We are going to focus on using your calculator when working with scientific notation. The first few practice problems we will do will "move the decimal" going back and forth from regular notation to scientific notation. However, as the picture represents, we will learn that scientific notation numbers are just like any other number and using your calculator correctly will insure success.

Scientific Notation Practice (pp 1-2):

A. Express in decimal notation:

1. 5.2 x 103

2. 9.65 x 10-4

3. 8.5 x 10-2

4. 2.71 x 104

5. 3.6 x 101

6. 6.452 x 102

7. 8.77 x 10-1

8. 6.4 x 10-3

B. Express in scientific notation:

1. 78,000

2. 0.00053

3. 250

4. 2,687

5. 16

6. 0.0043

7. 0.875

8. 0.012654

C. Compute the following:

1. (6.02 x 1023)(8.65 x 104)

2. (6.02 x 1023)(9.63 x 10-2)

3. 5.6 x 10-18/8.9 x 108

4. (-4.12 x 10-4)(7.33 x 1012)

5. 1.0 x 10-14/4.2 x 10-6

6. 7.85 x 1026/6.02 x 1023

7. (-3.2 x 10-7)(-8.6 x 10-9)

8. [(5.4 x 104)(2.2 x 107)]/4.5 x 105

9. [(6.02 x 1023)(-14.2 x 10-15)]/6.54 x 106

10. [(6.02 x 1023)(-5.11 x 10-27)]/-8.23 x 105

11. [(3.1 x 1014)(4.4 x 10-12)]/-6.6 x 10-14

12. [(8.2 x 10-3)(-7.9 x 107)]/7.3 x 10-16

13. [(-1.6 x 105)(-2.4 x 1015)]/8.9 x 103

14. (7.0 x 1028)(-3.2 x 10-20)(-6.4 x 1035)

You may have caught on through the lab that the equipment you use limits the accuracy of a calculation. For example, a metric ruler typically measures to 1 mm or 0.001 m (0.1cm). If you use a balance to measure the mass of a cube and the balance measures to 3 decimal places but the ruler you used to measure the length of the cube measures to 0.1 cm, then the density calculation can only be reported to one decimal place, not the 6 or 7 decimal places your calculator shows. The rules in the figure below point out some pretty picky "zero rules" but they can be summed it: If the number is greater than one and there is a DECIMAL, everything is significant. If the number is less than one, do not count the "placeholder" zeros after the decimal but count all digits after the zeros. If the number has no decimal and ends in zero(s), do not count the zero(s) at the end of the number.

Significant Figures (p 3) Practice: Give the number of sig figs in each of the following measurements

1. 1278.50

2. 120000

3. 90027.00

4. 0.0053567

5. 670

6. 0.00730

7. 8.002

8. 823.012

9. 0.005789

10. 2.60

11. 542000.

12. 2653008.0

13. 43.050

14. 0.147

15. 6271.91

16. 6

17. 3.47

18. 387465

Round to 3 sig figs:

19. 120000

20. 5.457

21. 0.0008769

22. 4.53619

The metric system is the only system used in scientific fields. The figure below shows the "prefixes" for the powers of ten that each represents. The figure on the right shows the most important examples of metric prefixes. Memorize these prefixes because you need to be completely comfortable with this system. We will work numerous problems to help you along.

Metric System, Practice (p 4-5):

1) 2000 mg = _______ g 2) 5 L = _______ mL 3) 16 cm = _______ mm

4) 104 km = _______ m 5) 198 g = _______ kg 6) 2500 m = _______ km

7) 480 cm = _____ m 8) 75 mL = _____ L 9) 65 g = _____ mg

10) 5.6 kg = _____ g 11) 50 cm = _____ m 12) 6.3 cm = _____ mm

13) 8 mm = _____ cm 14) 5.6 m = _____ cm 15) 120 mg = _____ g

You may have seen this method for solving problems before (picket fence, etc) but STOP RESISTING using this method because it makes everything so much easier! You can solve many problems in chemistry and physics by focusing on the UNITS instead of randomly punching numbers into your calculator until you get one of the multiple choice answers. By the way, mean instructors such as myself are on to this student method and incomplete answers (answers to part of the problem) are there to distract you on a test. So climb on the unit bandwagon and be more confident in your problem solving!

Dimensional Analysis, Practice (pp 5-6):

1. The average annual crude oil consumption in the US is 15,615 dm3. What is this value in cubic centimeters?

2. Perform each of the following conversions:

(a) 154 cm to in (b) 3.14 kg to g (c) 3.5 L to qt (d) 109 mm to in

3. A runner wants to run 10 km. She knows that her running pace is 7.5 mi per hr. How many minutes must she run?

4. A modest house has an area of 195 m2. What is the area in:

(a) km2 (b) dm2 (c) cm2

5. (a) The average U.S. farm occupies 435 acres. How many square miles is this (1 acre = 43,560 ft2; 1 mile = 5280 ft) (b) Total U.S. land area is 3,537 million square miles. What percentage of U.S. land is farmland?

You have probably used the density formula several times over the years you have been in school. It is not uncommon, however, for students to be in college and still not have a firm grasp of this concept. The picture above shows the particle view of density. Once students understand this concept AND then master the units for density, they do, then, understand density. Density is a ratio of mass and volume and as you may have found in the lab, the size of the sample for a particular material does not change the density because it is how those particles are arranged and the ratio, therefore, does not change. The units are g/mL if working with liquids; g/cubic centimeter if working with solids (mL = cubic centimeter).

Density Practice (p 7):

1) What is the volume of a tank that can hold 18 754 Kg of methanol whose density is

0.788g/cm3?

2) What is the density of a board whose dimensions are 5.54 cm x 10.6 cm X 199 cm and whose mass is 28.6 Kg?

3) CaCl2 is used as a de-icer on roads in the winter. It has a density of 2.50 g/cm3. What is the mass of 15.0 L this substance?

4) The volume of the aquarium in our classroom is 1890 L. The density of seawater is 1.03g/cm3. What is the mass of the water in our tank? Express your answer in the correct number of significant digits.

5) Spanish mahogany is a red, lustrous wood. It is prized over Honduras mahogany, which does have the characteristic red color and similar grain, but is the wood is not as compact and has a looser appearance than the Spanish mahogany. A hope chest made of Honduras mahogany has a volume of .648 m3 and a mass of 349 Kg. An identical hope chest made of Spanish mahogany has a mass of 550.8 Kg. What is the density of Honduras mahogany? What is the density of Spanish mahogany? Express your answer in the correct number of significant digits.

The Mole, practice (pp 1-4): By completing the following questions about eggs, you should be able to complete the chart below: One key piece of information is that one dozen eggs weighs 70 g on average (relationship = conversion factor).

Mass (g) <-------->Dozens<---------># of Eggs

1. How many eggs in 2 dozen? (what mathematical operation did you use?)

answer: you multiplied by 12 eggs/dozen, the number of eggs in a dozen. Put 12 eggs/dozen by (1) ABOVE.

2. How many dozen is 36 eggs? (insert the appropriate conversion factor above for (2))

3. What is the mass of 2 doz eggs? (insert the appropriate conversion above for (3))

4. How many doz is 210 g of eggs? (insert the appropriate conversion above for (4))

5. How many eggs in 210 g of eggs? (this is a repeat of which operation?)

6. What is the mass of 400 eggs? (this is a repeat of 2 steps above. Which two?)

Now we are going to change the egg chart into a mole chart.

Mass (g) <----------->Moles<------------># of Particles

Converting the chart:

# of eggs becomes # of particles. Particles can be atoms (if substance is made up of atoms), molecules (if substance is covalently bonded), or formula units (if substance is ionic).

Dozens becomes moles. The mathematical operations (and nature of conversions) we did for dozens stays the same. HOWEVER, the # of particles in a mole is not 12 but is 6.02 x 1023 (Avogadro's number). It is just a number like 12 is.

Mass is still mass but the mass of one mole will not be "constant" as it was for eggs. The mass of one mole of an element is found on the periodic table. We will add up the masses from the periodic table for compounds.

Just like you did for eggs, write in the conversion on the chart above for each new problem below:

1 (a) How many atoms in 2.0 moles of Cu? (The fact that it is Cu does not really matter as there is

the same number of atoms in one mole of any element)

(b) How many molecules in 3.0 mol of H2S? (Again, it does not really matter what compound it is as

long as the compounds are all made of molecules)

(c) How many formula units in 10.0 mol of NaCl? (Again, it does not really matter what compound it is as long as the compounds are all made of formula units)

2. (a) How many moles in 7.1 x 1024 atoms Cu?

(b) How many moles in 8.3 x 1015 molecules H2S?

(c) How many moles in 6.7 x 1028 formula units NaCl?

3. What is the mass of (a) 15. 3 mol Ag? (b) 3.7 mol NH3 USE THE PERIODIC TABLE TO GET MASS OF A MOLE!!

4. (a) How many mol in 250 g Au?

(b) How many mol in 28 g CO2?

THIS IS JUST THE OPPOSITE OPERATION FROM #3

5. How many particles in

(a) 45.0 g Mn (b) 60.0 g CuCl2

We have balances in the lab that measure mass but not moles or atoms. Mass is the "everyday" unit and the one we can actually measure. If you are trying to go from moles to mass for most elements, the molar mass is represented by the atomic mass on the PT. For compounds, you must add up the atomic masses to get the molar mass.

Practice (p 4):

1. Without doing any calculations, which of the following contains the most atoms?

a) one gram of iron b) one gram of boron c) one gram of mercury

2) A helium balloon contains 0.46 g of helium. How many moles of helium does it contain?

3) What is the mass of each elemental sample?

a) 6.64 mol W b) 0.581 mol Ba c) 68.1 mol Xe d) 1.57 mol S

4) What is the mass of 4.91 x 1021 platinum atoms?

You will have to select the mole slideshare from a list. It has a bunch of great examples from real life such as "how much quacamole can be made from a mole of avocados". There are also worked out examples.

We studied isotopes at the very beginning of this course. Now that we are getting into the math of chemistry, we can finally understand where those atomic masses that are on the PT come from and how they are related to the mass numbers we used before.

1. Three isotopes occur in nature: silicon-28 (92.23%), atomic mass 27.97693 amu; silicon-29 (4.68%), atomic mass 22.97649; and silicon-30 (3.09%), atomic mass 29.97377. Calculate the atomic weight of silicon. (practice problem #2, pg 51 in text)

problem done in video:

2. Silver has two naturally occurring isotopes (Ag-107 and Ag-109).

a. What is the atomic mass of silver from the PT?

b. If the natural abundance of Ag-107 is 51.84%, what is the natural abundance of Ag-109?

c. If the mass of Ag-107 is 106.905 amu, what is the mass of Ag-109?

See problems 2.33-2.39 (odd problems have answers in the back of the text).

It is a percentage-always going to "part"/"whole" x 100! In chemistry, this can be approached in two ways: from the formula as the picture represents or from actual data from the lab. We will do both in our practice and you will do an experiment determining the percentage of water in a hydrated compound.

1. A hydrocarbon is a compound that contains mostly carbon and hydrogen. Calculate the percent composition (by mass) of the following hydrocarbon: C 8 H 18 .

2. Calculate the percentage by mass of oxygen in the following compounds.

(a) carbon dioxide, CO2 (b) morphine, C17H19NO3 (c) digitoxin, C41H64O13

3. Silver chloride, often used in silver plating, contains 75.27% Ag. Calculate the mass of silver chloride in grams required to make 4.8 g of silver plating.

4. 37.5 g of copper II fluoride contains 14.0 g of fluorine. What is the percentage of copper and fluorine in the compound?

You can use mass percentages to arrive at "simplest ratio formula" or empirical formula. OR, you can use lab data as well just as you did for mass percentages.

The steps below are the latter approach and the figure to the right represents using mass pecentages-there is really no difference.

1. Convert mass of element to mol (whether the mass comes from lab data or mass %)

2. Divide each element amount in mol by the smallest molar amount (thereby setting subscript for the smallest element present at "one").

3. If the subscripts are not whole numbers, a multiplying factor can be used. Often, this produces the molecular formula.

4. If the molecular mass is known, the empirical mass and molecular mass can be compared to obtain the "factor".

1. A compound containing selenium and fluorine is decomposed in the laboratory and produces 2.231 g of selenium and 3.221 g of fluorine. Calculate the empirical formula of the compound.

2. Citric acid, the compound responsible for the sour taste in lemons, has a elemental composition of C, 37.51 %; H, 4.20 %; O, 58.29 %. Calculate the empirical formula for citric acid.

3. A 2.241-g sample of nickel reacts with oxygen to form 2.852 g of the metal oxide. Calculate the empirical formula of the oxide.

4. Caffeine, a stimulant found in coffee and soda, has the mass percent composition: C, 49.48%, H, 5.19%; N, 28.85%; O, 16.48%. Find the empirical formula and the molecular formula if the molar mass is 194.19 g/mol.

Chemistry is a subject that must be practiced everyday if possible. Work through the lecture examples stopping the video clips and then restarting to check yourself. Take advantage of the practice in Mastering Chemistry to give you the practice you need to be successful. DO NOT PROCRASTINATE! Check announcements on Blackboard everyday. Keep a printed copy of the most recent course calendar (included in syllabus) next to your work area. Email me with questions!!

Email: moliver@bishop.edu
Location: Bishop State CC Central Campus , Doctor Martin Luther King Junior Avenue, Mobile, AL, United States
Phone: (251)405-4504