AP Chemistry

Summer Review

Winston Churchill High School 2011-2012

Dear AP CHEM STUDENTS,

I’m looking forward to the school year and the work we’re going to do together. This packet is meant to refresh you on several things you already learned so that you have them firmly in mind for the start of school I recommend that you review your Honors Chemistry notes, and start looking at this material several weeks before school starts in the fall. This assignment will be graded

Good luck, and have a wonderful and Productive summer!

Sincerely,

J.Boppana

Chem Sheets to Memorize

SOLUBILITY CHART

Soluble / Exceptions / Insoluble / Exceptions
No / NO3- / --- / Smoking / S-2 / GroupIA, NH4+ Sr+2 Ba+2 Ca+2
CHeating / CH3COO- / --- / Counter / CO3-2 / GroupIA, NH4+
Cellphones / Cl- / Ag+ Pb+2 Hg2+2 / Productive / PO4-3 / GroupIA, NH4+
Bullying / Br- / '' / O (ZERO) / OH- / GroupIA, NH4+ Sr+2 Ba+2 Ca+2
Intimidating / I- / ''
Students / SO4-2 / Sr+2 Ba+2 Pb+2 Hg2+2

AcidsBases

HF – weak LiOH – strong

HCl – strong NaOH– strong

HBr – strong KOH – strong (all IA Metal Hydroxides)

HI – strong Ca, Ba and Sr Hydroxide too

H2SO4 –strong

HNO3 – strong

HClO3 – strong

HClO4 –strong

All else are weak

H2CO3 → H2O + CO2

(very weak acid-breaks down!)

Special Reactions

MetalNonmetal

metal + acid → salt + H2nonmetal + H2O → oxy-acid

metal oxide + H2O → metal hydroxideSO3 + H2O → H2SO4

metal oxide + CO2→ metal carbonateSO2 + H2O → H2SO3

metal chloride + O2→ metal chlorateN2O5 + H2O → 2 HNO3

N2O3 + H2O → 2 HNO2

P2O5 + 3H2O → 2 H3PO4

P2O3 + 3H2O → 2 H3PO3

Oxidizers

AcidBaseNeutral

MnO4- or MnO2→ Mn+2Cr2O7-2→ CrO4-2MnO4-→ MnO2

CrO4-2→ Cr+3

Cr2O7-2→ Cr+3

NO3- (dil) → NO

NO3- (conc) → NO2

metallic ions → metallous ions

free halogens → halide ions

Na2O2 → NaOH

HClO4→ Cl-

C2O4-2 → CO2

H2O2→ O2, H2O

Reducers

halide ions → free halogens

free metals → metal ions

metalous ions → metallic ions

SO3-2→ SO4-2

NO2-→ NO3-

free halogens (dil) → hypohalite ions

free halogens (conc) → halate ions

Colors of Complex ions

Ion / Color
[Cr(H2O)6]+2 / blue
[Cu(H2O)6]+2 / blue
[Cu(H2O)6]+3 / blue/violet
[Mn(H2O)6]+2 / very pale pink
[Co(H2O)6]+2 / pink
[Fe(H2O)6]+2 / pale green
[Ni(H2O)6]+2 / green
[Fe(H2O)6]+3 / yellow/brown

Flame Test Colors

Ion / Flame color
Li+, Sr2+, Ca2+ / Red (various shades)
Na+ / Yellow/Orange
K+ / Lilac
Ba2+ / Green
Cu2+ / Blue-green

TRANSITION METAL ION COLORS

+1 / +2 / +3 / +4 / +5 / +6 / +7
Sc / colorless
Ti / VIOLET / COLORLESS
V / VIOLET / GREEN / BLUE / YELLOW
Cr / BLUE / GREEN / YELLOW FOR CHROMATE
ORANGE FOR DICHROMATE ION
Mn / PALE PINK / BROWN / DARK GREEN / PURPLE
Fe / PALE GREEN / REDDISH BROWN
Co / PINK / ORANGE/
YELLOW
Ni / GREEN
Cu / COLORLESS / BLUE
Zn / Colorless

Common Precipitate colors:

WHITE / BLUE / YELLOW / BLACK / GREEN / REDDISH BROWN
AgCl / Many Copper
(II) ppt’s. / AgI / Many Sulfides / Many Fe(II)
ppt’s. / Many Fe(III)
ppt’s.
BaSO4 / PbI2
PbCl2
Many nontransition
metal
hydroxides
Many nontransition
metal
carbonates and sulfates

Common Tests for gases

GAS / TEST
Hydrogen gas / Squeaky pop with lighted splint
Oxygen gas / Re-lights glowing splint
Carbon Dioxide gas / Turns limewater (Calcium Hydroxide solution) milky
Ammonia gas / Pungent odor, turns red litmus paper blue,gives dense white fumes in contact with conc.HCl fumes

Common tests for cations and anions

ION / TEST
Carbonate and Hydrogen carbonate / Release CO2 gas with acids
Sulfate / White ppt of BaSO4 with barium ions
Chloride / White of AgCl with silver ions
Bromide / Cream ppt of AgBr with silver ions
Iodide / Yellow ppt of AgI with silver ions
Ammonium / NH3 released with hydroxide ions

HALOGENS:

Fluorine gas – pale yellow/green,

Chlorine gas – green,

Bromine liquid –orange/brown,

Iodine solid – dark purple

NO2 gas – orange/brown

Color Changes in REDOX reactions

1) MnO4-(aq) → Mn2+(aq)

(Dark Purple) →(Pale Pink)

2) Cr2O72-(aq) →Cr3+(aq)

(Orange) → (Green)

Acid/Base Indicator Color changes

INDICATOR / ACID / BASE
Methyl orange / RED / YELLOW
Methyl red / RED / YELLOW
Litmus / RED / BLUE
Universal / RED / BLUE/PURPLE
Phenolphthalein / COLORLESS / PINK

POLYATOMIC IONS

Group IIIB or 13 / Group IVB or 14 / Group VB or 15 / Group VIB or 16 / Group VIIIB or 17
Charge -3 / Charge -2 / Charge -1
BO3-3 borate / CO3-2 carbonate / NO3-1 nitrate
NO2-1 nitrite / O / F


one member in the ion family / SiO3-2 silicate / PO4-3 phosphate
PO3-3 phosphite / SO4-2 sulfate
SO3-2 sulfite / ClO4-1 perchlorate
ClO3-1 chlorate
ClO2-1 chlorite
ClO-1 hypochlorite
AsO4-3 arsenate
AsO3-3 arsenite / SeO4-2 selenate
SeO3-2 selenite / BrO4-1 perbromate
BrO3-1 bromate
BrO2-1 bromite
BrO-1 hypobromite
Remember:
  • ions with the greater # of oxygens: ATE
  • ions with the fewer # of oxygens: ITE
  • adding hydrogen in front makes BI and reduces charge by 1
/ Charge -3
two members in the ion family / TeO4-2 telurate
TeO3-2 telurite / IO4-1 periodate
IO3-1 iodate
IO2-1 iodite
IO-1 hypoiodite
Charge -2 / Charge -1
four members in the ion family

Other important polyatomic ions to remember:

acetate C2H3O2-1 / chromate CrO4-2 / bisulfite HSO3-1
hydroxide OH-1 / dichromate Cr2O7-2 / bisulfate HSO4-1
permanganate MnO4-1 / peroxide O2-2 / bicarbonate HCO3-1
cyanide CN-1 / oxalate C2O4-2 / biphosphite HPO3-2
hydronium H3O+1 / thiosulfate S2O3-2 / biphosphate HPO4-2
ammonium NH4+1 / tartrate C2H4O6-2 / hydrogen biphosphite H2PO3-1

I.NOMENCLATURE: NAMING AND WRITING FORMULAS OF CHEM.COMPOUNDS

FormulaName

1. P4O10
2. ZnAt2
3. SBr6
4. CaF2
5. P2S3
6. / carbon monoxide
7. / sodium hydride
8. / aluminum selenide
9. / xenon hexafluoride
10. / dinitrogen monoxide
11. KClO3
12. Pb(OH)2
13. Ca(MnO4)2
14. N2O4
15. Ti(HPO4)2
16. / manganese (VII) oxide
17. / francium dichromate
18. / copper (II) dihydrogen phosphate
19. / silver chromate
20. / ammonium oxalate
21. (NH4)2SO3
22. Ni3(PO4)2
23. Fe(IO2)3
24. NaBrO2
25. H3PO3
26. / tartaric acid
27. / hydrotellluric acid
28. / mercury (I) nitrate
29. / vanadium (V) oxide
30. / tetraphosphorous decaoxide

II. Significant Figures

1. Give the number of sig figs in each of the following numbers

a. 123 b. 0.078 c. 89007 d. 12,000 e. 1,000,000,000.0

f. 0.009 g. 23,000. h. 34,000 i. 34.89 j. 101

2. Do the following calculations giving the answer in the appropriate number of sig figs.

a. 1.23 + 75 b. 1.89 - .20 c. 45.6 x 8.2 d. 234/0.298

e. 0.887 + 0.3 f. 2340 - 100 g. 12.45 x 3 h. 25,600/ 3.0

3. Do the following calculations giving the answer in the appropriate number of sig figs

a. 45.0 x 9.0 + 89.22/ 75 b. (2.88 + .5) x ( 23,000 - 0.11)

c. 0.8897 x 2.15 + 0.002/.1 d. (8 + 9)/(34.0 – 20.)

III. Reactions

Please write net ionic balanced reactions (with states of matter included) for the following questions on a separate piece of paper. You’ll have reactions that are classified as precipitation, acid-base, or redox (reduction-oxidation…like, synthesis, decomposition, and single displacement/replacement).

  • Any ion has an aqueous state of matter.
  • For acid-base reactions, strong acids (HCl, HBr, HI, H2SO4, HClO4, and HNO3) and strong bases (metal ions in groups 1 and 2 paired with hydroxide (OH-) completely dissociate. Weak acids and bases do not.
  • For precipitation (and some redox) reactions, use the solubility rules below to determine which salts are soluble (aqueous) or insoluble (solid). Only aqueous solutions can dissociate…solids, liquids, and gases cannot.
  1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble.
  2. Salts containing nitrate ion (NO3-) and acetate ion (C2H3O2) are generally soluble.
  1. Salts containing Cl -, Br -, F-, and I - are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.
  2. Most silver salts are insoluble. AgNO3 and AgC2H3O2 are common soluble salts of silver; virtually anything else is insoluble.
  3. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4 and SrSO4 .
  4. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, and Co(OH)2 are not soluble.
  5. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, and Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
  6. Most chromates, phosphates, bicarbonates, and carbonates are frequently insoluble except those with alkali metals and ammonium.

Acid-Base Example: Hydrochloric acid is added to a solution of zinc hydroxide.

*First, write a molecular equation.

HCl + Zn(OH)2 ZnCl2 + H2O

*Next, you need to see what dissociates and what does not. Hydrochloric acid is a strong acid, so it will completely dissociate into its ions while zinc hydroxide is a weak base, so it will not dissociate. Zinc chloride is a soluble salt according to the solubility rules above, so it will also dissociate into its ions.

*Wait to balance the reaction until the end.

H+ + Cl- + Zn(OH)2 Zn+2 + Cl- + H2O

*Last, you need to see what can be cancelled out. Species that are identical on both sides of the reaction, called spectator ions, can be cancelled out. Cl- is present on both sides of the reaction and therefore can be cancelled out…giving you your net ionic reaction that you’ll now balance and put back on states of matter.

2 H+ (aq) + Zn(OH)2 (aq)  Zn+2 (aq) + 2 H2O (l )

Redox Example: Silver metal reacts with a solution of sodium nitrate.

Ag + NaNO3 Na + AgNO3

*Ag is a solid. NaNO3 is a soluble salt according to the solubility rules above, so it will dissociate into its ions. Na is a solid. AgNO3 is also a soluble salt and will dissociate.

Ag + Na+ + NO3- Na + Ag++ NO3-

*NO3- is a spectator ion.

Ag (s) + Na+ (aq)  Na (s) + Ag+ (aq)

Precipitation Example: Barium acetate is mixed with potassium sulfate.

Ba(C2H3O2)2 + K2SO4 BaSO4 + KC2H3O2

*According to the solubility rules, barium sulfate is the only insoluble salt. So, everything else will dissociate.

Ba+2 + C2H3O2- + K+ + SO4-2 BaSO4 + K+ + C2H3O2-

*The potassium ions and acetate ions can be cancelled out.

Ba+2 (aq) + SO4-2 (aq)  BaSO4 (s)

***Here are your questions. Please do these on a separate piece of paper.

  1. Solid sodium bicarbonate is mixed with copper (II) nitrate.
  2. Magnesium oxide is heated.
  3. Acetic acid is added to a solution of ammonia.
  4. Iron (III) chloride is mixed with silver sulfite.
  5. A solid piece of aluminum is put into a solution of nickel (II) chloride.
  6. A solution of lithium chloride is added to a solution of lead (IV) nitrite.
  7. Sulfuric acid is added to a solution of aluminum hydroxide.
  8. Cadmium nitrate is added to sodium sulfide.
  9. Chromium (III) sulfate is added to ammonium carbonate.
  10. Methane combusts in air.

In each of the equations below, the reactants are written correctly. You must write the correct products and then balance the equation. It might be useful to identify the type of chemical reaction before writing the products.

  1. CaCO3
  1. Al + O2
  1. Fe + CuSO4
  1. C6H12 + O2
  1. Zn + H2SO4
  1. Cl2 + MgI2
  1. NaOH 
  1. Fe + HCl 
  1. NaOH + H3PO4
  1. (NH4)2SO4 + Ca(OH)2
  1. AgNO3 + K2SO4
  1. Mg(OH)2 + H3PO4
  1. Na + H2O 
  1. KClO3
  1. Al2(SO4)3 + Ca3(PO4)2
  1. SO2 + H2O 
  1. (NH4)3PO4 + Ba(OH)2
  1. Ca(OH)2 + HNO3
  1. C3H8 + O2
  1. Li + S 

IV.Electron Structure and Periodicity

*Please do all questions on a separate piece of paper.

*You will need to know about valence electrons, electron shells, orbital notation, electron configuration, atomic radius, ionization energy, and electronegativity to do these questions.

  1. Draw the orbital notation for nickel.
  2. How many unpaired electrons are in arsenic?
  3. Write the electron configuration for palladium.
  4. How many valence electrons are in mercury?
  5. Write the electron configuration for uranium.
  6. Write the noble gas electron configuration for lead.
  7. Which is more electronegative, sulfur or chlorine, and why?
  8. Which has a larger atomic radius, potassium or bromine, and why?
  9. Which has the smaller ionization energy, nitrogen or phosphorus, and why?
  10. Write the noble gas electron configuration for copper.

Short Answer question from previous AP EXAM

11. Use the principles of atomic structure and/or chemical bonding to explain each of the following. In each part, your answer must include references to both substances.

  1. The atomic radius of Li is larger than that of Be.
  2. The second ionization energy of K is greater than the second ionization energy of Ca.
  3. The carbon-to-carbon bond energy in C2H4 is greater than it is in C2H6.
  4. The boiling point of Cl2 is lower than the boiling point of Br2.

Atomic Structure Sample problems:

12. Give the symbols for the isotopes of Carbon, nitrogen and uranium and Determine the number of protons, electrons and neutrons in each isotope.

13. Given the data below determine the average atomic mass

Isotope % Abundance Isotopic Mass

a. Sb-121 57.25% 120.9038 amu

Sb-123 42.75%122.0041 amu

b. Ag-107 51.82% 106.90509 amu

Ag-109 48.18% 108.9047 amu

Mole Concept Sample Problems

14. Convert each of the following to moles.

a. 12.64 g NaOHb. 3.00 x 1024 atoms Au c. 40.0 L of Ne gas

d. 800. g CaBr2 e. 3.011 x 1022 molecules H2O f. 6.78 L of Ar gas

15. Do the following

a. Given 0.250 moles of krypton determine

(i) the mass (ii) the number of atoms (iii) the volume at STP

b. Given 0.750 moles of oxygen determine

V. Bonding

*Please do all questions on a separate piece of paper.

*You will need to know about Lewis structures, covalent bonding, shape names, and bond angles to do these questions.

*For the following questions, draw the Lewis Structure, name the shape, and state the bond angle.

  1. SeCl2
  2. NO3-1
  3. OF2
  4. BF3
  5. SO4-2
  6. NH4+
  7. CO2
  8. CH3NH2
  9. HCOOH

10. HCN

VI.Stoichiometry

*Please do all questions on a separate piece of paper.

*You will need to be able to write molecular chemical reactions and do mole conversions to do these questions.

  1. 30.5 g of sodium metal reacts with a solution of excess lithium bromide. How many grams of lithium metal are produced?
  1. How many molecules are in 100. L of potassium hydroxide solution at STP?
  1. Propane, C3H8, undergoes combustion. How many grams of propane are needed to produce 45.9 g of water?
  1. How many moles are in 3.02 x 1026 molecules of water?
  1. Find the empirical and molecular formulas for a compound containing 11.66 g iron and 5.01 g oxygen if the molar mass of the compound is 320 g/mol.
  1. A solution of 3.50 g of sodium phosphate is mixed with a solution containing 6.40 g of barium nitrate. How many grams of barium phosphate can be formed?
  1. Find the empirical and molecular formulas for a compound containing 5.28 g of tin and 3.37 g of fluorine if the molar mass of the compound is 584.1 g/mol.
  1. Octane, C8H18, undergoes combustion. How many grams of oxygen are needed to burn 10.0 g of octane?
  1. Sodium azide, NaN3, decomposes into its elements. How many grams of sodium azide are required to form 34.8 g of nitrogen gas?
  1. Ammonia reacts with oxygen gas to form nitrogen monoxide and water. How many grams of nitrogen monoxide are formed when 1.50 g of ammonia react with 2.75 g of oxygen gas?

Short Answer Problems from Previous AP EXAMS

  1. The reaction between silver ion and solid zinc is represented by the following equation:

2Ag+ (aq) + Zn (s)  Zn+2 (aq) + 2Ag (s)

A 1.50 g sample of Zn is combined with 250 mL of 0.110 M AgNO3 at 25°C.

  1. Identify the limiting reagent. Show calculations to support your answer.
  1. On the basis of the limiting reactant that you identified in part (i), determine the value of [Zn+2] after the reaction is complete.
  1. Consider the hydrocarbon pentane, C5H12 (molar mass 72.15 g).
  2. Write the balanced equation for the combustion of pentane to yield carbon dioxide and water.
  1. What volume of dry carbon dioxide, measured at 25°C and 785 mmHg, will result from the complete combustion of 2.50 g pentane?

4) Find the mass percent of nitrogen in each of the following compounds:

a. NO

b. NO2

c. N2O4

d. N2O

5) Benzene contains only carbon and hydrogen and has a molar mass of 78.1 g/mol. Analysis shows the compound to be 7.74% H by mass. Find the empirical and molecular formulas of benzene.

6) Calcium carbonate decomposes upon heating, producing calcium oxide and carbon dioxide gas.

a. Write a balanced chemical equation for this reaction.

b. How many grams of calcium oxide will be produced after 12.25 g of calcium carbonate is completely decomposed?

c. What volume of carbon dioxide gas is produced from this amount of calcium carbonate, at STP?

7) Hydrogen gas and bromine gas react to form hydrogen bromide gas.

a. Write a balanced chemical equation for this reaction.

b. 3.2 g of hydrogen gas and 9.5 g of bromine gas react. Which is the limiting reagent?

c. How many grams of hydrogen bromide gas can be produced using the amounts in (b)?

d. How many grams of the excess reactant is left unreacted?

e. What volume of HBr, measured at STP, is produced in (b)?

8) When ammonia gas, oxygen gas and methane gas (CH4) are combined, the products are hydrogen cyanide gas and water.

a. Write a balanced chemical equation for this reaction.

b. Calculate the mass of each product produced when 225 g of oxygen gas is reacted with an excess of the other two reactants.

c. If the actual yield of the experiment in (b) is 105 g of HCN, calculate the percent yield.

9) When solutions of potassium iodide and lead (II) nitrate are combined, the products are potassium nitrate and lead (II) iodide.

a. Write a balanced equation for this reaction, including (aq) and (s).

b. Calculate the mass of precipitate produced when 50.0mL of 0.45M potassium iodide solution and 75mL of 0.55M lead (II) nitrate solution are mixed.

c. Calculate the volume of 0.50M potassium iodide required to react completely with 50.0mL of 0.50M lead (II) nitrate.