Parents Survival Guide to Chemistry

Parents’ Survival Guide to Chemistry

The Parents Survival Guide to Chemistry How do I help my kids with their Chemistry Homework? It has been years since I studied it and I don’t remember a thing about it. Sound familiar? Most parents feel that they cannot help their children with Chemistry. But they are wrong. Experience has shown that many students are unsuccessful in Chemistry because they do not take the time to learn the basics. Trying to understand Chemistry without a good grasp of the basics is like trying to be successful in High School English without having learned all the letters of the alphabet. But your child does know all the letters of the alphabet because you taught it to them. You can do the same with the basics of Chemistry. Much of the basic knowledge of Chemistry can be memorized. In this guide, items that can be memorized have been highlighted in yellow. You can use this guide to teach and quiz your child.

SJHS Chemistry Dept.

Page 1

Classification of Elements VII IA

VB

VIB

VII B

----------VIIIB---------

IB

IIB

VIA

Transition Metals

VII A

Metals

Lanthanum Series (rare earths) Actinium Series (transuraniums) Metal + Nonmetal = Binary Ionic Compounds 2 Nonmetals = Binary Molecular Compounds Hydrogen + Negative Ion = Binary Acids Students must memorize the parts of the periodic table. They must also learn the types of elements present in each of the types of compounds.

SJHS Chemistry Dept.

Page 2

Noble Gases

IVB

VA

Halogens

Alkaline Earths

Alkali Metals

III B

IVA

Metals

III A

IIA

Non

IA

Naming Compounds and Writing Formulas Students must learn to recognize the type of compound that they are dealing with before trying to name it or write its formula. This is absolutely essential. Binary Molecular Compounds Consist of 2 non-metal elements bonded covalently. Naming 1. Given the formula, determine the name of each element present. 2. List the name of the first element listed in the formula. If there is more than one of that type of atom, use the appropriate prefix*. 3. List the name of the second element listed in the formula. Use the appropriate prefix for it (including mono if there is just one of them). Change the ending to -ide eg CO2 becomes carbon dioxide Writing Formulas 1. Write the symbol for the first element listed in the name. If it has a prefix, use the number represented by that prefix as a subscript. 2. Write the symbol for the second element listed in the name. Convert its prefix to a subscript. eg dinitrogen tetraoxide becomes N2O4 *Prefixes mono=1 hexa=6

di=2 hepta=7

tri=3 octa=8

prefixes are not used for hydrogen

SJHS Chemistry Dept.

Page 3

tetra=4 nona=9

penta=5 deca=10

Binary Ionic Compounds Consist of a metal and a non-metal bonded ionically Naming 1. List the name of the metal ion first. Some metal ions have more than one charge. You have to figure out what its charge is by looking at the charge of the negative ion and how many of them there are to balance out the positive and negative charges. For metal ions with more than one charge, show the charge as a roman numeral following the name. (see example below) The sum of all the positive and negative charges has to add up to zero. If the metal ion has a subscript, it is not included in the name. 2. List the negative ion next. All simple negative ions end in -ide. eg. CaCl2 = calcium chloride CuI2= Copper(II) iodide V2O5= Vanadium(V) oxide Writing Formulas The charges of the positive and negative ions must add up to zero. 1. When given the name of a binary ionic compound first write the symbols for the ions involved (metal ion first) eg

Silver chloride Ag+ Calcium fluoride Ca2+

ClF-

2. Next, determine the lowest whole number ratio of ions which will provide an overall net charge of zero. For silver chloride, one Ag+ and one Cl- gives a formula AgCl For calcium fluoride, one Ca2+ and two F- gives a formula CaF2

SJHS Chemistry Dept.

Page 4

Ionic Compounds Containing Polyatomic Ions Usually consist of a metal, a nonmetal atom and oxygen atoms eg. K2CO3 Polyatomic ions are groups of atoms that have bonded together and then gained (or lost) electrons to become more stable. They have a naming system that we will not discuss now. Our periodic tables feature a list of the most common polyatomic ions. There is one common positive polyatomic ion that we will use. It is the ammonium ion NH4+. It should be memorized. Naming Ionic Compounds Containing Polyatomic Ions (except those containing ammonium) 1. Name the metal ion listed. 2. On the table of polyatomic ions, find the ion given in the formula. The most common ending for polyatomic ions is -ate. Next most common is -ite. One or two of them end in -ide. eg

Na2SO4 becomes sodium sulphate Ca(ClO2)2 becomes calcium chlorite KCN becomes potassium cyanide

SJHS Chemistry Dept.

Page 5

Molecular Elements

The following elements must form “free state”. Memorize them. hydrogen H2 nitrogen N2 oxygen O2 fluorine F2

molecules to remain stable when they are in the chlorine bromine iodine astatine

Nonsystematic Names for Certain ozone O3 water methane CH4 sucrose methanol CH3OH ethanol Propane

Cl2 Br2 I2 At2

phosphorus sulphur

Common Compounds (Memorize) H2O ammonia NH3 C12H22O11 glucose C6H12O6 C2H5OH hydrogen H2O2 peroxide

C3H8

Binary Molecular Compounds 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25

Molecular Formula CCl4 CO2 NO NO2 SO3 P4O10 P4O6 CH4 HCl H2O

SJHS Chemistry Dept.

P4 S8

Name

nitrogen sulphur dioxide carbon monoxide ozone ethanol sucrose sulphur chlorine dioxide methanol phosphorus ammonia dinitrogen oxide iodine glucose propane

Page 6

Binary Ionic Compounds Empirical Formula 1

CaCl2

2

MgO

3

NaBr

4

Al2O3

5

CaO

6

ZnO

7

Ag2S

8

CaF2

9

CaH2

Name

10

potassium iodide

11

aluminum chloride

12

lithium nitride

13

barium chloride

14

sodium chloride

15

silver bromide

16

magnesium hydride

17

magnesium chloride

18

zinc chloride

19

potassium chloride

20

sodium sulphide

21

zinc sulphide

22

aluminum chloride

23

scandium bromide

24

cesium iodide

25

strontium fluoride

SJHS Chemistry Dept.

Page 7

Multiple Ion Charges Empirical Formula 1

Cu2S

2

SnO2

3

NiBr2

4

AuCl3

5

TiO2

6

V2O5

7

HgO

8

FeS

9

MoS2

10

HgS

11

Sb2S3

12

Bi2O3

Name

13

uranium(IV) oxide

14

lead(IV) sulphide

15

manganese (IV) oxide

16

ferric oxide

17

cobaltous chloride

18

uranium (VI) fluoride

19

chromic oxide

20

stannous fluoride

21

lead(IV) oxide

22

copper(II) sulphide

23

platinum (IV) chloride

24

mercuric chloride

25

Titanium (IV) oxide

SJHS Chemistry Dept.

Page 8

Nomenclature Involving Polyatomic Ions Note: mixed in with these compounds possessing polyatomic ions are some molecular compounds.

Empirical Formula 1

K2CO3

2

(NH4)2S

3

Cr(NO3)3

4

SO3

5

NaNO2

6

K3PO4

7

K2Cr2O7

8

KMnO4

9

NaHCO3

10

NH4H2PO4

11

Na2SO4

Name

12

calcium hydroxide

13

magnesium silicate

14

iron(II) chlorite

15

ammonium dichromate

16

ammonium sulphate

17

calcium stearate

18

sodium nitrate

19

sodium thiosulphate

20

barium perchlorate

21

sodium hydrogen sulphide

22

potassium cyanide

23

sodium glutamate

24

potassium thiocyanate

25

lead(II) acetate

SJHS Chemistry Dept.

Page 9

Acids Rule 1

Hydrogen _____-ide becomes hydro_____-ic acid

Rule 2

Hydrogen _____-ate becomes _____-ic acid

Rule 3

Hydrogen _____-ite becomes _____-ous acid

Formula 1

H3BO3(aq)

2

H2SO4(aq)

3

H2CO3(aq)

4

H2S(aq)

5

H2SiO3(aq)

6

CH3COOH(aq)

7

H2SO3(aq)

8

HClO4(aq)

9

HCN(aq)

10

C6H5COOH(aq)

Name

11

thiosulphuric acid

12

nitrous acid

13

chromic acid

14

hydrofluoric acid

15

hypochlorous acid

16

nitric acid

17

stearic acid

18

phosphoric acid

19

oxalic acid

20

hydrochloric acid

Memorize the rules!

SJHS Chemistry Dept.

Page 10

SI Base Units Quantity Length Mass Time Electric Current Thermodynamic Temperature Amount of Substance Luminous Intensity Prefix tera giga mega* kilo* hecto deka

Symbol T G M k h da

Name metre kilogram second ampere Kelvin mole candela SI Prefixes Meaning trillion billion million thousand hundred ten

Symbol m kg s A K mol cd Multiplier 1012 109 106 103 100 10

deci* d tenth 0.1 centi* c hundredth 0.01 milli* m thousandth 0.001 micro* m millionth 10-6 nano* n billionth 10-9 pico p trillionth 10-12 femto f quadrillionth 10-15 atto a quintillionth 10-18 *These prefixes are commonly used and should be memorized. Rules for Counting Significant Digits Assuming that a measurement has been properly taken 1. All non-zero digits (1-9) are significant (eg. 13.4 s) 2. All zeroes between significant digits are significant (eg. 204 g) 3. Trailing zeroes in a decimal are significant (eg. 1l.230 m) 4. Zeroes to the left of all non-zeroes (leading zeroes) are not significant (eg. 0.0034) 5. Trailing zeroes in a whole number create a problem. There may be some circumstances when you want them to be significant and other circumstances where you do not want them to be. The best thing to do is to use scientific notation. For example, in the number 1390, if you want the 0 to be significant, write it as 1.390 x 103. If you do not want it counted as significant, write it as 1.39 x 103. As a local rule, for our convenience, we will use an additional notation. If we want the 0 counted as significant, we will put a decimal point in. Therefore, 1390. has 4 significant digits. 1390 has only 3.

SJHS Chemistry Dept.

Page 11

Rules for Rounding Off 1. If the first digit eliminated is less than 5, the last digit retained remains the same. 2. If the first digit eliminated is greater than 5, the last digit retained increases by 1. 3. If the first digit eliminated is 5 and it is followed at some point by another non-zero digit, the last digit retained increases by 1. 4. If the first and only digit eliminated is 5, the last digit retained increases by 1 if it is odd, but remains the same if it is even. Rules for Addition & Subtraction of Significant Digits Perform the addition or subtraction then round the answer off to the same number of decimal places as the least precise number in the calculation. Rules for Multiplication & Division of Significant Digits Perform the multiplication or subtraction then round the answer off to the same number of significant digits as the number (in the calculation) having the least number of significant digits.

These rules should be memorized. The SI base units and the prefixes with an asterisk next to them should also be memorized.

SJHS Chemistry Dept.

Page 12

SI Assignment Read Pages 28 to 38 of Addison-Wesley and answer the following questions 1. What one unit can all measurements of length be measured in? 2. What is it more convenient to sometimes do for extremely small or large measurements of length instead of using the unit mentioned in question 1? 3. List the approximations of SI distances or lengths shown in the table at the bottom of page 30. 4. Define the word volume. Why is it considered a derived unit? 5. What is the SI unit of volume? 6. What more convenient unit is often used? Define it. 7. When measuring extremely small volumes, what volume unit is used instead of the one mentioned in question 6? How big is it compared to that unit? 8. What measuring device is recommended for approximate measurements of volume in the lab? What is recommended for more precise mesurements? 9. What is mass? Define it. What is weight? How are they different? 10. What is the SI unit of mass? Define it. 11. Define density. 12. Define temperature. 13. When does heat transfer occur? 14. How do mercury thermometers work? 15. Who invented the Celsius scale? How did he define it? 16. What is the Kelvin Scale? 17. What is absolute zero?

SJHS Chemistry Dept.

Page 13

Sample Problems 1. How many significant figures are there in each of the following numbers? a) 143 g b) 0.074 cm c) 8.750 x 102 g d) 1.003 km e) 400 m f) 400. m 2. Round off each of these measurements to three significant figures a) 98.473L b) 0.00007333 g c) 543.4999 d) 14.25 kg e) 15.35 g f) 1.002 3. Perform each of the following calculations and round off to the appropriate number of significant figures. a) 5.3 x 10+4 g ÷0.2331 mol b) 8.7g + 15.43 g + 19g c) 38.742 kg ÷ 0.421 L d) 5.47 m + 11.1 m + 87.3000 m e) 853.2 L - 627.443 L 4. Perform the following conversions a) 4.23 m = _______ cm b) 14.239 nm = ________ mm c) 3.3 x 105 ng = _______ µg d) 4.75 kg/L to _________g/mL e) 1.42 g/cm3 = ________ µg/mm3 f) 14.3 dollars/kg = ________ cents/g 5. A mixture contains 14.3g sodium carbonate, 15.9 g aluminum sulphate and 22.3g sodium chloride. How much sodium chloride would be in 10.0g of the mixture.

SJHS Chemistry Dept.

Page 14

Mole ⇄ Gram Conversions Factor-Label Technique (memorize these steps) 1. Determine the relationship between the known and the unknown units. 2. Create two ratios from this relationship. 3. Set the given quantity (to be converted) equal to itself. 4. Multiply the right hand side of the equation by whichever ratio cancels out the old units and leaves you with the old units. eg. 14.2 mol H2 = _______ g 1. Relationship: 1 mol H2=2.02g 2. Ratios: 1 mol/2.02g or 2.02g/1 mol 3. Set given quantity equal to itself: 14.2 mol H2 = 14.2 mol H2 4. Multiply RHS by ratio: 14.2 mol H2 = 14.2 mol H2 x 2.02g/1 mol =28.7 g Exercise Perform the following conversions in your notebook. Show your work. 1. 2.02g NaCl = __________ mol

10. 50.0g NO2 = _______ mol

2. 14.2 g methane = _______mol 3. 100.g copper(II) chloride = ____ mol 4. 1000. g NaHCO3 = _________ mol 5. 1.298 mol sucrose = _________ g 6. 2.345 mol calcium chloride = ___ g

11. 1000. g CH4 = ___________L

7. 2.50 mol Ca = _______ atoms 8. 2.50 x 1025 atoms He = _____ mol 9. 1000. g CaCl2 = _______ mol

SJHS Chemistry Dept.

12. 2.22 x 1024 molecules CO2 = ____________ g 13. 100. L O2 = __________ g 14. 50.0g NO2 = _________ molecules 15. 50.0L Cl2 = __________ molecules

Page 15

Balancing Chemical Equations Balance the following equations 1. Ca + O2 → CaO 2. HI + NaOH → HOH + NaI 3. C3H8 + O2 → CO2 + H2O 4. HgO → Hg + O2 5. Zn + CuSO4 →

ZnSO4 + Cu

6. CH3OH + O2 → CO2 + H2O 7. Pb(NO3)2 + KI → PbI2 + KNO3 8. H2 + O2 → H2O 9. Al2O3 → Al + O2 10. HCl + Zn → ZnCl2 + H2

SJHS Chemistry Dept.

Page 16

Reaction Types Memorize these types and their characteristics. 1. Combination (also known as Simple Composition or Synthesis) element + element →compound example: 2Mg + O2 → 2MgO 2. Simple Decomposition binary compound → 2 (or more) elements example: 2H2O →2H2 + O2 3. Single Replacement element + compound → new element + new compound Cu + 2AgNO3(aq) → 2Ag + Cu(NO3)2(aq) 4. Double Replacement 2 compounds →2 new compounds (switch partners) HCl + NaOH →NaCl + H2O 5. Combustion compound containing C and H (and possibly N, O or S)

→

+ O2 CH4 + 2O2 → CO2 + 2H2O

SJHS Chemistry Dept.

Page 17

CO2 + H2O (+NO2 + SO2)

Predicting Chemical Equations Exercise Combination and Decomposition Predict the products in each of the following chemical equations. Balance each of the following. If an equation is already balanced, indicate so. Classify each one as combination or decomposition 1. Cu + O2 → 2. H2O → 3. Fe + S8 → 4. NaCl → 5. Al + O2 → 6. mercuric oxide decomposes 7. C + O2 → 8. Li + N2 → 9. S8 + O2 → 10. N2 + H2 → 11. Ag + S8 → 12. magnesium reacts with oxygen 13. CuI2

→

14. Fe + O2 → 15. H2 + Cl2 →

SJHS Chemistry Dept.

Page 18

Predicting Chemical Equations Exercise Single and Double Replacement Predict the products in each of the following chemical equations. Balance each of the following. If an equation is already balanced, indicate so. Classify each one as combination or decomposition 1. Na(s) + HOH(l) → + H2SO4(aq) →

2.

NaCl(s)

3.

Al(s)

4.

Ca(OH)2(aq) + Mg(HCO3)2(aq) →

5.

K(s)

6.

Cu(s)

7.

H2SO4(aq) + Ca3(PO4)2(s)

8.

Cl2(g)

9.

H2S(g)

+ Fe2O3(s)

+ AlCl3(s)

10. H2S(g)

→

→

+ Ag2SO4(aq) → →

+ MgBr2(aq) → + PbCrO4(s) + Ag(s)

→

→

11. Cl2(g)

+ KI(aq) →

12. Cu(s)

+ AgNO3(aq) →

13. Zn(s)

+ HCl(aq) →

14. Ca(s)

+ H2SO4(aq) →

15. HCl(aq) + KOH(aq) → 16. F2(g)

+ KCl(aq) →

17. Zn(s)

+ CuSO4(aq) →

18. CaCO3(s)

+ HCl(aq) →

19. NaOH(aq) + H3PO4(aq) →

SJHS Chemistry Dept.

Page 19

Predicting Chemical Equations Exercise Combustion Reactions Predict the products in each of the following chemical equations. Balance each of the following. If an equation is already balanced, indicate so. Classify each one as combination or decomposition 1. CH4(g) + O2(g) → 2. CH3OH(l) + O2(g) → 3. CS2(l) + O2(g) → 4. C3H8(g) + O2(g) → 5. Mg(s) + O2(g) → 6. C2H6(g) + O2(g) → 7. C2H5SH(l) + O2(g) → 8. CH3NO2(g) + O2(g) → 9. C12H22O11(s) + O2(g) → 10. C8H18(l) + O2(g) → 11. S8(s) + O2(g) → 12. H2S(g) + O2(g) → 13. C2H5OH(l) + O2(g) → 14. C4H10(g) + O2(g) → 15. C25H52(s) + O2(g) →

SJHS Chemistry Dept.

Page 20

Review Exercise for Predicting Reactions 1. Ca + O2 → 2. C3H8 + O2 → 3. CuCl2 + Zn → 4. HgO → 5. H2SO4 + NaOH → 6. Li + HOH → 7. C2H5OH + O2 → 8. H2O → 9. NaCl → 10. H3PO4 + Fe(OH)3 → 11. Al2O3 + Fe → 12. HCl + NH4OH → 13. FeCl3 + NH4OH → 14. Pb(NO3)2 + KI → 15. Zn + HCl → 16. Cl2 + KI → 17. C2H5SH + O2 → 18. Ca + HOH → 19. Mg + O2 → 20. CH3NO2 + O2 →

SJHS Chemistry Dept.

Page 21

Steps for Solving Stoichiometry Problems These steps should be memorized. 1. Obtain a balanced chemical equation and determine the known and unknown. 2. Convert the starting units to moles 3. Convert moles of the known to moles of the unknown using the mole ratio (which is obtained from the coefficients in the equation) 4. Convert moles of the newly-found quantity to the desired units. 5. Write a concluding statement.

SJHS Chemistry Dept.

Page 22

Chapter 12: Chemical Periodicity Read Chapter 12 in Addison-Wesley. As you read it, answer the following questions. You are responsible for them on the final exam. Write your answers in your notebook in complete sentences, so that they will make sense when you study them later. 1. Define the following: periods representative elements noble gases transition elements inner transition elements covalent atomic radius ionization energy electron affinity electronegativity 2. Who developed the periodic table upon which we base our modern periodic table? What led him to develop it? 3. In what order did he arrange the elements? (see p. 272 margin notes) 4. How did Moseley say the elements should be arranged? 5. How does the covalent atomic radius of an element change as you go from left to right on the periodic table (ie. as you go through a period)? 6. How does the covalent atomic radius of an element change as you go down a group? 7. How does ionization energy change as you go from left to right on the periodic table? 8. How does ionization energy change as you go down a group? 9. How does electron affinity change as you go from left to right on the periodic table? 10. How does electron affinity change as you go down a group? 11. How does ionic size change as you go from left to right on the table? 12. How does ionic size change as you go down a group? 13. How does the size of a positive ion compare with the size of the atom from which it is formed? 14. How does the size of a negative ion compare with the size of the atom from which it is formed? 15. How does electronegativity change as you go from left to right on the table? 16. How does electronegativity change as you go down a group? 17. In what field are Herzberg's main contributions?

SJHS Chemistry Dept.

Page 23

VSEPR Shapes Summary Chemistry 122 Number of Bonds 2 3 4 3 2

Number of Lone Pairs 0 0 0 1 2

5

0

4 3 2 6 5 4

1 2 3 0 1 2

Shape Name Linear trigonal planar tetrahedron pyramidal v-shape trigonal bipyramid SF4 shape T-shape linear octahedron square pyramid square planar

Type of Molecule AX2 AX3 AX4 AX3E AX2E2

Example BeCl2, CH2O BF3, C2H4 CH4 NH3 H2O

AX5

PCl5

AX4E AX3E2 AX2E3 AX6 AX5E AX4E2

SF4 ClF3 XeF2 SF6 IF5 XeF4

The AXE designation is a short-hand method for describing the number of bonds and lone pairs attached to a central atom. A refers to the central atom. X refers to the atoms bonded to the central atom. E refers to the lone pairs. Hence, AX2E2 refers to a central atom with 2 atoms bonded to it and 2 lone pairs on it.

The parts of this table highlighted in yellow are required for Chemistry 122, 121 and IB Chemistry courses. The items highlighted in green are needed for IB Chemistry only.

SJHS Chemistry Dept.

Page 24

Chapter 25 Self-Study Assignment: Hydrocarbons Chem 122/121 You are to work through the questions in this exercise. 1. Read Chapter 25. Define or explain all of the highlighted terms. In addition, define the term Organic Chemistry. 2. Prior to 150 years ago, what theory did scientists hold about the formation of carbon compounds? Who performed experiments to counter this theory? Briefly explain his experiment. How would we now define an organic compound? 3. What distinguishes alkanes from the other hydrocarbons? 4. How many bonds will carbon atoms almost always make? 5. Give the names, formulas and draw the structural formulas for the first ten straight-chain (or continuous chain) alkanes. Memorize them! 6. What organization recommends the use of the names that you found in question 5? 7. What is the difference between a continuous chain alkane and a branched chain alkane? 8. What is a substituent? What does it usually take the place of on a hydrocarbon chain? What are some typical substituents mentioned on page 595? 9. What is a methyl group? What is an ethyl group? What type of substituents are these examples of? Give the names and formulas for the next 5 alkyl groups. 10. How is:

different than

After all, both of them can be written as C4H10. 11. List all of the rules for naming branched-chain alkanes. Memorize them. 12. Draw structural formulas for a) 2-methyl pentane

SJHS Chemistry Dept.

Page 25

b) 3-methyl hexane c) 3-ethyl-2-methyl octane 13. Try questions 4 and 5 on p. 598. 14. What are the alkenes and alkynes collectively known as? Why? 15. How do you name an alkene by the IUPAC system? 16. How are the following two structures different? How would you name them? CH2=CH-CH2-CH3

and CH3-CH=CH-CH3

17. Look at the two versions of 2-butene at the top of page 602. How are they different from each other? What terms are used to distinguish between the two isomers? 18. Do question 7 p. 602. 19. What is an assymetric carbon (section 25-7)? What do these compounds possess? What does nonsuperimposable mean? What are stereoisomers? Do question 9 p. 604. 20. How are cyclic hydrocarbons different than alkanes, alkenes and alkynes?(sec. 25-8) 21. What are the arenes? What is the simplest one? What other term is applied to them? 22. Draw the resonance structures of benzene. 23. Do question 10 p. 606. 24. Do question 25 p. 611

SJHS Chemistry Dept.

Page 26

Chapter 26 Self-Study Assignment: Functional Groups and Organic Reactions Chem 122/121 1. Read Chapter 26 and define all of the bold print words. Omit sections 12 to the end. 2. What symbol is used to represent a carbon chain or ring attached to a functional group? (section 1) 3. Copy the table 26-1 found at the top of page 616 into your notes. 4. Do question 1 on p. 616. 5. What is a compound called when a hydrocarbon has a halogen substituted onto it in place of one of the hydrogens? What is the difference between alkyl halides and aryl halides? 6. Name the following: (See section 2)

7. According to section 3, what is one major reaction that is used to make alkyl halides? 8. Complete the following reaction: CH4 + Br2 → 9. Section 6 discusses a second type of reaction of hydrocarbons. It is called addition. What kind of hydrocarbon can take part in an addition reaction?

SJHS Chemistry Dept.

Page 27

10. According to section 6, what sorts of substances can add to a hydrocarbon? (see pink on pp. 624 & 625) For each type of compound that can add, what sort of product do you get? Copy down the examples illustrated on pp. 624 & 625. 11. Do question 6 a, b, c, d and 7a on p. 626. 12. What are alcohols? (section 4) Explain the differences between primary, secondary and tertiary alcohols (top of p. 621) 13. Make careful note of the examples of alcohols on p. 621. Where do you suppose the name for methanol comes from? Ethanol? Propanol? Why is there a 1-propanol and a 2-propanol? Where have the roots of each of these names come from? Where does the suffix come from? 14. What are glycols? What are phenols? Draw the formula for the parent of this group, which is, itself, named phenol. 15. Do question 4 and 5 on p. 622. 16. Why do small alcohols dissolve in water, even though hydrocarbons do not? Why do the larger ones not dissolve as readily in water? (see section 5) Explain why methanol has a much higher boiling point than methane. (types of intermolecular forces) 17. What is fermentation? Write an equation to describe it. 18. What is the old name for methanol? Why was it called this? Copy down the last two sentences on p. 623. 19. What are ethers? Write formulas for ethyl methyl ether, diethyl ether, dimethyl ether and ethyl propyl ether. For what was diethyl ether used for over 100 years. Why was it replaced? Why are ethers higher boiling than comparable sized hydrocarbons, but lower boiling points than the alcohols? Explain why they are more soluble in water than hydrocarbons but less soluble than alcohols. (Section 7) 20. What is a carbonyl group? What is the difference between aldehydes and ketones? Why is this compound known as 2-pentanone pentanone?

instead of just

21. Do #8 on p. 628

SJHS Chemistry Dept.

Page 28

22. Why are the aldehydes and ketones (particularly the smaller ones) soluble in water? Are aldehydes and ketones soluble in nonpolar solvents? 23. Note table 26-5 on p. 629. What is the common name for methanal? What is it used for? What is the common name for propanone? 24. What are carboxyl groups? What is the general formula for a carboxylic acid? Describe how carboxylic acids ionize in water? (see p. 630) 25. Will carboxylic acids dissolve in water? Why? Do they have high boiling points? Is there any connection between these two questions? 26. Make careful note of the table on p. 631. How do the names of the carboxylic acids compare to that of the alkanes? What is the common name of methanoic acid? What is the french word for ant? Any connection? What is the common name for ethanoic acid? 27. What is oxidation of a hydrocarbon also known as? (p. 632). What does an alkane produce when it is oxidized? What does an alkene produce when it is oxidized? 28. What does a primary alcohol produce when it is oxidized? What does a secondary alcohol produce when it is oxidized? Could a tertiary alcohol be oxidized? Why not? What does methanol produce when it is oxidized? How does this relate to the fact that methanol is one of the more toxic alcohols? 29. What will aldehydes oxidize to produce? 30. Using your answers to questions 28 and 29, explain why old wine (ethanol is the alcohol in all alcoholic beverages) turns to vinegar. 31. Can ketones generally be oxidized? 32. Do Questions 9 and 10 on p. 634. 33. What are esters? How are they produced? What are esters used for? Copy down the esterification reactions shown on p. 636. Do question 32 on p. 649.

SJHS Chemistry Dept.

Page 29

Answers to Exercises on pages 6 to 10, 14 to 16 and 18 to 21. p. 6 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 p. 7 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25

Molecular Formula CCl4 CO2 NO NO2 SO3 P4O10 P4O6 CH4 HCl H2O

Name Carbon dioxide Nitrogen monoxide Nitrogen dioxide Sulphur trioxide Tetraphosphorus decaoxide Tetraphosphorus hexaoxide Methane Hydrogen chloride water

nitrogen sulphur dioxide carbon monoxide ozone ethanol sucrose sulphur chlorine dioxide methanol phosphorus ammonia dinitrogen oxide iodine glucose propane

SO2 CO O3 C2H5OH C12H22O11 S8 ClO2 CH3OH P4 NH3 N2 O I2 C6H12O6 C3H8

Empirical Formula CaCl2 MgO NaBr Al2O3 CaO ZnO Ag2S CaF2 CaH2 KI Li3N BaCl2 NaCl AgBr MgH2 MgCl2 ZnCl2 KCl Na2S ZnS AlCl3 ScBr3 CsI SrF2

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

Carbon tetrachloride

N2

AlCl3

p. 8

16 17 18 19 20 21 22 23 24 25

Name Magnesium oxide Sodium bromide Aluminum oxide Calcium oxide Zinc oxide Silver sulphide Calcium fluoride Calcium hydride

potassium iodide aluminum chloride lithium nitride barium chloride sodium chloride silver bromide magnesium hydride magnesium chloride zinc chloride potassium chloride sodium sulphide zinc sulphide aluminum chloride scandium bromide cesium iodide strontium fluoride

SJHS Chemistry Dept.

PbS2 MnO2 Fe2O3 CoCl2 UF6 Cr2O3 SnF2 PbO2 CuS PtCl4 HgCl2 TiO2

1 2 3 4 5 6 7 8 9 10

NH4H2PO4

11 12 13 14

Na2SO4

15

(NH4)2Cr2O7

16 17 18 19 20 21 22 23

Page 30

UO2

Empirical Formula K2CO3 (NH4)2S Cr(NO3)3 SO3 NaNO2 K3PO4 K2Cr2O7 KMnO4 NaHCO3

p. 9

Calcium chloride

Empirical Formula Cu2S SnO2 NiBr2 AuCl3 TiO2 V2O5 HgO FeS MoS2 HgS Sb2S3 Bi2O3

Name Copper(II) sulphide Tin(IV) oxide Nickel (II) bromide Gold (III) chloride Titanium (IV) oxide Vanadium (V) oxide Mercury (II) oxide Iron (II) sulphide Molybdenum (IV) sulphide Mercury (II) sulphide Antimony (III) sulphide Bismuth (III) oxide

uranium(IV) oxide lead(IV) sulphide manganese (IV) oxide ferric oxide cobaltous chloride uranium (VI) fluoride chromic oxide stannous fluoride lead(IV) oxide copper(II) sulphide platinum (IV) chloride mercuric chloride Titanium (IV) oxide Name Potassium carbonate Ammonium sulphide Chromium (III) nitrate Sulphur trioxide Sodium nitrite Potassium phosphate Potassium dichromate Potassium permanganate Sodium hydrogen carbonate Ammonium dihydrogen phosphate Sodium sulphate

Ba(ClO4)2

calcium hydroxide magnesium silicate iron(II) chlorite ammonium dichromate ammonium sulphate calcium stearate sodium nitrate sodium thiosulphate barium perchlorate

NaHS

sodium hydrogen sulphide

KCN

potassium cyanide sodium glutamate potassium thiocyanate lead(II) acetate

Ca(OH)2 MgSiO3 FeCl2

(NH4)2SO4 Ca(C17H35COO)2

NaNO3 Na2S2O3

NaC5H8O4

24

KSCN

25

Pb(CH3COO)2

p. 10

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Formula H3BO3(aq) H2SO4(aq) H2CO3(aq) H2S(aq) H2SiO3(aq)

p. 17

Name Boric acid Sulphuric acid Carbonic acid Hydrosulphuric acid Silicic acid Acetic acid Sulphurous acid Perchloric acid Hydrocyanic acid Benzoic acid thiosulphuric acid nitrous acid chromic acid hydrofluoric acid hypochlorous acid nitric acid stearic acid phosphoric acid oxalic acid hydrochloric acid

CH3COOH(aq)

H2SO3(aq) HClO4(aq) HCN(aq) C6H5COOH(aq)

H2S2O3(aq) HNO2(aq) H2CrO4(aq) HF(aq) HClO(aq) HNO3(aq) C17H35COOH(aq)

H3PO4(aq) H2C2O4(aq) HCl(aq)

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. p. 19 1. 2. 3. 4.

p. 14

5. 6. 7.

1. a) 3 b) 2 c) 4 d) 4 e) 1 f) 3 2. a) 98.5 b) 0.0000733 c) 543 d) 14.2 e) 15.4 f) 1.00

8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

5

3. a) 2.27 x 10 g/mol b) 43g c) 92.0 kg/L d) 103.07m e) 225.8 L -5

4. a) 423 b) 14.239 x 10

c) 330 d) 4.75

e) 1420 f) 1.43 5. 4.25 g NaCl p. 15 1. 0.0346

2. 0.885 3. 0.744 4. 11.90 5. 444.4

6. 260.2 7. 1.50 x 10

24

8. 41.5 9. 9.011 10. 1.09

11. 1396 12. 0.162 13. 143 14. 6.54 x 10 15. 1.34 x 10

23

24

p. 16 1. 2Ca + O2 → 2CaO 2. HI + NaOH → HOH + NaI 3. C3H8 + 5O2 → 3CO2 + 4H2O 4. 2HgO → 2Hg + O2 5. Zn + CuSO4 →

ZnSO4 + Cu

6. 2CH3OH + 3O2 → 2CO2 + 4H2O 7. Pb(NO3)2 + 2KI → PbI2 + 2KNO3 8. 2H2 + O2 → 2H2O 9. 2Al2O3 → 4Al + 3O2 10. 2HCl + Zn → ZnCl2 + H2

SJHS Chemistry Dept.

2Cu + O2 → 2CuO 2H2O → 2H2 + O2 16 Fe + 3 S8 → 8Fe2S3 2NaCl → 2Na + Cl2 4Al + 3O2 → 2Al2O3 2HgO →2 Hg + O2 C + O2 → CO2 6Li + 3N2 → 2Li3N S8 + 8O2 →8SO2 N2 + 3H2 →2NH3 16Ag + S8 →8Ag2S 2Mg + O2 → 2MgO CuI2 → Cu + I2 4Fe + 3O2 → 2Fe2O3 H2 + Cl2 → 2HCl

Page 31

2Na(s) + 2HOH(l) →2NaOH + H2 2NaCl(s) + H2SO4(aq) →Na2SO4 + 2HCl 2Al(s) + Fe2O3(s) →Al2O3 + 2Fe Ca(OH)2(aq) + Mg(HCO3)2(aq) →Mg(OH)2 + Ca(HCO3)2 3K(s) + AlCl3(s) → Al + 3KCl Cu(s) + Ag2SO4(aq) → CuSO4 + 2Ag 3H2SO4(aq) + Ca3(PO4)2(s) → 2H3PO4 + 3CaSO4 Cl2(g) + MgBr2(aq) → MgCl2 + Br2 H2S(g) + PbCrO4(s) → PbS + H2CrO4 H2S(g) + 2Ag(s) → Ag2S + H2 Cl2(g) + 2KI(aq) → I2 + 2KCl Cu(s) + 2AgNO3(aq) → Cu(NO3)2 + 2Ag Zn(s) + 2HCl(aq) → ZnCl2 + H2 Ca(s) + H2SO4(aq) → CaSO4 + H2 HCl(aq) + KOH(aq) → HOH + KCl F2(g) + 2KCl(aq) → 2KF + Cl2 Zn(s) + CuSO4(aq) → ZnSO4 + Cu CaCO3(s) + 2HCl(aq) → CaCl2 + H2CO3 3NaOH(aq) + H3PO4(aq) → Na3PO4 + 3HOH

p. 20 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

11. 12. 13. 14. 15.

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) 2CH3OH(l) + 3O2(g) →2CO2(g) + 4H2O(g) CS2(l) + 3O2(g) → CO2(g) + 2SO2(g) C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) 2Mg(s) + O2(g) → 2MgO(s) 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(g) 2C2H5SH(l) + 9O2(g) → 4CO2(g) + 2SO2(g) + 6H2O(g) 4CH3NO2(g) + 8O2(g) → 4CO2(g) + NO2(g) + 6H2O(g) C12H22O11(s) + 12O2(g) → 12CO2(g) + 11H2O(g) 2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g) S8(s) + 8O2(g) → 8SO2(g) 2H2S(g) + 3O2(g) → 2SO2(g) + 2H2O(g) C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g) 2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(g) C25H52(s) + 38O2(g) → 25CO2(g) + 26H2O(g)

p. 21 1. 2Ca + O2 → 2CaO 2. C3H8 + 5O2 → 3CO2 + 4H2O 3. CuCl2 + Zn → ZnCl2 + Cu 4. 2HgO → 2Hg + O2 5. H2SO4 + 2NaOH → Na2SO4 + 2HOH 6. 2Li + 2HOH → 2LiOH + H2 7. C2H5OH + 3O2 → 2CO2 + 3H2O 8. 2H2O → 2H2 + O2 9. 2NaCl → 2Na + Cl2 10. H3PO4 + Fe(OH)3 → FePO4 + 3HOH 11. Al2O3 + 2Fe → Fe2O3 + 2Al 12. HCl + NH4OH → HOH + NH4Cl 13. FeCl3 + 3NH4OH → Fe(OH)3 + 3NH4Cl 14. Pb(NO3)2 + 2KI → PbI2 + 2KNO3 15. Zn + 2HCl → ZnCl2 + H2 16. Cl2 + 2KI → 2KCl + I2 17.

2C2H5SH + 9O2 → 4CO2 + 6H2O + 2SO2

18. Ca + 2HOH → Ca(OH)2 + H2 19. 2Mg + O2 → 2MgO 20. 4CH3NO2 + 7O2 → 4CO2 + 6H2O +

SJHS Chemistry Dept.

Page 32