UNIT ONE: Introduction to Chemistry and · PDF fileRegents Chemistry Mid-Term Exam Study Guide Revised: January 2009 UNIT ONE: Introduction to Chemistry and Measurements Content Outline

Regents Chemistry Mid-Term Exam Study Guide

Revised: January 2009

UNIT ONE: Introduction to Chemistry and Measurements Content Outline

a) Scientific Method a. Variables vs. Control b. Designing an experiment

b) Metric System

a. Conversions (Reference Table A, B, C and D)

c) Measurements a. Precision / accuracy b. Significant figures

i. All measured digits (certain) and one estimated (uncertain) ii. Addition / Subtraction = answer has as many places to the right of the

decimal as the number with the fewest places iii. Multiplication / Division = answer has as many significant figures as the

number with the fewest significant figures

Number Sig Figs 2077 4 20.77 4 20770 4 20770. 5 0.207 3 0.2070 4 0.00207 3

c. Scientific notation

i. 10,000,000 is equal to 1.0 x 10 7 ii. 0.00000001 is equal to 1.0 x 10 -8

d. Common lab equipment

Crucible Watch Glass Mortar and Pestle Graduated Cylinder

Erlenmeyer Test Tube Holder Clay Triangle Test Tube Clamp Flask

Beaker

e. Percent error

i. Reference Table T

d) Problem solving a. Steps of the scientific method b. Density (Reference Table T)

Student Objectives • Describe steps in scientific method • List basic safety procedures in science class • Identify metric units for length, width, and volume • Demonstrate differences between precision and accuracy • Name and explain use of significant figures • Determine and explain significant digits • Apply percent error to lab • Apply density to real-life situations • Use conversion factors • Name and describe four (4) steps to problem solving

UNIT TWO: Atomic Structure and the Periodic Table Content Outline

a) Development a. The periodic table is arranged in order of increasing atomic number b. Properties of elements differ most greatly from group to group

b) Basic definitions a. Element v. Compounds

i. Elements cannot be broken down ii. Compounds can be broken down

b. Physical v. Chemical properties c. Periods:

i. Horizontal rows (left and right), indicate principal energy level d. Groups / Families

i. Vertical columns (up and down) ii. Similar chemical properties due to same number of valence electrons

c) Atomic Structure a. Protons, electrons, neutrons b. Mass number

i. Protons plus neutrons c. Atomic number

i. Equal to the number of protons d. Energy levels (quantum number)

i. Equal to the period (row) which an element is found e. Valence numbers

i. Lewis dot diagrams ii. Number of electrons in outer shell

f. Octect rule i. All elements (except for hydrogen and helium) want 8 electrons in their

valence shell

d) Classification of elements a. Metals

i. Make up 2/3 of the atoms ii. Have low ionization energies and low electronegativities iii. Lose electrons to form positive ions iv. All solid at room temperature except for Hg v. Are malleable, ductile, good conductors, have a shiny luster

b. Nonmetals i. Upper right side of the periodic table ii. High ionization energies and high electronegativities iii. Fluorine has the highest electronegativity iv. Gain electrons to form negative ions v. Poor conductors of heat and electricity

c. Metalloids i. Properties of both metals and non-metals ii. Found on both sides of the “staircase” on the PT

d. Solids, liquids, and gases

e) Groups and Families a. Alkali metals

i. Group 1 ii. Very reactive iii. Low ionization energies iv. Low electronegatvities v. Found only in compounds

b. Alkaline earth metals i. Group 2 ii. Same as alkali metals

c. Transition metals i. Groups 3-11 ii. Give compounds their color

d. Carbon, nitrogen, oxygen families

e. Halogens i. Group 17 ii. F and Cl are gases iii. Br is a liquid iv. I and At are solids

f. Noble gases i. Group 18 ii. Unreactive

f) Forms of elements a. Ions

i. Charged atoms or molecules b. Isotopes

i. Atoms of the same element with different atomic masses ii. C-12 & C-14 iii. Average atomic mass

g) Trends

a. Ionization energy i. Increases from left to right ii. Decreases from top to bottom

b. Electronegativity i. Increases from left to right ii. Decreaes from top to bottom

c. Atomic radius i. Decreases from left to right ii. Increases top to bottom

Student Objectives

• State the periodic law • Explain the basic structure of an atom • Define mass number, atomic number, stable octet, valence number, and energy level • Explain why elements in a group have similar properties • Define and describe trends in the periodic table • Define: ion, isotopes • Relate position in periodic table to the number of energy levels and valence electrons • Describe the physical and chemical characteristics of each family or group in the periodic table • Calculate the average atomic mass of an element, given the percent of naturally occurring isotopes

for that element • Describe general physical and chemical properties of metals, nonmetals, and metalloids

Content Outline

a) Historical contributions of a. Rutherford’s gold foil experiment

b) Basic atomic structure a. Protons, neutrons, electrons b. Atomic mass unit c. Valence electrons d. Lewis-Dot diagrams

c) Bohr Model a. Energy levels

i. Electron configuration (ie: Br = 2-8-18-7) b. Ground state v. Excited state

i. Ground state is normal electron configuration ii. Excited state is not (ie: Br = 2-8-17-8) iii. Energy is absorbed when electrons become excited iv. Energy is released (heat and light) when it returns to ground state

c. Sublevels: s, p, d, f

d) Electron cloud a. Orbitals

Student Objectives:

• Define atom • Explain organization of atomic models • Explain the contributions of Rutherford’s experiment to atomic theory • Name and describe three (3) subatomic particles • Explain Bohr’s model of the atom • Distinguish between ground and excited state • Describe sublevels represented in energy levels of atoms • Define orbitals and electron cloud model

UNIT THREE: The Mathematics of Chemistry Content Outline

a) Types of chemical formulas a. Empirical b. Molecular c. Structural

b) Nomenclature / Formula writing a. Binary compounds v. Ternary compounds

i. Binary = only two elements ii. Ternary = three or more

b. Prefixes i. Mono- , di-, tri-, etc

c. Roman numerals i. Copper (II) oxide

d. Polyatomic Ions i. Reference Table E

e. Acids i. Begin with a hydrogen ii. Polyatomic ions ending in “-ate”, form acids ending in “-ic” iii. Polyatomic ions ending in “-ite”, form acids ending in “-ous”

c) Formula calculations

a. Moles (Table T) b. Gram to mole conversions (Table T, use gram formula mass) c. Gram formula mass d. Percent composition (Table T) e. Empirical formulas

i. Convert a % composition to grams ii. Divide grams by gram formula mass to get moles iii. Divide smallest number of mules into all the moles to get a ratio of elements iv. Write the empirical formula using the mole ratios as the subscripts for each

element Example: 75% Carbon and 25% H Step 1: 75 g carbon 25 g Hydrogen Step 2: 75 g / 12 g 25 g / 1 g Step 3: 6.25 moles C 25 moles H Step 4: 6.25 / 6.25 = 1 25 / 6.25 = 4 Final Answer: CH4

f. Molecular formulas from the molecular mass of the empirical formula i. Determine the mass of the molecular formula ii. Divide the molecular mass by the mass of the empirical formula to get number

of formula units iii. Empirical formula multiplied by # formula units = molecular formula

d) Equations

a. Information in an equation b. Types of equations

i. Addition ii. Decomposition / analysis iii. Single replacement iv. Double replacement

c. Laws of conservation of mass and energy i. Mass and energy can neither be created nor destroyed

d. Balancing equations

e) Stoichiometry (Reference Table T) a. Moles b. Mass c. Liters

Student Objectives

• Name binary compound, ternary compounds, and acids • Demonstrate use of Stock system / prefixes in naming compounds • Write formulas, given the name of the compound • Explain the difference between molecular, empirical, and structural formulas • Describe the meanings of symbols used in chemical equations • Identify the four (4) types of equations • Explain how the Laws of Conservation of Mass and Energy applies to balanced equations • Define a mole and describe its importance to chemistry • Relate gram formula mass, Avogadro’s number, moles, and molar volume to each other • Find the percent composition of a given formula • Define hydrates and find the percent of water in a hydrate • Use percent composition to determine the empirical formula of an unknown sample • Given the molecular mass and percent composition or empirical formula, determine the molecular

formula

UNIT FOUR: Bonding Content Outline

a) Bond formation a. Stable octet (Nobel gas configuration) b. Energy transferred c. Lewis-dot diagrams for ionic and covalent bonds d. Electronegativity predicting bond type

i. The ability to attract electrons ii. Difference of 1.7 and higher is ionic iii. Difference of 1.6 and lower is covalent

e. Ionization energy i. Amount of energy needed to take an electron away

b) Types of bonds a. Ionic

i. Ion formation ii. Give and take electrons

b. Covalent i. Share electrons ii. Multiple bonds

1. single, double, and triple iii. Unsaturated compounds

1. have a double or triple bond iv. Saturated compounds

1. have all single bonds c. Metallic

i. “sea of electrons” d. Coordinate covalent bonds

i. Pair of shared electrons come from one atom only

c) Polarity of molecules a. Dipoles b. Symmetry / Asymmetry

i. Symmetrical = non-polar (equal pulling) 1. higher electronegativity wins the electron

ii. Asymmetrical = polar (unequal pulling) 1. WATER is a polar molecule

c. Shapes i. Linear

ii. Tetrahedral

iii. Bent

iv. Pyramidial

d) Intermolecular forces a. Hydrogen bonding

i. Attraction between dipole molecules b. Van der waals forces

i. “gravity” the bigger the molecule, and the closer they are causes more attraction

ii. Attraction between non-polar molecules Student Objectives

• Describe characteristics of ionic bonds • Describe properties of ionic compounds • Explain octet rule • Draw Lewis dot structures to show electron configuration • Define covalent bond • Distinguish between molecular and structural formula • Explain difference between single, double, and triple covalent bonds • Compare and contrast polar and non-polar covalent bonds • State two (2) factors that determine polarity of a molecule • Discuss metallic bonding • Discuss network solids • Describe different types of intermolecular forces

UNIT FIVE: Matter and Energy Content Outline

a) Energy types

a. Potenital (energy of position) i. Chemical

b. Electrical c. Kinetic

i. Mechanical (energy of motion) ii. Thermal (energy of heat)

d. Radiant energy (light energy) e. Nuclear energy

b) Law of conservation of energy a. Definition of temperature b. Temperature scales

i. Absolute zero (-273 °C) ( 0 K)

c) Heating and Cooling curves a. Joules b. Endothermic / Exothermic Reactions c. Terms and energy involved with:

i. Vaporization ii. Fusion iii. Sublimation iv. Condensation v. Deposition

d. Changes in Potential / Kinetic Energy

AB, CD, & EF: all have temperature changes and increase in kinetic energy BC, & DE: no temperature change, but change of phase, increase in potential energy

d) Phase changes a. Changes in energy b. Changes in intermolecular distance c. Properties of solid, liquid, gas, and plasma d. Unique properties of water

i. Viscosity ii. Surface tension iii. Universal solvent iv. Vapor pressure (Reference Table H)

e) Matter

a. Physical vs. Chemical properties b. Elements vs. Compounds c. Physical vs. Chemical change d. Mixtures

i. Homogeneous 1. alloy

ii. Heterogeneous iii. Separation of mixtures

1. Physical properties: density, particle size, molecular polarity, boiling point, freezing point

2. Processes: distillation, chromatography, filtration

Student Objectives • Name and describe four (4) basic types of energy • Explain the law of conservation of energy • Compare °C / K temperature scales • Explain absolute zero • Define activation energy • Describe the difference between endothermic and exothermic reactions • Describe how energy and chemical reactions are related • Explain how different intermolecular forces influence properties of solid, liquid, and gases • Describe major events on a heating and cooling curve • Interpret phase diagrams • Solve calorimetry problems using phase diagrams • Describe four (4) states of matter, including examples • Describe differences between chemical and physical changes • Explain the law of conservation of matter • Define substance, mixture, homogeneous, heterogeneous • Describe four (4) techniques and properties to separate mixtures

1. In a sample of pure copper, all atoms haveatomic numbers which are(1) the same and the atoms have the same number of

electrons(2) the same but the atoms have a different number of

electrons(3) different but the atoms have the same number of

electrons(4) different and the atoms have a different number of

electrons

2. Which formula represents a binary compound?(1) NH4NO3(2) CH4(3) CH3COCH3(4) CaCO3

3. Which of the following represents a mixture?(1) NaCl(aq)(2) NaCl(s)(3) NH3(g)(4) NH3(…)

4. The compound whose molecules have the highest average kinetic energy is(1) NO(g) at 25°C(2) N2O(g) at 15°C(3) NO2(g) at 30°C(4) N2O3(g) at 20°C

5. Which kind of energy is stored within a chemical substance?(1) free energy (3) kinetic energy(2) activation energy (4) potential energy

6. Which change of phase is exothermic?(1) H2O(s) → H2O(g)(2) CO2(s) → CO2(…)(3) H2S(g) →

H2S(…)(4) NH3(…) → NH3(g)

7. Base your answer to the following question on the graphs shown below.

Which graph best represents how the volume of a given mass of a gas varies with the pressure exerted on it at constant temperature?(1) 1 (3) 3(2) 2

8. A sample of gas A was stored in a container at a temperature of 50ºC and a pressure of 0.50 atmosphere. Compared to a sample of gas B at STP, gas A had a(1) higher temperature and a lower pressure(2) higher temperature and a higher pressure(3) lower temperature and a lower pressure(4) lower temperature and a higher pressure

9. A sample of a gas occupies 6.00 liters at a temperature of 200. K. If the pressure remains constant and the temperature is raised to 600. K, the volume of the gas sample would be(1) 18.0 L (3) 3.00 L(2) 2.00 L (4) 12.0 L

10. A 2.5 liter sample of gas is at STP. When the temperature is raised to 273ºC and the pressure remains constant, the new volume of the gas will be(1) 1.25 L (3) 5.0 L(2) 2.5 L (4) 10. L

11. The boiling point of water at standard pressure is(1) 0.000 K (3) 273 K(2) 100 K (4) 373 K

12. Base your answer to the following question on the graph below, which represents uniform cooling of a sample of a pure substance, starting as a gas.

The boiling point of the substance is(1) 10°C (3) 120°C(2) 60°C (4) 180°C

13. The heat of fusion is defined as the energy required at constant temperature to change 1 unit mass of a(1) gas to a liquid (3) solid to a gas(2) gas to a solid (4) solid to a liquid

14. What occurs when a substance melts?(1) It changes from solid to liquid, and heat is absorbed.(2) It changes from solid to liquid, and heat is released.(3) It changes from liquid to solid, and heat is absorbed.(4) It changes from liquid to solid, and heat is released.

15. The chart below shows the change in vapor pressure of four pure liquids with increasing temperature. Which liquid has the lowest normal boiling point?

(1) A (3) C(2) B (4) D

16. The graph below represents the uniform heating of a substance, starting with the substance as a solid below its melting point.

Which segment of the graph represents a time when both the solid and liquid phases are present?(1) AB (3) DE(2) BC (4) EF

17. What is the mass number of a 31H atom?(1) 1 (3) 3(2) 2 (4) 4

18. What is the total number of protons contained in the nucleus of a carbon-14 atom?(1) 6 (3) 12(2) 8 (4) 14

19. What is the total number of electrons in a neutral atom of fluorine?(1) 9 (3) 19(2) 10 (4) 28

20. What is the total number of electrons in a Mg2+ ion?(1) 10 (3) 12(2) 2 (4) 24

21. Which pair of atoms are isotopes?(1) 14

6 C and 147 N

(2) 4019K and 40

18Ar(3) 222

88 Ra and 222 86 Rn

(4) 4019K and 42

19K

22. Compared to an atom of 126 C, an atom of 14

6 C has(1) more protons (3) more neutrons(2) fewer protons (4) fewer neutrons

23. The principal quantum number of the outermost electron of an atom in the ground state is n = 3. What is the total number of occupied principal energy levels contained in this atom?(1) 1 (3) 3(2) 2 (4) 4

24. Which electron configuration represents an atom of an element having a completed third principal energy level?(1) 2-8-2 (3) 2-8-10-2(2) 2-8-6-2 (4) 2-8-18-2

25. Which is the atomic number of an atom with six valence electrons?(1) 6 (3) 10(2) 8 (4) 12

26. Which is the correct electron-dot symbol for a boron atom in the ground state?

(1) (3)

(2) (4)

27. Which of the following atoms has the greatest tendency to gain electrons?(1) Al (3) I(2) Rb (4) F

28. Which compound would have the greatest degree of ionic character?(1) Na2O(2) H2O(3) CO2(4) NO2

29. Which pair of elements form a bond with the most ionic character?(1) KCl (3) PCl(2) ICl (4) HCl

30. Which type of bond is present in a water molecule?(1) polar covalent (3) ionic(2) nonpolar covalent (4) electrovalent

31. Which molecule contains a nonpolar covalent bond?

(1) (3)

(2) (4)

32. Which could form a coordinate covalent bond?

(1) (3)

(2) (4)

33. Which formula represents a molecular substance?(1) CaO(2) CO(3) Li2O(4) Al2O3

34. Which type of bond is present in copper wire?(1) covalent (3) electrovalent(2) ionic (4) metallic

35. Which molecule is polar?

(1)

(3)

(2)

(4)

36. The unusually high boiling point of water is due to the(1) network bonds between the molecules(2) hydrogen bonds between the molecules(3) linear structure of the molecules(4) nonpolar character of the molecules

37. Which of the following liquids has the weakest van der Waals forces of attraction between its molecules?(1) Kr(…) (3) Ar(…)(2) Ne(…) (4) He(…)

38. Which compound has the empirical formula CH?(1) CH4(2) C2H4(3) C6H6(4) C3H8

39. The correct name of the compound with the formula PbO2 is(1) lead (I) oxide (3) lead (III) oxide(2) lead (II) oxide (4) lead (IV) oxide

40. The correct formula for calcium phosphate is(1) CaPO4(2) Ca2(PO4)3(3) Ca3P2(4) Ca3(PO4)2

41. When the equation

__H2 + __N2 → __NH3

is completely balanced using smallest whole numbers, the sum of all the coefficients will be(1) 6 (3) 3(2) 7 (4) 12

42. Given the unbalanced equation:

__CaSO4 + __AlCl3→Al2(SO4)3 + __CaCl2

What is the coefficient of Al2(SO4)3 when the equation is completely balanced using the smallest whole-number coefficients?(1) 1 (3) 3(2) 2 (4) 4

43. Which element in Period 3 has the least tendency to lose an electron?(1) argon (3) phosphorus(2) sodium (4) aluminum

44. Which element in Period 4 of the Periodic Table exhibits the most nonmetallic properties?(1) Ca (3) Ga(2) Cr (4) Br

45. Which is the electron configuration of a transition element in the ground state?(1) [Ar]4s1 (3) [Ar]3d104s24p1

(2) [Ar]3d104s1 (4) [Ar]3d104s24p6

46. Which element exists as a diatomic molecule at STP?(1) bromine (3) sulfur(2) argon (4) rubidium

47. Which group contains elements in three phases of matter at STP?(1) Group 2 (3) Group 15(2) Group 6 (4) Group 17

48. Which is an example of a metalloid?(1) sodium (3) silicon(2) strontium (4) sulfur

49. Which two elements have chemical properties that are most similar?(1) Cl and Ar (3) K and Ca(2) Li and Na (4) C and N

50. A chloride dissolves in water to form a colored solution. The chloride could be(1) HCl(2) KCl(3) CaCl2(4) CuCl2

51. Which element is a noble gas?(1) antimony (3) gold(2) krypton (4) francium

52. What is the gram formula mass of Ca(OH)2?(1) 29 g (3) 57 g(2) 34 g (4) 74 g

53. What is the gram formula mass of Na2CO3 • 10H2O?(1) 106 g (3) 266 g(2) 142 g (4) 286 g

54. What is the volume of 1.50 moles of an ideal gas at STP?(1) 11.2 L (3) 33.6 L(2) 22.4 L (4) 44.8 L

55. The percent by mass of oxygen in Na2SO4 (formula mass = 142) is closest to(1) 11% (3) 45%(2) 22% (4) 64%

56. Given the equation:

Zn + 2 HCl → ZnCl2 + H2

How many moles of HCl would be required to produce a total of 2 moles of H2?(1) 0.5 (3) 3(2) 2 (4) 4

57. Given the reaction:

2 NaOH + H2SO4 → Na2SO4 + 2 H2O

What is the total number of moles of NaOH needed to react completely with 2 moles of H2SO4?(1) 1 (3) 0.5(2) 2 (4) 4

58. A sample of gas occupies 15.0 liters at a pressure of 2.00 atmospheres and a temperature of 300. K. If the pressure is lowered to 1.00 atmosphere and the temperature is raised to 400. K, the volume of the gas sample would be(1) 5.63 L (3) 22.5 L(2) 10.0 L (4) 40.0 L

59. If a student pours a mixture of sand and salt water through a filter paper into a beaker, what will be found in the beaker after filtering?(1) salt, only (3) salt and water(2) sand, only (4) salt and sand

60. Which piece of laboratory equipment is represented by the diagram below?

(1) crucible tongs (3) test tube clamp(2) beaker tongs (4) pinch clamp

61. The diagram below shows the upper part of a laboratory burner.

Which letter represents the hottest part of the burner flame?(1) A (3) C(2) B (4) D

62. Which measurement contains three significant figures?(1) 0.05 g (3) 0.056 g(2) 0.050 g (4) 0.0563 g

63. The mass of a solid is 3.60 grams and its volume is 1.8 cubic centimeters. What is the density of the solid, expressed to the correct number of significant figures?(1) 2 g/cm3 (3) 0.5 g/cm3

(2) 2.0 g/cm3 (4) 0.50 g/cm3

64. A student determined that the percent of H2O in a hydrate was 39.0%. The percent of H2O in this hydrate is 36.0% according to an accepted chemistry reference. What is the student's percent of error?(1) 9.1% (3) 3.0%(2) 8.3% (4) 11%

1

1 Short Answer Review

Short Answer Review Directions: Predict the products and balance each chemical equation

1. potassium carbonate + sodium

2. manganese (II) fluoride + lithium bromide

3. beryllium oxide + phosphorus

4. zinc iodide + ammonium chloride

5. lead (IV) bromide + silicon

6. barium hydroxide + hydrogen

7. magnesium chlorate + copper (I) sulfate

8. cobalt (II) cyanide + iron (II) chromate

9. sulfuric acid + sodium hydroxide

10. calcium hydroxide + chromic acid

2


Directions: Complete all questions below given the data below showing the solubility of salt x:

Temperature (°C)

Mass of Solute per 100 g of H2O (g)

10 22 25 40 30 48 60 107 70 135

a. On the graph provided below, scale and label the y – axis including appropriate units. [ b. Plot the data from the data table. Surround each point with a small circle and draw a best-fit curve for

the solubility of salt X.

c. Using your graph, predict the solubility of salt X at 50°C.

3


Directions: Answer all questions. Round all gram formula masses to the nearest tenth. Use significant figures and units. Show your work.

(a) How many moles are present in 284.2 grams of copper (II) sulfate _____________________

(b) How many grams would be present in 16.38 moles of barium nitrate? _____________________

(c) How many moles are present in 27.0 grams of hydrochloric acid? _____________________

4


(d) How many moles of NaCl would be produced if 0.73 moles of NaF reacts according to the following equation?

2 NaF + CuCl2 2 NaCl + CuF2

___________________

(e) How many moles of aluminum nitrate would be produced when 5.32 moles of aluminum bromide reacts with nitric acid?

___________________

(f) Carbon will react with zinc oxide to produce zinc and carbon dioxide. How many moles of carbon dioxide will be produced if 0.38 mol of ZnO is completely reacted?

___________________

(g) Phosphorus will react with bromine to produce phosphorus tribromide. How many moles of phosphorus will be consumed if 0.78 mol of bromine is reacted?

____________________

5


Directions: Answer each question. Show all work.

(1) What is the quantity heat that is released when 50. grams of water is cooled from 343 K to 333 K? ___________________

(2) What is the final temperature of 42.3 g of water after 12 kJ of heat is added to it and it’s starting

temperature is 4°C? ___________________

(3) When 10.0 grams of water at 293 K absorbs 420. joules of heat, the temperature of water increases by: ___________________

(4) How much heat energy is absorbed by the melting of 275 g of ice at 0°C? ___________________

6


Directions: Draw the appropriate particle diagrams based on the following elements. Element A Element B

(1) Draw a particle diagram that shows no less than 5 molecules of Element A in the solid phase

(2) Draw a particle diagram that shows no less than 5 molecules of Element B in the gaseous phase

(3) Draw a particle diagram that shows no less than 5 molecules of Compound AB in the liquid phase

7


Directions: Use Reference Table H to answer the following problems (1) For water, what is the difference in vapor pressure (in kPa) between 75°C and 100°C? (2) At what Kelvin temperature will ethanol boil at, if the atmospheric pressure is 150. kPa? (3) Which boils at the highest temperature: ethanol at 1.3 atm or water at 0.80 atm?

Directions: Solve each of the following problems as directed.

(1) List three types of intermolecular forces of attraction:

__________________________________________________________________________________ (2) Which one of these forces is present in F2 and Cl2? ___________________________ (3) Explain the effect that larger molecular size has on the strength of the intermolecular forces.

_________________________________________________________________________________

(4) Which molecule is larger, F2 or Cl2? _______________________________________ (5) Which substance has the stronger intermolecular force? _______________________

(6) What is the boiling point of F2 in Kelvin? __________________________________

(7) What is the boiling point of Cl2 in Kelvin? __________________________________

8


Directions: Solve each of the following problems as directed. Show all work.

(1) On the graph of a heating curve for water given below, label 0°C and 100°C on the temperature axis. On the likes provided, indicate the states of matter represented by A, C, and E. Then indicate what changes of state are occurring in sections B and D.

(2) Explain what is happening to the average kinetic energy of the molecules during the changes of state

marked by B and D on the graph.

(3) It takes 6 kJ to convert a mole of ice into water. It takes 41 kJ to convert a mole of water into steam. Calculate the total amount of energy absorbed through the two phase changes.

9


Directions: On the line at the left, write the term that matches each description below. Vaporization Melting Equilibrium vapor pressure Condensation Phase change Volatile Freezing point Sublimation Deposition ____________________________ 1. conversion of a solid directly into a gas

____________________________ 2. opposite of vaporization ____________________________

3. temperature at which the solid and liquid forms of a substance exist in equilibrium

____________________________

4. conversion of a substance from one of the three physical state if matter into another

____________________________ 5. change of state from liquid to gas

____________________________ 6. pressure exerted by a constant number of gas molecules above a liquid or a solid

____________________________ 7. description of a liquid that evaporates easily

____________________________ 8. transformation of gas directly into a solid ____________________________

9. phase change from solid to liquid

10


Directions: (1) Determine if the bonding is IONIC or COVALENT (2) If the bond type is IONIC, draw the transfer of electrons (3) If the bond type is COVALENT, draw the molecule (4) When appropriate, determine the shape of the molecule (5) When appropriate, determine whether the molecule is POLAR or NON-POLAR (6) When appropriate, determine the degrees of bond angles in the molecule (7) When appropriate, label and place the DIPOLE CHARGES

(1) BaF2 Draw the transfer of e- or the shape of the molecule below Electronegativity Difference __________ Bond Type ________________________ Bond Angle _______________________ Shape ____________________________ Polarity ___________________________ (2) H2S Draw the transfer of e- or the shape of the molecule below Electronegativity Difference __________ Bond Type ________________________ Bond Angle _______________________ Shape ____________________________ Polarity ___________________________ (3) CCl4 Draw the transfer of e- or the shape of the molecule below Electronegativity Difference __________ Bond Type ________________________ Bond Angle _______________________ Shape ____________________________ Polarity ___________________________

1. A student obtained the following data to determine the percent by mass of water in a hydrate.

What is the approximate percent by mass of the water in the hydrated salt?(1) 2.5% (3) 88%(2) 12% (4) 98%

2. Which piece of laboratory equipment should be used to remove a heated crucible from a ring stand?

(1) (3)

(2)

(4)

3. Which diagram represents a crucible?

(1) (3)

(2)

(4)

4. Which measurement has the greatest number of significant figures?(1) 6.060 mg (3) 606 mg(2) 60.6 mg (4) 60600 mg

5. An 8.24-gram sample of a hydrated salt is heated until it has a constant mass of 6.20 grams. What was the percent by mass of water contained in the original sample?(1) 14.1% (3) 32.9%(2) 24.8% (4) 75.2%

6. What is the sum of 0.0421 g + 5.263 g + 2.13 g to the correct number of significant digits?(1) 7 g (3) 7.44 g(2) 7.4 g (4) 7.435 g

7. The following weighings were made during a laboratory exercise:

Mass of evaporating dish.....59.260 gMass of sugar sample ..........1.61 g

What is the total mass of the evaporating dish plus the sample, expressed to the proper number of significant figures?(1) 60.870 g (3) 60.9 g(2) 60.87 g (4) 61 g

8. The solid object shown below has a mass of 162.2 grams.

What is the density of the object to the correct number of significant figures?(1) 0.22 g/cm3 (3) 4.5 g/cm3

(2) 0.2219 g/cm3 (4) 4.505 g/cm3

9. A student in a laboratory determined the boiling point of a substance to be 71.8°C. The accepted value for the boiling point of this substance is 70.2°C. What is the percent error of the student's measurement?(1) 1.60% (3) 2.23%(2) 2.28% (4) 160.%

10. In a sample of pure copper, all atoms haveatomic numbers which are(1) the same and the atoms have the same number of

electrons(2) the same but the atoms have a different number of

electrons(3) different but the atoms have the same number of

electrons(4) different and the atoms have a different number of

electrons

11. The diagram below shows a section of a 100-milliliter graduated cylinder.

When the meniscus is read to the correct number of significant figures, the volume of water in the cylinder would be recorded as(1) 75.7 ml (3) 84.3 ml(2) 75.70 ml (4) 84.30 ml

12. The atomic mass of an element is defined as the weighted average mass of that element's(1) most abundant isotope(2) least abundant isotope(3) naturally occurring isotopes(4) radioactive isotopes

13. What is a possible mass number of a sodium atom, Na?(1) 1 (3) 12(2) 11 (4) 23

14. What is the mass number of a 31H atom?(1) 1 (3) 3(2) 2 (4) 4

15. What is the atomic number of an element whose atoms each contain 47 protons, 60 neutrons, and 47 electrons?(1) 13 (3) 60(2) 47 (4) 107

16. Which ion has the same electron configuration as an Mg2+ ion?(1) Ca2+ (3) Na+

(2) Cl– (4) S2–

17. The nucleus of a fluorine atom (F) has a charge of(1) 1+ (3) 19+(2) 9+ (4) 0

18. An experiment in which alpha particles were used to bombard thin sheets of gold foil led to the conclusion that an atom is composed mostly of(1) empty space and has a small, negatively charged

nucleus(2) empty space and has a small, positively charged

nucleus(3) a large, dense, positively charged nucleus(4) a large, dense, negatively charged nucleus

19. Which pair must represent atoms of the same element?(1) 14

6 X and 147 X

(2) 126 X and 13

6 X(3) 2

1X and 42X(4) 13

6 X and 147 X

20. In which two atoms do both nuclides contain the same number of neutrons?(1) 20

10Ne and 4018Ar

(2) 6529Cu and 65

30Zn(3) 24

12Mg and 2612Mg

(4) 146 C and 16

8 O

21. Which particle has approximately the same mass as a proton?(1) alpha (3) electron(2) beta (4) neutron

22. Which principal energy level of an atom contains an electron with the lowest energy?(1) n = 1 (3) n = 3(2) n = 2 (4) n = 4

23. Which is the electron-dot symbol for a chlorine atom in the ground state?

(1) (3)

(2) (4)

24. Which is the correct electron-dot symbol for thefluoride ion?

(1) (3)

(2) (4)

25. Which atom has the strongest attraction for electrons?(1) Cl (3) Br(2) F (4) I

26. The bonds in all network solids are(1) covalent (3) metallic(2) ionic (4) nonpolar

27. Which element requires the least amount of energy to remove its most loosely bound electron?(1) Li (3) Ba(2) Mg (4) Ca

28. As a chemical bond forms between two hydrogen atoms the potential energy of the atoms(1) decreases (3) remains the same(2) increases

29. Which compound would most likely have the greatest ionic character?(1) CO (3) CaO(2) KF (4) LiH

30. Which compound would have the greatest degree of ionic character?(1) Na2O(2) H2O(3) CO2(4) NO2

31. Which element would most likely form an ionic bond with chlorine?(1) O (3) S(2) N (4) K

32. The bond between which two elements is the least ionic in character?(1) H-F (3) H-I(2) H-Cl (4) H-O

33. Which molecule contains a nonpolar covalent bond?

(1) (3)

(2) (4)

34. In a nonpolar covalent bond, electrons are(1) shared equally by two atoms(2) shared unequally by two atoms(3) transferred from one atom to another(4) located in a mobile "sea" shared by many atoms

35. Which formula represents a substance that contains covalent bonds?(1) LiCl(2) CaCl2(3) K2O(4) CO2

36. Which molecule contains a polar covalent bond?

(1) (3)

(2) (4)

37. Which could form a coordinate covalent bond?

(1) (3)

(2) (4)

38. Which element consists of positive ions immersed in a "sea" of mobile electrons?(1) sulfur (3) calcium(2) nitrogen (4) chlorine

39. Which compound has the lowest melting point?(1) HCl (3) NaCl(2) KCl (4) LiCl

40. Which two compounds contain only polar bonds?(1) CCl4 and CH4 (2) HCl and Cl2(3) HCl and NH3(4) CO and O2

41. The abnormally high boiling point of HF as compared to HCl is primarily due to intermolecular forces of attraction called(1) network bonds (3) van der Waals forces(2) electrovalent forces (4) hydrogen bonds

42. Which of the following liquids has the weakest van der Waal's forces of attraction between its molecules?(1) Xe(…) (3) Ne(…)(2) Kr(…) (4) He(…)

43. The attraction that nonpolar molecules have for each other is primarily caused by the presence of(1) hydrogen bonding(2) high ionization energy(3) electronegativity differences(4) van der Waals forces

44. Which formula represents a nonpolar compound?(1)

(2)

(3)

(4)

45. In which system are molecule-ion attractions present?(1) NaCl(s) (3) NaCl(…)(2) NaCl(g) (4) NaCl(aq)

46. Which is an empirical formula?(1) N2O4(2) P4O10(3) C6H12O6(4) Al2O3

47. What is the correct name of Fe2O3?(1) iron (I) oxide (3) iron (III) oxide(2) iron (II) oxide (4) iron (V) oxide

48. Which formula represents lead (II) phosphate?(1) PbPO4(2) Pb4PO4(3) Pb3(PO4)2(4) Pb2(PO4)3


__C3H8(g) + __O2(g) → __H2O(g) + __CO2(g)

When the equation is completely balanced using smallest whole numbers, the coefficient of O2 is(1) 5 (3) 3(2) 2 (4) 10

50. More than two-thirds of the elements of the Periodic Table are classified as(1) metalloids (3) nonmetals(2) metals (4) noble gases


__SiO2 + __C → __SiC + __CO

is correctly balanced using whole-number coefficients, the sum of all the coefficients is(1) 6 (3) 8(2) 7 (4) 9


__ C2H4 + __ O2 → __ CO2 + __ H2O

is balanced using smallest whole numbers, what is the coefficient of the O2?(1) 1 (3) 3(2) 2 (4) 4

53. The highest ionization energies in any period are found in Group(1) 1 (3) 17(2) 2 (4) 18

54. Which element in Period 2 has the greatest tendency to gain electrons?(1) Li (3) F(2) C (4) Ne

55. Which element has the highest ionization energy?(1) barium (3) calcium(2) magnesium (4) strontium

56. Which of the following Period 3 elements has the least metallic character?(1) Na (3) Al(2) Mg (4) Si

57. At STP, which of the following elements has the most metallic character?(1) iodine (3) chlorine(2) bromine (4) fluorine

58. Which element exists as a diatomic molecule at STP?(1) bromine (3) sulfur(2) argon (4) rubidium

59. In which area of the Periodic Table are the elements with the strongest nonmetallic properties located?(1) lower left (3) lower right(2) upper left (4) upper right

60. Which represents the electron configuration of a metalloid in the ground state?(1) 2-3 (3) 2-8-5(2) 2-5 (4) 2-8-6

61. What is the mass in grams of 2.0 moles of NO2?(1) 92 (3) 46(2) 60. (4) 30.

62. Which sequence of atomic numbers represents elements which have similar chemical properties?(1) 19, 23, 30, 36 (3) 3, 12, 21, 40(2) 9, 16, 33, 50 (4) 4, 20, 38, 88

63. A white anhydrous powder that dissolves in water to form a blue aqueous solution could be(1) MgSO4(2) BaSO4(3) CuSO4(4) CaSO4

64. Which salt forms a colored aqueous solution?(1) Mg(NO3)2(2) NaNO3(3) Ca(NO3)2(4) Ni(NO3)2

65. Which three elements have the most similar chemical properties?(1) Ar, Kr, Br (3) B, C, N(2) K, Rb, Cs (4) O, N, Si

66. Which substance has the greatest molecular mass?(1) H2O2(2) NO(3) CF4(4) I2

67. The percent by mass of nitrogen in NH4NO3 (formula mass = 80) is approximately(1) 18% (3) 32%(2) 23% (4) 35%

68. The temperature of 100 grams of water changes from 16ºC to 20ºC. What is the total number of calories of heat energy absorbed by the water?(1) 25 (3) 100(2) 40 (4) 400

69. A compound contains 40% calcium, 12% carbon, and 48% oxygen by mass. What is the empirical formula of this compound?(1) CaCO3(2) CaC2O4(3) CaC3O6(4) CaCO2

70. Vitamin C has an empirical formula of C3H4O3 and a molecular mass of 176. What is the molecular formula of vitamin C?(1) C3H4O3(2) C6H8O6(3) C9H12O9(4) C10H8O3


2 NaOH + H2SO4 → Na2SO4 + 2 H2O

What is the total number of moles of NaOH needed to react completely with 2 moles of H2SO4?(1) 1 (3) 0.5(2) 2 (4) 4


CH4 + 2 O2 → CO2 + 2 H2O

What amount of oxygen is needed to completely react with 1 mole of CH4?(1) 2 moles (3) 2 grams(2) 2 atoms (4) 2 molecules


__N2(g) + __H2(g) →_NH3(g)

When the equation is balanced using smallest whole-number coefficients, the ratio of moles of hydrogen consumed to moles of ammonia produced is(1) 1:3 (3) 2:3(2) 3:1 (4) 3:2

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UNIT ONE: Introduction to Chemistry and · PDF fileRegents Chemistry Mid-Term Exam Study Guide Revised: January 2009 UNIT ONE: Introduction to Chemistry and Measurements Content Outline