? In meters per second (m/s)? (round answers to the nearest tenth)

______ km/h ______ m/s

10. The speed of light is 3.0x108 meters per second (m/s). What is the speed in miles per hour (miles/h)? (give the answer in scientific notation with two decimal places)

______ miles/h

Answers: (1) 193.2 km (2) 250 inches, 6.35 m (3) 5.51 pounds, 2500 g, 88 ounces (4) 910 cm3 , 0.91 L, 0.2 gallons (5) 9.32x107 miles (6) 0.4 gallons , 1820 mL, 1820 cm3,

1.82x10−3 m3 (7) 3x106 m2 , 3 km2 , 300 hectares, 741 acres (8) 93.5 inches squared, 603.2 cm2 (9) 128.8 km/h, 35.8 m/s (10) 6.71x108 miles/h

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1-11. Read metric scales to the correct number of significant figures and with an appropriate unit.

Sample Problem: Read the balance to the appropriate number of digits.

Sample Problem: Read the graduated cylinder to the appropriate number of digits.

Sample Problem: Read the ruler to the appropriate number of digits.

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WS #8 (Learning Target 1-11. Read metric scales to the correct number of significant figures and with an appropriate unit.)

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1.12 Report values and round numbers to the correct number of significant figures.

1. What are significant figures

a. Significant Figures refer to measured numbers only.

i. Counting numbers have infinite sig figs. (example: 12 eggs, 10 people, circumference = 2𝜋r)

b. For the purposes of significant figures there are two major categories:

i. Nonzero digits: 1,2,3,4,5,6,7,8,9

ii. Zero digits: 0

2. Rules for counting significant figures

a. All nonzero digits are significant

b. Zero digits

a. Leading zeros (Ex. 0.0123)

i. Leading zeros are never significant.

b. Trapped zeros (Ex. 303)

i. Trapped zeros are ALWAYS significant.

c. Trailing zeros (Ex. 200)

i. Trailing zeros are significant ONLY when the number has a decimal.

a. No decimal (Ex. 200 has 1 sig fig)

b. Decimal (Ex. 200. has 3 sig figs)

Sample Problem. How many significant digits in each of the following: 1.027 s 0.0076 g 12 oranges 3000 feet 73 people

Sample Problem. How many significant digits in each of the following: 3.0800 0.00418 7.09 x 10-5 91,600 0.003005 3.200 x 109 250 780,000,000 0.0101 0.00800

3. To round to the correct number of sigfigs, count sigfigs from the left, and see if the digit just to the right of the sigfigs

rounds the previous digit up.

Sample Problem: Write 168 in one significant figure

Sample Problem: Round to 2 sigfigs: 0.368

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Sample Problem: Write 3,284 in 2 significant figures.

Sample Problem: Write 3,284 in 1 significant figure.

Sample Problem: Write the following in 2 significant digits.

345 380 521 183 189 275 9652 925 18600 1345

Sample Problem: Write the following in 3 significant digits.

43256 5889 2346 3586 6745 1693 2943 134,565 6755 6.890 x 102

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WS #9 (Learning Target 1-12. Report values and round numbers to the correct number of significant figures.)

1. How many significant figures are in each of the following numbers?

a. 225.0 b. 1000.0

c. 0.0003210 d. 0.0067

e. 1,000,000. f. 2.00001

2. Round off each of the following numbers to three significant figures

a. 15.9994 b. 1.0080

c. 0.6654 d. 4885

e. 87,550 f. 0.027225

3. Put the following numbers into scientific notation and write down the number of significant digits in each:

a. 0.225 b. 2.5

c. 44,163 d. 20,190

e. 0.00000000000991 f. 7,000

5. Express the following fractions in decimal form to 2 significant digits.

a. ¼ b. 2/3

c. 5/3 d. 13/5

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1.12 Report values and round numbers to the correct number of significant figures.

4. Reporting Calculations in significant figures

a. Multiplying/Dividing

i. Record your answer in the least number of sigfigs.

b. Addition/Subtraction

i. Record your answer in the least number of decimal places.

Sample Problem: 7.8 x 1000 = ?

Sample Problem: 3.20 x 0.5 = ?

Sample Problem: 7.8 – 3.75 = ?

Sample Problem: 23.1 + 16.01 = ?

Sample Problem: 7.1 x 10-3 + 3.2503 x 10-2 = ?

Sample Problem: How many inches are in 7.86 feet?

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WS #10 (Learning Target 1-12. Report values and round numbers to the correct number of significant figures.)

1. Give the number of significant figures in each of the following.

a) 10.0005 g ______

b) 0.003423 mm ______

c) 2900 + 100 ft ______

d) 8.9 x 105 L ______

2. Determine the answer for each of the following. Be sure to use the correct number of significant figures.

a) 27.34 b) 2.8023

6.90 - 4.762

+ 13.124

c) 0.32 x 14.50 x 120 = d) 24.1 / 0.005 =

3. Round each of the following to 3 significant figures.

707.5 ____________ 2,300.2 ______________

0.0003350 _____________ 10.26730 _______________

18.95 x 1021 ______________

4. Convert each of the following into correct scientific notation.

1747 _________________________________

0.00000984 ___________________________

3200.0 x 102 ____________________________

0.002014 x 102 __________________________

25600000000000000 ______________________

5. Calculate the following using the correct number of significant figures.

a) 2.34 x 1047 b) 9132.0

+ 9.2 x 1046 - 1.6 x 103

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6. Calculate the following using the correct number of significant figures.

a) (1.54 x 1058)(3.5 x 1060)

b) (7.9 x 1034) / (8.32 x 1023)

7. Express the following numbers in scientific notation.

a). 810,000 g b). 0.000634 g c). 60,000,000 g

8. State the number of significant digits in the following measurements.

a). 3218 kg b). 60.080 kg c). 0.000534 kg

9. Add/Subtract as indicated and round the answer using the correct number of significant digits.

a). 85.26 g + 4.7 g

b). 1.07 km + 0.608 km

c). 186.4 kg – 57.83 kg

10. Multiply/Divide as indicated and round the answer using the correct number of significant digits.

a). ( 5,108 m )( 4.2107 m )

b). ( 1.67 x 10-2 km )( 8.5 x 10-6 km )

c). ( 2.6 x 104 cm )( 9.4 x 103 cm )

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1.13 Define density using a mathematical equation and give the appropriate units.

𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =𝑚𝑎𝑠𝑠

𝑣𝑜𝑙𝑢𝑚𝑒

𝑚𝑎𝑠𝑠 =

𝑣𝑜𝑙𝑢𝑚𝑒 =

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1.14 Perform calculations involving density, mass, and volume, giving answers with the appropriate units.

Sample Problem: Calculate the density in g/mL of aluminum if a 13.7 mL block weighs 37.0 g.

Sample Problem: Calculate the mass of a 100. cm3 block of silver with a density of 10.5 g/cm3.

Using the density equation:

Sample Problem: Calculate the mass of a 100. cm3 block of silver with a density of 10.5 g/cm3.

Using dimensional analysis:

Sample Problem: Calculate the volume of a 10.0 g cube of iron with a density of 7.80 g/cm3.

Using the density equation:

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Sample Problem: Calculate the volume of a 10.0 g cube of iron with a density of 7.80 g/cm3.

Using dimensional analysis:

Sample Problem: You perform an experiment in which you place a rod of an unknown metal weighing 5.40 g in a graduated cylinder full of water. The rod displaces the water 2.0 mL. Identify the unknown metal. On the next slide are densities of various substances.

Substance Density (g/mL)

Water 1.00

Aluminum, Al 2.70

Iron, Fe 7.80

Gold, Au 19.30

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WS #11 (Learning Target 1-14. Perform calculations involving density, mass, and volume, giving answers with the appropriate

units.)

Directions: Answer each of the following questions in the space provided.

𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =𝑚

𝑉

𝑚 = 𝐷𝑉

𝑉 =𝑚

𝐷

Substance Density (g/mL or g/cm3 )

water 1.00

Ethanol 0.800

Aluminum 2.70

Iron 7.86

Lead 11.34

Gold 19.30

Tin 7.31

Silver 10.50

Chromium 7.20

Copper 8.95

(1) (a) Calculate the density of a metal having a mass of 108 g and a volume of 15.0 cm3 . (b) Identify the metal.

D = ____________ metal = ____________

(2) (a) Find the density of aluminum. (b) Calculate the mass of 16.0 cm3 of aluminum.

D = ____________ m = ____________

(3) (a) Find the density of silver. (b) Calculate the volume of a silver block with a mass of 25.2 g.

D = ____________ V = ____________

(4) (a) Find the density of copper. (b) Calculate the mass of 40.0 cm3 of copper. (c) What is the mass of the copper in kg?

D = ____________ m = ____________ m = ____________kg

(5) (a) Find the density of iron? (b) What volume will 393 g of iron occupy? (c) Calculate the volume in Liters.

D = ____________ V = ____________ V = ____________L

(6) (a) A block of lead measures 2.00 cm by 5.00 cm by 5.00 cm. What is the volume of the block? (b) Find the density of lead? (c) What is the mass of the block? (d) What is the mass in kilograms?

V = ____________ D = ____________ m = ____________ m = ____________kg

(7) (a) An ice cube has sides with a length of 2.00 cm. Calculate the volume of the ice cube. (b) If the mass of the ice cube is 7.36 g, calculate the density of ice. (c) Find the density of water. (d) Ice will ________ in water because the density of ice is ________ than the density of water.

V = ____________ D = ____________ D = ____________

(8) (a) A ring weighing 38.6 g is placed in a graduated cylinder. The water in the graduated cylinder rises from 2.00 mL to 4.00 mL. What is the volume of the ring? (b) Calculate the density of the ring. (c) Of what metal is the ring composed?

V = ____________ D = ____________ metal = ____________

(9) (a) An empty beaker weighs 200 g. When the beaker is filled with ethanol it weighs 420 g. What is the mass of the ethanol? (b) Find the density of ethanol. (c) What is the volume of the ethanol in mL? (d) What is the volume in cm3

m = ____________ D = ____________ V = ____________ ? = ____________cm3

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(10) (a) An empty graduated cylinder has a mass of 150.0 g. When 40.0 mL of acetone are poured into the graduated cylinder, the mass increases to 181.6 g. What is the mass of the acetone? (b) Calculate the density of acetone. (c) A piece of titanium weighing 22.7 g is placed in a graduated cylinder. The water in the graduated cylinder rises from 5.00 mL to 10.00 mL. What is the volume of the titanium? (d) Calculate the density of titanium. (e) Titanium will ________ in acetone because the density of titanium is ________ than the density of acetone.

m = ____________ D = ____________ V = ____________ D = ____________

Answers: (1) (a) 7.20 g/cm3 (b) chromium (2) (a) 2.70 g/cm3 (b) 43.2 g (3) (a) 10.50 g/cm3 (b) 2.40 cm3 (4) (a) 8.95 g/cm3 (b) 358 g (c) 0.358 kg

(5) (a) 7.86 g/cm3 (b) 50.0 mL (c) 0.0500 L (6) (a) 50.0 cm3

(b) 11.34 g/cm3 (c) 567 g (d) 0.567 kg

(7) (a) 8.00 cm3 (b) 0.920 g/cm3 (c) 1.00 g/cm3 (d) float, less (8) (a) 2.00 mL (b) 19.3 g/mL (c) gold

(9) (a) 220 g (b) 0.800 g/mL (c) 275 mL (d) 275 cm3 (10) (a) 31.6 g (b) 0.790 g/mL (c) 5.00 mL (d) 4.54 g/mL (e) sink, greater

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WS #12 (Unit 1 Review)

Directions: Complete the questions in the space provided. Show all work, with any equations used, and answer with appropriate

units and significant figures.

(1) Convert the numbers between scientific notation and standard form.

(a) 45600000 = _________________ (b) 8.3x10–4 = _________________

(c) 0.0000571 = _________________ (d) 2.8x103 = _________________

(2) Complete the following metric conversions.

(a) 0.065 L = _________________ mL (b) 0.00900 Mg = _________________ g

(c) 1.5x105 cs = _________________ ks (d) 890 nm = _________________ m

(3) Complete the following temperature conversions.

(a) 90.0 C = ____ F = ____ K (b) ____ C = –155 F = ____ K

(4) A bag holds 20.0 ounces of coffee. What is the mass of the coffee in grams? In kilograms? In pounds?

_______ g _______ kg _______ pounds

(5) A container measures 9.10 cm by 12.50 cm by 6.40 cm. What is the volume in centimeters cubed? in milliliters? in inches cubed?

in Liters? in Gallons?

_______ cm3

_______ mL

_______ inch3

_______ L

_______ Gallons

(6) An Olympic sprinter can run at 11 m/s. What is the speed in miles per hour? _______ mph

(7) Give the correct number of significant figures in each value.

(a) 5175 (b) 1.49 x 10–6 (c) 0.00075 (d) 0.720

(8) Round each value to the number of significant figures indicated.

(a) 25.34 (3) (b) 50300 (2) (c) 0.008598 (3) (d) 2.98 x 103 (2)

(9) Complete the following addition/subtraction/multiplication/division operations and give the answer to the correct number of

significant figures.

(a) 8.2 0.0952 (b) 18.918 – 0.25 (c) 0.362 x 11.65 x 2.37 (d) 12.35 + 30.278 + 16.1

(e) (2.213 + 4.1) x (9.8516 – 7.328) (f) 10.5 2.5 – 5.50 x 0.225

(10) (a) A gold bar measures 20.00 cm by 8.00 cm by 4.50 cm. What is the volume of the bar? Find the density of gold. What is the

mass of the bar in g? in kg?

volume: ___________

density: ___________

___________ g

___________ kg

32

(b) A necklace weighing 23.1 g is placed in a graduated cylinder. The water in the graduated cylinder rises from 21.15 mL to 23.35

mL. What is the volume of the necklace? What is the density of the necklace in g/mL? What metal is the necklace composed of?

What is the density of the metal in kg/L?

volume: ___________

___________ g/mL

metal: ___________

___________ kg/L

(c) An empty beaker has a mass of 70.00 g. When some acetone is added to the beaker, the mass of the beaker and the acetone is

109.30 g. What is the mass of the acetone? The density of acetone is 0.786 g/mL. What is the volume of the acetone in the beaker in

mL? in μL?

mass: _________

___________ mL

___________ μL

Answers:

(1) (a) 4.56x107 (b) 0.00083 (c) 5.71 x 10–5 (d) 2800

(2) (a) 65 mL (b) 9.00x103 g (c) 1.5 ks (d) 0.890 m

(3) (a) 194 F, 363 K (b) –104 ºC, 169 K

(4) 567 g, 0.567 kg, 1.25 pounds

(5) 728 cm3, 728 mL, 44.4 inch3 , 0.728 L, 0.160 Gallons

(6) 25 miles/hour

(7) (a) 4 (b) 3 (c) 2 (d) 3

(8) (a) 25.3 (b) 5.0 x 104 (c) 0.00860 (d) 3.0 x 103

(9) (a) 86 (b) 18.67 (c) 10.0 (d) 58.7 (e) 16 (f) 3.0

(10) (a) 720 cm3 , 19.30 g/cm3, 1.39 x 104 g, 13.9 kg (b) 2.20 mL, 10.5 g/mL, silver, 10.5 kg/L (c) 39.30 g, 50.0 mL, 5.00x104 μL