Unit II

Atomic Theory

Worksheets

Lesson Date Topic WS #

1. Early Atomic Theory 1

2. Bohr Evidence 2

3. Bohr Diagrams 3

4. Quantum Theory 4

5. Mass Spectrometer/ Elegant Universe-1 5

6. Elegant Universe-2/Periodic Chem 6

7. Ionic Theory 7

8. Classifying Matter Lab 8

9. Classifying and Naming Formulas 1 9

10. Classifying and Naming Formulas 2 10

11. Electron Dot Diagram Structural Formula 1 11

12. Electron Dot Diagram 2 12

13. Practice Test 1 13

14. Practice Test 2 14

15. Test

Worksheet # 1 Early Atomic Theory

Briefly describe each atomic theory listed below. Include a diagram.

1. The Four-Element Theory

    The Four Element Theory

(a) evidence (b) explanation within theory

No evidence Nonscientific Theory

2. Dalton’s Atomic Theory

(a) Evidence (b) Explanation within theory

Conservation of mass Atoms are indestructible

Law of Constant composition           Elements combine in simple ratios

3. The Thompson Atom

(a) Evidence (b) Explanation within theory

Electrical Nature of Matter          Positive and negative particles

Worksheet # 2 Early Atomic Theory

1. The Rutherford Atom

(a) Evidence (b) Explanation within theory

A few alphas are radically deflected           Small dense nucleus

Most alphas are not deflected Most of atom is empty space

2. The Bohr Atom

The Bohr Atom

(a) evidence (b) explanation within theory

Line spectrum of discharge tubes Electrons are in orbitals

Worksheet # 3 Bohr Diagram

Draw Bohr atomic diagrams for the following atoms. Be sure to include protons, neutrons and electrons.

1. Oxygen 7. Calcium

2. Silver 8. Barium

3. Cs 9. I

4. Na 10. V

5. Cl^- 11. Al³⁺

6. Se^2- 12. Ca²⁺

Worksheet #4 Quantum Mechanics

1. What is the main difference between the Bohr Theory of the atom and the Quantum Mechanical Theory?

Electrons are waves in Quantum Theory and particles in the Bohr Theory.

2. How many electrons will fill the smallest orbital in quantum mechanical theory?

Two

3. How is a 3s orbital different than a 2s orbital in terms of shape and distance from the nucleus?

They are both spherical in shape but the 3s is further from the nucleus.

4. Explain what happens to the energy when an electron falls from a 3s orbital to a 2s orbital.

Energy is emitted in the form of a photon of light with energy corresponding to the difference between the two orbitals.

Use your Quantum Periodic Table to write quantum electron configurations for each element below.

5. F 1s²2s²2p⁵

6. K 1s²2s²2p⁶3s²3p⁶4s¹

7. C 1s²2s²2p²

8. Kr 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶

9. S 1s²2s²2p⁶3s²3p⁴

10. Rb 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶5s¹

11. Co 1s²2s²2p⁶3s² 3p⁶3d⁷4s²

12. P 1s²2s²2p⁶3s² 3p³

13. Ca 1s²2s²2p⁶3s² 3p⁶4s²

14. Al 1s²2s²2p⁶3s² 3p¹

15. Ag 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶4d⁹5s²

16. 1s²2s²2p⁶3s¹ Na

17. 1s²2s²2p⁶3s²3p⁵          Cl

18. 1s²2s²2p⁶3s²3p⁶3d⁹4s²          Cu

19. 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰5s²5p⁵ I

20. 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰4f¹⁴5s²5p⁶6s² Ba

21. Give the formula of four chemical species that are isoelectronic (same electron configuration) as Ar.

S^2- P^3- Cl^- K⁺ Ca²⁺

Worksheet # 5 Mass Spectrometry

Calculate the average atomic mass for each element. Round off to the correct number of sig figs. Write down the atomic mass from the periodic table rounded off to the same number of sig figs.

Isotope Mass % Abundance Average Mass Atomic Mass (table)

1.

¹⁴N 14.0030744 99.6340 14.0067 amu 14.0067 amu

¹⁵N 15.000108 0.366001

0.996340(14.0030744) + 0.00366001(15.000108) = 14.0067 amu

2.

²⁰Ne 19.992404 90.92 20.2 amu 20.1798 amu

²¹Ne 20.993849 0.257

²²Ne 21.991385 8.82

0.9092(19.992404) + 0.00257(20.993849) + 0.0882(21.991385) = 20.2 amu

3.

⁴⁶Ti 45.952633 7.93 47.9 amu 47.90 amu

⁴⁷Ti 46.95176 7.28

⁴⁸Ti 47.947948 73.94

⁴⁹Ti 48.947867 5.51

⁵⁰Ti 49.944789 5.34

You will lose marks if you don’t show the work!

4.

⁵⁴Fe 53.93962 5.8202 55.847 amu 55.845 amu

⁵⁶Fe 55.93493 91.660

⁵⁷Fe 56.93539 2.1901

⁵⁸Fe 57.93327 0.33001

You will lose marks if you don’t show the work!

5. Silver has two common isotopes. One is 106.90508 amu and 51.35 % and the other is 48.65 %. If the average atomic mass is 107.9730 amu, what is the atomic mass of the other isotope?

106.90508 (.5135) + X (0.4865) = 107.9730 amu

NOW SOLVE FOR X

109.1 amu

6. Copper has two common isotopes. One is 62.92959 amu and 69.09 % and the other is 30.91 %. If the average atomic mass is 63.5472 amu, what is the atomic mass of the other isotope.

62.92959 (.6909) + X (0.3091) = 63.5472 amu

NOW SOLVE FOR X

64.92 amu

7. Complete the chart below.

	protons	electrons	neutrons
^28Si	14	14	14
^29Si	14	14	15
^30Si	14	14	16

8. Write a quantum electron configuration for each of the following.

a) Ne 1s²2s²2p⁶

b) Mg

c) Ti 1s²2s²2p⁶3s² 3p⁶3d²4s²

d) Cr

e) Sr 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶5s²

f) Ag

g) Br 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁵

9. What was the first atomic theory to account for the Law of Conservation of Mass?

Explain how the theory accomplished this.

Dalton. Atoms are indestructible.

10. What was the first atomic theory to account for electromagnetic radiation (light)?

Explain how the theory accomplished this.

Bohr Theory. Electrons are in orbitals.

11. What was the first atomic theory to account for the small, dense nucleus?

Explain how the theory accomplished this.

Rutherford Atom. A few alpha particles were radically deflected.

12. What was the first atomic theory to have a wave theory for the electron?

Explain how the theory accomplished this.

Quantum Theory. Electrons vibrate around the nucleus in 3 dimensional wavelike orbitals.

13. What was the first atomic theory to account for positive and negative charges in matter? Explain how the theory accomplished this.

Thomson Atom. The matter in the atom was positive with negative particles throughout.

Worksheet # 6 Periodic Chemistry

1. Define the following:

a) Oxidation Loss of electrons

b) Reduction   Gain of electrons

c) Anion Negative ion

d) Cation    Positive ion

e) Atom  Neutral particle of an element

f) Chemical family Column on Periodic Table

g) Period Row on Periodic Table

2. Why are noble gases stable? Full outer or valence shells

3. Why are non-noble gases un-stable or reactive? Incomplete outer or valence shells

4. Draw Bohr diagrams for the following chemical species.

a) He b) K

c) K⁺ d) S^2-

e) P^3- f) Li⁺

5. Fill in the chart below.

symbol	atom, cation or anion	protons	neutrons	electrons	valence electrons	stable or reactive?
Mg^2+	cation	12	12	10	8	stable
Mg	atom	12	12	12	2	unstable
F	atom	9	10	9	7	unstable
F^-	anion	9	10	10	8	stable
Ne	atom	10	10	10	8	stable
C	atom	6	6	6	4	unstable
Be	atom	4	5	4	2	unstable
Be^2+	cation	4	5	2	2	stable
N^3-	anion	7	7	10	8	stable

Worksheet # 7 Ionic Chemistry

1.

symbol	atom, cation or anion	protons	neutrons	electrons	valence electrons	stable or reactive?
Ga	Atom	31	39	31	3	Reactive
Ga^3+	Cation	31	39	28	18	Stable
Br	Atom	35	45	35	7	Unstable
Br^-	Anion	35	45	36	8	Stable
Kr	Atom	36	48	36	8	Stable
Ca	Atom	20	20	20	2	Unstable
Ca^2+	Cation	20	20	18	8	Stable
P	Atom	15	16	15	5	Unstable
P^3-	Anion	15	16	18	8	Stable

2. What happens to protons, electrons and neutrons as you move form left to right within a row on the periodic table?

Protons, Electrons, and Neutrons all increase.

3. Write half-reactions to show how each atom forms an ion. Label each as oxidation or reduction. The first two are done for you.

a) K → K⁺ + 1e^- oxidation

b) N₂ + 6e^- → 2N^3-  reduction

c) P + 3e^- → P^3- reduction

d) O₂ + 4e^- → 2O^2-        reduction

e) Ca → Ca²⁺ + 2e^-       oxidation

f) Br₂ + 2e^- → 2Br^-      reduction

g) I₂ + 2e^- → 2I^- reduction

h) Al → Al³⁺ + 3e^-        oxidation

i) Ba → Ba²⁺ + 2e^-       oxidation

j) Cs → Cs⁺ + 1e^- oxidation

k) Mg → Mg²⁺ + 2e^- oxidation

l) Zn → Zn²⁺ + 2e^-        oxidation

m) Ga → Ga³⁺ + 3e^- oxidation

n) Cl₂ + 2e^- → 2Cl^- reduction

o) F₂ + 2e^- → 2F^- reduction

4. Describe five properties of:

a) Metals

Shiny Conductors Malleable Ductile Lose electrons Left side of periodic table

b) Non-metals

Dull Nonconductors Brittle Gain electrons Right side of periodic table

5. Draw Bohr diagrams for each of the following.

a) Na b) Na⁺

c) O d) O^2-

e) Ca f) Ca²⁺

Worksheet # 8

1. Complete the table.

	Salt	Base	Acid	Covalent Nonacid
Litmus	Neutral	Blue	Red	Neutral
Conductivity	Good	Good	Good	Non

2. Put each formula into the table below.

Ca(OH)₂ NH₄OH CH₃OH C₁₂H₂₄O₁₁

HCl PI₃ K₂SO₄ RbOH

H₃PO₄ NaOH CaCl₂ Li₂SO₄

H₂SO₃ BaF₂ BCl₅ CH₃COOH

H₂CO₃ CsOH S₂Cl₂ Fr₂S

Salt	Base	Acid	Covalent Nonacid
K2SO4	Ca(OH)2	HCl	CH3OH
CaCl2	NH4OH	H3PO4	C12H24O11
Li2SO4	RbOH	H2SO3	PI3
BaF2	NaOH	CH3COOH	BCl5
Fr2S	CsOH	H2CO3	S2Cl2

3. Draw Bohr diagrams for each of the following.

a) K⁺ b) P^3-

4. Write half-reactions to show how each atom forms an ion. Label each as oxidation or reduction. The first two are done for you.

a) Ca → Ca²⁺ + 2e^- oxidation

b) O₂ + 4e^- → 2O^2- reduction

c) I₂ + 2e^- → 2I^-      reduction

d) N₂ + 6e^- → 2N^3- reduction

e) Cs → Cs⁺ + e^-      oxidation

f) Ba → Ba²⁺ + 2e^- oxidation

g)      Al      →      Al³⁺   +       3e^-                     oxidation

h)      F₂ +   2e^-     →      2F^-                             reduction

i)       H₂     →      2H⁺      +      2e^-                    oxidation

j)       Na⁺      +   1e^- →   Na_(s)                     reduction

k)      N^3- →      N₂      +       6e^-   oxidation

l)       Ca²⁺    +   2e^-   →      Ca                          reduction

m)     Ba²⁺ +   2e^- → Ba reduction

Worksheet # 9

1. Complete the following table by classifying and naming each compound.

Formula	Classification	Name
CuS(s)	Salt	Copper II sulphide
H3PO4(s)	covalent nonacid (acids must be aq)	Hydrogen phosphate
P2O5(s)	Nonacid	Diphosphorus pentoxide
NH4OH(s)	Base	Ammonium hydroxide
Al2O3(s)	Salt	Aluminum oxide
MgSO4(s)	Salt	Magnesium sulphate
HCl(g)	covalent nonacid (acids must be aq)	Hydrogen chloride
HCl(aq)	Acid	Hydrochloric acid
H2SO4(l)	covalent nonacid (acids must be aq)	Hydrogen sulphate
H2SO4(aq)	Acid	Sulphuric acid
NI3(s)	Nonacid	Nitrogen triiodide
N3Cl3(s)	Nonacid	Trinitrogen trichloride
CO(g)	Nonacid	Carbon monoxide
K2CrO4(s)	Salt	Potassium chromate
H2Cr2O7(aq)	Acid	Dichromic acid
H2O(l)	Nonacid	Water
CrCO3(s)	Salt	Chromium II carbonate
HBr(g)	covalent nonacid (acids must be aq)	Hydrogen bromide
P3O5(s)	Nonacid	Triphosphorus pentoxide

Worksheet # 10

2. Complete the following table by classifying and naming each compound.

Formula	Classification	Name
HI (aq)	Acid	Hydroiodic acid
(NH4)3PO4(s)	Salt	Ammonium phosphate
NCl3(l)	Nonacid	Nitrogen trichloride
Ba(OH)2(s)	Base	Barium hydroxide
Rb2SO4(s)	Salt	Rubidium sulphate
CuCl2(s)	Salt	Copper II chloride
Al2O3(aq)	Salt	Aluminum oxide
N3Cl3(aq)	Nonacid	Trinitrogen trichloride
CO(g)	Nonacid	carbon monoxide
H2SO3(aq)	Acid	Sulphurous acid
CuSO4 ^.6H2O(aq)	Salt	Copper II sulphate hexahydrate
H3PO3(s)	covalent nonacid (acids must be aq)	Hydrogen phosphite
Mg3(PO4)2(aq)	Salt	Magnesium Phosphate
HCH3COO(aq)	Acid	Acetic Acid or Ethanoic acid
HF(aq)	Acid	Hydrofluoric acid
N2O5(aq)	Nonacid	Dinitrogen pentoxide
Na3PO4^. 5H2O(aq)	Salt	Sodium phosphate pentahydrate
Ni(NO3)3(aq)	Salt	Nickel III nitrate
SO(g)	Nonacid	Sulphur monoxide

Use your Quantum Periodic Table to write quantum electron configurations for each element below.

3. Sr 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶5s²

4. V 1s²2s²2p⁶3s²3p⁶3d³4s²

5. Mg 1s²2s²2p⁶3s²

6. P 1s²2s²2p⁶3s²3p³

7. Cr 1s²2s²2p⁶3s²3p⁶3d⁴4s²

Pick the best answers. Answers can be used more than once.

Answers: Four Element Theory

Dalton’s Atomic Theory

Thomson’s Atomic Theory

Rutherford’s Atomic Theory

Bohr’s Atomic Theory

Quantum Mechanical Theory

8. Rutherford’s Atomic Theory The 1^st model of the atom to explain the gold foil experiment

9. Bohr’s Atomic Theory The 1^st model to explain light

10. Dalton’s Atomic Theory The 1^st model to account for the Law of Constant Composition

11. Rutherford’s Atomic Theory The 1^st model to have a small, dense nucleus

12. Quantum Mechanical Theory The 1^st model to have an electron as a wave

13. Four Element Theory Non-scientific Theory

14. Thomson’s Atomic Theory The 1^st model to have electrons

15. Dalton’s Atomic Theory The 1^st model to account for the Law of Conservation of Mass

16. Quantum Mechanical Theory Modern theory of the atom

17. Rutherford’s Atomic Theory The 1^st model to claim the atom is mainly “empty space”

Worksheet # 11 Electron Dot Diagrams

Draw structural and electron-dot diagrams for each.

	Structural	Dot-Diagram
CH4
CI4
S2
P2
C2Cl6
C2F4
NF3
CS2
N2Cl2
HCN
CH4N2O (symmetrical)
C6H6 (cyclic)
CF4		.. : F : .. .. .. : F ׃ C ׃ F : .. .. .. : F : ..
N2Cl4	Cl – N – N – Cl | | Cl Cl
NBr3	Br – N – Br | Br	.. .. .. : Br : N : Br : .. .. .. : Br : ..
N2	N ≡ N	: N : : : N :
O2	O = O	.. .. : O : : O :
I2	I - I	.. .. : I : I : .. ..
CO2	O = C = O	.. .. : O : : C : : O :
COBr2	Br | O = C | Br	.. : Br : .. : O : : C .. : Br : ..
CNCl2F (symmetrical)		.. : Cl: .. .. : F : C : : :N .. .. : Cl : ..

Name each compound

1. CH₃COOH_(aq) Acetic or ethanoic acid

2. HBr_(aq) Hydrobromic acid

3. HF_(g) Hydrogen fluoride

4. HNO_3(aq) Nitric acid

5. HClO_4(aq) Perchloric acid

Write the quantum electron configurations for the following.

6. Cl^- 1s²2s²2p⁶3s²3p⁶

7. Sr²⁺ 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶

8. I 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶4d¹⁰5p⁵

Write a dissociation equation for each to show how each ionizes in water.

9. CH₃COOH_(l) → H⁺_(aq) + CH₃COO^-_(aq)

10. HNO_3(l) → H⁺_(aq) + NO₃^-_(aq)

11. Al₂(SO₄)_3(s) → 2Al³⁺_(aq) + 3SO₄^2-_(aq)

12. Co₃(PO₄)_2(s) → 3Co²⁺_(aq) + 2PO₄^3-_(aq)

Name each compound above.

13. Hydrogen acetate

14. Hydrogen nitrate

15. Aluminum sulphate

16. Cobalt II phosphate

17. Classify the following compounds.

NaOH             BaF₂ BCl₅ CH₃COOH

H₂CO₃ CsOH S₂Cl₂ BaCl₂

Salt	Base	Acid	Covalent Nonacid
BaF2	NaOH	CH3COOH	S2Cl2
BaCl2	CsOH	H2CO3	BCl5

Worksheet # 12 Electron Dot Diagrams

Draw structural and electron-dot diagrams for each.

ClO3^-		PO4^3-	.. 3- : O : .. .. .. : O : P : O : .. .. .. : O : ..
IO3^-	- .. : O : I : O : .. : O : ..	BrO3^-	- .. .. .. : O :Br: O : .. .. .. : O : ..
CN^-		NO3^-	- .. .. : O : N :: O : .. .. : O : ..
SO4^2-	2- .. : O : .. .. .. : O : S : O : .. .. .. : O : ..	CaCO3	2+       Ca
Li2SO4	+ 2 Li 2- .. : O : .. .. .. : O : S : O : .. .. .. : O : ..	CCl4	.. : Cl : .. .. .. : Cl ׃ C ׃ Cl : .. .. .. : Cl : ..
NI3	.. .. .. : I ׃ N ׃ I : .. .. .. : I : ..	NSCl	.. .. : Cl : N : : S : ..
NH4^+	H + .. H : N : H .. H	H3O^+
NaCl	+ − .. .. : Na : : Cl : .. ..	ClO3^-

Draw structural and electron-dot diagrams for each.

BrO4^-	3- .. : O : .. .. .. : O : Br : O : .. : O : ..	PO3^3-	3- .. : O : .. .. .. : O : P : O : ..
IO4^-	- .. : O : .. .. .. : O : I : O : .. .. .. : O : ..	NO3^-	.. - : O : .. .. .. : O : N : : O : ..
HCN	H : C : : : N :	SO3^2-	2- .. : O : .. .. .. : O : S : O : .. .. ..
CO3^2-		CaS	2+     Ca 2- .. : S : ..
Na2SO4	+ 2 Na                 2- .. : S : ..	NCl3
N2		O2	.. .. : O : : O :
Cl2	.. .. : Cl : Cl : .. ..	C2H6
C2H4		C2H2

Draw electron dot diagrams for each ionic compound

LiCl	.. [ Li ]^+ [ :Cl: ]^- ..
Na2O	.. [ Na ]^+[ :O: ]^2-[ Na ]^+   ..
K2S	.. [K ]^+[ : S : ]^2-[ K ]^+   ..
BaO	.. [ Ba ]^2+ [ :O: ]^2-  ..
GaH3	[ H: ]^- [ H: ]^-[ Ga ]^3+[ H: ]^-

Worksheet # 13 Practice Test # 1

1. Classify as stable or reactive.

Na N⁺ Ne Cl^- S^2- S^3-

P P^3- Ca Ca²⁺ NaCl N^3-

2. Describe a metal and a nonmetal in terms of gaining or losing electrons.

Metals lose electrons and nonmetals gain electrons.

3. Why are noble gases always stable?

Full outer shells

4. Determine the number of valence electrons for:

Ca 2 Ca²⁺ 8 Cl 7 Cl^- 8

O 6 O^2- 8 Al 3

5. Draw a Bohr diagram for

Ca Ca²⁺ N N^3-

6. Determine the number of protons, neutrons, and electrons in each.

	Protons	Neutrons	Electrons
S	16	16	16
S^2-	16	16	18
Al	13	14	13
Al^3+	13	14	10
^42Ca	20	22	20

7. Classify as ionic or covalent compounds.

HCl CH₃OH H₂O NH₄OH

NaCl MgSO₄ CoCl₂ H₃PO₄

NH₃ P₂O₅ Ba(OH)₂

8. Classify the above compounds into acids, non-acids, salts and bases.

Acids H₃PO₄ HCl

Bases NH₄OH Ba(OH)₂

Non-acids H₂O NH₃

Salts MgSO₄ NaCl

9. Calculate the average atomic mass for magnesium using the following percentage abundance data.

²⁴Mg 78.70% (24.00 amu)

²⁵Mg 10.13% (25.00 amu)

²⁶Mg 11.17% (26.00 amu)

0.7870(24.00) + .1013(25.00) + .1117(26.00) = 24.32 amu

10. Write the formulas for each ionic compound.

Magnesium chloride MgCl₂

Silver phosphate Ag₃PO₄

Cobalt III oxide Co₂O₃

Zinc phosphate Zn₃(PO₄)₂

Calcium nitride Ca₃N₂

Copper I nitrate CuNO₃

11. Name each ionic compound.

Fe₂(SO₄)₃ iron III sulphate CoCl₃ col

Na₂O sodium oxide AgCl

Na₃PO₄ sodium phosphate CaF₂

NH₄OH ammonium hydroxide Ca(NO₃)₂

K₂Cr₂O₇ potassium dichromate MgCrO₄

12. Name each covalent compound.

P₂O₅ diphosphorus pentoxide         N₂O₃ dinitrogen trioxide

CO carbon monoxide CO₂     carbon dioxide

SO₂ sulphur dioxide P₃O₅ triphosphorus pentoxide

C₆H₆ hexacarbon hexahydride H₂SO_4(l) hydrogen sulphate

HCl_(l) hydrogen chloride HNO_3(l) hydrogen nitrate

H₂CrO_4(aq) chromic acid HF_(aq)    hydrofluoric acid

H₃PO_4(l) hydrogen phosphate H₂CO_3(aq) carbonic acid

13. Complete the chart below.

	Protons	Neutrons	Electrons	Reactive or stable?	# of valence electrons
Li^+	3	4	2	stable	2
S^2-	16	16	18	stable	8
^22Mg^2+	12	10	10	stable	8

14. Write dissociation equations showing how each salt or acid dissolves in water and

forms ions. The first one is done for you.

MgCl₂ → Mg²⁺ + 2Cl^-

FeCl₃ → Fe³⁺ + 3Cl^-

Co₂(SO₄)₃ → 2Co³⁺ + 3SO₄^2-

HClO₃ → H⁺ + ClO₃^-

H₂Cr₂O₇ → 2H⁺ + Cr₂O₇^2-

16. Name each acid:

HCl_(aq) hydrochloric acid HClO_3(aq) chloric acid

HNO_3(aq) nitric acid HBr_(aq) hydrobromic acid

17. Classify each as formula units or molecules.

NaF H₂O

CH₄ NH₄Cl

Ionic compounds start with metals and have formula units.

Covalent compounds start with nonmetals and have molecules.

CaSO₄ CH₃OH

H₂SO₄

18. Indicate the solutions that conduct electricity.

NaCl_(s) solid NaCl_(aq)

Ca(OH)_2(aq) HCl_(aq)

NH₄OH_(aq) C₆H₁₂O_6(aq)

19. Consider the following electron configuration 1s²2s²2p⁶. Determine the element and

some ions that have the above electron configuration.

Element Ne

Cation Na⁺ or Mg²⁺

Anion F^- O^2- N^3-

20. Describe why NaCl_(s) doesn’t conduct electricity. Describe what happens to

NaCl_(s) when it is dissolved in water. Why does it conduct electricity?

The ions in NaCl_(s) are not free to move and conduct electricity.

When it is dissolved in water, the NaCl dissociates into ions that conduct electricity.

NaCl_(s) → Na⁺_(aq)  +   Cl^-_(aq)

21. Write the quantum electron configuration for each (1s²2s²p.....)

He 1s² Ar 1s²2s²2p⁶3s²3p⁶

Na 1s²2s²2p⁶3s¹ Na⁺ 1s²2s²2p⁶

Cl 1s²2s²2p⁶3s²3p⁵ Cl^- 1s²2s²2p⁶3s²3p⁶

K 1s²2s²2p⁶3s²3p⁶4s¹ K⁺ 1s²2s²2p⁶3s²3p⁶

Br 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁵ Br^- 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶

22. What two particles make up most of the mass within an atom? Protons and neutrons.

23. I am an atom with 35p 40n. Who am I? Br

24. I am a cation with 56p 81n & 54e. Who am I? Ba²⁺

25. Define isotopes. Elements that have the same atomic number but different

atomic mass because of having different amounts of neutrons.

26. In Rutherford’s Gold foil experiment some particles were completely

un-deflected and others were radically deflected. Describe the significance of each in terms of the structure of the atom.

There is a small dense positive nucleus in the center of the atom with most of the mass.

27. Define ionic and covalent bonding.

Ionic bonding transfers an electron from the metal, which becomes a cation to the nonmetal,

which becomes the anion.

Covalent bonding occurs between two nonmetals and involves shared electrons.

28. How many valence electrons are in the calcium ion?

8

29. How many valence electrons are in the fluoride ion?

8

30. What is the name of the family that has and electron configuration of:

a) s²p⁵ Halogens

b) s¹ Alkali Metals

c) s²p² Carbon Family

31. Consider the following electron configuration 1s²2s²2s²2p⁶3s²3p⁶ . Determine the element and some ions that have the above electron configuration. These are called isoelectronic.

Element: Ar

Cations: K⁺ Ca²⁺

Anions: Cl^- S^2- N^3-

Draw electron dot diagrams for the following using brackets for ions. Write a dissociation equation first.

32. NaCl

..

[ Na ]⁺ [ :Cl: ]^-

..

33. Li₂O

..

[ Li ]⁺ [ :O: ]^2- [ Li ]⁺

..

34. CaF₂

    .. ..

[ :F: ]^- [ Ca ]²⁺ [ :F: ]^-

.. ..

Name and classify each compound as an acid, molecular, salt, or base.

35. CuSO_4(aq) Salt Copper II sulphate

36. P₂O_4(s) Molecular Diphosphorus tetroxide

37. H₂SO_4(aq) Acid Sulphuric acid

38. H₂Cr₂O_7(aq) Acid Dichromic acid

39. H₂Cr₂O_7(l) covalent nonacid Hydrogen dichromate

40. Ca(OH)₂ ^. 5H₂O Base Calcium hydroxide pentahydrate

41. HBr_(aq) Acid Hydrobromic acid

42. Calculate the average atomic mass for neon if there are three naturally occurring isotopes and they are:

²⁰Ne mass = 19.9924404 amu abundance = 90.92 %

²¹Ne mass = 20.993849 amu abundance = 0.2570 %

²²Ne mass = 21.991385 amu abundance = 8.820 %.

Show some work if you want some marks. Round to an appropriate number of sig figs.

0.9092(19.9924404) + 0.002570(20.993849) + 0.08829(21.991385) = 20.17 amu

Ca(OH)₂ calcium hydroxide                  NH₄OH           ammonium hydroxide

CH₃OH            methanol                         C₁₂H₂₂O₁₁         sucrose

HCl hydrochloric acid                    PI₃ phosphorus triiodide

K₂SO₄ potassium sulphate RbOH rubidium hydroxide

H₃PO₄ phosphoric acid NaOH sodium hydroxide

CaCl₂                 calcium chloride                     Li₂SO₄                 lithium sulphate

SiO₂ silicon dioxide                        BaF₂ barium fluoride

BCl₅ boron pentachloride              CH₃COOH      acetic acid

H₂CO₃ carbonic acid CsOH cesium hydroxide

S₂Cl₂ disulphur dichloride               Fr₂S                 francium sulphide

Fe₂(SO4)₃ iron (III) sulphide                   ZnCl₂ zinc chloride

Co₃(PO₄)₂ cobalt (II) phosphate             Ag₂Cr₂O₇ silver dichromate

Worksheet # 14 Practice Test # 2

Balance each equation.

1. 2C₁₆H₃₄ + 49O₂ → 32CO₂ + 34H₂O

2. 2Ga + 3H₂SO₄ → 3H₂ + 1Ga₂(SO₄)₃

Write a balanced equation including phase symbols.

3. Solid carbon reacts with chlorine gas to produce liquid tetracarbon decachloride.

4C(s) + 5Cl_2(g) → C₄Cl_10(l)

Write chemical formulas for each ionic or molecular compound.

4. Strontium sulphide SrS

5. triphosphorous hexoxide P₃O₆

6. Osmium IV sulphide OsS₂

Name each chemical formula

7. Sn(CO₃)₂^.5H₂O Tin IV carbonate pentahydrate

8. Si₃F₈ Trisilicon octafluoride

9. NaHCO₃ Sodium bicarbonate

Classify the following as acids, bases, salts, and molecular non-acids. Name each.

10. Sn(SO₄)₂ Salt Tin IV sulphate

11. Ca(OH)₂ Base Calcium hydroxide

12. CH₃COOH Acid Acetic Acid

13. S₂O₅ Molecular Disulphur pentoxide

Round off each measured number to three significant figures.

14. 0.0056349 0.00563

15. 539663 540000 or 5.40 x 10⁵

Add or subtract the measured quantities.

16. 153.267 + 0.53493 153.802

17     ( 4.5631 x 10²⁴ ) ( 2.36 x 10^{- 23} )                    108

Simplify the following rounding to the correct number of significant figures.

18. (5.6 x 10 ^-24) (5.37894 x 10^-25)(6.532 x 10¹⁵) = 1.9 x 10^-67

(2.059378 x 10²⁴)(5.23 x 10²²)(9.37894 x 10^-13)

Use unit analysis and the conversion factors to perform the following conversions:

2.210 lb = 1.000 kg 14 lb = 1 stone (defined)

2000 lb = 1 ton (defined) 1.61 km = 1.00 mile

4.54 L = 1.00 gallon 16 oz = 1 lb (defined)

19. 236 oz to stone

236 oz x 1 lb x 1 stone = 1.05 stone

16 oz 14 lb

20. 8.53 stone to oz

8.53 stone x 14 lb x 16 oz = 1.9 x 10³ oz

1 stone 1 lb

21. 25.6 mi/h to km/s

25.6 mi x 1.61 km x 1 h = 0.0114 km/s

h 1.00 mi 3600s

State the model of the atom is best described by each statement below.

22. First model to account for The Law of Conservation of Mass. Dalton

23. The first theory to explain the emission of photons.           Bohr

24. First model to account for positive and negative charges. Thomson

25. First model to account for the wave properties of electrons.           Quantum

26. First model to include a small dense nucleus.          Rutherford

27. Describes the atom as a small dense nucleus surrounded with electrons, which are not in orbitals. Rutherford

28. Describes the atom as a small dense nucleus surrounded with electrons, which are in spherical orbitals.Bohr

29. Describes the atom as a spherical atom that is indestructible and combines in simple whole number ratios to form compounds. Dalton

30. Describes electrons as being contained in a complex 3D orbitals as negative clouds of vibrational energy. Quantum

31. Non-scientific theory that delayed modern theories of the atom for 1800 years and was shown to be incorrect. Aristotle’s

What did the evidence tell us about the nature of the atom?

32. 99 % of alphas in the gold foil experiment were completely un-deflected.

Most of the atom is empty space.

33. 1 % of alphas in the gold foil experiment were radically deflected.

There is a small dense positive nucleus.

34. Flame spectroscopy of an element produces an emission spectrum consisting of 4 photons.

Electrons are in orbitals

35. A beam of negative particles is produced in a Crooke’s tube.

There are electrons.

36. There are five naturally occurring isotopes of Germanium. Complete the chart to show the number of protons neutrons and electrons.

protons neutrons electrons At. Mass Abundance

⁷⁰Ge 32 38 32 69.92428 20.52%

⁷²Ge 32 40 32 71.92174 27.43%

⁷³Ge 32 41 32 72.9234 7.760%

⁷⁴Ge 32 42 32 73.92115 36.54%

⁷⁶Ge 32 44 32 75.9214 7.760%

37. Calculate the average atomic mass of Germanium. Show some work if you want some marks. Round to an appropriate number of significant figs.

0.2052(69.92428) + 0.2743(71.92174) + .07760(72.9234) + 0.3654(73.92115) + .07760(75.9214) = 72.64 amu

38. Write the quantum electron configurations for the following atoms or ions.

39. F 1s²2s²2p⁵

40. Ga 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p¹

41. Br^- 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶

42. Rb⁺ 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶

Determine the element that has the following electron configuration.

43. 1s²2s²2p⁶3s² Mg

Determine a cation that has the following electron configuration.

44. 1s²2s²2p⁶3s²3p⁶ K⁺ Ca²⁺ Ga³⁺

Determine an anion that has the following electron configuration.

45. 1s²2s²2p⁶3s²3p⁶            P^3- S^2- Cl^-

Complete the following chart.

Symbol p e n valance el. stable/unstable atom/cation/anion

46. Na 11 11 12 1 unstable atom

47. P^-3 15 18 16 8 stable anion

48. Xe 54 54 77 8 stable atom

49. Sr⁺² 38 36 50 8 stable cation

Name and classify each compound as an acid, molecular, salt, or base.

50. CuSO_4(aq) Salt      Copper II sulphate

51. P₂O_4(s) Molecular Diphosphorus tetroxide

52. H₂SO_4(aq) Acid Sulphuric acid

53. H₂CO_3(aq) Acid      Carbonic acid

54. H₂CO_3(l) Nonacid Covalent Hydrogen carbonate

55. Ba(OH)₂^. 2H₂O Base       Barium hydroxide dihydrate

56. HF_(aq) Acid       Hydrofluoric acid

Write structural diagrams and electron-dot diagrams for each.

57. CCl₄

  ..

: Cl :

.. .. ..

: Cl ׃ C ׃ Cl :

.. .. ..

: Cl :

..

58. S₂

.. ..

: S : : S :

59. NH₃

..

H : N : H

..

H

60. CO₂

.. ..

: O : : C : : O :

Write electron-dot diagrams for each.

61. NaCl

.. ..

[ : Na : ]⁺ [ : Cl : ]^-

.. ..

62. SO₄^2-

2-

..

: O :

.. .. ..

: O : S : O :

.. .. ..

: O :

..

63. Na₃PO₄

64. NO₃^-

.. .. -

: O : N :: O :

.. ..

: O :

..

65. NH₄⁺

66. IO₄^-

      -

..

: O :

.. .. ..

: O : I : O :

.. .. ..

: O :

..