Worksheets - Answers Grade 11 University Level

Unit 1Nomenclature

Worksheet 1 – Nomenclature 1 Read Pgs 52 – 65 Q# Pg 66 2 – 6Worksheet 2 – Nomenclature 2Worksheet 3 – Nomenclature 3Worksheet 4 – Nomenclature 4Worksheet 5 –Nomenclature 5

Types of ReactionsNo Worksheet – Chemical Equations Read Pgs 72 – 78 Q# Pg 79 1 – 6Worksheet 6 – Balancing Chemical EquationsWorksheet 7 – Types of Reactions Read Pgs 79 – 86, Q# Pg 86 1 – 5Worksheet 8 – Types of Reactions Read Pgs 87 – 102 Q# Pg 103 3 – 5Worksheet 9 – Net Ionic Equations

Atomic StructureWorksheet 10 – Aufbau Diagrams Read Pgs 4 – 14, Pg 14, Q# 1 – 6Worksheet 11 – Blank Periodic TableWorksheet 12 – Energy Level DiagramsWorksheet 13 – Electron ConfigurationsWorksheet 14 – Orbital and Energy Level DiagramsWorksheet 15 – Quantum NumbersWorksheet 16 – Atomic Structure Review

Unit 2Periodicity

Worksheet 17 – Periodic Trends Read Pgs 15 – 25 Pg 26 Q# 1 - 7Worksheet 18 – Alkali Metals and Halogens Pg 30 Q# 1 – 15, 17

BondingNo Worksheet – Ionic Compounds Read Pgs 34 – 40, Pg 41 Q# 1,2,4,5No Worksheet – Covalent Compounds Read Pgs 41 – 51, Pg 51 Q# 1 – 6No Worksheet – Ionic/Covalent Structural FormulaWorksheet 19 – Polarity Read Pgs 48 – 50, Pgs 51 Q# 7No Worksheet – Unit 2 Review Pgs 70 – 71 Q# 1 – 18,

Worksheet 1 – Nomenclature 1

1. Write the formulas for each of the following compounds:

a) calcium fluoride CaF2 b) sodium sulfide Na2S

c) aluminum nitride AlN d) aluminum chloride AlCl3

e) potassium oxide K2O f) calcium chloride CaCl2

g) copper (II) sulfide CuS h) lead (II) bromide PbBr2

i) silver iodide AgI j) barium nitride Ba3N2

k) iron (II) fluoride FeF2 l) manganese (IV) oxide MnO2

m) mercury (II) sulfide HgS n) ferric oxide Fe2O3

o) cuprous nitride Cu3N p) antimony (V) iodide SbI5

q) cesium carbide Cs4C r) strontium phosphide Sr3P2

2. Write the names for the following compounds:

a) lime, CaO(s) calcium oxide

b) road salt, CaCl2(s) calcium chloride

c) magnesia, MgO(s) magnesium oxide

d) bauxite, Al2O3(s) aluminum oxide

e) zinc ore, ZnS(s) zinc sulfide

f) cassiterite, SnO2(s) tin (IV) oxide

3. Write the IUPAC and Classical names for the following compounds if needed:

a) Na2O(s) sodium oxide

b) SnCl4(s) tin (IV) chloride stannic chloride

c) ZnI2(s) zinc iodide

d) SrCl2(s) strontium chloride

e) AlBr3(s) aluminum bromide

f) PbCl4(s) lead (IV) chloride plumbic chloride

4. Write the chemical formulas and names for the compounds formed by the pairs of elements below:

a) strontium and oxygen SrO strontium oxide

b) sodium and sulfur Na2S sodium sulfide

c) barium and fluorine BaF2 barium flrouide

d) silver and iodine AgI silver iodide

e) calcium and bromine CaBr2 calcium bromide

f) lithium and chlorine LiCl lithium chloride

Worksheet 2 – Nomenclature 2

1. For each of the following compounds use stock and classical naming systems when needed:

a) NaNO3(s) sodium nitrate -found in tobacco

b) NaNO2(s) sodium nitire -a meat preservative

c) Cu(NO3)2(s) copper (II) nitrate cupric nitrate -blue solution in water

d) CuNO3(s) copper (I) nitrate cuprous nitrate -green solution in water

e) Al2(SO3)3(s) aluminum sulfite -a food additive in pickles

f) Ca(OH)2(s) calcium hydroxide -firming agent in fruit products

g) PbCO3(s) lead (II) carbonate plumbous carbonate -cerussite, a mineral

h) Sn3(PO4)2(s) tin (II) phosphate stannous phosphate -use to fix paints to silk

i) Fe2(SO4)3(s) iron (III) sulfate ferric sulfate -a mineral found on Mars

2. Write the chemical formula for each of the following compounds:

a) calcium carbonate CaCO3 -active ingredient in antacids

b) sodium bicarbonate NaHCO3 -a foaming agent added to foods

c) sodium hypochlorite NaClO -a component of bleach

d) calcium sulfate CaSO4 -plaster of Paris

e) ammonium nitrate NH4NO3 -used in fertilizer

f) ammonium phosphate (NH4)PO4 -a leavening agent added to foods

g) copper (II) sulfate CuSO4 -used as a fungicide

h) sodium hydroxide NaOH -a strong base used as a washing agent

i) potassium permanganate KMnO4 -a traditional antiseptic

3. Use IUPAC and Archaic systems (if necessary) to name each of the following compounds:

a) LiClO3(s) lithium chlorate

b) BaSO4(s) ) barium sulfate

c) Hg2CO3(s) mercury (II) carbonate mercurous carbonate

d) Mg(NO3)2(s) magnesium nitrate

e) Fe(BrO3)3(s) iron (III) bromate ferric bromate

f) Na3PO4(s) sodium phosphate

g) NH4IO3(s) ammonium iodate

h) AuC2H3O2(s) gold (I) acetate aurous acetate

i) Zn3(PO4)2(s) zinc phosphate

j) Sb(ClO3)5(s) antimony (V) chlorate antimonic chlorate

k) MnSO4(s) manganese (II) sulfate manganous sulfate

l) KBrO(s) potassium hypobromite

m) AlPO5(s) aluminum perphosphate

4. Write the chemical formulas for each of the following molecules:

a) nitrogen N2 b) carbon dioxide CO2

c) carbon monoxide CO d) nitrogen dioxide NO2

e) nitrogen monoxide NO f) dinitrogen monoxide N2O

g) dinitrogen tetroxide N2O4 h) sulfur dioxide SO2

i) diiodine pentoxide I2O5 j) silicon tetrafluoride SiF4

k) boron trifluoride BF3 l) phosphorus triiodide PI3

m) diphosphorus pentoxide P2O5 n) hexanitrogen heptachloride N6Cl7

5. Rename each of the following compounds using the Stock system and give the chemical formula for each compound:

a) ferrous sulfide FeS iron (II) sulfide

b) plumbic bromide PbBr4 lead (IV) bromide

c) stannous chloride SnCl2 tin (I) chloride

d) cuprous hypophosphite Cu3PO2 copper (I) hypophosphite

e) stannic chlorite Sn(ClO2)4 tin (IV) chlorite

f) ferrous bromate Fe(BrO3)2 iron (II) bromate

g) ferric chlorite Fe(ClO2)3 iron (III) chlorite

h) plumbic sulfate Pb(SO4)2 lead (IV) sulfate

i) titaniumic arsenite Ti3(AsO3) titanium (IV) arsenite

j) arsenous nitrite As(NO2)3 arsenic (III) nitrite

k) cobaltic acetate Co(C2H3O2)3 cobalt (III) acetate

6. Name the following compounds:

a) SF6(g) sulfur hexafluoride

b) N2O3(g) dinitrogen trioxide

c) NO2(g) nitrogen dioxide

d) PCl3(l) phosphorus trichloride

e) PCl5(s) phosphorus pentachloride

f) S5P4(s) pentasulfur tetraphosphide

g) Se2F(s) diselenium monoflrouide

h) N2I7(s) dinitrogen heptaiodide

Worksheet 3 – Nomenclature 3

1. Name the following hydrated compounds:

a) CuSO4•5H2O copper (II) sulfate pentahydrate

b) Na2SO4•10H2Osodium sulfate decahydrate

c) MgSO4•7H2O magnesium sulfate heptahydrate

d) LiNO3•3H2O lithium nitrate trihydrate

e) CaSO3•8H2O calcium sulfite octahydrate

2. Write the chemical formulas for the following hydrates:

a) iron(III) oxide trihydrate Fe2O3•3H2O -rust

b) aluminum chloride hexahydrate AlCl3•6H2O -component of antiperspirant

c) sodium thiosulfate pentahydrate NaS2O3•5H2O -photographic “hypo”

d) cadmium (II) nitrate tetrahydrate Cd(NO3)2•4H2O -photographic emulsion

e) lithium chloride tetrahydrate LiCl•4H2O -in fireworks

f) calcium chloride dihydrate CaCl2•2H2O -deicer

3. Write the names of the following bases:

a) KOH(aq) potassium hydroxide

b) Ca(OH)2(aq) calcium hydroxide

4. Write the formulas of the following bases:

a) aqueous magnesium hydroxide Mg(OH)2(aq)

b) aqueous sodium hydroxide NaOH(aq)

c) aqueous aluminum hydroxide Al(OH)2(aq)

5. Write the chemical formulas for the following compounds:

a) aqueous hydrogen chloride HCl(aq)

b) hydrochloric acid HCl(aq)

c) aqueous hydrogen sulfate H2SO4(aq)

d) sulfuric acid H2SO4(aq)

e) aqueous hydrogen acetate CH3COOH(aq)

f) acetic acid CH3COOH(aq)

g) aqueous hydrogen nitrite HNO3(aq)

h) nitric acid HNO3(aq)

i) hydrobromic acid HBr(aq)

j) hyposulfurous acid H2SO2(aq)

k) hydroiodic acid HI(aq)

l) aqueous hydrogen perchlorate HClO4(aq)

6. Name each of the following compounds using both the “classical” and the “IUPAC” nomenclature systems:

a) H2SO3(aq) aqueous hydrogen sulfite sulfurous acid

b) H3PO4(aq) aqueous hydrogen phosphate phosphoric acid

c) HCN(aq) aqueous hydrogen cyanide hydrocyanic acid

d) H2CO3(aq) aqueous hydrogen carbonate carbonic acid

e) H2S(aq) aqueous hydrogen sulfide hydrosulfuric acid

f) HCl(aq) aqueous hydrogen chloride hydrochloric acid

g) H2SO4(aq) aqueous hydrogen sulfate sulfuric acid

h) H3PO3(aq) aqueous hydrogen phosphite phosphorous acid

Worksheet 4 – Nomenclature 4

1. Write the chemical names for each of the following compounds.

- be sure to include both Stock and Classical Names for any compounds with polyvalent ions.

- also if it is a binary or ternary acid be sure to include both IUPAC and classical system names.

a) MgCl2(aq) magnesium chloride

b) FeO(s) iron (II) oxide ferrous oxide

c) Cs2S2O3 cesium thiosulfate

d) Cu3(AsO4)2 copper (II) arsenate cupric arsenate

e) Cu3AsO4 copper (I) arsenate cuprous arsenate

f) SrCl2•3H2O strontium chloride trihydrate

g) HI(g) hydrogen iodide

h) HI(aq) aqueous hydrogen iodide hydroiodic acid

i) O8Cl4 octaoxygen tetrachloride

j) H3PO4(aq) aqueous hydrogen phosphate phosphoric acid

k) CaHPO4(aq) calcium hydrogen phosphate

l) P3F9 triphosphorus nonafluoride

m) BeS beryllium sulfide

n) AuI3 gold (III) iodide auric iodide

o) KMnO4(s) potassium permanganate

p) (NH4)OH(aq) ammonium hydroxide

q) HCl(aq) aqueous hydrogen chloride hydrochloric acid

r) HBrO2(aq) aqueous hydrogen bromite bromous acid

s) HCN(aq) aqueous hydrogen cyanide hydrocyanic acid

t) HBrO2(g) hydrogen bromite

u) LiH2PO4(aq) lithium dihydrogen phosphate

v) Al(H2PO2)3(aq) aluminum dihydrogen phosphate

w) AlCl3•8H2O aluminum chloride octahydrate

x) S3Cl5 trisulfur pentachloride

y) Ti3N4 titanium (IV) nitride titaniumic nitride

z) Sr(ClO)2(s) strontium hypochlorite

aa) HNO(aq) aqueous hydrogen hyponitrite hyponitrous acid

bb) FrIO2 francium iodite

cc) (NH4)3AsO3 ammonium arsenite

dd) HgMnO4 mercury (I) permanganate mercurous permanganate

ee) HBr(g) hydrogen bromide

ff) H2SO4(aq) aqueous hydrogen sulfate sulfuric acid

gg) N2O2 dinitrogen dioxide

hh) XeF6 xenon hexaflouride

ii) H2SO3(aq) aqueous hydrogen sulfite sulfurous acid

jj) HBr(aq) aqueous hydrogen bromide hydrobromic acid

kk) BeHSO3(aq) beryllium hydrogen sulfite

ll) RbClO4 rubidium perchlorate

mm) H2CO3(aq) aqueous hydrogen carbonate carbonic acid

nn) H2CO4(aq) aqueous hydrogen percarbonate percarbonic acid

oo) H2Se(aq) aqueous hydrogen selenide hydroselenic acid

pp) BaSO2•10H2O barium hydrosulfite decahydrate

qq) Pb(HCO)4 lead (II) hypocarbonite plumbous hypocarbonite

rr) Fe2(HCO4)3 iron (III) percarbonate ferric percarbonate

ss) H2O dihydrogen monoxide water

Worksheet 5 –Nomenclature 5

1. Write the chemical formulas of the following compounds:

a) radium oxide RaO

b) hexanitrogen difluoride N6F2

c) tin (IV) permanganate Sn(MnO4)4

d) cupric carbonite CuCO2

e) aqueous hydrogen iodide HI(aq)

f) magnesium dihydrogen phosphate Mg(H2PO4)2

g) cuprous nitrite dehydrate CuNO2∙10H2O

h) hypobromous acid HBrO(aq)

i) aqueous hydrogen hypobromite HBrO(aq)

j) hydrogen bromide HBr(g)

k) aqueous hydrogen bromide HBr(aq)

l) hydrobromic acid HBr(aq)

m) ammonium phosphate (NH4)3PO4

n) titaniumic chloride TiCl4

o) oxygen dichloride OCl2

p) aluminum nitride AlN

q) potassium dichromate K2Cr2O7

r) phosphoric acid H3PO4(aq)

s) phosphorous acid H3PO3(aq)

t) hydrogen cyanide HCN

u) lithium iodite LiIO2

v) ammonium thiosulfate (NH4)2S2O3

w) antimonic persulfate Sb2(SO5)5

x) aqueous hydrogen sulfate H2SO4(aq)

y) sulfuric acid H2SO4(aq)

z) aluminum hydrogen hypocarbonite Al(HCO)3

aa) octasulfur decachloride S8Cl10

bb) pernitric acid HNO4(aq)

cc) aqueous hydrogen pernitrate HNO4(aq)

dd) cuprous chloride hexahydrate CuCl∙6H2O

ee) nickelic phosphide NiP

ff) lithium iodate LiIO3

gg) diphosphorus tetraiodide P2I4

hh) potassium thiosulfate K2S2O3

ii) calcium hydrogen phosphate CaHPO4

jj) aqueous magnesium hydroxide Mg(OH)2(aq)

kk) hydrochloric acid HCl(aq)

ll) acetic acid CH3COOH(aq)

mm) cobalt (III) dichromate Co2(Cr2O7)3

nn) mercury (I) carbide Hg4C

oo) titaniumous hydrogen sulfide Ti(HS)3

pp) cesium cyanate CsOCN

qq) nonanitrogen trifluoride N9F3

rr) nickel (III) oxide Ni2O3

ss) hypophosphorous acid H3PO2(aq)

tt) perphosphoric acid H3PO5(aq)

uu) aqueous hydrogen perphosphate H3PO5(aq)

vv) persulfuric acid H2SO5(aq)

Worksheet 6 - Balancing Chemical Equations

1. Write balanced equations for the following reactions:

a) The reaction of ammonia with iodine forms aqueous nitrogen triiodide and hydrogen gas.

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b) When heated mercury (II) oxide decomposes into mercury and oxygen gas.

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c) The reaction of iron (III) oxide with carbon produces carbon monoxide and iron.

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d) The combustion of ammonia produces nitrogen monoxide and water

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e) Magnesium reacts with sulfuric acid, forming magnesium sulfate and releasing hydrogen gas.

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f) Ethane, CH3CH3, burns in oxygen gas to form carbon dioxide and water.

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g) Potassium chlorate, a good oxidizing agent, will readily produce potassium chloride and oxygen gas.

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h) Sulfuric acid is produced by reacting sulfur dioxide with oxygen gas, and water.

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2. Balance the following equations

a) Li2O(s) + H2O(l) LiOH(aq)

b) Na(s) + Cl2(g) NaCl(s)

c) Al2O3(s) + H2(g) Al(s) + H2O(l)

d) Fe(s) + O2(g) Fe2O3(s)

e) Al(s) + HCl(aq) AlCl3(aq) + H2(g)

f) Al(s) + O2(g) Al2O3(s)

g) KMnO4(s) + HCl(aq) Cl2(g) + MnCl2(aq) + KCl(aq) + H2O(l)

h) MgCl2(aq) + NH4NO3(aq) Mg(NO3)2(aq) + NH4Cl(aq)

i) Pb(C2H5)(l) + O2(g) PbO(s) + CO2(g) + H2O(l)

j) Ca(s) + H2O(l) H2(g) + Ca(OH)2(aq)

Worksheet 7 – Types of Reactions

For the following reactions identify:a) the type of reactionb) predict the products (be sure to include states of matter)c) balance the equation

Write “no reaction” if nothing happens. Type of Reaction

1. CaSO4(s) CaO(s) + SO3(g) decomposition

2. Cl2(aq) + NaI(aq) I2(aq) + NaCl(aq) single-displacement

3. Zn(s) + CuCl2(aq) Cu(s) + ZnCl2(aq) single-displacement

4. KOH(aq) + MgCl2(aq) KCl(aq) + Mg(OH)2(s) double-displacement

5. Li2O(s) + CO2(g) Li2CO3(s) synthesis

6. Al(s) + H2O(l) Al(OH)3(s) + H2(g) single-displacement

7. Mg(s) + Zn(NO3)2(aq) Mg(NO3)2(aq) + Zn(s) single-displacement

8. Sn(s) + HCl(aq) SnCl2(aq) + H2(g) single-displacement

9. Hg2(NO3)2(aq) + NaCl(aq) NaNO3(aq) + HgCl(s) double-displacement

10. Na2CO3(aq) + H3PO4(aq) Na3PO4(aq) + H2O(l) + CO2(g) double/decomposition

11. Zn(s) + O2(g) ZnO(s) synthesis

12. Br2(aq) + MgI2(aq) MgBr2(aq) + I2(aq) single-displacement

13. S(s) + O2(g) SO2(g) synthesis

14. CaO(s) + H2O(l) Ca(OH)2(aq) synthesis

15. Cu(s) + Hg(NO3)2(aq) No Reaction

16. KOH(aq) + HClO(aq) KClO(aq) + H2O(l) double-displacement

17. Ni(s) + CuSO4(aq) NiSO4(aq) + Cu(s) single-displacement

18. H2(g) + Cl2(g) HCl(g) synthesis

19. NH4NO3(aq) + KOH(aq) KNO3(aq) + H2O(l) + NH3(g) double/decomposition

20. NaOH(aq) + H2SO4(aq) Na2SO4(aq) + H2O(l) double displacement

21. F2(aq) + AlCl3(aq) AlF3(aq) + Cl2(aq) single-displacement

22. Al(NO2)3(s) Al2O3(s) + NO(g) decomposition

23. CsNO3(aq) + Na2S(aq) No Reaction

24. BaO(s) + ClO2(g) BaClO3(s) synthesis

25. Ca(s) + HIO3(aq) CaIO3(aq) + H2(g) single-displacement

26. NH4Cl(aq) + Al(OH)3(aq) AlCl3(aq) + H2O(l) + NH3(g) double/decomposition

27. Na(s) + CuCl(aq) NaCl(aq) + Cu(s) single-displacement

Worksheet 8 – Types of Reactions

In each of the following identify the type of reaction and then state the products.

1. Na3PO4(aq) + CaBr2(aq) NaBr(aq) + Ca3(PO4)2(s)

2. Pb(NO3)2(aq) + HCl(aq) PbCl2(s) + HNO3(aq)

3. Na2CO3(aq) + KF(aq) No Reaction

4. AgNO3(aq) + CuSO4(aq) Cu(NO3)2(aq) + Ag2SO4(s)

5. AgF(aq) + NiCl2(aq) NiF(aq) + AgCl(s)

6. Pb(s) + FeSO4(aq) No Reaction

7. CaCO3(s) CaO(s) + CO2(g)

8. P4(s) + O2(g) P2O3(g)

9. RbNO3(aq) + BeF2(aq) No Reaction

10. AgNO3(aq) + Cu(s) Cu(NO3)2(aq) + 2 Ag(s)

11. Li2SO3(aq) + HNO3(aq) LiNO3(aq) + H2O(l) +SO2(g)

12. Na2O(s) + H2O(l) NaOH(aq)

13. HNO3(aq) + Ca(OH)2(aq) Ca(NO3)2(aq) + H2O(l)

14. NaOH + (NH4)3PO4 Na3PO4(aq) + H2O(l) +NH3(g)

15. Al(s) + Sn(NO3)2(aq) Al(NO3)3(aq) + Sn(s)

16. C3H6(g) + O2(g) CO2(g) + H2O(l)

17. Na(s) + CaSO4(aq) No Reaction

18. BeO(s) + SO2(g) BeSO3(s)

19. Pb(s) + Br2(l) PbBr2(s) or PbBr4(s)

20. Ti3(PO3)4(s) TiO2(s) + PO2(g)

21. H2O(l) + BrO2(g) HBrO3(aq)

22. IO3(g) + H2O(l) HIO4(aq)

23. H3PO5(aq) H2O(l) + PO4(g)

24. RbOH(aq) Rb2O(s) + H2O(l)

25. CoBr3(aq) + I2(s) No Reaction

26. CoBr3(aq) + Cl2(g) CoCl3(aq) + Br2(l)

27. K2S(aq) + HNO3(aq) KNO3(aq) + H2S(g)

28. H2O(l) + Ni(s) Ni(OH)2(aq) + H2(g)

29. MgCl2(s) Mg(s) + Cl2(g)

30. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Worksheet 9 – Net Ionic Equations

1. Write the balanced chemical equation, total ionic equation, and net ionic equation for the reaction of sodium iodide solution and aqueous bromine.

Balanced Equation: 2NaI(aq) + Br2(aq) 2NaBr(aq) + I2(aq)

Total Ionic Equation: 2Na+(aq) + 2I-

(aq) + Br2(aq) 2Na+(aq) + 2Br-

(aq) + I2(aq)

2Na+(aq) + 2I-

(aq) + Br2(aq) 2Na+(aq) + 2Br-

(aq) + I2(aq)

Net Ionic Equation: 2I-(aq) + Br2(aq) 2Br-

(aq) + I2(aq)

2. Write the balanced chemical equation, total ionic equation, and net ionic equation for the reaction of lead (II) nitrate solution and potassium chloride solution.

Balanced Equation: Pb(NO3)2(aq) + 2KCl(aq) 2KNO3(aq) + PbCl2(s)

Total Ionic Equation: Pb2+(aq) + 2NO3

-(aq) + 2K+

(aq) + 2Cl-(aq) 2K+

(aq) + 2 NO3-(aq) + PbCl2(s)

Pb2+(aq) + 2NO3

-(aq) + 2K+

(aq) + 2Cl-(aq) 2K+

(aq) + 2 NO3-(aq) + PbCl2(s)

Net Ionic Equation: Pb2+(aq) + 2Cl-

(aq) PbCl2(s)

3. Write the balanced chemical equation, total ionic equation, and net ionic equation for the reaction of sodium carbonate solution and hydrochloric acid.

Balanced Equation: Na2CO3(aq) + 2HCl(aq) 2NaCl(aq) + H2O(aq) + CO2(aq)

Total Ionic Equation: 2Na(aq)+ CO32-

(aq)+2H+(aq)+2Cl-

(aq) 2Na+(aq) + 2Cl-

(aq) + H2O(aq) + CO2(aq)

2Na(aq)+ CO32-

(aq)+2H+(aq)+2Cl-

(aq) 2Na+(aq) + 2Cl-

(aq) + H2O(aq) + CO2(aq) Net Ionic Equation: CO3

2-(aq) + 2H+

(aq) H2O(aq) + CO2(aq)

4. Write the net ionic equation for each of the following reactions:

a) When aqueous sodium carbonate solution is added to a solution of calcium chloride, solid calcium carbonate is formed and sodium chloride remains in solution.

Balanced Equation: Na2CO3(aq) + CaCl2(aq) 2NaCl(aq) + CaCO3(s)

Total Ionic Equation: 2Na+(aq)+ CO3

2-(aq) + Ca2+

(aq) +2Cl-(aq) 2Na+

(aq)+2Cl-(aq)+ CaCO3(s)

2Na+(aq)+ CO3

2-(aq) + Ca2+

(aq) +2Cl-(aq) 2Na+

(aq)+2Cl-(aq)+ CaCO3(s) Net

Ionic Equation: CO32-

(aq) + Ca2+(aq) CaCO3(s)

b) When solid magnesium metal is added to a solution of zinc chloride, solid zinc metal is formed, leaving a solution of magnesium chloride

Balanced Equation: ZnCl2(aq) + Mg(s) MgCl2(aq) + Zn(s)

Total Ionic Equation: Zn2+(aq) + 2Cl-

(aq) + Mg(s) Mg2+(aq) + 2Cl-

(aq) + Zn(s)

Zn2+(aq) + 2Cl-

(aq) + Mg(s) Mg2+(aq) + 2Cl-

(aq) + Zn(s)

Net Ionic Equation: Zn2+(aq) + Mg(s) Mg2+

(aq) + Zn(s)

c) When sodium iodide solution is mixed with aqueous lead (II) nitrate, a yellow lead (II) iodide precipitate is formed, leaving a solution of sodium nitrate.

Balanced Equation: 2NaI(aq) + Pb(NO3)2(aq) 2NaNO3(aq) + PbI2(s)

Total Ionic Equation: 2Na+(aq) +2I-

(aq)+ Pb2+(aq)+ 2NO3

-(aq)2Na+

(aq) + 2NO3-(aq) + PbI2(s)

2Na+(aq) +2I-

(aq)+ Pb2+(aq)+ 2NO3

-(aq)2Na+

(aq) + 2NO3-(aq) + PbI2(s)

Net Ionic Equation: 2I-(aq) + Pb2+

(aq) PbI2(s)

5. Write the net ionic equation for each of the following reactions. a) An aqueous solution of sodium sulphide reacts with hydrochloric acid to produce

hydrogen sulphide gas and a solution of sodium chloride.

Balanced Equation: Na2S(aq) + 2HCl(aq) 2NaCl(aq) + H2S (g)

Total Ionic Equation: 2Na+(aq) + S2-

(aq)+ 2H+(aq) + 2Cl-

(aq)2Na+(aq) +2Cl-

(aq) +H2S(g)

2Na+(aq) + S2-

(aq)+ 2H+(aq) + 2Cl-

(aq)2Na+(aq) +2Cl-

(aq) +H2S(g)

Net Ionic Equation: S2-(aq)+ 2H+

(aq) H2S(g) b) An aqueous solution of potassium hydroxide reacts with nitric acid to give a

solution of potassium nitrate and water.

Balanced Equation: KOH(aq) + HNO3(aq) KNO3(aq) + H2O(l)

Total Ionic Equation: K+(aq) + OH-

(aq) + H+(aq) + NO3

- (aq) K+

(aq) + NO3- (aq) + H2O(l)

K+(aq) + OH-

(aq) + H+(aq) + NO3

- (aq) K+

(aq) + NO3- (aq) + H2O(l)

Net Ionic Equation: OH-(aq) + H+

(aq) H2O(l)

Worksheet 13 – Electron Configurations of Atoms

1) Write the full and shorthand (condensed) electron configurations for each of the following atoms.

a) phosphorus 1s22s22p63s23p3 [Ne]3s23p3

b) beryllium 1s22s2 [He]2s2

c) fluorine 1s22s22p5 [He]2s22p5

d) calcium 1s22s22p63s23p64s2 [Ar]4s2

e) zinc 1s22s22p63s23p64s23d10 [Ar]4s23d10

f) tellurium, 1s22s22p63s23p64s23d104p65s24d105p4 [Kr] 5s24d105p4

g) sodium 1s22s22p63s1 [Ne]3s1

h) nickel 1s22s22p63s23p64s23d8 [Ar]4s23d8

i) oxygen 1s22s22p4 [He]2s22p4

j) iron 1s22s22p63s23p64s23d6 [Ar]4s23d6

k) magnesium 1s22s22p63s2 [Ne]3s2

l) aluminum 1s22s22p63s23p1 [Ne]3s23p1

2) Write shorthand (condensed) electron configuration for each of the following atoms.

a) holmium [Xe]6s24f11

b) osmium [Xe]6s24f145d6

c) lead [Xe]6s24f145d106p2

d) radium [Rn]7s2

e) bohrium [Rn]7s25f146d5

3) Determine which elements have the following electron configurations:

a) 1s22s22p63s23p4 S b) 1s22s22p63s23p64s23d104p65s1 Rb

c) [Kr] 5s24d105p3 Sb d) [Xe] 6s24f145d6 Os

e) [Rn] 7s25f11 Es

4) Determine which electron configurations are not valid.

a) 1s22s22p63s23p64s24d104p5 no b) 1s22s22p63s33d5 no

c) [Ra] 7s25f8 no d) [Kr] 5s24d105p5 yes

e) [Xe] no – better to put [Kr]5s24d105p6

5) Which block of the periodic table are each of the following elements located in?

a) einsteinium f b) polonium p

c) rutherfordium d d) francium s

6) Without looking at the periodic table, identify the group number, period number, and block of an atom that has each of the following electron configurations.

a) [Ne]3s1 group 1, period 3, s block

b) [He]2s2 group 2, period 2, s block

c) [Kr]5s24d105p5 group 17, period 5, p block

7) What is the number of orbitals in each of the following subshells?

a) 2s 1 b) 4p 3 c) 3d 5 d) 1p 0

e) 4f 7 f) 2p 3 g) 2d 5 h) 6g 9

8) What is the maximum number of electrons in each of the following subshells?

a) 2s 2 b) 4p 6 c) 3d 10 d) 1p 0

e) 4f 14 f) 2p 6 g) 2d 10 h) 6g 18

Worksheet 14 – Orbital and Energy Level Diagrams

1) Identify and correct the errors in each of the following valence shell orbital diagrams: Element s orbital p orbitals

a) carbon

Correct:

b) sulfur

Correct:

c) aluminum

Correct:

2) Consider the valence level electron configurations of the following five elements:Element s orbital p orbitals

1

2

3

4

5

a) Which of the above elements are metals? 2, 3, 4 could be metals

b) Which of the above elements would form an ion with a +2 charge? 3, possibly 4

c) Which of the above elements is a noble gas? 1, possibly 3 a halogen? 5

d) Which of the above elements has the same configuration as S2- ? 1

3) Write the orbital occupied by the last electron of each of the following elements:

As W Li U O Rn V

4p 5d 2s 5f 2p 6p 3d

4) Draw an orbital diagram for each of the following.

a) a cobalt atom (Co)

1s 2s 2p 3s 3p 4s 3d

b) a sulfide ion S2–

1s 2s 2p 3s 3p 4s

c) a cadmium ion (Cd2+)

1s 2s 2p 3s 3p 4s 3d

4p 5s 4d

5) Draw a complete energy level diagram (vertical) for an atom of iron, Fe.

Worksheet 15 – Quantum Numbers

1. State the four quantum numbers and the possible values they may have.

n = 1,2,3,4,5,6,7,8 , l = 0,1,2,3 , ml = -l...+l , ms = +½ , -½

2. Name the orbitals described by the following quantum numbers

    a. n = 3, L = 0 , 3s    b. n = 3, L = 1 , 3p    c. n = 3, L = 2 , 3d    d. n = 5, L = 0 , 5s

3. Give the n and L values for the following orbitals

    a. 1s , n = 1 l = 0     b. 3s , n = 3 l = 0    c. 2p , n = 2 l = 1    d. 4d , n = 4 l = 2    e. 5f , n = 5 l = 3

4. Place the following orbitals in order of increasing energy:

    1s, 3s, 4s, 6s, 3d, 4f, 3p, 7s, 5d, 5p , 1s, 3s, 3p, 4s, 3d, 5p, 6s, 5d, 7s

5. What are the possible mL values for the following types of orbitals?

    a. s , 0    b. p , -1, 0, +1    c. d , -2, -1, 0, +1, +2    d. f , -3, -2, -1, 0, +1, +2, +3

6. How many possible orbitals are there for n =

    a. 4 , 16    b. 10 , 100

7. How many electrons can inhabit all of the n=4 orbitals? 32

8. Tabulate all of the possible orbitals (by name, i.e. 4s) for n=4 and give the three quantum numbers which define each orbital.4s n = 4, l = 0, ml = 0 4p n = 4, l = 1, ml = -1, 0, +1 4d n = 4, l = 2, ml = -2, -1, 0, +1, +2 4f n = 4, l = 3, ml = -3, -2, -1, 0, +1, +2, +3

9. Write electron configurations for the following atoms:    a. H 1s1 d. F 1s22s22p5

    b. Li 1s22s1 e. Br 1s22s22p63s23p64s23d104p5

    c. N 1s22s22p3

   

Worksheet 16 - Atomic Structure Test Review

1) Which one of the following statements regarding Rutherford’s gold foil experiment is TRUE?a) Rutherford predicted that most of the alpha particles would pass straight through the

foil because atoms are mainly empty space.b) Rutherford predicted that the alpha particles that hit negative particles would bounce

back towards the radioactive source.c) The actual experimental results led Rutherford to conclude that atoms have a central

positive core because most of the alpha particles were deflected at large angles.d) Since most of the alpha particles passed straight through the foil without deflection,

Rutherford concluded that atoms are mainly empty space.e) The most surprising thing about the experimental results was that only some of the

alpha particles were deflected at large angles.

2) In Thomson’s model of the atom:a) negative particles are embedded within a sphere of positive chargeb) atoms are tiny, indivisible spheresc) positive particles are embedded within a sphere of negative charged) there is a positive nucleus surrounded by empty space

3) Which one of the following could be the electron configuration for an element with the following Lewis diagram?

a) 1s22s1 b) 1s22s22p3

c) 1s22s22p1 d) [Ar]4s23d3

e) more than one are correct

4) The species with the following energy level diagram is:

a) Ne b) Na

c) Na+ d) K

e) Cl–

5) In the emission spectrum of hydrogen, which electronic transition would produce a line in the visible region of the electromagnetic spectrum?

a) n = 2 n = 1b) n = 3 n = 2c) n = 2 n = 3d) n = n = 1

X

6) How many unpaired electrons are in the ground state energy level diagram for an atom of sulfur?

a) 1 b) 2 c) 3 d) 4 e) 5

7) Which of the following particles has the same electron configuration as krypton?a) Ga3+ b) Ga3– c) Sn4+ d) Sn2+ e) As3–

8) Complete the following table:

Name of Element

Symbol of

Atom or Ion

Atomic Number

Mass Number

Number of

Protons

Number of

Neutrons

Number of

Electrons

Overall Charge

a beryllium Br 4 9 4 5 4 0

b fluorine F- 9 19 9 10 10 1-

c iron Fe3+ 26 56 26 30 23 3+

d selenium Se2- 34 79 34 45 36 2-

9) Write the full electron configuration for each of the following atoms:a) strontium 1s22s22p63s23p64s23d104p65s2

b) germanium 1s22s22p63s23p64s23d104p2

10) Write the shorthand electron configuration for each of the following atoms:a) silicon [Ne]3s23p2

b) molybdenum [Kr]5s24d4

c) terbium [Xe]6s24f9

d) gold [Xe]6s24f145d9

e) seaborgium [Rn]7s25f146d4

11) Draw orbital diagrams for each of the following ions (shorthand), indicate which electrons have been gained/lost:a) Cu+

4s ↑ 3d ↑↓ ↑↓

↑↓ ↑↓

↑↓

b) Pb2+

6s ↑↓

5d ↑↓ ↑↓

↑↓ ↑↓ ↑↓ 6p ↑ ↑

c) Br–

4s ↑↓

3d ↑↓ ↑↓

↑↓ ↑↓ ↑↓ 4p ↑↓ ↑↓ ↑↓

d) Sb3–

5s ↑↓

4d ↑↓ ↑↓

↑↓ ↑↓ ↑↓ 5p ↑↓ ↑↓ ↑↓

e) Ba2+

6s ↑↓

f) Sn4+

5s ↑↓

4d ↑↓ ↑↓

↑↓ ↑↓ ↑↓ 5p ↑ ↑

12) What is the maximum number of electrons in:a) n = 3 , 18 b) 3d , 10 c) 6p , 6

d) 4f , 14 e) 2s , 2 f) n = 5 , 50

g) all d orbitals of curium, Cm , 30

13) State the four quantum numbers for each of the following:a) the 31st electron of Ga , 4, 1, -1, +½

b) the 12th electron of magnesium , 3, 0, 0, -½

c) the 80th electron of mercury , 4, 2, +2, -½

d) the 54th electron of xenon , 5, 1, +1, -½

e) the 68th electron of holmium , 4, 3, 0, +½

Worksheet 17 - Periodic Trends

1) Which of the following elements has the largest atomic radius?a) beryllium b) carbon c) nitrogen d) oxygen

2) Which of the following elements has the smallest atomic radius?a) sulfur b) selenium c) oxygen d) tellurium

3) Which of the following elements has the smallest first ionization energy?a) strontium b) calcium c) barium d) magnesium

4) Which of the following elements has the largest first ionization energy?a) bromine b) potassium c) arsenic d) calcium

5) Which of the following elements has the highest electron affinity?a) chlorine b) silicon c) sodium d) phosphorus

6) Which element in the following sets should have the largest atomic radius, and why?a) boron, lithium, or fluorine b) potassium, lithium, or sodium

7) Which element in the following pairs has the higher first ionization energy?a) cesium or gold b) sulfur or phosphorusc) magnesium or aluminum d) neon or kryptone) oxygen or selenium f) barium or calcium

8) Which element in the following sets loses an electron most readily, and why?a) boron, lithium, or fluorine b) potassium, lithium, or sodium- lowest ionization energy/largest size - lowest ionization energy/largest size- electron farther from attractive positive force of nuceleus

9) List each of the following sets in order of increasing radius:a) Ne, Na+, O2– Na+, Ne, O2–

b) Mg2+, Na+, Al3+ Al3+, Mg2+, Na+ c) N3–, O2–, F1– N3–, O2–, F1–

10) The ion Na+ and the atom Ne have the same electron configuration. To remove an electron from gaseous neon atoms requires 2081 kJ/mol. To remove an electron from a gaseous Na+ ion requires 4562 kJ/mol. Why are these values not the same?

Both have 10 electrons (2 electrons in first shell, 8 in second), but Ne has 10 protons in nucleus and Na+ has 11 protons in nucleus. (ENC of Ne is 0, ENC of Na+ is +1) The 11 protons are holding on to the electrons more strongly than neon’s 10 protons, therefore it would take more energy to remove an electron from Na+. (Note: You should draw/write configurations, orbital diagram or Bohr-Rutherford diagrams in your answer.)

11) The second ionization energy of magnesium is only about twice as great as the first ionization energy. However, the third ionization energy is about ten times as great as the first. Why does it take so much more energy to remove the third electron from magnesium?

A magnesium atom has 2 electrons in the valence shell. Once the first two electrons are removed, the third electron is a whole shell closer to the nucleus, therefore it would require significantly more energy to remove this electron. (You should draw/write configurations, orbital diagrams, or a Bohr-Rutherford diagram in your answer.)

12) The following graph shows the melting points (and boiling points) of the period 3 elements.

a) Why is there a general increase in melting point for Na, Mg, and Al?Smaller radius and increased nuclear charge cause increased strength of metallic bonding.

b) Why does Si have the highest melting point?Si is the only element in the period that exists as a network covalent structure.

c) Why is there a general decrease in melting point for P4, S8, and Cl2?Molecules with larger mass/size experience stronger intermolecular forces (van der Waals’ forces), thus S8 has slightly higher m.p. than P4, and Cl2 is lower than both of the others.

d) Why do the noble gases have the lowest melting point?Noble gases exist as single, diatomic elements, therefore very low van der Waals forces.

13) Although the first ionization energy of K is smaller than that of Ca, the second ionization energy of K is much higher than that of Ca. Why is this so?

Potassium has one valence electron, while calcium has two valence electrons. Once the first electrons have been removed, the second electron of potassium is in the third shell, while the second electron of calcium is still in the fourth shell. It would require much more energy to remove the second electron from potassium because it is much closer to the nucleus. (You should draw/write configurations, orbital diagrams, or Bohr-Rutherford diagrams in your answer.)

14) Element I1 I 2 I 3 I 1 = 1st ionization energy (kJ/mol) AA 531 1087 6270 I 2 = 2nd ionization energy (kJ/mol) BB 2090 3135 4180 I 3 = 3rd ionization energy (kJ/mol) CC 523 8360 11704

a) Which of the above elements would most likely be an alkali metal?CC would most likely be the alkali metal (group 1) because there is a big jump in ionization energy from I1 to I2 indicating that the second electron is a shell closer to the nucleus than the first.

b) Which of the above elements would most likely be a noble gas?BB would most likely, of the three elements above, to be a noble gas because there is no big jump in ionization energy. AA is a Group 2 element.

Worksheet 18 - Alkali Metals and Halogens

1) State whether each of the following properties increase or decrease going down the alkali metal group.

a) atomic radius increases

b) ionic radius increases

c) ionization energy decreases

d) electronegativity decreases

e) reactivity decreases

2) State whether each of the following properties increase or decrease going down the halogen group.

a) atomic radius increases

b) ionic radius increases

c) ionization energy decreases

d) electronegativity decreases

e) reactivity decreases

3) What is the state of matter at room temperature for each of the following?

a) F2 gas

b) Cl2 gas

c) Br2 liquid

d) I2 solid

4) For each of the following, circle the correct element.

Li Si S metal

N P As smallest ionization energy

K Ca Sc largest atomic mass

S Cl Ar member of the halogen family

Al Si P greatest electron affinity

Ga Al Si largest atomic radius

V Nb Ta largest atomic number

Te I Xe member of noble gases

Si Ge Sn 4 energy levels

Li Be B member of alkali metals

As Se Br 6 valence electrons

H Li Na nonmetal

Hg Tl Pb member of transition metals

Na Mg Al electron distribution ending in s2 p1

Pb Bi Po metalloid

B C N gas at room temperature

Ca Sc Ti electron distribution ending in s2 d2

Worksheet 19 – Polarity

In each of the following problems, rank the molecules from lowest to highest polarity:

1) PF3, LiOH, SF2, NF3

NF3 < SF2 < PF3 < LiOH

2) Ni(OH)3, N2H2, CH3OH, C2H5OHN2H2 < C2H5OH < CH3OH < Ni(OH)3

3) B2F4, H2C2O4, CuCl2, CF2ONO2 < SO < SeCl2 < PbCl2

4) PH3, PF3, NH3, NF3

PH3 < NH3 < NF3 < PH3

5) H2O, H2S, HF, H2

H2 < H2S < H2O < HF