Exam on monday

bring your own pen, pencil, eraser, and liquid--no sharing rule is being implemented

exam news

I will answer any questions you have until Saturday at 12 noon. do not barrage me with questions on monday or sunday night etc.

ep 6 exam topics

ch 19 pages 586-629
focus on acid base theories problems with ph, pOH, ka, kb, and kw.
indicators and buffers plus neutralization reactions

EP 5 exam topics

ch 13 states of matter 384-411 states of matter, kinetic theory, colloids
ch 14 behavior of gases page 412-434 focus on gas law problems properties of ideal gases
ch 15 pages 444-458 water--polarity and homogeneous aqueous systems

ep 4 exam topics and PAGES

ch 1 pages 1-20--areas of chemistry
ch 2 page 38-61 matter and change

ch 3 pages 73-99 focus on density, SI, and conversion between metrics
ch 4 atomic structure--pages 100-125 focus on the neutrons, electrons, and protons plus atomic theories of rutherford, thomson, millikan, THERE ARE NO CALCULATIONS WITH ATOMIC MASS ON TEST page 117

ch 5 126-132 no electron configuration on the exam

ep 4 test coverage

ch 1-4.4
p.8, p. 22 sci method, p 41-states of matter, p 57 vocabulary, p. 92 density problems, and ch 4.1-4.2 thomson, rutherfrod, and millikan. also do review p 37 #1-10 and p. 61 # 1-14

ep 5 test review problems

problem samples p 419, 421, 424, 427, 428 and 429, and 434 also know vocab page 406.

sample exam ep 4

I would print this out, try it without the answers, and then, look at the answers

EP 4 Test Review Ch 1-4

Multiple Choice
Identify the letter of the choice that best completes the statement or answers the question.

____ 1. Which field of science studies the composition and structure of matter?
a. physics c. chemistry
b. biology d. geology


____ 2. The study of chemicals that, in general, do not contain carbon is traditionally called what type of chemistry?
a. bio c. physical
b. inorganic d. analytical


____ 3. Which of the following best describes an example of pure chemistry?
a. testing the effects of lower concentrations of a drug on humans
b. studying chemicals containing carbon
c. developing a cure for osteoporosis
d. finding an antidote for a new strain of virus


____ 4. Which of the following are considered physical properties of a substance?
a. color and odor c. malleability and hardness
b. melting and boiling points d. all of the above


____ 5. A vapor is which state of matter?
a. solid c. gas
b. liquid d. all of the above


____ 6. Which state of matter has a definite volume and takes the shape of its container?
a. solid c. gas
b. liquid d. both b and c


____ 7. Which state of matter takes both the shape and volume of its container?
a. solid c. gas
b. liquid d. both b and c


____ 8. Which state of matter is characterized by having an indefinite shape, but a definite volume?
a. gas c. solid
b. liquid d. none of the above


____ 9. Which of the following is a physical change?
a. corrosion c. evaporation
b. explosion d. rotting of food


____ 10. Which of the following is a heterogeneous mixture?
a. air c. steel
b. salt water d. soil


____ 11. Which of the following is true about homogeneous mixtures?
a. They are known as solutions.
b. They consist of two or more phases.
c. They have compositions that never vary.
d. They are always liquids.


____ 12. Separating a solid from a liquid by evaporating the liquid is called ____.
a. filtration c. solution
b. condensation d. distillation


____ 13. The diameter of a carbon atom is 0.000 000 000 154 m. What is this number expressed in scientific notation?
a. 1.54 10 m
c. 1.54 10 m

b. 1.54 10 m
d. 1.54 10 m



____ 14. The expression of 5008 km in scientific notation is ____.
a. 5.008 10 km
c. 5.008 10 km

b. 50.08 10 km
d. 5.008 10 km



____ 15. What is the SI unit of mass?
a. liter c. candela
b. joule d. kilogram


____ 16. What is the temperature of absolute zero measured in C?
a. –373 C
c. –173 C

b. –273 C
d. –73 C



Commonly Used Metric Prefixes
Prefix Meaning Factor
mega (M) 1 million times larger than the unit it precedes 10

kilo (k) 1000 times larger than the unit it precedes 10

deci (d) 10 times smaller than the unit it precedes 10

centi (c) 100 times smaller than the unit it precedes 10

milli (m) 1000 times smaller than the unit it precedes 10

micro ( )
1 million times smaller than the unit it precedes 10

nano (n) 1000 million times smaller than the unit it precedes 10

pico (p) 1 trillion times smaller than the unit it precedes 10



____ 17. What is the quantity 0.0075 meters expressed in centimeters? Use the table above to help you.
a. 0.075 cm c. 7.5 cm
b. 0.75 cm d. 70.5 cm


____ 18. What is the quantity 7896 millimeters expressed in meters? Use the table above to help you.
a. 7.896 m c. 789.6 m
b. 78.96 m d. 789,600 m


____ 19. What is the quantity 987 milligrams expressed in grams? Use the table above to help you.
a. 0.000 987 g c. 9.87 g
b. 0.987 g d. 98,700 g


____ 20. Which of the following equalities is NOT correct? Use the table above to help you.
a. 100 cg = 1 g c. 1 cm = 1 mL

b. 1000 mm = 1 m d. 10 kg = 1 g


____ 21. Who was the man who lived from 460B.C.–370B.C. and was among the first to suggest the idea of atoms?
a. Atomos c. Democritus
b. Dalton d. Thomson


____ 22. Which of the following was NOT among Democritus’s ideas?
a. Matter consists of tiny particles called atoms.
b. Atoms are indivisible.
c. Atoms retain their identity in a chemical reaction.
d. Atoms are indestructible.


____ 23. The smallest particle of an element that retains the properties of that element is a(n) ____.
a. atom c. proton
b. electron d. neutron


____ 24. Dalton's atomic theory included which idea?
a. All atoms of all elements are the same size.
b. Atoms of different elements always combine in one-to-one ratios.
c. Atoms of the same element are always identical.
d. Individual atoms can be seen with a microscope.


____ 25. Why did J. J. Thomson reason that electrons must be a part of the atoms of all elements?
a. Cathode rays are negatively-charged particles.
b. Cathode rays can be deflected by magnets.
c. An electron is 2000 times lighter than a hydrogen atom.
d. Charge-to-mass ratio of electrons was the same, regardless of the gas used.


____ 26. Who conducted experiments to determine the quantity of charge carried by an electron?
a. Rutherford c. Dalton
b. Millikan d. Thomson


Short Answer

27. What is the temperature 128 K expressed in degrees Celsius?

28. What is the density of an object having a mass of 4.0 g and a volume of 39.0 cubic centimeters?

29. What is the volume of an object with a density of 7.73 g/cm3 and a mass of 5.40 10 g?

Essay

30. Explain how Dalton improved upon atomic theory more than 2000 years after Democritus’s hypotheses about atoms.

31. What observations by Rutherford led to the hypothesis that atoms are mostly empty space, and that almost all of the mass of the atom is contained in an atomic nucleus?

EP 4 Test Review Ch 1-4
Answer Section

MULTIPLE CHOICE

1. ANS: C DIF: L1 REF: p. 7 OBJ: 1.1.1

2. ANS: B DIF: L2 REF: p. 8 OBJ: 1.1.1

3. ANS: B DIF: L2 REF: p. 9 OBJ: 1.1.2

4. ANS: D DIF: L2 REF: p. 40 OBJ: 2.1.2

5. ANS: C DIF: L1 REF: p. 42 OBJ: 2.1.3

6. ANS: B DIF: L1 REF: p. 41 OBJ: 2.1.3

7. ANS: C DIF: L1 REF: p. 42 OBJ: 2.1.3

8. ANS: B DIF: L1 REF: p. 41 OBJ: 2.1.3

9. ANS: C DIF: L2 REF: p. 42 OBJ: 2.1.4

10. ANS: D DIF: L1 REF: p. 45 OBJ: 2.2.2

11. ANS: A DIF: L1 REF: p. 45 OBJ: 2.2.2

12. ANS: D DIF: L2 REF: p. 46 OBJ: 2.2.3

13. ANS: D DIF: L1 REF: p. 63 OBJ: 3.1.1

14. ANS: A DIF: L1 REF: p. 63 OBJ: 3.1.1

15. ANS: D DIF: L1 REF: p. 76 OBJ: 3.2.1

16. ANS: B DIF: L1 REF: p. 77 OBJ: 3.2.1

17. ANS: B DIF: L1 REF: p. 84 OBJ: 3.3.2

18. ANS: A DIF: L1 REF: p. 84 OBJ: 3.3.2

19. ANS: B DIF: L1 REF: p. 84 OBJ: 3.3.2

20. ANS: D DIF: L2 REF: p. 84 OBJ: 3.3.2

21. ANS: C DIF: L2 REF: p. 101 OBJ: 4.1.1

22. ANS: C DIF: L2 REF: p. 101 OBJ: 4.1.1

23. ANS: A DIF: L1 REF: p. 101, p. 102
OBJ: 4.1.1, 4.1.2

24. ANS: C DIF: L2 REF: p. 102 OBJ: 4.1.2

25. ANS: D DIF: L2 REF: p. 105 OBJ: 4.2.1

26. ANS: B DIF: L2 REF: p. 105 OBJ: 4.2.1

SHORT ANSWER

27. ANS:
C = K – 273 = 128 – 273 = –145 C

DIF: L2 REF: p. 77, p. 78 OBJ: 3.2.3

28. ANS:
Density = mass/volume = 4.0 g/39.0 cm = 0.10 g/cm

DIF: L2 REF: p. 91 OBJ: 3.4.1

29. ANS:
Volume = mass / density = (5.40 10 g) / 7.73 g/cm = 69.9 cm

DIF: L2 REF: p. 91 OBJ: 3.4.1

ESSAY

30. ANS:
John Dalton used experimental methods to hypothesize that atoms of the same element are identical, and that they differ from atoms of other elements. Dalton also noted that atoms of different elements can form compounds in which the elements combine in whole-number ratios. He noted that individual atoms are not transformed into different atoms as a result of a chemical reaction. What Dalton did not contribute was the fact that individual atoms are divisible into subatomic particles, which came from the discovery of the electron and the atomic nucleus by Thomson and Rutherford, respectively.

DIF: L3 REF: p. 101, p. 102 OBJ: 4.1.1, 4.1.2

31. ANS:
Rutherford’s gold-foil experiment led to this hypothesis. Alpha particles were observed to mostly pass through a gold foil, which suggests that the volume of individual gold atoms consists mainly of empty space. The observation that some alpha particles were scattered at large angles led to the suggestion that the gold atom has a central core, or nucleus, composed of a concentrated mass capable of deflecting the alpha particles.

DIF: L3 REF: p. 107 OBJ: 4.2.2

sample exam ep 6 old questions

Match each item with the correct statement below.
a. acid dissociation constant h. Lewis acid
b. diprotic acid i. pH
c salt hydrolysis j. equivalence point
d. end point k. buffer capacity
e. oxidation number l. oxidizing agent
f. half-reaction m reducing agent
g. hydrogen-ion donor n titration


____ 1. can accept an electron pair

____ 2. acid with two ionizable protons

____ 3. Brønsted-Lowry acid

____ 4. negative logarithm of the hydrogen ion concentration

____ 5. ratio of the concentration of the dissociated to the undissociated form


____ 6. process of adding a known amount of solution of known concentration to determine the concentration of another solution

____ 7. The number of moles of hydrogen ions equals the number of moles of hydroxide ions.

____ 8. Indicator changes color.

____ 9. Cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to water.

____ 10. the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs


____ 11. substance that accepts electrons

____ 12. substance that donates electrons

____ 13. integer related to the number of electrons under an atom's control

____ 14. reaction showing either the reduction or the oxidation reaction












Match each item with the correct statement below.
a. substituent j asymmetric carbon
b. structural isomers k. trans configuration
c. geometric isomers l. cis configuration
d. aromatic compound m. lignite
e. aliphatic hydrocarbon n bituminous coal
f. condensed structural formula o. saturated compound
g. homologous series p. complete structural formula
h. unsaturated compound q Stereoisomers
i. anthracite coal


____ 15. atom or group of atoms that can take the place of a hydrogen in a parent hydrocarbon molecule

____ 16. compounds that have the same molecular formula, but the atoms are joined in a different order

____ 17. arrangement in which substituted groups are on the same side of a double bond

____ 18. molecules in which atoms are joined in the same order but differ in the arrangements of their atoms in space

____ 19. arrangement in which substituted groups are on opposite sides of a double bond

____ 20. compounds that differ in the orientation of groups around a double bond

____ 21. carbon atom to which four different atoms or groups are attached


____ 22. group of compounds in which there is a constant increment of change in molecular structure from one compound in the series to the next

____ 23. formula showing all the atoms and bonds in a molecule

____ 24. structural formula in which some bonds and/or atoms are left out

____ 25. organic compound that contains the maximum number of hydrogens per carbon atom

____ 26. organic compound that contains at least one double or triple carbon-carbon bond


____ 27. any straight-chain or branched-chain alkane, alkene, or alkyne

____ 28. any hydrocarbon compound in which a ring has bonding similar to benzene

____ 29. hard coal, having a carbon content of over 80%

____ 30. brown coal, having a carbon content of approximately 50%

____ 31. soft coal, having a carbon content of 70–80%

EP 6 exam questions

1. what is the formula for sulfuric acid, nitrous acid, and phosphorous acid?
2. what is the difference between chemical and physical properties?
3. Compare an ester, ketone, and aldehyde structures.
4. using the scientific method, outline how the gas collection experiment combining ammonium sulfate and sodium nitrate works or doesn't work? what are sources of error?
5. You have an unknown gas that burns blue in a flame test what is it?
6. in the neutralization of HCl and NaOH, what is your product?

EP 6 Exams topics

Acid naming on exam
Chemical/physical properties
Dilutions
Vocab ch 19-23
Hess’s law


Labs—Mg+HCl
Zn+H2SO4
Ammonium sulfate + sodium nitrate=
NaOH+ HCl=
Know your neutralization reactions
Acid + base= water + salt


Flame test
Hydrogen
Carbon dioxide
Oxygen
Methane
Flame test
Calcium
Sodium
copper

EP 5 Review Questions for exams

Ep 5 review exam
Vocab
Ch 16, 17
Page 498 p. 534

Solubility page 477 # 3-7
Molarity page 481 # 8, 9
Find the moles of solute in solution page 482 # 10-11
Dilutions page 484 # 12-13
Page 486 #16-17, 19-21
Molality page 492 # 29-30
thermochemistry
Page 510 # 3-11
Page 513 Enthalpy #12-13
Page 517 # 16-17, 20
Page 521 #21, 22
Page 523 Heating curve
Page 524 heat of vaporization #23-24
Page 526 # 27-31
Page 535 #38-55
labs
Dilution m1v1=m2v2
Hess’s law
Flame test—hydrogen pops, carbon dioxide puts the fllame out, oxygen, reignites the flame, methane burns blue
Lab equipment
Gas collecting apparatus

Ep 4 exam things

All vocab ch 9, 10
All lab equipment
Video will be posted to acchemistry.blogspot site

Page 291—number of atoms to moles
Page 292 5, 6 moles to atoms
Page 296 molar mass of a compound 5, 6 -15
Page 298 moles to mass 16, 17
Page 299 mass to moles 18, 19

Page 301 volume of gas stp 20-21
Page 302 molar mass of a gas 22-23
Page 303 24-31
Page 306 percent composition 32-33
Page 307 34-35 percent composition from forula
Page 310 36-37 empirical formula
312 molecular formula 38-46


Page 315-317 more practice

Flame Test Hydrogen

Ch 16 Review Questions

molarity
Calculate the molarity of each of the following solutions.
a. 0.40 mol of NaCl dissolved in 1.6 L of solution

20.2 g of potassium nitrate, KNO3, in enough water to make 250.0 mL of
solution
dilution
You must prepare 300.0 mL of 0.750M NaBr solution using 2.00M NaBr stock solution. How many milliliters of stock solution should you use?

5. In order to dilute 1.0 L of a 6.00M solution of NaOH to 0.500M solution
how
much water must you add?
v/v
6. What is the concentration in percent by volume, %(v/v), of the following
solutions?
a. 60.0 mL of methanol in a total volume of 500.0 mL
b. 25.0 mL of rubbing alcohol (C3H7OH) diluted to a volume of 200.0 mL
with water


7. How many grams of solute are needed to prepare each of the following
solutions?
a. 1.00 L of a 3.00% (m/m) NaCl solution?
b. 2.00 L of 5.00% (m/m) KNO3 solution?


How many particles in solution are produced by each formula unit of potassium carbonate, K2CO3?

3. How may moles of particles would 3 mol Na2SO4 give in solution?


Calculate the mole fraction of solute in each of the following solutions.
a. 3.0 moles of lithium bromide, LiBr, dissolved in 6.0 moles of water

Find the molality of each of the following solutions.
a. 2.3 moles of glucose dissolved in 500.0 g of water
b. 131 g of Ba(NO3)2 dissolved in 750.0 g of water

Chem review EP 5

Find the percent by mass of water in
NiCl2 * 6H2O.

In your own words, explain hydrogen bonds.
Draw a diagram of the hydrogen bonding between three water molecules.
Explain why the density of ice at 0°C is less than the density of liquid water at 0°C.

Draw a diagram of the hydrogen bonding between three water molecules.


Explain why the density of ice at 0°C is less than the density of liquid water at 0°C.

Identify the solute and solvent in a dilute aqueous solution of potassium chloride.

Which of the following compounds are soluble in water? Which are insoluble?
a. CaCl2
b. N2
c. HBr
d. NH4C2H3O2

Write the formulas for the following hydrates.
a. Calcium sulfate dihydrate
b. Cobalt(II) chloride hexahydrate

ep 6 Powerpoint review text only

Identify the following acids as monoprotic, di-protic, or triprotic. Explain your reasoning. H2CO3

What is true about the relative concentrations of hydrogen ions and hydroxide ions in each kind of solution?
basic

How many moles of HCl are required to neutralize aqueous solutions of these bases?
2 mol NH3

Write complete balanced equations for the following acid-base reactions?
H2SO4(aq) + KOH(aq) →

What substances are combined to make a buffer?


What is the molarity of sodium hydroxide if 10.0 mL of the solution is neutralized by each of the following 2.00M solutions?
56.0 mL of HCl
predict
Acidic solution?

Naming acids
Hydrochloric
HBr
Acid Nomenclature Flowchart

EP 4 Review Powerpoint text only

How many valence electrons are in each atom?
Carbon
Know your Bonds
Ionic
Covalent
Polar covalent
Coordinate covalent
Network solid

Write the correct chemical formula for the compounds formed from each pair of ions.
K+,S2−

Write formulas for each compound.
barium chloride
VSEPR
Two bonding atoms, one lone pair

Properties of metals
Sea of electrons


formula
Write the formula for these binary compounds.beryllium chloride

Naming acids

Give the names of these acids. HNO2

Test Review ep 5 text only

Test Review EP 5 Sept 2010
Gary Sakuma
Like dissolves like
What is a polar molecule?
Water
Dissolves ionics, sugars, not lipids, methane, hydrocarbons
Formulae and percent
calcium chloride dihydrate
Percent by water
Solubility is affected by
Temperature
Pressure
size
Mole fraction problem
Calculate the mole fraction of each component in a solution of 1.50 mol ethanoic acid (CH3COOH) in 12.00 mol of water.
Colligative properties
What is the freezing point of a solution of 20.0 g of CCl4 dissolved in 500.0 g of benzene? The freezing point of benzene is 5.48°C; Kf is 5.12°C/m.
m/m problem
Calculate the grams of solute required to make the following solutions.
1100 g of saline solution (0.90% NaCl (m/m))
molarity
How many grams of NaCl would you need to make 500 ml of 3.0M

Concentration of Solute
The amount of solute in a solution is given by its concentration.
Molality
What is the molality of 5 moles of KCl in 500 ml of water
Two Other Concentration Units
V/V problems
What is the concentration (in % (v/v)) of the following solutions?
55 mL of ethanol (C2H5OH) is diluted to a volume of 250 mL with water.
Calculating Concentrations
Calculate molality






Boiling Point Elevation and Freezing Point Depression
∆T = K•m•i
i = van’t Hoff factor = number of particles produced per molecule/formula unit. For covalent compounds, i = 1. For ionic compounds, i = the number of ions present (both + and -)
Compound Theoretical Value of i
glycol 1
NaCl 2
CaCl2 3
Ca3(PO4)2 5
Boiling Point Elevation and Freezing Point Depression
∆T = K•m•i
Change in Boiling Point
Dissolve 62.1 g of glycol (1.00 mol) in 250. g of water. What is the boiling point of the solution?
Kb = 0.52 oC/molal for water (see Kb table).
Solution ∆TBP = Kb • m • i

1. Calculate solution molality = 4.00 m
2. ∆TBP = Kb • m • i
∆TBP = 0.52 oC/molal (4.00 molal) (1)
∆TBP = 2.08 oC
BP = 100 + 2.08 = 102.08 oC (water normally boils at 100)

Freezing Point Depression
Calculate the Freezing Point of a 4.00 molal glycol/water solution.
Kf = 1.86 oC/molal (See Kf table)
Solution
∆TFP = Kf • m • i
= (1.86 oC/molal)(4.00 m)(1)
∆TFP = 7.44
FP = 0 – 7.44 = -7.44 oC (because water normally freezes at 0)


Freezing Point Depression
At what temperature will a 5.4 molal solution of NaCl freeze?
Solution
∆TFP = Kf • m • i
∆TFP = (1.86 oC/molal) • 5.4 m • 2
∆TFP = 20.1 oC
FP = 0 – 20.1 = -20.1 oC
Solution Molality
How many grams of potassium iodide must be dissolved in 500.0 grams of water to produce a 0.060 molal KI solution
.5 kg H20 * .6 mole KI/1 kg H2O *
166.0 gKI/1 mol KI
= answer
Types of problems on exam
Ch 15 and 16 vocabulary and reading
Mass/Mass m/m
Volume/volume v/v
Hydrates percent water and writing formulae
Molarity
Molality
Freezing point depression/boiling point elevation T=k m I

EP 5 and 6 Matching Sample

Match each item with the correct statement below.
a. anode d. half-cell
b. battery e. cathode
c. fuel cell


____ 1. the electrode at which oxidation occurs

____ 2. one part of a voltaic cell in which either oxidation or reduction occurs

____ 3. the electrode at which reduction occurs

____ 4. a group of cells that are connected together

____ 5. a voltaic cell in which a fuel substance undergoes oxidation and from which electrical energy is obtained continuously

Match each item with the correct statement below.
a. electrode d. voltaic cell
b. electrolysis e. dry cell
c. salt bridge


____ 6. a tube containing a conducting solution

____ 7. a conductor in a circuit that carries electrons to or from a substance other than a metal

____ 8. an electrochemical cell that is used to convert chemical energy to electrical energy

____ 9. a voltaic cell in which the electrolyte is a paste

____ 10. a process in which electrical energy is used to bring about a chemical change

Match each item with the correct statement below.
a. activity series of metals c. combustion reaction
b. single-replacement reaction d. decomposition reaction


____ 24. a reaction in which a single compound is broken down into simpler substances

____ 25. a reaction in which oxygen reacts with another substance, often producing heat or light

____ 26. a reaction in which the atoms of one element replace the atoms of a second element in a compound

____ 27. a list of metals in order of decreasing reactivity

matching answers
MATCHING

1. ANS: A DIF: L1 REF: p. 665 OBJ: 21.1.2

2. ANS: D DIF: L1 REF: p. 665 OBJ: 21.1.2

3. ANS: E DIF: L1 REF: p. 665 OBJ: 21.1.2

4. ANS: B DIF: L1 REF: p. 668 OBJ: 21.1.4

5. ANS: C DIF: L1 REF: p. 669 OBJ: 21.1.5

6. ANS: C DIF: L1 REF: p. 665 OBJ: 21.1.2

7. ANS: A DIF: L1 REF: p. 665 OBJ: 21.1.2

8. ANS: D DIF: L1 REF: p. 665 OBJ: 21.1.2

9. ANS: E DIF: L1 REF: p. 667 OBJ: 21.1.3

10. ANS: B DIF: L1 REF: p. 678 OBJ: 21.3.1


24. ANS: D DIF: L1 REF: p. 332 OBJ: 11.2.1

25. ANS: C DIF: L1 REF: p. 336, p. 337
OBJ: 11.2.1

26. ANS: B DIF: L1 REF: p. 333 OBJ: 11.2.1

27. ANS: A DIF: L1 REF: p. 333 OBJ: 11.2.2

28. ANS: C DIF: L1 REF: p. 294 OBJ: 10.1.3

29. ANS: B DIF: L1 REF: p. 294, p. 295
OBJ: 10.1.3, 10.1.4

30. ANS: A DIF: L1 REF: p. 300 OBJ: 10.2.2

p. 681 OBJ: 21.3.3