KEY330 2014-15 Mid-Year Exam - REVIEW

CHAPTERS 1 & 2

1. Define chemistry: {the study of matter, its properties and changes}_____________________________2. Identify the quantity, unit and abbreviation represented by each of the following:

Quantity Unit Abbrev. Quantity Unit Abbrev.mass grams g volume liters Lenergy joules J amount of matter mol moldensity grams/millileter g/mL time seconds stemperature kelvin K molar mass grams/mole g/mol

3. Fill in the following table:

PropertyState of Matter

Solid Liquid GasMolecular motion Fixed Slow Very fastMolecular distance Close Close Far apart

4. Physical/Chemical Changes:

A. Define physical change: change is substance identity that doesn’t change its identity

B. Define chemical change: change in which one or more substances’ identities are changedC. Identify each as a chemical or physical change, and briefly explain.

i. paper burning: chemical (complete combustion would lead to CO2 & H2O)

ii. iron rusting: chemical (mix of iron oxides)

iii. ice melting: physical (still H2O)

iv. leaves turning color in the autumn: chemical

v. copper roof turning from orange to green: chemical (copper + oxygen copper oxide)

5. Define matter: something that has mass and takes up space

6. Define and give an example of each:

A. atom: smallest unit of an element that has the properties of that element)

B. compound: substance made of atoms of 2 or more elements chemically bonded

C. mixture: blend of two or more kinds of matter, each retains its own properties

D. heterogeneous: mixture that is not uniform throughout

E. pure substance: has fixed composition – every sample is same

F. which of the following is(are) a pure substance: element compound mixture solution

7. Define:

A. quantitative data: observation that includes definite number

B. qualitative data: observation that doesn’t include definite number

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8. Give one example of each:

A. direct relationship:_two quantities are directly related if dividing one by the other gives a constant; if one goes up, the other does too.

B. inverse relationship:___________two quantities are indirectly related if their product is a constant; if one goes up, the other goes down.

9. Convert:

A. 25 oC to K:________________________________________________________25 + 273 = 298 K

B. 293 g/cm3 to kg/L:293 g 1 kg 1 cm3

1000 mL =

293 kg

1 cm3 1000

g 1 mL 1 L L

10. Express 639,000,000 in scientific notation: 6.39 x 108

11. Periodic Table: Label metals, nonmetals and metalloids. Label s-, p-, d- and f-blocks Name the group identified by each letter:

A. s-block

B. d-block or transition metals

C. p-block

D. f-block orinner transition metals

E. alkali metalsF. alkaline earth

metals

G. halogens

H. noble gases

I. lanthanides

J. actinides

12. A student performs an experiment to determine the most effective chemicals to use in a cold pack. S/he added water and three different chemicals to four cups as given in the below table and recorded the temperature when the chemicals were first added (initial) and 5.0 minutes later (final).

Cup 1 Cup 2 Cup 3 Cup 4water: 100 mL 100 mL 100 mL 100 mLchemical (mass): NaCl (5.03 g) CaCl2 (4.98 g) NH4NO3 (5.00g) (none)initial temp: 21.5oC 22.0 oC 19.2 oC 20.7 oCfinal temp: 22.0 oC 29.3 oC 14.7 oC 21.3 oC

A. What was independent variable? chemical identity

B. What was the dependent variable? solution temperature

C. What was(were) the control? no chemical added (water only; Cup 4)

D. From this experiment, which chemical would be most effective to use in a cold pack? Why?

Ammonium nitrate (T = -4.5oC). The other chemicals raised (CaCl2 T = +7.3oC) by or essentially kept constant (NaCl T = +0.5oC) the temperature.

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E. What could be done differently to improve the validity of the findings?

1. replicates; 2. all initial temperatures the same as well as the temperatures of the chemicals, 3. other?

13. Define the law of conservation of mass:__ mass is neither created nor destroyed in a chemical reaction

14. According to the law of conservation mass, what is the relationship between the masses of the reactant(s) and product(s) in a chemical reaction? they’re equal

15. Identify the contribution of each to science:

A. Democritus: atomos (indivisible particles)

B. Dalton: 5 postulates for first modern atomic theory

C. Rutherford:_____discovered nucleus, proton, nucleus is positively charged & atom is mostly space(Solar System model of the atom)

D. Thomson: discovered electron (Plum-Pudding model)

E. Moseley: arranged elements on periodic table by atomic number (number of protons) not mass

F. Mendeleev modern periodic table

16. Most of the volume of an atom is occupied by: empty space/ electron cloud

17. What are the three subatomic particles comprising an atom? proton, neutron, electron

18. Which paerticle determines the identity of the atom? number of protons

19. Define mass number: (number of protons) + (number of neutrons)

20. Define atomic number: number of protons

21. Fill in the following spaces for the properties of the subatomic particles:

Property Neutron Proton Electron

symbol no p+ e–

location nucleus nucleus electron cloud

relative mass 1 1 1/2000

charge 0 + –

22. Define isotope:__atoms of the same element but with different masses (different numbers of neutrons)

23. Fill in the following spaces for isotopes:

Element IsotopeAtomic Number

Mass Number

Number Protons

Number Neutrons

Number Electrons

Oxygen-16 8 16 8 8 8

Magnesium-221222 Mg 12 22 12 10 12

Iron-562656 Fe 26 56 26 30 26

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24. Convert the following:

A. 293 g calcium to moles 293g Ca 1 mol Ca =

7.31

mol Ca

40.08 g Ca

B. 1.45 x 1024 atoms of zinc to moles

1.45E+24

atoms Zn 1 mol Zn =

2.41

mol Zn

6.02E+23

atoms Zn

C. The number of atoms in 12.0 g of gold._______________________________________________1.20E+0

1g Au 1

mol Au

6.02E+23

atom Au =

3.67E+22

atoms

1196.966

5 g Au 1mol Au

25. Name three types of electromagnetic energy: e.g., FM radio, TV, X-rays, microwaves

26. What problem was addressed by the development of the Bohr atom? line-emission of hydrogen

27. Draw a picture of the atom using the Bohr model. Identify the ground and excited states: (lowest = ground; others: excited)

28. Using the Bohr model of the atoms, describe how the line-emission spectrum of hydrogen is produced.

A. B. C.

Figure 1. Bohr model of the atom used to explain the line-emission spectrum of hydrogen. A. In this model, the hydrogen is shown initially at the ground state (electron is at n=1). B. Energy (e.g., photon of light) hits the electron with sufficient energy to excite it to the third excited state (n=4). C. However, the attraction between the negatively-charged electron and the positive nucleus causes the electron to move back towards the nucleus. In this case, it only returns to the first excited state (n=2) whereby it emits energy in order to do this. D. Eventually, the electron does return to the ground state (n=1) by further emitting energy.

29. What are wave functions? solutions to the Schrödinger wave equation

30. What does the Heisenberg uncertainty principle describe?__ that location & velocity of electron cannot

be simultaneous known.

31. What is the difference between an orbit and an orbital? orbit = fixed; orbital = probable

32. What is the highest occupied energy level in a bromine atom? 4p

33. Label the following orbital shapes.

A. B. C.

34. Which is a higher energy level: 2s or 2p? 2s or 3s? 4s or 4d?

35. How many electrons are each of the following energy sublevels? s ___2 p__6 d__10 f__14

36. What is the orbital notation for zinc?

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37. What is the noble gas notation for yttrium? [Kr] 4d1 5s2

38. What is the electron configuration for selenium? [Ar] 3d10 4s2 4p4

39. What is the electron configuration for barium? [Xe] 6s2

40. What is the noble gas configuration for iodine? [Kr] 4d10 5s2 5p5

41. For what element is the noble gas configuration [Xe]4f145d106s26p6? radon (Rn)

42. What are the horizontal rows on the periodic table called? periods

43. What are the vertical columns on the periodic table called? groups/families

44. Atomic radius:

A. What is the trend for atomic radius as one goes from left to right across a period? smaller

B. Explain.The basis for this results from the balance between (1) the attraction between electrons and the positively charged nucleus and (2) the repulsion between electrons. As one moves left to right across a period, one adds not only electrons but also protons. The increased pull by the more highly charged nucleus on the electrons is stronger than the repulsion between the electrons. Thus, the atomic radii across a period decrease.

C. What is the trend for atomic radius as one goes from top to bottom in a group? larger

D. Explain.

The radius of each atom gets larger as one goes down a group. This is expected: each row adds another shell of electrons and electrons repel electrons, making each shell larger,

45. Define valence electrons. Outer most s and p electrons

46. Define octet rule.______atoms tend to gain, lose, or share electrons in order to have a full set of eight valence electrons

47. Why is it called a “periodic table”? physical & chemical properties repeat

48. What is the most probable ion for each of the following?

A. calcium: Ca2+ B. phosphorus: P3–

C. iodine I– D. oxygen O2–

49. What is electronegativity? _____It is the measure of the ability of an atom in a chemical compound to attract electrons.

50. What is the relationship between atomic radius and electronegativity?____radius electronegativity

51. List five properties of metals:

A. conduct heat & electricity B. ductile

C. malleable D. have luster

E. high tensile strength

52. What are properties of each group?

A. alkali metals: most reactive metals, unstable in nature

B. alkaline earth metals: 2nd most reactive metals; unstable in nature

C. halogens: most reactive nonmetals; react vigorously with most metals to form salts

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D. noble gases: unreactive; gasses (@ room temp)

53. What two subatomic particles are involved in a chemical bond? protons & electrons

54. What are the diatomic molecules? a molecule containing two identical atoms (HONClBrIF)

55. What type of bond always connects two atoms in a diatomic molecule? nonpolar covalent

56. What type of bond is between each of the following pairs of atoms?

A. carbon and hydrogen:

(2.5-2.1) = 0.4polar covalent B. nitrogen and phosphorus:

(3.0-2.1) = 0.9polar covalent

C. lithium and fluorine

(4.0-1.0) = 3.0ionic D. chlorine and iodine

(3.0-2.5) = 0.4polar covalent

57. How many valence electrons are in each of the following atoms?

A. carbon: 4 B. silicon 4

C. sodium 1 D. bismuth 5

58. What is the formula unit for the compound formed by calcium and phosphorus? Ca3P2

59. How many electrons are in a double bond?______________________________________________4

60. Compare the differences in physical properties (including the ability to conduct electricity when either dissolved in water or in the molten state) of molecules and ionic compounds.

PHYSICAL PROPERTY IONIC COMPOUNDS MOLECULESMELTING POINT HIGH LOWBOILING POINT HIGH LOW

BRITTLE YES NOHARD YES NO

CONDUCTS ELECTRICITY MOLTEN STATE OR DISSOLVED IN WATER

NO

61. Which has the most energy – a single, a double, or triple bond?___________________________triple

62. Draw Lewis structures for following molecules:A. CH4

B. NH3

63. Formula units are used for: ionic compounds or molecules? ionic compounds

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64. What is, and give an example of, a polyatomic ion?_sulfate, nitrate, phosphate, hydroxide, acetate, etc

65. What is the correct formula for the compound formed by magnesium and phosphate? Mg3(PO4)2

66. What is the correct formula for the compound formed by aluminum and fluorine? AlF3

67. Name Cu2SO4 copper(I) sulfate

68. Name Ba3(PO4)2 barium phosphate

69. Name N2O10 dinitrogen decaoxide

70. How many nitrogen atoms are in NH4NO3? 2

71. What is the formula for magnesium acetate? Mg(C2H3O2)2

72. The formula for octane is C8H18. This compound is a(n) molecule, ionic compound, or metal?__________________________________________________________________________molecule

73. What does a subscript in a molecular formula indicate?____number of atoms immediately preceding it

74. Write a balanced equation for:

A. Magnesium and iron (III) oxide react to form magnesium oxide and iron.

3Mg(s) + Fe2O3(s) 3MgO(s) + 2Fe(s)

B. Propane gas (C3H8) burns.

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

C. Aqueous barium nitrate and sodium carbonate react to form barium carbonate and sodium nitrate.

Ba(NO3)2(aq) + Na2CO3(aq) 2NaNO3(aq) + BaCO3(s)

75. Write Ionic and Net ionic equations for # 74 (C)

76. Indicate whether the following are electrolytes:A. C3H8 No- moleculeB. Al(NO3)3 Yes- soluble ionic compoundC. Sr3(PO4)2 No- solid ionicD. H2SO4 Yes- acidE. CO2 No- gas

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77. Indicate whether the following are precipitates or aqueous:A. CaSO4 SolidB. Al(HCO3)3 AqC. NH4OH AqD. MgS SolidE. PbBr2 Solid

78. 15.24 grams Magnesium react with 23.54 grams of nitric acid. What volume of gas is produced?

79. 220.65 grams of phosphoric acid react with 46.00 grams magnesium. 156.87 grams of solid product are made. What is the % Yield?

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80. 245.01 liters of carbon dioxide gas are collected after 167.98 grams of propane (C3H8) are burned. What is the % Yield?

81. What is the % Water in copper (II) chloride hexahydrate?

82. What is the molecular formula of CH2O if the molecule has a molar mass of 540 grams/mole?

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83. What is the hybridization state of both carbons in acetic acid?

84. What is the Formal Charge for each atom in nitric acid?

85. What is the energy (Joules) of a radio station at 95.1 MHz?

86.

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87. What is the empirical formula for a compound consisting of 32% sodium, 23% sulfur, and the remainder oxygen?

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