Unit 5, Worksheet 1— Relative Mass Relative Mass From...

Modeling Chemistry TN Modeling Curriculum Committee Pope John Paul II High School

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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 1— Relative Mass

Relative Mass From Gases We have established that the combining ratio of gases can be explained if two

assumptions are made:

1. Equal volumes of gases contain the same number of molecules at the same pressure and temperature.

2. Some pure elemental gases are clustered into pairs to form diatomic molecules.

Now, you should remember that back in Unit 1 we found that iron was more

dense than aluminum. Two possible models arose to account for this difference.

A. The masses of Al and Fe atoms are about the same, but there are more atoms of Fe than atoms of Al in each cm3 sample.

B. One cm3 samples of Fe and Al contain about the same number of atoms, but the Fe atoms are more massive.

A third possibility – that both the size and the mass of the atoms of these two

elements were different – also came up. At the time, we did not have enough evidence to make a decision about these possible models.

While the reason for density variation between particle types is difficult to

determine for liquids and solids, a conclusion can be reached more easily for gases due to the fact that particles in a gas are widely spaced. This means that particle size does not have an effect on the volume that a given number of gaseous particles occupy.

This makes the determination of the relative mass of individual particles in a gas

fairly simple. If one liter of gas A weighs 5 times as much as one liter of gas B (at the same T and P) we assume that each particle of gas A weighs 5 times as much as each particle of gas B. Work through the following example to test your understanding of this concept.

1. The density of oxygen gas at standard temperature and pressure is 1.43 g/liter,

whereas the density of hydrogen gas under these conditions is 0.089 g/liter. How many times more massive is one molecule of oxygen than one molecule of hydrogen? Explain your reasoning.

You shouldn’t conclude that chemists were able to determine the molar masses of

all the elements using this technique. Measurements of the density of the gaseous phase of many of the elements would be difficult, if not impossible.


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However, we are going to see that chemists could use another tool – the percent composition of compounds to determine molar masses.

Relative Mass From Compounds Many substances combine with oxygen to form a type of compound called an

oxide. We have already seen that such combinations often occur in multiple proportions. John Dalton made the assumption that the lowest ratio was a 1:1 combination of elements. For now, we will make a similar assumption. We may have to re-examine this assumption later.

2. Based on the % composition of each substance in the table, calculate the mass of each of the following elements that would combine with 100 grams of oxygen.

Mass of element

Mass of oxygen

Mass of element that would combine with

100g of oxygen

Answers from Question 4

11.11 g of H 88.89 g of O 1.0

42.86 g of C 57.14 g of O

46.67 g of N 53.33 g of O

77.72 g of Fe 22.28 g of O

92.59 g of Hg 7.41 g of O

93.10 g of Ag 6.90 g of O

3. If these elements combine in a 1:1 ratio in these compounds, we would now have

relative masses for these elements. One could conclude that each atom of oxygen

is

€

10012.5

= 8.0 times as heavy as an atom of hydrogen. Water presents a problem,

however; you have seen evidence that two atoms of hydrogen combine with one atom of oxygen in water. This makes the relative mass of oxygen twice as great as this value. Explain.

4. If we choose 1.0 g of hydrogen as the standard weighable amount – 1 mole – then

the mass of a mole of oxygen atoms must be 16.0 g. Now, using proportional reasoning, use the mass of the elements relative to 100 g of oxygen you have calculated in the table above to determine the mass of one mole of each these elements. Record these values in the table above. How do these molar masses compare with the values you find in the periodic table? Is the assumption that we made at the outset valid for all of these compounds? Explain.

Name_______________________________________


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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 2— Relative Mass Activity

Purpose The purpose is to determine the relative mass of different kinds of hardware

and to learn to count by massing.

Data Hardware Mass (g) Empty vial

Vial + Washers

Vial + Hex Nuts

Vial + Bolts

Calculations Work must be shown and quantities labeled with units and appropriate sig figs.

1. A box of hardware contains 100 pieces. Assuming there are 4 pieces in each vial, calculate the mass of a box of each kind of hardware. Express these values in units of g/box.

Washers: Nuts: Bolts: 2. If you had 1.00 kg of each kind of hardware, how many boxes of each would you

have? Washers: Nuts: Bolts:


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3. You learned that a barrel of the 1” bolts had a mass of 65.2 kg. The mass of the barrel was 9.6 kg. How many boxes of bolts are in the barrel?

4. Someone at the Home Depot tells you that a 2” bolt is 6.75 times as heavy as a

washer. What would be the mass of a box of such bolts? 5. Suppose that you were given the job of shipping 25,000 hex nuts to a customer.

How many boxes of hex nuts would this be? All you have is a hanging scale and a barrel of hex nuts. Describe how you could determine the proper number of pieces without physically counting them out.

Conclusion Do you agree or disagree with the following statement? Support your answer. “You can count by weighing.” Extension

Each vial contains the same number of pieces of hardware. Calculate the relative mass of each kind of hardware. Divide each mass by the mass of the smallest. (The smallest will be 1.00)

Relative mass: Nuts _____ : Bolts _____ : Washers _____ Suppose that the washer represented an atom of the element carbon. From your

relative masses, determine the elements that would be represented by the nut and the bolt.


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What is a Mole in Chemistry?

Author: Dr.Badruddin Khan, teaches Chemistry in the University of Kashmir, Srinagar, India

Molecules and atoms are extremely small objects, both in size and mass. Consequently, working with them in the laboratory requires a large collection of them. How large does this collection need to be? A standard needs to be introduced. In chemistry, a mole is a certain number of particles, usually of atoms or molecules. This number is very special in chemistry and is given the name Avogadro's number, in honor of Italian chemist and physicist Aamadeo Avogadro, who first suggested the concept of a molecule. In theory, one could use any number of different terms for counting particles. For example, one could talk about a dozen (12) particles or a gross (144) of particles. The problem with these terms is that they describe far fewer particles than one usually encounters in chemistry. A unit like the mole is needed because of the way chemists work with and think about matter. A "unit" is the smallest measurable entity in the substance, generally either an atom or a molecule. A mole is the quantity of a substance that contains 6.02 x 1023 units. One mole of a substance is equal to the substance's atomic weight (the average weight of an atom of an element) or molecular weight, in grams.

When chemists work in the laboratory, they typically handle a few grams of a substance. They might mix 15 grams of sodium with 15 grams of chlorine. But when substances react with each other, they don't do so by weight. That is, one gram of sodium does not react exactly with one gram of chlorine. Instead, substances react with each other atom-by-atom or molecule-by-molecule. In the above example, one atom of sodium combines with one atom of chlorine. This ratio is not the same as the weight ratio because one atom of sodium weighs only half as much as one atom of chlorine. A mole (symbol Mol) is the base unit of quantity of a substance in the metric system.

The standard mole is based upon the carbon-12 isotope. Careful measurements yield a value for NA = 6.0221367x1023. This is an incredibly large number, almost a trillion trillion. We may say that a convenient name is given when there is an Avogadro's number of objects; it is called a "mole". Now the mole concept is no more complicated than the more familiar concept of a dozen. The mass of a mole of objects will be huge if we consider a mole of objects of appreciable size such as pennies; however a mole of atoms or molecules is a different story. We know that the atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom. Consequently we have the relation: NA x 12 amu = 12 g. Thus, a mole of carbon-12 atoms has a mass of just 12 g.


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The mole unit, then, acts as a bridge between the level on which chemists actually work in the laboratory (by weight, in grams) and the way substances actually react with each other (by individual particles, such as atoms). One mole of any substance—no matter what substance it is—always contains the same number of particles: the Avogadro number of particles. Let us think of what this means in the reaction between sodium and chlorine. If a chemist wants this reaction to occur completely, then exactly the same number of particles of each must be added to the mixture. That is, the same number of moles of each must be used. One can say: 1 mole of sodium will react completely with 1 mole of chlorine. It's easy to calculate a mole of sodium; it is the atomic weight of sodium expressed in grams. And it's easy to calculate a mole of chlorine; it is the molecular weight of chlorine expressed in grams. This conversion allows the chemist to weigh out exactly the right amount of sodium and chlorine to make sure the reaction between the two elements goes to completion. Even the tiniest speck of sodium chloride (table salt), for example, contains trillions and trillions of particles. The term mole, by contrast, refers to 6.022137 × 1023 particles which when written out in the long form, is 602,213,700,000,000,000,000,000 particles.

The term mole involves the acceptance of two dictates, the scale of atomic masses and the magnitude of the gram. Both have been established by international agreement. Formerly, the connotation of "mole" was "gram molecular weight." Current usage tends to apply the term "mole" to an amount containing Avogadro's number of whatever units are being considered. Thus, it is possible to have a mole of atoms, ions, radicals, electrons, or quanta. This usage makes unnecessary such terms as "gram-atom," "gram-formula weight," etc. All stoichiometry essentially is based on the evaluation of the number of moles of substance. The most common involves the measurement of mass. The convenient on gases are pressure, volume, and temperature. Use of the ideal gas law constant R allows direct calculation of the number of moles: n=P V/R T. T is the absolute temperature, R must be chosen in units appropriate for P, V, and T. The acceptance of Avogadro's law is inherent in this calculation.


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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 3— What is a mole? Reading Questions

1. What is a mole in chemistry?

2. Provide another example from personal experience of a counting term.

3. Why are chemists not able to use the term “gross” (144) for counting particles?

4. Why are mass ratios not a good way of determining the definite ratio of elements in a compound?

5. What is Avogadro’s number? What does it represent?

6. What is another way of expressing or measuring a mole?

7. How is one mole of aluminum the same as one mole of iron? How are they different?


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Veritas:______________________________________

Unit 5, Worksheet 4— Size of a Mole

To help you better visualize the enormous size of Avogadro's number, 6.02 x 1023,

consider the following analogies: 1. If we had a mole of rice grains, all the land area of the earth would be covered

with rice to a depth of about 75 meters! 2. One mole of rice grains is more grain than the number of all grain grown since

the beginning of time. 3. One mole of marshmallows (standard 1 in3 size) would cover the United States to

a depth of 650 miles. 4. If the Mount St. Helens eruption had released a mole of particles the size of sand

grains, the entire state of Washington would have been buried to a depth equal to the height of a 10-story building.

5. A mole of basketballs would just about fit perfectly into a ball bag the size of the

earth.

Your turn: Show your work including units. Use dimensional analysis and show the cancellation of units. Keep 2

sf’s in your answers. 1. Assuming that each human being has 60 trillion body cells (6 x 1013) and that the

earth's population is 7 billion (7 x 109), calculate the total number of living human body cells on this planet. Is this number smaller or larger than a mole? Divide the larger value by the smaller to determine the relative size of the two values.


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2. One of the fastest supercomputers can perform about 12 teraflops (1 teraflop is 1012 calculations per second). Determine how many seconds it would take this computer to count a mole of things. Convert this figure into years.

3. If you started counting when you first learned how to count and then counted by

ones, eight hours a day, 5 days a week for 50 weeks a year, you would be judged a 'good counter' if you could reach 4 billion by the time you retired at age 65. If every human on earth (about 7 x 109) were to count this way until retirement, what fraction of a mole would they count?


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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 5— Gram – Mole – Particles Conversions

1. An old (pre-1987) penny is nearly pure copper. If such a penny has a mass of 3.3 g, how many moles of copper atoms would be in one penny?

2. Four nails have a total mass of 4.42 grams. How many moles of iron atoms do

they contain? 3. A raindrop has a mass of 0.050 g. How many moles of water does a raindrop

contain? 4. What mass of water would you need to have 15.0 moles of H2O? 5. One box of Morton’s Salt contains 737 grams. How many moles of sodium

chloride is this? 6. A chocolate chip cookie recipe calls for 0.050 moles of baking soda (sodium

bicarbonate, NaHCO3). How many grams should the chef mass out?

7. Rust is iron(III) oxide (Fe2O3). The owner of a l959 Cadillac convertible wants

to restore it by removing the rust with oxalic acid, but he needs to know how many moles of rust will be involved in the reaction. How many moles of iron(III) oxide are contained in 2.50 kg of rust?


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8. First-century Roman doctors believed that urine whitened teeth and also kept them firmly in place. As gross as that sounds, it must have worked because it was used as an active ingredient in toothpaste and mouthwash well into the 18th century. Would you believe it’s still used today? Thankfully, not in its original form! Modern dentists recognized that it was the ammonia that cleaned the teeth, and they still use that. The formula for ammonia is NH3. How many moles are in 0.75 g of ammonia? How many molecules?

9. Lead (II) chromate, PbCrO4, was used as a pigment in paints. How many moles

of lead chromate are in 75.0 g of lead (II) chromate? How many atoms of oxygen are present?

10. The diameter of the tungsten wire in a light bulb filament is very small, less

than two thousandths of an inch, or about 1/20 mm. The mass of the filament is so very small – 0.0176 grams – that it would take 1,600 filaments to weigh an ounce! How many tungsten atoms are in a typical light bulb filament?

11. Two popular antacids tablets are Tums and Maalox. The active ingredient in

both of these antacids is calcium carbonate, CaCO3. Tums Regular Strength

tablets contain 0.747 g and Maalox tablets contain 0.600 g of calcium carbonate. Compare the number of formula units of calcium carbonate in both Tums and Maalox.


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Unit 5, Lab 1

Title

Introduction In this experiment, a measured amount of magnesium will be allowed to react with oxygen in the air. A product of the reaction is magnesium oxide. You will obtain data that will enable you to determine the empirical formula of magnesium oxide, MgxOy. “Empirical” means “based on experimental evidence”.

Question

Safety 1. Wear safety goggles at all times. 2. Handle magnesium ribbon with forceps, not your hands 3. If magnesium begins glowing very bright, do not look directly at it. It

could damage your eyes. 4. Never turn your back to Bunsen burner flame. Long hair should be

pulled back. 5. Handle hot crucible and lid with tongs. 6. Crucible and lid will break if dropped. Be careful when handling!

Procedure 1. Clean your crucible (used crucibles will not become perfectly clean) and

then rinse with dH20 (blue bottle). 2. Heat crucible for a couple of minutes to dry. Remove from flame and

allow crucible to return to room temperature. 3. Record the mass of the clean, dry crucible and lid. 4. Polish, with steel wool, 0.15-0.20g of magnesium ribbon.

The ribbon should be a bright grey color when finished. 5. Loosely curl ribbon so it will lie in the bottom of the

crucible. Do not wad the ribbon or curl it too tightly. 6. Record the mass of the crucible, lid, and magnesium

ribbon. 7. Place the crucible in a clay triangle (on a tripod) as

shown in the picture to the right. 8. Using a Bunsen burner, slowly heat the sample. Be

sure to leave the lid slightly ajar. (See picture) 9. Occasionally, lift the lid to allow more air into the

crucible (see picture). Too much air will cause the ribbon to glow brightly. If this happens, replace the lid immediately.


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10. Continue to heat until no noticeable change is observed in the ash at the bottom of the crucible.

11. Remove the lid and heat for an additional 30 seconds. 12. Stop heating and allow the crucible to return to room temperature. 13. Record the mass of the crucible, lid, and ash on the same balance you

used previously. 14. Add 2-3 drops of water (blue bottle) and reheat the crucible for 1-2

minutes. No water should be remaining in crucible. 15. Once the crucible has cooled to room temperature, record the mass of

the crucible, lid, and ash. 16. If the mass has changed, reheat the crucible for another 1-2 minutes.

Allow to cool and mass again. 17. Clean crucible and lid. Make sure all materials are put back in their

appropriate locations. Wipe down lab table. Check table and floor for trash.

Data Create a data table for the data recorded in lab. Observations ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________


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Calculations Show all work. All numbers should have units and the appropriate number of sig figs. 1. Determine the mass of magnesium reacted. 2. Determine the mass of magnesium oxide (product). 3. Determine the mass of oxygen in the magnesium oxide. 4. Determine the number of moles of magnesium 5. Determine the number of moles of oxygen. 5. Determine the mole-to-mole ratio of Mg to O:

Conclusion 1. We know that atoms of compounds combine in simple, whole number

ratios (ex. H2O). What do you think is the likely ratio for the compound containing magnesium and oxygen?


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2. What is the empirical formula of magnesium oxide? 3. How did you decide on that formula? 4. What are some possible lab errors that could lead to an incorrect formula?

Your lab errors must be logical given your mole-to-mole ratio. How, if possible, could those errors be avoided?


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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 6— Empirical and Molecular Formulas

Show all your work when solving the following problems. Circle the final answer. Be sure to include units and the correct number of significant figures. 1. Find the empirical formula of a compound containing 32.00 g of bromine and 4.90

g of magnesium. 2. What is the empirical formula of a carbon-oxygen compound, given that a 95.2 g

sample of the compound contains 40.8 g of carbon and the rest oxygen? 3. A compound was analyzed and found to contain 9.8 g of nitrogen, 0.70 g of

hydrogen, and 33.6 g of oxygen. What is the empirical formula of this compound? 4. A compound composed of hydrogen and oxygen is found to contain 0.59 g of

hydrogen and 9.40 g of oxygen. The molar mass of this compound is 34.0 g/mol. Find the empirical and molecular formulas.


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5. A sample of iron oxide was found to contain 1.116 g of iron and 0.480 g of oxygen. Its molar mass is roughly 5 times as great as that of oxygen gas. Find the empirical formula and the molecular formula of this compound.

6. Find the percentage composition of a compound that contains 17.6 g of iron and

10.3 g of sulfur. The total mass of the compound is 27.9 g. 7. Find the percentage composition of a compound that contains 1.94 g of carbon,

0.48 g of hydrogen, and 2.58 g of sulfur in a 5.00 g sample of the compound. 8. What is the % by mass of oxygen in Mg(NO3)2 ?


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Name:______________________________________

Veritas:______________________________________

Unit 5, Worksheet 7— More Empirical and Molecular Formulas

Show all your work when solving the following problems. Circle the final answer. Be sure to include units and the correct number of significant figures. 1. The compound benzene has two formulas, CH and C6H6. Which of these is

the empirical formulas and which is the molecular formula? _______________ empirical formula _______________ molecular formula 2. There are two common oxides of sulfur. One contains 32 grams sulfur and 32

grams oxygen. The other oxide contains 32 grams sulfur and 48 grams oxygen. What are the empirical formulas for the two oxides?

3. A form of phosphorus called red phosphorus is used in match heads. When

0.062 grams of red phosphorus burns, 0.142 grams of phosphorus oxide is formed. What is the empirical formula of this oxide?


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4. A certain compound is composed of 7.20 grams carbon, 1.20 grams hydrogen, and 9.60 grams oxygen. The molecular mass of the compound is 180.0 g/mol. What is the empirical formula and the molecular formula for this compound?

5. Oxalic acid is a compound used in cosmetics and paints. A 0.725 gram

sample of oxalic acid was found to contain 0.194 grams carbon, 0.016 grams hydrogen, and 0.516 grams oxygen. If the molecular mass of oxalic acid is 90.04 g/mol, what is the molecular formula for oxalic acid?


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Name:______________________________________

Veritas:______________________________________

Unit 5 — More Practice Problems 1. Definitions

a. mole

b. molar mass

c. Avogadro’s number

d. empirical formula

e. molecular formula 2. Find the molar mass of the following (include units):

a. KNO3 ______________ h. UF6 ______________

b. (NH4)2CO3 ______________ h. UF6 ______________

c. Ag2CrO4 ______________ j. H3PO4 ______________

d. oxygen gas ______________ k. (NH4)2SO4 ______________

e. Ca(NO3)2 ______________ l. CH3COOH ______________

f. PbSO4 ______________ m. Pb(NO3)2 ______________

g. Mg(OH)2 ______________ n. Ga2(SO3)3 ______________

3. Consider the masses of various hardware below.

Type Mass (g) Relative mass

Washer 1.74

Hex nut 3.16

Anchor 3.00

Bolt 7.64 a. Do the calculations necessary to complete the table.

b. Explain the connection between these calculations and the atomic masses in

the Periodic Table.


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4. Convert from g à moles or from moles à g. Show units cancelling. a. 12.0 g Fe x = moles b. 25.0 g of Cl2 gas x = moles

c. 0.476 g of (NH4)2SO4 x = moles

d. 0.15 moles NaNO3 x = g e. 0.0280 moles NO2 x = g f. 0.64 moles AlCl3 x = g

g. Convert 30 grams of H3PO4 to moles.

h. Convert 25 grams of HF to moles. i. Convert 110 grams of NaHCO3 to moles.

j. Convery 4 moles of Cu(CN)2 to grams.

k. Convert 5.6 moles of C6H6 to grams.

l. Convert 21.3 moles of BaCO3 to grams.

m. If you had 2.50 moles of oxygen gas, what mass of the gas would be in the sample?


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5. Use Avogadro’s number to do the following conversions.

a. How many atoms are there in 0.00150 moles Zn?

b. A 4.07 g sample of NaI contains how many atoms of Na?

c. How many atoms of chlorine are there in 16.5 g of iron (III) chloride, FeCl3?

d. What is the mass of 100 million atoms of gold? Could you mass this on a

balance?

e. How many molecules are there in 24 grams of FeF3?

f. How many molecules are there in 450 grams of Na2SO4?

g. How many grams are there in 2.3 x 1024 atoms of silver?

h. How many grams are there in 7.4 x 1023 molecules of AgNO3?

i. How many grams are there in 7.5 x 1023 molecules of H2SO4?

j. How many grams are there in 4.5 x 1022 molecules of Ba(NO2)2?

k. How many molecules are there in 9.34 grams of LiCl?


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6. Empirical and Molecular Formulas

a. What is the molecular formula of each compound? Empirical Formula Actual Molar Mass of Molecular Formula Compound CH 78 g/mole NO2 92 g/mole

b. Calculate the empirical formula of a compound that contains 4.20 g of nitrogen and 12.0 g of oxygen.

c. When 20.16 g of magnesium oxide reacts with carbon, carbon monoxide forms and 12.16 g of Mg metal remains. What is the empirical formula of magnesium oxide?

d. A compound is composed of 7.20 g of carbon, 1.20 g of hydrogen and 9.60 g of oxygen. The molar mass of the compound is 180 g/mole. Determine the empirical and molecular formulas of this compound.


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e. What is the % by mass of oxygen in water?

f. A compound of iron and oxygen is found to contain 28 g of Fe and 8.0 g of O. What is the % by mass of each element in the compound?

g. A 5.438 gram sample, was found to contain 2.549 grams of iron, 1.947 grams of oxygen, and 0.9424 grams of phosphorus. What is its empirical formula?

h. Aniline, a starting material for urethane plastic foams, consists of C, H, and N. Combustion of such compounds yields CO2, H2O, and N2 as products. If the combustion of 9.71 g of aniline yields 6.63 g H2O and 1.46 g N2, what is its empirical formula? The molar mass of aniline is 93 g/mol. What is its molecular formula?


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i. When 2.5000 g of an oxide of mercury, (HgxOy) is decomposed into the

elements by heating, 2.405 g of mercury are produced. Calculate the empirical formula.

j. A 5.438 gram sample, was found to contain 2.549 grams of iron, 1.947 grams of oxygen, and 0.9424 grams of phosphorus. What is its empirical formula?

7. Find the molecular formula of the following compounds.

a. A compound with an empirical formula of CFBrO and a molar mass of 254.7 grams per mole.

b. A compound with an empirical formula of C2H8N and a molar mass of 46 grams per mole.

c. A compound with an empirical formula of C2OH4 and a molar mass of 88 grams per mole.

d. A compound with an empirical formula of C4H4O and a molar mass of 136 grams per mole.


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Practice Problems for Semester Exam (Units 1-5)

Part I: Measurement and Calculations 1. Demonstrate understanding of the use of measurements in science. You should

be able to apply the rules of significant figures to choose the answer with the correct number of significant figures.

a. State the number of significant figures in each of the following numbers.

1) 5.432 ________ 8) 17.20 ________ 15) 300.10 ________ 2) 31.2 ________ 9) 0.450 ________ 16) 4000.5 ________ 3) 304.1 ________ 10) 4.560 ________ 17) 2.30 x106 ________

4) 20.9 ________ 11) 3.4x10¯4 ________ 18) 1.20 x10¯8 ________

5) 0.56 ________ 12) 2.7x105 ________ 19) 22.0030 ________ 6) 0.032 ________ 13) 500 ________ 20) 4.00900 ________ 7) 34.0 ________ 14) 360 ________ 21) 34,000 ________

b. Rewrite the following numbers rounded to the indicated number of sig figs.

22) 5.67 (2 SF) ___________ 28) 31.8 (2 SF) ___________

23) 30.53 (3SF) ___________ 29) 221.351 (4 SF) ___________

24) 24.25 (3 SF) ___________ 30) 16.05 (3 SF) ___________

25) 0.04500 (3 SF) ___________ 31) 3456 (2 SF) ___________

26) 40,001 (3SF) ___________ 32) 3000.34 (2 SF) ___________

27) 77 (1 SF) ___________ 34) 44 (4 SF) ___________


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c. Express the answer in the correct number of significant figures. Label with appropriate units. a. 21.3 g = 16.384615 __________________

1.3 cm3

b. 6.34 cm2 x 1.2 cm = 6.251437 __________________ 1.217 cm

c. 13.21m x 61.5 m = 812.415 __________________

d. 21.50 cm = 2.529411765 __________________

8.50 in

e. 334.54 grams + 198 grams __________________

f. 34.1 grams / 1.1 mL __________________

g. 2.11 x 103 joules / 34 seconds __________________

h. 0.0010 meters – 0.11 m __________________

i. 349 cm + 1.10 cm + 100 cm __________________

j. 450 meters / 114 seconds __________________ k. 298.01 kilograms + 34.112 kg __________________

l. 4 m/s x 31.221 s __________________

d. You should also be able to measure the length of an object or volume of a liquid to the appropriate number of significant figures based upon the measuring instrument.

Which of the following

best expresses the width of the business card?

a. 5 cm b. 5.0 cm c. 5.05 cm d. 5.50 cm


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Which of the following best expresses the volume of the liquid in the graduated cylinder?

a. 40 mL

b. 43 mL c. 43.0 mL c. 44.0 mL d. 43.01 mL

Below each cylinder, record the volume of the liquid in the graduated cylinder using appropriate significant figures?

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2. Demonstrate proficiency in the use of scientific notation and use of dimensional analysis in metric conversions. Know the meaning of the following metric prefixes and be able to make conversions utilizing them: milli-, centi-, kilo-

Example:

€

150mm ×1m

1000mm= 0.15m

Complete the indicated conversions: a. 37 g x = mg b. 4.7 kg x = g c. 138 m x = km d. 4021 mm x = m f. 1000. cL = ? L

g. 2.66 cm = ? mm

h. 64 mm = ? cm

i. 4.32 kg = ? mg

3. Be able to convert standard (decimal) notation to scientific notation and vice

versa. Standard: Scientific: Standard: Scientific: 1300 ___________ ___________ 24,212,000 0.00155 ___________ ___________ 0.000665 ___________ 1.68 x 106 0.00332 ___________

___________ 2.73 x 10-2 314159 ___________


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4. What is the Law of Conservation of Mass? Study the mass and change lab.

a. Under what conditions did you observe a decrease in mass?

b. Explain how could the mass decrease if matter is conserved.

c. Under what conditions did you observe an increase in mass?

d. Explain how mass can increase if matter is conserved. 5. Determine the density of an object from a data table or from a graph of Mass v

Volume.

Volume (cm3) Mass (g) 1.5 11.7 3.0 24.0 4.5 35.1 6.0 48.0 7.5 58.5 9.0 70.0

a. Plot the data above.

b. Determine the density of the

substance.

c. What volume would 150g of the substance occupy? Show work; use labels. d. What mass would 5 cm3 of the substance have?

Show on the graph above how you could answer the question.

Show below how you could answer the question mathematically.


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7. Complete density calculations and conversion problems.

a. Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From this information, calculate the density of mercury.

b. A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block weighs 1587 g. From this information, calculate the density of lead.

c. What is the mass of the ethanol that exactly fills a 200.0 mL container? The density of ethanol is 0.789 g/mL.

d. Find the mass of 250.0 mL of benzene. The density of benzene is 0.8765 g/mL.

e. 28.5 g of iron shot is added to a graduated cylinder containing 45.50 mL of water. The water level rises to the 49.10 mL mark, From this information, calculate the density of iron.

Part II: The Role of Energy in Physical Change 1. Describe the ways energy is stored in solids, liquids and gases (thermal, phase,

chemical). Also describe ways energy is transferred (working, heating, radiating).

Thermal energy – energy of motion – related to the absolute temperature • Hotter molecules move more rapidly than slower ones; for a given volume

the gas will have a greater pressure due to the greater number of collisions

Phase energy – energy due to attractions between molecules; the stronger the attractions, the lower the energy of the system of particles

• Lowest for solids, greater in liquids, greatest in gas phase

2. Heat is the transfer of energy into or out of a system due to molecular collisions. Energy is transferred from hotter (faster) molecules to colder (slower) molecules.

a. Explain why the alcohol level in a thermometer rises when it is placed in a warmer fluid. (3-step process)


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b. Describe what happens (at the molecular level) when a glass of cold water warms up to room temperature.

3. Be able to draw energy bar graphs to account for energy storage and transfer in all sorts of changes.

Complete the energy bar chart for the following scenario: An ice cube tray of water at room temperature is placed into the freezer and the liquid changes to solid.

Motion before:

Motion after:

Arrangement before:

Arrangement after:

4. When energy is transferred to a sample of matter, either the particles speed up (temperature increases) or they get pulled apart (phase change), but not both at the same time. This helps account for the shape of the warming curve you got in the Icy Hot lab.

a. On the graph above label which phases are present in each portion of the curve.

b. Label the sections in which the thermal energy (Eth) of the sample is

changing (indicating an increase or a decrease). Label the sections where the phase energy (Eph) is changing (indicating an increase or a decrease).


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5. Use the equations

€

Q = mH, where Hf = 334 J g , Hv = 2260 J g to determine

the energy transferred to or from water during a phase change. Use

€

Q = mcΔt and c = 4.18 J g˚C to determine the energy transferred to or from water

during heating or cooling. Draw a heating curve and mark on the curve the beginning and ending points.

Using the appropriate equations and values answer the following:

a. How much energy would be required to bring 100g of ice at 0°C to its boiling point?

b. Suppose that during the Icy Hot lab that 65 kJ of energy were transferred to 450 g of water at 20.˚C. What would have been the final temperature of the water?

c. An ice cube tray full of ice (235g) at –7.0˚C is allowed to warm up to room

temperature (22˚C). How much energy must be absorbed by the contents of the tray in order for this to happen?


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Part III: Gases and Kinetic Theory Demonstrate knowledge of the relationships that exist among the pressure, volume, temperature, and number of molecules of a gas. You should be able to determine the pressure of a sample of gas in a flask connected to a manometer.

1. You should be able to identify the correct graphic representation of the

relationships between volume, temperature, and pressure.

a. Which graph describes the relationship between gas pressure and volume?

Explain.

b. Which graph describes the relationship between gas pressure and the Kelvin

temperature? Explain. 2. Solve problems given volume, pressure, or temperature of gases (PVTn charts).

a. A sample of carbon dioxide has a volume of 2.0 L at a temperature of –10˚C. What volume will this sample have when the temperature is increased to 110˚C. Assume that the pressure does not change and that no carbon dioxide leaks from the sample.

b. A 12.7 L sample of gas is under a pressure of 740 mm Hg at 20°C. What will be the volume of the gas if the pressure increases to 1.00 atm and the temperature drops to 0.0°C?


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3. Based on the height of a column of mercury in an open U-tube manometer, be able to compare pressures on both sides of the tube.

a. Determine the pressure in each of the flasks. Part IV: Matter and Atomic Theory

1. Describe how matter is organized.

You should be able to identify diagrams and distinguish between pure substances, atoms, molecules, elements, compounds, mixtures, gases, and solids.

a. Which diagram(s) show only molecules?

b. Which diagram(s) show pure substances?

c. Which diagram(s) show mixtures?

d. Which diagram(s) show only atoms?

e. Does Figure B represent a compound or an element? Explain.

f. Which diagram(s) show only an element?

g. Which diagram(s) show only a compound?


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2. Demonstrate knowledge of separation techniques for mixtures and compounds.

a. Distinguish between the separation techniques required for mixtures and compounds.

b. List and describe a few examples below.

c. How would you separate a mixture of water and ethanol? Describe the process and include a temperature-time graph.

d. How would you separate a mixture of salt and sand? Describe the process(es) involved.

e. How would you separate the compound water into its elemental components?

Describe the process.

Identify the apparatus used. Label which substance is found on the anode side and which substance is found on the cathode side.


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3. Demonstrate knowledge of the meaning of a chemical formula in terms of atoms and molecules. Given the formula of a compound, identify the number of atoms present. Example: Identify the number of atoms of each kind in each compound. Pb(NO3)2 Na3PO4 Al2(SO4)3

Part V. The mole concept and chemical reactions Recognize that atoms are too small to count directly. We determine how many there are in a sample by finding their mass. We use the mole to determine the number of atoms and molecules. Molar mass (on Periodic Table) is relative mass, based originally on hydrogen (lightest element).

1. Be able to determine the molar mass of a compound. Determine the molar masses (using correct SFs and giving unit of measurement!). a. Pb(NO3)2 _____________ c. BaSO4 _____________ b. MgCl2 _____________ d. oxygen gas _____________

2. Determine the number of atoms or moles using Avogadro’s number and the molar

mass of a compound. a. 12 g MgCl2 x = moles MgCl2

b. 3 moles Cl x = atoms Cl

c. How many moles are 1.20 x 1025 atoms of phosphorous?

d. How many atoms are in 0.750 moles of zinc?

e. Find the grams in 1.26 x 10-4 mol of HC2H3O2.

f. Find the number of moles of argon in 452 g of argon.


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3. Determine the empirical and molecular formulas from data. a. A compound is composed of 7.20 g of carbon, 1.20 g of hydrogen, and 9.60 g of oxygen. The molar mass of the compound is 180 g. Find the empirical and molecular formulas for this compound.

b. Aniline, a starting material for urethane plastic foams, consists of C, H, and N. Combustion of such compounds yields CO2, H2O, and N2 as products. If the combustion of 9.71 g of aniline yields 6.63 g H2O and 1.46 g N2, what is its empirical formula?

The molar mass of aniline is 93 g/mol. What is its molecular formula?

4. Percent Composition a. What is the % by mass of oxygen in water?

b. A compound of iron and oxygen is found to contain 28 g of Fe and 8.0 g of O. What is the % by mass of each element in the compound?


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Part VI. Additional Terms and Concepts

Law of Definite Proportions

Definintion:

Examples:

Law of Multiple Proportions

Definintion:

Examples:

Empedocles Importance:

Democritus Importance?

Dalton’s Atomic Theory

Components:

a. b. c. d.

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Unit 5, Worksheet 1— Relative Mass Relative Mass From...

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