Atomic Structure I: Nuclear Chem



2013-2014 Review Packet for Regents Chemistry Midterm

Date: Wednesday, January 24, 2104

PRACTICE: ATOMIC STRUCTURE

1. What is the definition of each of the following terms:

a. Atomic number number of protons

b. Mass number number of protons plus number of neutrons

c. Isotope same element with different masses (due to different number of neutrons)

d. Nucleon collective term for protons and neutrons

2. What is the charge, mass, location, and symbol of each of the following subatomic particles?

|Particle |Charge |Mass |Location |symbol |

|proton |+1 |1 amu |in nucleus |11p |

|neutron |none |1 amu |in nucleus |10n |

|electron |-1 |0 amu |outside nucleus |0-1e |

3. The number of electrons in a neutral atom is equal to the number of protons

4. What is the charge on the nucleus of an atom? positive

5. Most of the volume of an atom is empty space

6. Describe Rutherford’s famous experiment. What were the two results of the experiment and what conclusion did Rutherford draw from each result?

1. atom consists of mostly empty space

2. the center of the atom is extremely dense and is positively charged

7. What is the difference between an isotope and an ion?

Isotope: neutral atoms of same element with different masses

Ion: atoms that have gained/lost electrons and become charged

8. Given that chlorine is 75% Cl-35 and 25% Cl-37, calculate the average atomic mass of chlorine.

(35)(0.75) + (37)(0.25) = average atomic mass of chlorine

9. Complete the following table:

|Atom |# of protons |# of neutrons |# of electrons |Atomic # |Mass # |

|126C |6 |6 |6 |6 |12 |

|73Li |3 |4 |3 |3 |7 |

|2311Na |11 |12 |11 |11 |23 |

|31H |1 |2 |1 |1 |3 |

|42He |2 |2 |2 |2 |4 |

10. What radioactive isotope is used to date organic material (wool, linen, wood, bones, etc)?

carbon - 14

11. What is the absolute temperature scale?

Kelvin

12. In the Bohr model of the atom, describe the type of electron transition which will result in the production of a spectral line. Is this transition endothermic or exothermic?

An electron transition in which an electron drops from a higher energy level to a lower energy level will result in the production of a specific wavelength of color (a spectral line). This is an exothermic process since energy is released by the electron during the above described transition.

13. How many atoms are represented in the formula (NH4)2SO4?

2 N; 8 H; 1 S; 4 O

14. What is sublimation? What are two common substances that sublime?

Substances that change from a solid directly to a liquid are said to undergo sublimation. CO2(s) and are two substances that sublime.

PRACTICE: NUCLEAR CHEMISTRY

1. We use two laws to balance nuclear equations. These two laws are the conservation of atomic number and the conservation of atomic mass number.

2. In the symbol 31H what does the 3 represent? atomic mass number

3. In the symbol 126C what does the 6 represent? atomic number

4. What nuclide is a radioisotope used to date rocks? U-238

Identify each of the following as natural transmutation (NT) or artificial transmutation (AT); if AT identify if can be fission or fusion. Explain how to identify which type of reaction it is.

| |Nuclear reaction |type |explanation |

|5 |23892U → 23490Th + 42He |NT |one reactant |

|6 |21H + 31H → 42He + 10n + energy |AT - fusion |two reactants; fusion because 2 small atoms unite|

| | | |to form one slightly larger one |

|7 |23994Pu + 10n → 9038Sr + 14756Ba + 3 10n |AT - fission |two reactants; fission because one very large |

| | | |atom is split into 2 smaller atoms |

|8 |2311Na + 10n → 2411Na |AT |two reactants |

|9 |147N + 42He → 11H + 178O |AT |two reactants |

|10 |23592U + 10n → 9236Kr + 14156Ba + 3 10n + energy |AT - fission |two reactants; fission because one very large atom|

| | | |is split into 2 smaller atoms |

|11 |21283Bi → 21284Po + 0-1e |NT |one reactant |

12. Find the missing term in the following nuclear equations:

a. 209F → 0-1e + X X = 2010Ne

b. 18079Au → 17677Ir + X X = 42He

c. X → 0-1e + 3719K X = 3718Ar

d. 94Be + X →126C + 10n X = 42He

13. For fusion reactions to occur, high temperatures are required because both of the reacting nuclei have

a. a large mass c. a positive charge

b. many neutrons d. a negative charge

14. Name one beneficial use and one destructive use of nuclear fission reactions?

Beneficial: large amounts of energy created from small amount of reactants

Destructive: nuclear waste is created

15. Explain why it is more difficult to cause an artificial transmutation reaction with an alpha

particle than with a neutron. Alpha particles are positively charged, therefore have to be

accelerated to great speeds to overcome repulsion forces for collisions to occur

16.Where in the universe do fusion reactions take place? What conditions have to be present

(in terms of pressure and temperature)? Place: sun; need high temp and high pressure

17. Where do fission reactions take place in a controlled fashion? In an uncontrolled one?

controlled fission reactions: nuclear power plant; uncontrolled fission reactions: nuclear bomb

18. What is the total number of neutrons in an atom of Pb-207? 125 [207-82]

19. What is the nuclear charge of fluorine? +9

20. What happens to the half-life of a radioactive substance as the temperature of the

substance increases? Half-life remains the same

21.Based on Reference Table N, what fraction of a sample of gold-198 remains radioactive

after 2.69 days?

HL Au-198 is 2.69 days after 1 HL: ½ of the original sample is left

22. How many days are required for 200.0 grams of radon-222 to decay to 50.0 grams?

HL Rn-222 is 3.82 days 200.0g → 100.0g → 50.0g 2HL = (3.82 days)(2) = 7.64 days

23.How much time must elapse before 16 grams of potassium-42 decays, leaving 2 grams of

the original isotope?

HL K-42 is 12.4h 16g → 8g → 4g → 2g (12.4)(3HL) = 37.2h

24. The half-life of iodine-131 is 8.07 days. What fraction of a sample of I-131 remains

after 24.21 days?

24.21 days = 3HL ½ ( ¼ ( ⅛

8.07 days

25. What was the original mass of a radioactive sample that decayed to 25 grams in 4 half-lives?

1 2 3 4

400g ( 200g ( 100g ( 50g ( 25g 400g was the original mass

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PRACTICE: MATTER

Using the symbol to represent atoms of elements A and the symbol to represent atoms of element B, draw the following particle diagrams.

PRACTICE: Energy Changes: Identify the Q equation needed to solve each of the following energy word problems below, then solve each problem.

1. What are the three equations for calculating an energy change? Describe when to use each equation.

Q = mcΔT change in temperature (no phase change) ΔT = Tf - Ti

Q = mHf change in phase (constant temperature) between solid and liquid (melting/freezing)

Q = mHv change in phase (constant temperature) between liquid and gas (boiling/condensation)

2. What is the Law of Conservation of Energy?

Energy is neither lost or gained

3. What is the Law of Conservation of Matter?

Matter (mass/atoms) is neither lost nor gained

4. Given the reaction: 2X + 3Y ( Q + 4Z

If 5g of X react completely with 30g of Y to produce 4g of Q, how much Z will be produced?

5g + 30g = 4g + Z Z = 31g

5. How much energy is required to heat a 35.0 g sample of water from 35ºC (Ti) to 75ºC (Tf)?

Equation: Q = mcΔT Q = (35.0g)(4.18J/g°C)( 40°C) = 5852J

6. An ice cube at 0˚C with a mass of 175.0g melts. How much energy does the ice absorb to melt?

Equation: Q = mHf Q = (175.0g)(334J/g) = 58,50J

7. An 85.0g sample of water at 100.0˚C boils. How much energy is required to convert the water to steam at 100.0˚C?

Equation: Q = mHv Q = (85.0g)(2260J/g) = 192,100J

8. A sample of water with a mass of 160.0g is heated from 15.0˚C (Ti) to 70.0˚C (Tf). How much energy is required to heat the water?

Equation: Q = mcΔT Q = (160.0g)(4.18J/g°C)(55°C) = 36,784J

9. 450.0g of water at 0.0˚C is frozen to ice at 0.0˚C. How much energy is released to the environment?

Equation: Q = mHf Q = (450.0g)(334J/g) = 150,300J

10. 200.0g of steam at 100.0˚C condenses to water at 100.0˚C. How much energy is released to the environment?

Equation: Q = mHv Q = (200.0g)(2260J/g) = 452,000J

11. A 225g sample of water is cooled from 95.0˚C (Ti) to 40.0˚C (Tf). How much energy is released to the environment?

Equation: Q = mcΔT Q = (225g)(4.18J/g°C)(-55 °C)

12. A block of ice at 0.0˚C with a mass of 1.0kg melts. How much energy is absorbed from the environment to convert it to water at 0.0˚C?

Equation: Q = mHf Q = (1000g)(334J/g) = 334,000J

13. A 200.0g sample of water at 100.0˚C is boiled and converted to steam at 100.0˚C. How much energy is required?

Equation: Q = mHv Q = (200.0g)(2260J/g) = 452,000J

14. A 345.0g sample of steam at 100.0˚C condenses to water at 100.0˚C. How much energy is released to the environment?

Equation: Q = mHv Q = (345.0g)(2260J/g) = 779,700J

15. If you carry out a reaction in a Styrofoam cup calorimeter and

a. the temperature of the water increases, is the reaction endothermic or exothermic?

b. the temperature of the water decreases, is the reaction endothermic or exothermic?

c. How would you calculate the energy change? Q = mcΔT

PRACTICE: Phase Changes

|boiling endothermic heat of fusion phase change condensation temperature evaporation |

|heat of vaporization solid constant exothermic increasing sublimation |

|decreasing fusion liquid deposition gas melting |

|vaporization |

16. temperature is a measure of the average kinetic energy of the particles of a substance.

17. vaporization is another word for boiling; fusion is another word for melting.

18. Heat of fusion is the amount of energy required to melt one gram of a solid at its melting point.

19. One type of physical change is a phase change.

20. Solid to liquid or gas is an example of an endothermic process.

21. evaporation is the spontaneous change from liquid to gas at any temperature.

22. sublimation is the change from solid phase to gas phase.

23. The average kinetic energy of the particles of a substance is increasing when the temperature is increasing.

24. gas takes the shape and volume of its container. solid has a definite shape and a definite volume.

25. boiling is the change from liquid phase to gas phase at a constant temperature.

26. condensation is a change from gas to liquid phase which is noticeable on glass.

27. melting is the change from solid to liquid phase at a constant temperature.

28. Heat of vaporization is the amount of energy required to convert one gram of a liquid to the gas phase at its boiling point.

29. deposition is the change from gas phase to the solid phase.

30. Gas to liquid or solid is an example of an exothermic process process.

31. When the temp. is constant, the average kinetic energy of the particles in a substance is constant.

32. liquid has definite volume and indefinite shape.

33. When temperature decreases, the average kinetic energy of the particles of a substance is decreasing.

PRACTICE: Gas Laws

1. What are the 4 assumptions of the KMT of gases?

1. gas particles move in random, straight-line motion

2. gas particles undergo “elastic” collisions

3. volume of gas particles is considered negligible

4. gas particles do not attract to repel each other

2. What is an ideal gas?

gas that behaves according to the KMT all the time

3. Under which two conditions do real gases behave ideally?

high temperature and low pressure

4. What is Dalton’s Law of Partial Pressures?

The total pressure of all the gases in a system is equal to the sum of all the partial pressures of the individual gases in the system

5. A 150 ml container that is filled with a mixture of nitrogen, oxygen, and CO2 has a total pressure of 175 kPa. The partial pressure of oxygen is 75 kPa. The partial pressure of nitrogen is 75 kPa. What is the partial pressure of CO2?

Ptotal = P1 + P2 + P3… 175kPa = 75kPa + 75kPa + PCO2

25kPa = PCO2

6. What is Avogadro’s hypothesis?

One mole of any gas occupies 22.4 liters of space

7. What is the combined gas law equation?

P1V1 = P2V2

T1 T2

8. What happens to the pressure if a given mass of gas at 50.6 kPa and 546 K is changed to STP?

P1V1 = P2V2 (50.6kPa) = P2 P2 = 25.3kPa

T1 T2 546K 273K

9. What volume will a 200 ml sample of a gas at STP occupy when the pressure is doubled at constant temperature?

P1V1 = P2V2 (1atm)(200ml) = (2atm)(V2) V2 = 100ml

T1 T2

10. A gas occupies a volume of 444 ml at 273 K and 79 kPa. What is the final Kelvin temperature when the volume of the gas is changed to 1880 ml and the pressure is changed to 38.7 kPa.

P1V1 = P2V2 (79kPa)(444ml) = (38.7kPa)(1880ml) X = 566.27K

T1 T2 273K X

11. At STP, 1 liter of O2 would have the same number of molecules as

a. 1 liter of H2 c. 3 liter of CO2

b. 2 liters of CO d. 0.5 liter of Ne

12. What is the formula to calculate percent error?

% error = (actual value – experimental value( x 100

actual value

13. What temperature scale must you use in gas law word problems? WHY?

Kelvin; there is no zero value to complicate the math calculations in word problems

PRACTICE: Periodic Table

1. What is nuclear charge? What is effective nuclear charge?

nuclear charge is the charge of the particles in the nucleus (equal to the # of protons; + sign)

effective nuclear charge is the charge felt by the valence electrons in an atom

• To calculate: subtract the # of kernel electrons from the # of protons

2. What is ionization energy? WHERE CAN THESE VALUES BE FOUND?

amount of energy needed to remove the most loosely held valence electron in an atom;

values are found in Table S

3. What is electro negativity?

ability of an atom to hold onto its own electrons and also attract its neighbor’s electrons (relates to strength of PPP)

4. What are the trends for atomic size, ionization energy, & electro negativity in the Periodic Table?

• Atomic size:

o ( as go down a column from top to bottom and ( as go left to right across a row

• Ionization energy:

o ( as go down a column from top to bottom and ( as go left to right across a row

• electronegativity:

o ( from bottom left corner (Fr) to upper right corner

PRACTICE: IMF

1. What is vapor pressure?

pressure exerted by a vapor above a liquid; occurs when molecules at surface of a liquid vaporize and the resulting gas molecules then exert pressure on the liquid

2. How is vapor pressure related to boiling point?

normal boiling of liquid is temperature at which vapor pressure equals standard atmospheric pressure

3. If a substance has a high vapor pressure and a low boiling point what can you say about the IMF?

IMF are weak

4. What variable does vapor pressure depend on?

temperature

5. What are the three IMF? Which force is weakest? Which force is strongest?

dispersion forces (weakest), dipole-dipole forces, hydrogen bonds (strongest)

6. Describe the kind of molecule that exhibits each kind of intermolecular force (ex: polar, non-polar, etc)

• dispersion forces: monatomic atoms (noble gases), diatomic elements, non-polar molecules (symmetrically shaped), pure hydrocarbons (compounds containing only C and H)

• dipole-dipole forces: polar molecules (non-symmetrical)

• hydrogen bonds: H in one molecule interacting with F, O or N in neighboring molecule

7. Why does water have a high boiling point? IMF are strongest: H-bonds

8. What kind of IMF exists between each of the following molecules?

a. CH4 dispersion

b. HCl dipole-dipole

c. CH3OH H-bonds

d. CO2 dispersion

e. Xe dispersion

f. NH3 H-bonds

g. He dispersion

h. N2 dispersion

i. HF H-bonds

j. H2 dispersion

k. C5H12 dispersion

l. F2 dispersion

m. HBr dipole-dipole

9. If a substance has strong IMF what phase is it likely to be in at room temperature? solid/liquid

10. If a substance has weak IMF what phase is it likely to be in at room temperature? gas

PRACTICE: Mole Math

1. Calculate the formula mass of the following:

a. NH3 1(14) + 3(1) = 17g/mole d. CH4 1(12) + 4(1) = 16g/mole

b. NaCl 1(23) + 1(35) = 58g/mole e. C2H4 2(12) + 4(1) = 28g/mole

c. CO 1(12) + 1(16) = 28g/mole

2. Perform the following conversions:

a. 2 moles NH3 to grams 34g f. 59.5g NH3 to moles 3.5moles

b. 0.4 moles NaCl to grams 23.2g g. 3 moles CO to liters 67.2 L

c. 0.8 moles CO to grams 22.4moles h. 4.5 moles C2H4 to liters 100.8 L

d. 234.0g NaCl to moles 4.03moles i. 11.2 liters CO to moles 0.5moles

e. 32.0g CH4 to moles 2moles j. 89.6 liters CH4 to moles 4moles

3. What is the definition of percent?

part x 100 = %

whole

4. Calculate the percent composition of the following:

a. N and H in NH3 FM = 17g/mole

N 14 x 100 = 82.35% H 3 x 100 = 17.65%

17 17

b. C and H in CH4 FM = 16g/mole

C 12 x 100 = 75% H 4 x 100 = 25%

16 16

c. Na and Cl in NaCl FM = 58g/mole

Na 23 x 100 = 39.66% Cl 35 x 100 = 60.34%

58 58

5. Calculate the percent by mass of water in the following:

d. BeSO4•2H2O FM = 141g/mole 36 x 100 = 60.34%

141

e. CaO2•8H2O FM = 216g/mole 144 x 100 = 66.7%

216

5. What is the empirical formula of a substance with the following molecular formulas:

a. P4O10 P2O5

b. N2O4 NO2

c. H2C2O4 HCO2

6. What is the molecular formula of a substance with the empirical formula C4H5N2O and a molecular mass of 194.0g/mol. FM = 4(12) + 5(1) + 2(14) + 1(16) = 97g/mole

194.0 = 2 2(C4H5N2O) is C8H10N4O2

97

7. What is the molecular formula of a substance with the empirical formula CH2 and a molecular mass of 84.0g/mol. FM = 1(12) + 2(1) = 14g/mole

84 = 6 6(CH2) is C6H12

14

PRACTICE: Significant Figures

How many sig figs are in each of the following measurements?

1. 6.25x10-6 g 3 6. 86270 mm 4

2. 5.505x102 cm 4 7. 86270.0 mm 6

3. 0.0081 L 2 8. 140 m/hr 2

4. 87.0˚C 3 9. 729 s 3

5. 0.00750 dm3 3 10. 13.90 ml 4

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e) mixture of compounds

a) pure substance – diatomic element

h) a gas

g) a liquid

f) a solid

d) mixture of elements

c) mixture of elements and compounds

b) pure substance – compound

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