Westside High School Summer Review Packet for 20162017 · PDF fileAP Chem PAP Chemistry...

AP Chem PAP Chemistry Review.notebook

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Westside High School

Summer Review Packet

for 20162017 AP Chemistry

This review packet must be printed, completed and submitted for a doubleminor grade on 08/26/16 (Fri). All 34 exercise questions (indicated by ) must be answered with all work shown for a full credit. You will have the opportunity to ask questions during the first week of school in order to complete the packet; however, the instructor will proceed with the curriculum under the assumption that you have thoroughly reviewed the materials in this packet before the first day of class.


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PreAP Chemistry: Review

• Matter> takes up space, has mass and exhibits inertia> composed of atoms

• Water> made up of one oxygen and two hydrogen atoms> can be decomposed by electrolysis (separation by passing

through an electric current); reactions are reversible.

2 H2O 2H2 + O2


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• Chemistry> the study of matter, energy and the changes> important for safety> provides better understanding in all areas of science

• Robert Boyle> performed many experiments with air> created the first vacuum pump

> Boyle's Gas Law: P1V1 = P2V2

• Scientific Laws> Law of Conservation of Mass

Mass is neither created nor destroyed.> Law of Conservation of Energy

Energy is neither created nor destroyed.


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• Units of Measure> a quantitative observation or measurement> consists of a number and an unit> SI system (French: Le Systeme International )– international fundamental units of measurement agreed upon in 1960

– mostly the Metric system

The Fundamental SI Units:

Physical Quantity Name of Unit Abbreviation

Mass kilogram kgLength meter mVolume liter LTime second s (or sec)Temperature Kelvin KAmount of substance mole molElectric current Ampere ALuminous intensity Candela cd


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Common Prefixes Used in the SI System:

Prefix Symbol Meaning Exponential Notation

mega M 1,000,000 106

kilo k 1,000 103

hecto h 100 102

deka da 10 101

1 100

deci d 0.1 101

centi c 0.01 102

milli m 0.001 103

micro m 0.000001 106

nano n 0.000000001 109


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• Volume> derived from length> a cube with 1m length on each edge = 1.0 m3

> (1m)3 = (10 dm)3 = 103 dm3 = 1,000 dm3

> (1m)3 = (100 cm)3 = 1003 cm3 = 1,000,000 cm3

> 1 L = 1,000 mL = 1,000 cm3 > 1g of H2O = 1 mL = 1 cm3 (at 4 oC)

• Mass vs. Weight> mass (g or kg) is a measure of the resistance of an object to a

change in its state of motion; the quantity of matter present> weight (N for Newtons) is a force; the response of mass to

gravity.> often used interchangeably since all of our measurements are

made on Earth and the acceleration due to gravity is considered a constant


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• Precision vs. Accuracy> accuracy is the correctness or agreement of a measurement with

the true value> precision is the reproducibility or consistency among several

measurements

• Random vs. Systematic Error> random (or indeterminate) errors cause the measurement to be

scattered randomly; degrades precision.> systematic (or determinate) errors cause the measurement to be

displaced by a particular amount from the true value; degrades accuracy.


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Exercise 1 Precision and Accuracy

To check the accuracy of a graduated cylinder, a student filled the culinder to the 25mL mark using water delivered from a buret and then read the volume delivered. Following are the results of five trials:

Volume shown by Volume showTrial Graduated Cylinder by the Buret

1 25 mL 26.54 mL

2 25 mL 26.51 mL

3 25 mL 26.60 mL

4 25 mL 26.49 mL

5 25 mL 26.57 mL

Average 25 mL 25.54 mL

Is the graduated cylinder accurate?


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• Significant Figures

> Nonzero digits are always significant.

> A zero is significant IF AND ONLY IF:

– the zero is a trailing zero after the decimal

– the zero is a captive zero between two nonzero digits

> Exact, counting numbers and fundamental constants have an infinite amount of significant figures

Exercise 2 Significant Figures

Give the number of significant figures for each of the following experimental results.

a) A student's extraction procedure on a sample of tea yields 0.0105 g of caffeine.

b) A chemist records a mass of 0.050080 g in analysis.

c) In an experiment, a span of time is determined to be 8.050 x 103 s.


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• Reporting the Result of a Calculation to the Proper Number of Significant Figures

> When multiplying or dividing, the term with the LEAST number of SIG FIGS determines the number of SIG FIGS in the final answer.

4.56 x 1.4 = 6.38 6.4

> When adding or subtracting, the term with the LEAST number of DECIMAL PLACES determines the number of DECIMAL PLACES in the final answer.

12.11 18.0

+ 1.013

31.123 31.1

> When reporting pH, the number of SIG FIGS in [H+] determines the number of DECIMAL PLACES in pH.

corrected to

corrected to

limiting term with only 1 decimal place


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• Dimensional Analysis

Example: Consider a straight pin measuring 2.85 cm in length. Calculate its length in inches.

Start with a conversion factor such as 1 inch = 2.54 cm.

You can write TWO different conversion factors from this:

1 in 2.54 cm

2.54 cm 1 in

Why is this legal? Because both quantities represent the exact same "thing" so the conversion factor is actually equal to "1."

To convert the length of the pin from cm to inches, simply multiple your given quantity by a conversion factor set up to "cancel" the undesired unit and places the desired unit where you want it.

2.85 cm 1 in

2.54 cm

or

= 1.12 in


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Exercise 3

A pencil is 7.00 inches long. Calculate the length in centimeters.

Exercise 4

You want to order a bicycle with a 25.5in frame, but the sizes in the catalog are given only in centimeters. What size should you prder?


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Exercise 5

A student has entered a 10.0km run. How long is the run in miles?

Note: 1 km = 1000 m

1 m = 1.094 yd

1760 yd = 1 mi


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Exercise 6

The speed limit on many highways in the United States is 55 mi/h. What number would be posted if expressed in kilometers per hour?


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Exercise 7

A Japanese car is advertised as having a fuel economy of 15 km/L. Convert this rating to miles per gallon.


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• Temperature

boiling point of pure water

freezing point of pure water

absolute zero

• K = oC + 273.15• oC = K 273.15• oC = (oF 32)/1.8 • oF = (1.8 x oC) + 32


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• Density> density =

mass

volume

Exercise 8 Determining Density

A chemist, trying to identify the main component of a compact disc cleaning fluid, determines that 25.00 cm3 of the substance has a mass of 19.625 g at 20 oC. Use the information in the table below to identify which substance may serve as the main component of the cleaning fluid. Justify your answer with a calculation.

Compound Density (g/cm3) at 20 oC

Chloroform 1.492Diethyl ether 0.714Ethanol 0.789Isopropyl alcohol 0.785Toluene 0.867


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• Classification of Matter> Changes of state involve altering IMFs (Intermolecular Forces)

and not altering actual chemical bonds.> Solids have rigid, defined shape and volume; molecules are

close together and vibrate. Virtually incompressible.> Liquids have defined volume, but take on the shape of the

container; molecules are still close together and vibrate, rotate or slide around one another. Slightly compressible.

> Gases have no defined volume and take on the shape of the container; molecules are very far apart, and move independently of each other. Highly compressible.

> Vapor refers to the gas phase of a substance that is normally a solid or a liquid at room temperature.

> Fluid refers to a substance that flows; including liquids, gases, plasmas and some plastic solids.


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• Pure Substances> compounds and elements> compounds can be separated into elements by chemical means> electrolysis is a common method of chemical separatation of

compounds into elements> elements can be broken down into atoms> atoms can be broken down into nuclei and electrons> nuclei can be broken into protons and neutrons> protons and neutrons can be broken into quarks

Matter

Mixtures

Pure Substances

Homogeneous Mixtures

Heterogeneous Mixtures

Elements

Compounds

AtomsNuclei

Electrons

protons

neutrons

quarks

quarks


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• The Early History of Chemistry> 4,000 B.C. Processing of ores to produce metals for weapons

and ornaments; use of embalming fluids> 3,500 B.C. Alchemy was a pseudoscience where people sought

to turn metals into gold.> 400 B.C. Greeks proposed all matter was made up of 4

"elements" fire, earth, water and air; Democritus first used the term "atoms" to describe the smallest particles of matter.

> 1662 Robert Boyle (English) was the first chemist to perform quantitative experiments on pressure vs. volume.

> 17th century J.J. Becher (German) suggested "phlogiston" flowed out of burning material. "Phlogisticated" substances burn and become "dephlogisticated."

> 18th century Joseph Priestley (English) and Carl Scheele (Swedish) discovered oxygen.

https://www.youtube.com/watch?v=thnDxFdkzZs

https://www.youtube.com/watch?v=thnDxFdkzZs


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• Fundamental Chemical Laws> 18th century Antoine Lavoisier (French) explained the true

natur of combustion, published the first modern chemistry textbook and stated the Law of Conservation of Mass by being the first to perform quantitative experimentation.

> 1808 John Dalton (English) stated the Law of Definite Proportions. He later went on to develop the Atomic Theory of Matter.

The Law of Definite Proportions:

A given compound always contains exactly the same proportions of elements by mass.

The Law of Multiple Proportions:

When 2 elements combine to form a compound, the ratios of the masses of the second element that combine with 1g of the first element can always be reduced to small whole numbers.

https://www.youtube.com/watch?v=ZgmGxN66rTQ

https://www.youtube.com/watch?v=ZgmGxN66rTQ


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Dalton considered compounds of C and O and determined:Mass of O that combines with 1 g of C

Compound I 1.33 g

Compound II 2.66 g

Therefore, Compound I may be CO while Compound II may be CO2.

Postulates of Dalton's Atomic Theory:

1) All matter is made of atoms that are indivisible and indestructible.

2) All atoms of a given element are identical. However, atoms of different elements have different weights and chemical properties.

3) Compounds are formed by the combination of different atoms in the ratio of small whole numbers.

4) A chemical reaction involves only the combination, separation or rearrangement of atoms; atoms are neither created nor destroyed in the course of ordinary chemical reactions.

subatomic particles have been discoveredisotopes were discovered


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• 1809 Joseph GayLussac (French) stated that if the mass and volume of a gas are held constant, then gas pressure increases linearly as the temperature rises. P1/T1 = P2/T2

• 1811 Avogadro (Italian) proposed his hypothesis: At the same temperature and pressure, equal volumes of different gases contain the same number of particles. V1/n1 = V2/n2

• 1897 J.J. Thomson (English) discovered the electron in his cathoderay tube experiment. He called his model of an atom the "plum pudding" model.

• 1909 Robert Millikan (American) determined the charge (1.6 x 1019 Coulomb) and the mass (9.11 x 1031 kg) of an electron through his Oil Drop experiment.

https://www.youtube.com/watch?v=XMfYHag7Liw

https://www.youtube.com/watch?v=XMfYHag7Liw


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• 1896 Henri Becquerel (French) discovered radioactivity.• 1911 Ernest Rutherford (English) carried out the Gold Foil experiment to test Thomson's plum pudding model; he discovered the positively charged nucleus.

• 3 Types of radioactive emission:> alpha (α) He> beta (β) e or β > gamma (γ) pure energy; no particle

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01

01


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The Modern View of Atomic Structure:• Elements> composed of only one type of atom> 92 naturally occurring elements; all others are synthetic

• Atoms> the smallest particle of an element that retains the chemical

properties> nucleus contains the protons and neutrons and is very dense> protons (p+) are positively charged particle that defines the

atomic number (the identity of the element)> neutrons (no) have no charge; have the same size and mass as

p+; responsible for isotopes> electrons (e) are negatively charged; same size as p+ or no, but

1/1,800 the mass of p+ or no

> atomic number (Z) is the number of p+ in an atom> mass number (A) is the sum of the number of p+ and no


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AZ X

mass number

atomic number

element symbol

Exercise 9 Writing the Symbols for Atoms

Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and how many neutrons does this atom have?


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• Isotopes> atoms having the same number of p+ but a different number of

no > most elements have at least two stable isotopes> Hydrogen isotopes have special names:– H1 (with no no) is called Hydrogen– H2 (with one no) is called Deuterium– H3 (with two no) is called Tritium

• Molecules and Ions> chemical bonds are the forces that hold atoms together> in covalent bonds, atoms share electrons and make molecules;

examples include H2, CO2, H2O, NH3, O2, CH4, etc.> molecules are the smallest units of a compound that retains the

chemical characteristics of the compound> molecular formula uses symbols and subscripts to represent the

composition of the molecule> in a structural formula, bonds are shown by lines to represent

shared e pairs; may or may not indicate shape.> ions are formed when e are lost or gained– metals often form cations that are (+) ions by losing e

– nonmetals often form anions that are () ions by gaining e

– polyatomic ions are units of atoms behaving as one entity


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The Periodic Table


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• Metals> malleable, ductile and shiny> most exist as cations in a sea of electrons> form oxides readily

• Group A = representative elements• Group B = transition metals with various oxidation states

• Nomenclature> Naming (+) ions– for monatomic metal cations, simply name the metal off the periodic table (ex. Al3+ is the aluminum ion)

– transition metals form more than one ion, so use Roman Numerals in parenthesis following the ion's name (ex. Cu2+ is copper(II)

– Mercury(I) is an exception: it is Hg22+ (two Hg+ associated together)

– NH4+ is ammonium– No Roman Numeral is used with Al3+, Ag+, Cd2+ and Zn2+ because they only form one oxidation state

> Naming () ions– for monatomic anions, add the suffix ide to the stem of the nonmetal's name. Halogens are called the halides.

– for polyatomic anions: Examples

« hypoite the least amount of Oxygen hypochlorite ClO

« ite 1 more oxygen than hypoite chlorite ClO2

« ate 1 more oxygen than ite chlorate ClO3

« perate the most oxygen perchlorate ClO4


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Exercise 10 Naming Type I Binary Compounds

Name each binary compound.

a) CsF

b) AlCl3

c) LiH

Exercise 11 Naming Type II Binary Compounds

a) CuCl

b) HgO

c) Fe2O3

d) MnO2

e) PbCl2


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Metal present?

No Yes

Naming Binary Compounds: a Flowchart

Type III: Use prefixes.

Does the cation formvarious oxidation states?

No Yes

Type I:Use the element name for the cation.

Type II:Determine the charge of the cation; use a Roman numeral in parenthesis after the cation name.


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Common Polyatomic Ions To Memorize:

Cations NameHg22+ mercury(I)NH4+ ammonium

Anions NameNO2 nitriteNO3 nitrateSO32 sulfiteSO42 sulfateHSO4 Hydrogen sulfate (or bisulfate)OH hydroxideCN cyanidePO43 phosphateHPO42 Hydrogen phosphateH2PO4 dihydrogen phosphateNCS thiocyanateCO32 carbonateHCO3 Hydrogen carbonate (or bicarbonate)


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Common Polyatomic Ions To Memorize: (continued)

Anions NameClO hypochloriteClO2 chloriteClO3 chlorateClO4 perchlorateBrO2 bromiteBrO3 bromateC2H3O2 acetateMnO4 permanganateCr2O72 dichromateCrO42 chromateO22 peroxideC2O42 oxalateBO33 borateS2O32 thiosulfate


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Exercise 12 Naming Binary Compounds

Give the systematic name for each of the following compounds.

a) CoBr2 b) CaCl2

c) Al2O3 d) CrCl3

Exercise 13 Naming Compounds with Polyatomic Ions


a) Na2SO4 f) Na2CO3

b) KH2PO4 g) NaHCO3

c) Fe(NO3)3 h) CsClO4

d) Mn(OH)2 i) NaOCl

e) Na2SO3 j) KBrO3


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Exercise 14 Naming Type III Binary Compounds

Name each of the following compounds.

a) PCl5

b) PCl3

c) SF6

d) SO3

e) SO2

f) CO2


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Naming Acids:• If oxygen is absent, "hydro" is listed first followed by a suffix "ic" and the word "acid."

• If the negative ion's name ends in:> ide, then hydro (anion root name)ic acid

ex. hydrosulfuric acid, H2S> ate, then ic acid

ex. sulfuric acid, H2SO4> ite, then ous acid

ex. sulfurous acid, H2SO3


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Naming Acids: a Flowchart

Does the anioncontain oxygen?

No Yes

hydro+anion root+ ichydro(anion root)ic acid

Check the ending of the anion

ite ate

anion or element root+ ous(root)ous acid

anion or element root+ ic(root)ic acid


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The Exceptions to the Naming Rules:You only need to memorize the first two; the AP exam either avoids the rest or gives you the formula.

• water H2O• ammonia NH3• hydrazine N2H4• phosphine PH3• nitric oxide NO• nitrous oxide (aka "laughing gas") N2O


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Exercise 15 Naming Acids

Name each of the following acids.

a) HBr

b) HI

c) HBrO2

d) HBrO3

e) HClO4

f) HNO2

g) HNO3


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Exercise 16 Naming Various Types of Compounds


a) P4O10

b) Nb2O5

c) Li2O2

d) Ti(NO3)4

Exercise 17 Writing Compound Formulas from Names

Given the following names, write the formula for each compound.

a) Vanadium(V) fluoride

b) Dioxygen difluoride

c) Rubidium peroxide

d) Gallium(III) oxide


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• The Mole> the number of Carbon atoms in exactly 12.0 g of C12.> Avogadro's number 6.022 x 1023

An Alternative to Dimensional Analysis: The Mole Map

Multiply by the conversion factor on the arrow when traveling in the direction of the arrow; divide when traveling against an arrow.

Concentration of Solution (mol/L)

Mass(grams)

Number ofMoles

Number ofParticles

Gas Volume(Liters)

Molar Mass(g/mol)

6.022 x 1023

Volume

of Solution(Liters)

22.4 L of gas @

STP


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Exercise 18 Determining the Mass of a Sample of Atoms

Americium is an element that does not occur naturally. It can be made in very small amounts in a device known as a particle accelerator. Calculate the mass in grams of a sample of americium containing six atoms.


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Exercise 19 Determining Moles of Atoms

Aluminum is a metal with a high strengthtomass rati and a high resistance to corrosion; thus it is often used for structural purposes. Calculate both the number of moles and the number of atoms in a 10.0g sample of aluminum.


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Exercise 20 Calculating the Number of Moles and Mass

Cobalt (Co) is a metal that is added to steel to improve its resistance to corrosion. Calculate both the number of moles in a sample of cobalt containing 5.00 x 1020 atoms and the mass of the sample.


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• Molar Mass (MM) is the sum of all the atomic masses in a given chemical formula in units of g/mol; aka molecular weight or formula weight

• Empirical Formula is the simplest ratio of elements in a compound.• Stick with 2 decimal places for all MM calculated from the Periodic Table.


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Exercise 21 Calculating Molar Mass I

Juglone, a dye known for centuries, is produced from the husks of black walnuts. It is also a natural herbicide (weed killer) that kills off competitive plants around the black walnut tree but does not affect grass and other noncompetitive plants. The formula for juglone is C10H6O3.

a) Calculate the molar mass of juglone.

b) A sample of 1.56 x 102 g of pure juglone was extracted from black walnut husks. Calculate the number of moles of juglone present in this sample.


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Exercise 22 Calculating Molar Mass II

Calcium carbonate, also called calcite, is the principal mineral found in limestone, marble, chalk, pearls and the shells of marine animals such as clams.

a) Calculate the molar mass of calcium carbonate.

b) A certain sample of calcium carbonate contains 4.86 moles. Calculate the mass in grams of this sample. Also calculate the mass of the carbonate ions present.


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Exercise 23 Molar Mass and Numbers of Molecules

Isopentyl acetate (C7H14O2), the compound responsible for the scent of bananas, can be produced commercially. Interestingly, bees release about 1 mg of this compound when they sting. The resulting scent attracts other bees to join the attack.

a) Calculate the number of molecules of isopentyl acetate released in a typical bee sting.

b) Calculate the number of carbon atoms present.


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Diatomic molecules Br2 I2 N2 Cl2 H2 O2 F2 the "gens"• pure hydrogen, nitrogen, oxygen and the halogens exist as diatomic molecules under normal conditions.

• Others to be aware of:> P4 tetratomic form of elemental phosphorous;

an allotrope (different physical form of the same element)> S8 sulfur's elemental form; an allotrope > Carbon allotropes include diamond and graphite


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Percent Composition (by mass)

% Composition = x 100mass of desired element

total mass of compound

Example: Consider ethanol, C2H5OH

C (2)(12.01)=24.02H (6)( 1.01)= 6.06O 16.00

46.08 g/mol

% C = x 100 = 52.14% C

% H = x 100 = 13.2% H

% O = x 100 = 34.72% O

24.02

46.08

6.06

46.08

16.00

46.08


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Exercise 24 Calculating Mass Percent I

Carvone is a substance that occurs in two allotropes with the same formula (C10H4O) and mass. One type of carvone gives caraway seeds their characteristic smell, and the other type is responsible for the smell of spearmint oil. Calculate the mass percent of each element in carvone.


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Exercise 25 Calculating Mass Percent II

Penicillin, the first of a now large number of antibiotics, was discovered accidentally by the Scottish bacteriologist Alexander Fleming in 1928, but he was never able to isolate it as a pure compound. This and similar antibiotics have saved millions of lives that might have been lost to infections. Penicillin F has the formula C14H20N2SO4. Calculate the mass percent of each element.


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Determining the Formula of a Compound

When faced with an unknown hydrocarbon compound, one of the most common technique is to combust it with oxygen to produce oxides of the nonmetals CO2 and H2O which are then collected and weighed.

• Calculating Empirical and Molecular Formulas:> Empirical Formula represents the simplest or smallest whole

number ratio of elements within a compound> Molecular Formulas represent the actual number of elements

within a compound> example: CH2O is the empirical formula for glucose whose

molecular formula is C6H12O6.• Hydrates: water of hydration or "dot waters." They count in the calculation of molar masses for hydrates and used to "cement" crystal structures together. ex. CuSO4.5H2O, MgSO4.7H2O, etc.

• anhydrous: means without water. Hydrates can be heated to become anhydrous. ex. CuSO4.5H2O CuSO4 + 5H2O

hydrate anhydrous salt


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Example: Determination of Empirical Formula

A compound is composed of carbon, nitrogen and hydrogen. When 0.1156 g of this compound is reacted with oxygen, 0.1638 g of carbon dioxide and 0.1676 g of water are collected. Determine the empirical formula of the compound, and the molecular formula given that the MM of the compound is 31.06 g/mol.

Compound + O2 CO2 + H2O + N2

0.1638 g CO2 1 mol CO2 1 mol C

44.01 g 1 mol CO2

0.1676 g H2O 1 mol H2O 2 mol H

18.02 g 1 mol H2O

0.003781 mol C 12.01g

1 mol C

0.01860 mol H 1.01 g

1 mol H

0.1156 g of compound (0.04540 g C + 0.01879 g H) = 0.05141 g N

0.05141 g N 1 mol N

14.01 g

= 0.003781 mol C

= 0.01860 mol H

= 0.04540 g C

= 0.01879 g H

= 0.003670 mol N


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Example: Determination of Empirical Formula (continued)

0.003781 mol C

0.003670

0.01860 mol H

0.003670

0.003670 mol N

0.003670

≅ 1

≅ 5

= 1

The empirical formula is CH5N

C 12.01H (5)(1.01) = 5.05N 14.01

31.07 g/mol

MM of Molecular formula 31.06 g/mol

MM of Empirical formula 31.07 g/mol

The molecular formula is also CH5N

= ≅ 1 x 1


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Exercise 26

Determine the empirical and molecular formula for a compound that gives the following analysis in mass percents:

71.65% Cl 24.27% C 4.07% H

The molar mass is known to be 98.96 g/mol.


58

Exercise 27

A white powder is analyzed and found to contain 43.64% phosphorus and 56.36% oxygen by mass. The compound has a molar mass of 283.88 g/mol. What are the compound's empirical and molecular formula?


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Exercise 28

Caffeine, a stimulant found in coffee, tea and chocolate, contains 49.48% carbon, 5.15% hydrogen, 28.87% nitrogen and 16.49% oxygen by mass and has a molar mass of 194.2 g/mol. Determine the molecular formula of caffeine.


60

• Writing Chemical Equations> solid (s), liquid (l), gas (g) and aqueous solution (aq)> metals and halogens combine to make metal halides

ex. Sn + 2Cl2 SnCl4> combustion of hydrocarbons yields CO2 and H2O

ex. CH4 + O2 CO2 + H2O> any time H2CO3 is produced, it spontaneously decomposes into

CO2 and H2Oex. 2HCl + CaCO3 CaCl2 + CO2 + H2O

> metal carbonates decomposes into metal oxides and CO2at very high temperaturesex. Na2CO3 (s) Na2O (s) + CO2 (g)


61

Exercise 29

Solid ammonium dichromate, a vivid orange compound, decomposes into solid chromium(III) oxide, nitrogen gas and water vapor. Write a balanced equation for this reaction.

Exercise 30

At 1000 oC, ammonia gas reacts with oxygen gas to form gaseous nitric oxide and water vapor. Write a balanced equation for this reaction.


62

Stoichiometric Calculations

Example:

Calculate the mass of oxygen that will react completely with 96.1 grams of propane.

step 1) Write a balanced equation.

C3H8 + 5O2 3CO2 + 4H2O

step 2) Convert given grams to moles.

96.1 g C3H8 1 mol C3H8

44.11 g

step 3) Use the moletomole ratio indicated by the coefficients in the balanced equation to solve for the unknown moles.

2.18 mol C3H8 5 mol O2

1 mol C3H8

step 4) Convert the calculated moles to grams.

10.9 mol O2 32.00 g

1 mol O2

= 2.18 mol C3H8

= 10.9 mol O2

= 349 g O2


63

Exercise 31

Solid lithium hydroxide is used in space vehicles to remove exhaled carbon dioxide from the living environment by forming solid lithium carbonate and liquid water. What mass of gaseous carbon dioxide can be absorbed by 1.00 kg of lithium hydroxide?


64

Exercise 32

Baking soda (NaHCO3) is often used as an antacid. It neutralizes excess hydrochloric acid secreted by the stomach:

NaHCO3 (s) + HCl (aq) NaCl (aq) + H2O (l) + CO2 (aq)

Milk of magnesia, which is an aqueous suspension of magnesium hydroxide, is also used as an antacid:

Mg(OH)2 (s) + 2HCl (aq) 2H2O (l) + MgCl2 (aq)

Which is the more effective antacid per gram, NaHCO3 or Mg(OH)2 ? Justify your answer.


65

Calculations Involving a Limiting Reagent

Example:

Examine the diagrams below and explain the differences between the two situations with regards to what is (or isn't) reacting and total yield of ammonia.

N2 (g) + 3H2 (g) 2NH3 (g)

Given 5 mol of N2 and 15 moles of H2 , 10 mol of NH3 are produced with no reactants left over.

Given 5 mol of N2 and 9 mol of H2 , 6 mol NH3 are produced and 2 mol of N2 are left over. All H2 were used up. (i.e., H2 was the limiting reagent)


66

Example: Haber Process (Production of Ammonia)

Suppose 25.0 kg of nitrogen reacts with 5.00 kg of hydrogen to form ammonia. What mass of ammonia can be produced? Which reactant is the limiting reactant? What is the mass of the reactant that is in excess?

N2 + 3H2 2NH3

25.0kg 5.00kg ?kg

25.0 kg N2 1000 g 1 mol N2

1kg 28.02g

5.00kg H2 1000 g 1 mol H2

1kg 2.02g

892 mol N2 2 mol NH3

1 mol N2

2475 mol H2 2 mol NH3

3 mol H2

2475 mol H2 1 mol N2

3 mol H2

= 892 mol N2

= 2475 mol H2

= 1784 mol NH3

= 1650 mol NH3 17.04g 1kg

1 mol NH3 1000g = 28.1 kg

ammonia produced

H2 is the limiting reactant

= 825 mol N2 was used


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Example: Haber Process (Production of Ammonia) continued

892 mol N2 825 mol N2 used = 67 mol N2 left over

67 mol N2 28.02g 1kg

1 mol N2 1000g= 1.88 kg N2 is in excess


68

Exercise 33

Nitrogen gas can be prepared by passing gaseous ammonia over solid copper(II) oxide at high temperatures. The other products of the reaction are solid copper and water vapor. If a sample containing 18.1g of ammonia is reacted with 90.4g of CuO, which is the limiting reactant? How many grams of nitrogen gas will be formed?


69

• Theoretical Yield> the amount of product formed when a limiting reactant is

completely consumed.> assumes perfect conditions and the maximum amount

(unrealistic)• Actual Yield> more realistic, the actual amount of product formed

• Percent Yield> the ratio of actual yield to theoretical yield

actual yield

theoretical yieldx 100 = percent yield


70

Exercise 34

Methanol (CH3OH), also called methyl alcohol, is the simplest alcohol. It is used as a fuel in race cars and is a potential replacement for gasoline. Methanol can be manufactured by combination of gaseous carbon monoxide and hydrogen. Suppose 68.5 kg CO (g) is reacted with 8.60 kg H2 (g). Caclulate the theoretical yield of methanol. If 3.57x104 g CH3OH is actually produced, what is the percent yield of methanol?

Westside High School Summer Review Packet for 20162017 · PDF fileAP Chem PAP Chemistry...

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Transcript of Westside High School Summer Review Packet for 20162017 · PDF fileAP Chem PAP Chemistry...

Westside High School Summer Review Packet for 2016­2017 · PDF fileAP Chem PAP Chemistry...

Documents

Transcript of Westside High School Summer Review Packet for 2016­2017 · PDF fileAP Chem PAP Chemistry...

Westside High School Summer Review Packet for 20162017 · PDF fileAP Chem PAP Chemistry...

Transcript of Westside High School Summer Review Packet for 20162017 · PDF fileAP Chem PAP Chemistry...