Advanced Placement (AP) Chemistry Summer AssignmentSummer 2014

To the AP Chemistry Student:

Welcome to my AP Chemistry class!

The Advance Placement Chemistry experience is designed to provide a full year of college-level chemistry, so it places heavy demands on the student, especially in terms of the time commitment required. In fact, the College Board suggests that students devote a minimum of five hours per week for individual study outside of the classroom. The ultimate objective, of course, is to prepare you to take the AP Chemistry test on May 4, 2015, and in order to accomplish this, topics are covered very quickly. For this reason, most students take AP Chemistry after they already completed a year of high school chemistry, since that provides them with a solid foundation. In order to ensure the best start for you next fall, I prepared a Summer Assignment that reviews many basic chemistry concepts. This is a required assignment, and your first assessment in September will be drawn from topics in the Summer Assignment.

If you have already taken a high school chemistry course, you will find much of the material in the Summer Assignment to be very familiar. It is important that everyone come to the first day of class well prepared! Extensive remediation is not an option as we work towards our goal of becoming completely prepared for the AP Exam in early May, so seek help early if you are struggling.

To help you understand what is expected of you as you begin AP Chemistry, I prepared an “Expectation” sheet which is found on my website. This document describes basic skills and concepts, and it also has some tables of information that are absolutely essential for you to know.

It is also important that you realize up front how your performance in this course will be measured. The course grade will mainly depend on your assessment scores, although some lab reports will also be assigned and evaluated. Do not expect any grade curves or ‘fluff’ assignments this year! Assessments will be graded as if they are AP exams.

There is a vast amount of chemistry available on the Internet, including sources cited on my own website. With ready access to these websites in your home, school or at the local library, I am confident that you will have everything you need to learn chemistry at the AP level.

Finally, I recommend that you spread out the Summer Assignment, rather than trying to complete it in the final week of the summer! It takes time for a student to process, practice and subsequently learn chemistry at the level necessary for success in AP Chemistry. Remember, AP Chemistry is an equivalent course to an Introductory Chemistry college course, a full year program. Taking a college level course in high school is difficult, and it requires commitment, hard work and time, but completion of a class like this is a great investment in your education. Prepare yourself and arrive ready to learn!

You may contact me by email over the summer, and I will do my best to respond quickly. (ccumming@hatboro-horsham.org)

Have a great summer !

Mrs C Cummings

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Required Knowledge

At the end of this assignment, you are expected to know how to the following topics:

1. Knowing Basic element names and symbols - page 4 of packet2. Knowing the oxidation numbers of the monatomic ions included on page 5 of this

packet3. Knowing the names and oxidation numbers of the polyatomic ions listed on page 6

of this packet4. Significant digits calculations5. Converting among units with dimensional analysis6. Determining number of protons, neutrons and electrons in an atom of an element

using the periodic table7. Calculating the average atomic mass given the percent abundance and mass of each

isotope of that element8. Naming and Writing formulas for Ionic compounds9. Naming and writing formulas for molecular (covalent) compounds10. Naming and writing formulas for acids11. Calculating molar masses12. Converting among moles, grams and molecules (formula units) in a compound13. Calculating percent composition14. Determining an empirical formula when percent composition is given15. Determining a molecular formula if an empirical formula and a molar mass is

known16. Balancing Chemical equations17. Predicting products for the following types of reactions:

i. single displacementii. double displacement

iii. synthesisiv. decompositionv. combustion

18. Doing a mole-mole and mass-mass stoichiometry problems

Summer Assignment June 2014Use Internet resources, textbooks and my own reference material as you complete these problems. The URLs represent a fraction of the available chemistry addresses available. Please feel free to expand the list and find other web sites that help prepare you for the coming year. Take detailed notes, and practice the items indicated in the packet.

Useful links: http://highschoolhub.org/hub/chemistry.cfm ; http://www.chemistrycoach.com/home.htmhttp://www.collegeboard.com/ap/students/chemistry/index.htmlwww.chemmybear.com

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Topic 1:Common elements and their symbols

Remember Upper case U has no tail and lower case u does. Upper case M and N are pointed and lower case is not. Upper case I has a line at the top and the bottom so that it is not confused with a lower case l

Aluminum AlAntimony SbArgon ArArsenic AsBarium BaBeryllium BeBismuth BiBoron BBromine BrCadmium CdCalcium CaCarbon CCesium CsChlorine ClChromium CrCobalt CoCopper CuFluorine FGermanium GeGold AuHelium HeHydrogen HIodine IIron FeKrypton KrLanthanum LaLead Pb

Lithium LiMagnesium MgManganese MnMercury HgNeon NeNickel NiNitrogen NOxygen OPhosphorus PPotassium KRadium RaRadon RnSilicon SiSilver AgSodium NaStrontium SrSulfur STellurium TeTin SnUranium UXenon XeZinc Zn

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Topic 2:Common Oxidation Numbers

Group 1 Alkali Metals = +1 example – sodium ion

Group 2 Alkaline Earth Metals = +2 example – calcium ion

Group 13 = + 3 example – aluminum ion

Group 15 = -3 example – nitride ion

Group 16 = -2 example – oxide ion

Group 17 Halogens = -1 example – chloride ion

Silver Ag +1 silver ion

Zinc Zn +2 zinc ion

Hydrogen H+1 hydrogen ion

Hydride H-1 hydride ion

Multivalent metals

Copper Cu +1 = Copper (I) copper (I) ion Cu +2 = Copper (II) copper(II) ion

_______________________________________________________________________________________Iron Fe +2 = Iron (II) iron (II) ion

Fe +3 = Iron (III) iron (III) ion

cobalt Co +2 = coalt (II) cobalt (II) ion Co +3 = cobalt (III) cobalt(III) ion

chromium Cr +2 = Iron (II) chromium (II) ion Cr +3 = Iron (III) chromium (III) ion

Nickel Ni +2 = Nickel (II) nickel (II) ion Ni +3 = Nickel (III) nickel(III) ion

______________________________________________________________________________

Lead Pb +2 = Lead (II) lead (II) ion Pb +4 = Lead (IV) lead (IV) ion

Tin Sn +2 = Tin (II) tin(II) ionSn +4 = Tin (IV) tin (IV) ion

TOPIC 3: POLYATOMIC IONS+1 oxidation number

NH4+1 ammonium

-1 oxidation number C2H3O2 acetate OH hydroxide

ClO4 perchlorate CN cyanide

ClO3 chlorate MnO4 permanganate

ClO2 chlorite ClO hypochlorite

NO2 nitrite HCO3 hydrogen carbonate/bicarbonate

NO3 nitrate

-2 oxidation number

CO3 carbonate

SO3 sulfite

SO4 sulfate

CrO4 chromate

Cr2O7 dichromate

-3 oxidation number

PO4 phosphate

PO3 phosphite

AsO4 arsenate

Topic 4. Significant digit CalculationsDIRECTIONS: Solve the problems below, showing your setup, calculations, units and correct Significant Figures (SigFigs). Neatly-written solutions are acceptable. (Points are deducted from AP tests if correct SigFig rules and labels are ignored, so get into the habit of using them all of the time!)

1. Demonstrate that you know the correct use of significant figures (SigFig digits) by completing

the following:

a. 738.90 m has _____ SigFgs.

b. 0.0304 g has _____ SigFigs.

c. 1.40 X 104 joules has _____SigFigs.

d. 1 dozen donuts has _____ SigFigs.

c. 40 mL has _____ SigFigs.

f. 800.0 m has _____ SigFigs.

2. A cylinder rod formed from silicon is 21.3 cm long and has a mass of 5.00 kg. The density of silicon is 2.33 g/cm3. What is the diameter of the cylinder? (The volume of cylinder is given by V = πr2h, where r is the radius and h is the length.)

3. Calculate the following to the correct number of significant figures.a. 1.27 g / 5.296 cm3 = ________________________

b. 12.235 g / 1.01 L = __________________________

c. 12.2 g + 0.38 g = _________________________

d. 17.3 g + 2.785 g = _______________________

e. 2.1m x 3.215m = _____________________

f. 200.1mi x 120 min = __________________

g. (17.6 + 2.838 + 2.3 + 110.77)g = _________________

4. A solid white substance A is heated strongly in the absence of air. It decomposes to form a new white substance, B, and a gas, C. The gas has exactly the same properties as the product obtained when carbon is burned in an excess of oxygen. Based on these observations, can we determine whether solids A and B and the gas C are elements or compounds? Explain your conclusions for each substance.

Topic 5- Converting among metric units

1. Convert 1.09 ng to g

2. Convert 90.1 ms to s

3. Convert 145 pm to m

4. Convert 1.55 kg/m3 to g/L

5. Determine the number of L in a box that measures 4cm x 8cm x 10cm. (1 ml= 1 cm3)

Topic 6: Determining the number of protons, neutrons and electrons in an element using the periodic table

is the general symbol for an isotope where A is the mass number (number of protons + number

of neutrons), X is the element’s symbol, and Z is the atomic number (number of protons). This information can also be found on the periodic table. Use this information to fill in the chart for the following elements

Symbol Atomic Number

Atomic Mass Protons Electrons Neutrons Element Name

copper

nickel

lead

zinc

Topic 7: Calculating the average atomic mass of an element given the percent abundance and mass of each isotope

Isotopes are atoms of the same element with different numbers of neutrons, and therefore, different masses

Directions:

1. Write the isotopes.2. Write the masses.3. Multiply the masses by the decimal equivalent of the relative abundance.4. Add the products. This is the average atomic mass.

1. Naturally occurring chlorine molecules (chlorine is diatomic) have masses of 69.996, 721.9987and 73.999 amu. They occur in the percentages of 56.25%, 37.50% and 6.250%. What is the average atomic mass of chlorine atoms?

2. In Nature, strontium consists of four isotopes with masses and percent abundance of 83.9134 amu (0.50%), 85.9094 amu (9.9%) , 86.9089 amu (7.0 %) , and 87.9056 amu (82.6 %). Calculate the atomic mass of Sr.

3. Calculate the average atomic mass of Si with the following isotopic composition: 92.23% Si-28, with a mass of 27.97693 amu, 4.68% Si-29 with a mass 28.97649, and 3.09% Si-30 with a mass of 29.97377 amu.

4. Calculate the average atomic mass of Copper with the following isotopic composition: 69.17 % of Cu-63 with a mass of 62.9296 amu , and 30.83% Cu-65 with a mass of 64.9278 amu.

Topics8-10: Naming and writing formulas for ionic compounds, molecular (covalent) compounds, and acids

Case 1: Ionic compounds containing monatomic ions (i.e. ions that can only have one charge) Name of Compound = name of metal + name of non-metal w/ide suffix or name of

polyatomic ion. No prefixes are used! e.g. NaF = sodium fluoride; Na3PO4 = sodium phosphate; (NH4)3PO4 = ammonium phosphate

Case 2: Ionic compounds containing a metal that can form more than one ion Name of Compound = name of metal, followed by charge of metal in Roman numerals in

parentheses, followed by name of non-metal w/ -ide suffix or name of polyatomic ion. No prefixes are used!

e.g. PbCl2 = Lead (II) chloride; Cu(NO3)2 = copper (II) nitrateCase 3: Binary molecular compounds:

Name of Compound = name of first element + name of second element with -ide suffix. Use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, nona-, deca-) to indicate the number

of atoms. The mono prefix is not used with the first element. e.g. CO = carbon monoxide; NO2 = nitrogen dioxide; N2O = dinitrogen monoxide; P2O5 =

diphosphorus pentoxide Case 4: Binary acid solutions (i.e. binary acids dissolved in water = binary acids in aqueous solution)

Name of Compound = hydro + name of halogen w/ -ic suffix e.g. HF(aq) = hydrofluoric acid; HCl(aq) = hydrochloric acid

Unless stated otherwise assume the formula of a binary acid is for the acid dissolved in water. E.g. assume HCl = HCl(aq)

Naming Oxoacids (i.e. compound with the general formula HxMOy, where M = nonmetal) The name of an oxoacid is based on the name of the polyatomic ion from which the acid is derived.Case 5: -ate -ic If the name of the polyatomic ion ends in “-ate,” the name of the corresponding acid ends in “-ic acid.”

Polyatomic ion (-ate) Acid (-ic) sulfate = SO4

2- H2SO4 = sulfuric acid Chlorate = ClO3

1- HClO3 = chloric acidCase 6: -ite -ous If the name of the polyatomic ion ends in “-ite,” the name of the corresponding acid ends in “-ous acid.”

Polyatomic ion (-ite) Acid (-ous) sulfite = SO3

2- H2SO3 = sulfurous acid Chlorite = ClO2

1- HClO2 = chlorous acid

Write the formula.

1. Ammonium Nitrate

2. Calcium Bicarbonate or Calcium Hydrogen

Carbonate

3. Barium Chlorate

4. Dimercury (I) Iodide

5. Hydronitric acid

6. Plumbic Oxide or Lead (IV) Oxide

7. Potassium Thiocyanide

8. Perchloric acid

9. Sulfur hexafluoride

10. Zinc Hydroxide

11. Potassium Sulfite

12. Copper (I) Sulfide

13. Potassium bisulfate

14. Zinc Bromide

15. Ferric Chromate

16. Sodium Perchlorate

17. Potassium Hypochlorite

18. Magnesium Nitride

19. Sodium Permanganate

20. Potassium Permanganate

Write the name.

21. Hg2(CN)2

22. H2SO4

23. Fe(C2H3O2)2

24. KClO3

25. PbF2

26. HBr

27. N2O4

28. HgCrO4

29. Ag3PO4

30. K2Cr2O7

31. Ba3N2

32. NF3

33. CuBr2

34. (NH4)2S

35. Ca(NO3)2

36. Zn(OH)2

37. NaHCO3

38. PbO2

39. KClO4

40. Hg2I2

Topics 11-13: Calcuating Molar masses, Converting among moles, grams and atoms, and calculating percent composition

1. Calculating Molar (also formula ) mass.find the mass of a mole of an element by using the atomic mass on the periodic table. Now we are applying it to the formula for a molecular or ionic compound. Here are the steps for calculating molar formula mass:

1. Write the formula2. Write each element below it individually3. Write how many there are of that atom (look for the subscript and use that – if there is none, it is one)4. Multiply the number of that atom by its molar mass (use 2 decimal places).5. Find the total of these masses.

Example:H2OH = 2 x 1.01 g = 2.02 gO = 1 x 16.00 g = 16.00 g

18.0 2 g/mol

1. Calculate the molar mass ( g/ mol) of a. Ammonia ( NH3)

b. Baking soda ( NaHCO3)

c. Osmium Metal (Os)

2. The molecular formula of morphine, a pain-killing narcotic, is C17H19NO3.

a. What is its molar mass?

b. What fraction of atoms in morphine is accounted for by carbon?

c. Which element contributes least to the molar mass?

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3. Use dimensional analysis to Calculate the mass in grams of each of the following: (there are 6.022 x 1023 molecules or formula untis of a compond in one mole of that compound) (there are 6.022 x 1023 atoms of an element in one mole of that element)

example: 1. How many molecules of glucose (C6H12O6) in 15.0 grams of glucose?

a. 6.02 x 1023 atoms of Mg

b. 3.01 x 1023 formula units of CaCl2

c. 12.4 x 1015 molecules of formaldehyde (CH2O)

2. Mercury has an atomic mass of 200.59 amu. Calculate the a. Mass of 3.0 x 1010 atoms of mercury

b. What is this mass in nanograms?

3. a. Determine the number of molecules present in a 4.56 mol sample of methane gas (CH4) ;

b. then determine the number of hydrogen atoms in the sample.

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4. In an experiment, a student gently heated a hydrate of a copper compound to remove the water .The following data was recorded:

1. Mass of crucible, cover, and contents before heating = 23.4 g.2. Mass of empty crucible and cover = 18.82 g.3. Mass of crucible, cover, and contents after heating to constant mass = 20.94 g.

Calculate the percent by mass of water in the copper compound.

Percent composition

When you calculate the percent composition of a compound (for example – water), you are really finding out what percentage of the mass of water is hydrogen and what percentage is oxygen. To do this, you must:

1. Calculate the molar mass2. Divide each elements mass contribution by the molar mass3. Multiply this answer by 100 to make it into a percentage4. Check your answer by adding the percentages. They should add up to 100.

Example

H2O Answer:H = 2 x 1.0 g = 2.02 g / 18.02 g = 0.111 x 100 = 11.21% HO = 1 x 16.0 g = 16.00 g / 18.02 g = 0.8879 x 100 = 88.79% O

18.02 g

1. Calculate the percent composition of:

a. Aluminum hydroxide

b. iron II nitrate

2. White gold is an alloy that typically contains 45.0% by mass gold and the remainder is platinum. If 154 g of gold are available, how many grams of platinum are required to combine with the gold to form this alloy?

3. The hormone, thyroxine is secreted by the thyroid gland, and has the formula: C15H17NO4I4. How many milligrams of iodine can be extracted from 15.0 grams of thyroxine?

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Topic 14 and 15 Empirical and Molecular Formulas

1. Assume that you have 100 g of the compound and change your percentages to g.2. Convert the g to moles.3. Divide these answers by the smallest number of moles.4. Use these whole numbers as the subscripts for the formula.5. If they aren’t whole numbers, you may need to multiply each answer by a 2, 3, or 4 to make it a whole number. Use

2 if it ends close to .5, use 3 if it ends close to .33 or .67, and use 4 if it ends close to .25 or .75.

Example

A compound is 63% Mn and 37% O. What is the empirical formula?

63 g Mn x 1 mol Mn = 1.15 mol Mn 54.9 g Mn

37 g O x 1 mol O = 2.31 mol O 16.0 g O

The lowest number of moles is 1.15 mol Mn

Mn 1.15 mol Mn = 1 mol Mn 1.15 mol

O 2.31 mol O = 2 mol O 1.15 mol

The formula is MnO2

1. What is the empirical formula of a compound that contains 53.73% Fe and 46.27% of S ?

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2. Determine the empirical formula of the compounds with the following compositions by mass:a. 10. 4 % C, 27. 8% S , 61. 7 % Cl

b. 21.7 % C, 9.6 % O, and 68.7 % F

3. Arsenic reacts with chlorine to form a chloride. If 1.587 g of arsenic reacts with 3.755 g of chlorine, what is the empirical formula of the chloride?

4. Vanillin, a flavoring agent, is made up of carbon, hydrogen, and oxygen atoms. When a sample of vanillin weighing 2.500g burns in pure oxygen, 5.79 g of carbon dioxide and 1.18 g of water are obtained. Calculate the empirical formula of vanillin?

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Molecular Formulas

If you know the empirical formula of a compound and know the true molecular mass, then it is easy to find its true formula (remember that the empirical formula is just he lowest ratio of atoms in a compound).

A compound has the percent composition of 68.54% C, 8.63% H and 22.83% O. Its true molecular mass is 140 g. What is the molecular formula.

1. Find the empirical formula

68.54 g C x = 5.712 molC / 1.427 = 4.003

8.63 g H x = 8.63 molH / 1.427 = 6.05

22.83 g O x = 1.427 mol O / 1.427 = 1

Empirical formula is C4H6O

2. Find the molar mass of the empirical formula.

C4H6OC = 4 x 12.0 = 48.0H = 6 x 1.0 = 6.0O = 1 x 16.0 = 16.0

70.0

3. Divide the true molecular mass by the mass of the empirical formula. Multiply the subscripts of the empirical formula by the factor just computed. This is the molecular formula.

= = =2

2 (EF) = 2 (C4H6O) = C8H12O2

molecular formula = C8H12O2

problems:

1. What is the molecular formula of each of the following compounds?a. Empirical formula CH2 , molar mass =84g/mol

b. Empirical formula NH2Cl, molar mass = 51.5 g/mol

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2. Determine the empirical and molecular formula of each of the following substances:a. Ibuprofuren, a headache remedy contains 75.6 % C, 8.80 % H , and 15.5 % O by mass and a molar mass about 206 g/mol.

b. Epinerphine (adrenaline), a hormone secreted into the bloodstream in times of danger or stress contains 59% C, 7.1% H, 26.2% O, and 7.7% N by mass; its molar mass is about 180 amu.

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Topics 16 and 17 Balancing and writing Chemical Equations

Types of Chemical ReactionsAll chemical reactions can be grouped into one of six categories. These categories are based on how the atoms in the molecules react with each other.

Synthesis – A synthesis reaction is hwen two or more simple compounds combine to form a more complicated one. These reactions come in the general form of:

A + B → AB

One example of a synthesis reaction is the combination of iron and sulfur to form ferrous sulfide

8 Fe + S8 → 8 FeS

Decomposition – A decomposition reaction is the opposite of a synthesis reaction. A complex molecule breaks down to make simpler ones. These reactions come in the general form:

AB → A + B

One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen gas:

2 H2O 2 H2 + O2

Single Displacement/Replacement – This is when one element trades places with another element in a compound. These reactions come in the general form of:

A + BC → AC + B or X + YZ → YX + Z

One example of a single displacement reaction is when magnesium replaces hydrogen in water to make magnesium hydroxide and hydrogen gas:

Mg + 2 H2O → Mg(OH)2 + H2

Double Displacement/Replacement – This is when the anions and cations of two different molecules switch places, forming two entirely different compounds. These reactions are in the general form:

AB + CD → AD + CB

One example of a double displacement reaction is the reaction between plumbous nitrate with potassium iodide to form plumbous iodide and potassium nitrate:

Pb (NO3)2 + 2 KI → PbI2 + 2 KNO3

Acid-base/Neutralization – This is a special kind of double displacement reaction that takes place when an acid and base react with each other. The H+ ion in the acid reacts with the OH- ion in the base, causing the formation of water. Generally, the product of this reaction is some ionic salt and water:

HA + BOH → H2O BA

One example of an acid-base reaction is the reaction of hydrobromic acid with sodium hydroxide:

HBr + NaOH → NaBr + H2O

Combustion – A combustion reaction is when oxygen combines with another compound (usually a hydrocarbon) to form water and carbon dioxide. These reactions are usually exothermic (meaning they produce heat). An example of this kind of reaction is the burning of naphthalene:

C10H8 + O2 → 10 CO2 + 4 H2O

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1. Write a balanced chemical equation for each of the following:

a. The reaction of magnesium oxide with iron to form iron(III) oxide and magnesium.

b. The decomposition of dinitrogen oxide gas to its elements.

c. The reaction of solid calcium carbide with water to form calcium hydroxide and

acetylene (C2H2) gas.

d. The reaction of solid calcium cyanamide (CaCN2) with water to form calcium

carbonate and ammonia gas.

e. Propane (C3H8) burns in excess air (oxygen).

f. Nitrogen gas reacts with hydrogen to form ammonia.

g. Hydrogen reacts with Iodine gas to form hydrogen iodide.

h. Sodium reacts with iodine gas to form sodium iodide.

i. Sodium oxide reacts with water to form sodium hydroxide and hydrogen.

j. Magnesium and nitrogen gas combine to form magnesium nitride.

k. Concentrated hydrochloric acid reacts with concentrated sodium hydroxide to form

sodium chloride and water.21

2. Predict the products, and balance the equations:

a. NaCl + MgF2 →

b. C6H12 + O2 →

c. AgNO3 + Cu → (assume cupric)

d. Cr(OH)3 + HF →

e. Na2CO3 + MnO →

f. HF + NH4OH →

g. . C2H2 + O2 →

h. Ca + AlCl3 →

i. Au2S3 + H2 →

j. Hg(OH)2 + H3PO4 →

k. Au2O3 →

l. C4H10 + O2 →

2. Write the formulas for all reactants and products and balance:

a. Solid potassium chlorate is decomposed

b. Aqueous solutions of sulfuric acid and magnesium hydroxide react

c. piece of Iron reacts with oxygen gas (assume ferric in product)

d. Solid nickel II carbonate decomposes

e. Aqueous magnesium chloride is reacted with bromine gas

f. Aqueous aluminum nitrate is reacted with a piece of magnesium

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Topics 18 Stoichiometry: Mole-mole, mass-mass problems

Mole-Mole Problems

These problems are used to calculate the number of moles of substance that will react with or be produced by a given number of moles of substance A. The general form of the math equation will be:

Moles A → Moles B

Lest go back to the shuttle fuel reaction to see this type of problem:

N2H4 + 2 H2O2 → N2 ↑ + 4 H2O ↑

1. If 275 moles of N2H4 are required to fuel a space shuttle, how many moles of H2O2 will you need?

Staying with our space theme, here’s another problem … In a spaceship, the carbon dioxide exhaled by an astronaut is often removed by its reaction with lithium hydroxide, which forms lithium carbonate and water as its products.

2. Write the balanced equation for this reaction.

3. The average person exhales 20 moles of carbon dioxide. How many moles of lithium hydroxide would be needed to react with this amount of carbon dioxide?

4. The unbalanced equation for the synthesis of ammonia is N2 + H2 → NH3. How many moles of H2 are required to produce 6 moles of NH3?

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Mass-Mass Problems

In this type of problem, you will determine the mass of a substance that will react with, or be produced by a given mass of a second substance. The general form of the math equation will be:

Mass A → Moles A → Moles B → Mass B

This is a 3 step problem that will require 3 conversion factors – the molar mass of A, the mole ratio of substance B to substance A, and the molar mass of substance B.

Mass Mass Problems1. Write the balanced chemical equation.2. Write the amount given under the formula for that substance.3. Identify what is being sought.4. Write a “?” under the formula for that substance.5. Convert g of the given compound to moles by multiplying by a conversion factor.6. Convert moles of the given compound to moles of the compound sought by multiplying by a mole ratio.7. Convert moles of the compound sought to g by multiplying by a conversion factor.

Remember The equation must be balanced. The units in the conversion factors/mole ratios must cancel. The numbers in the mole ratio come from the balanced chemical equation. The number of moles in the conversion factor is always 1. The mass in the conversion factor is the molar mass.

Example – Calcium chloride reacts with sodium carbonate. If 5.00 g of calcium chloride reacts, how any grams of calcium carbonate will form?

1. a. Write the balanced equation that occurs when potassium bromide and fluorine react in a single displacement reaction.

b. How many grams of bromine are formed from 1.72 g of KBr?

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1. How many grams magnesium sulfide is formed when 35.0 g of magnesium reacts with an excess of sulfur in a synthesis reaction?

2. How many grams of beryllium hydroxide is needed to react completely with 100.0 g of hydrochloric acid?

3. How many grams of ferrous nitrate is formed when 25.0 g of pure iron is dissolved in pure nitric acid?

4. How many grams of plumbous sulfide is precipitated in the reaction between 21.0 g of plumbous bromide and an excess of hydrogen sulfide?

5. How many moles of lithium chloride is produced by the reaction of 7.45 moles of lithium metal and excess chlorine gas?

6. How many moles of silver chloride are produced from 5.0o moles of silver nitrate reacting with an excess of barium chloride?

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