Name_____________________



AP Chemistry 1: Structure of Matter Name __________________________

A. Measurement (1.4 to 1.6)

1. science knowledge is advanced by observing patterns (laws) and constructing explanations (theories), which are supported by repeatable experimental evidence

a. theory lasts until disproven

b. theory is never 100 % certain

2. uncertainty in measurements

a. precision and accuracy

1. precise = consistent (even if incorrect)

2. accurate = correct (even if inconsistent)

[pic]

precise precise & accurate accurate

b. data analysis

1. accuracy is measured by percent difference

percent Δ = 100|mean – true|/true

2. precision is measured by percent deviation

% Δ = 100Σ|trial – mean|/N(mean)

(N is number of trials)

• absolute Δ = |trail – mean|

• average Δ = Σ(absoluteΔ)/N

• % Δ = 100(averageΔ)/(mean)

c. significant figures (sf) indicate level of certainty

[pic]

measurement includes all certain (numbered) plus one estimated value ∴ 7.5 cm (2 sf)

d. rules for counting significant figures

1. all nonzero digits are significant

2. zero is sometimes significant, sometimes not

a. example: 0.00053000021000

never always ?

b. (?) decimal vs. no decimal

1. significant with decimal: 120. (3 sf)

2. not significant w/o decimal: 120 (2 sf)

3. exact numbers (metric conversions, counting or written numbers) have infinite number of sf

4. scientific notation: C x 10n

a. C contains only significant figures

b. 1200 with 3 sf: 1.20 x 103

e. rules for rounding off calculations

1. limited by least accurate measurement

2. x, ÷: answer has the same number of sf as the measurement with the fewest

3. +, –: answer has same end decimal position as measurement with left most end position

3. SI measuring system

a. summary chart

|Measurement |SI standards |Chemistry |

|mass |kilogram (kg) |gram (g) |

|volume |cubic meter (m3) |liter (L) |

|temperature |kelvin (K) |Celcius (oC) |

|time |second (s) |varies |

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b. prefixes system (x 10X)

1. k3, c-2, m-3, µ-6, n-9

2. squared/cubed prefix: 1 cm2 = 1 x (10-2)2 m2

3. 1 mL = 1 cm3

4. 455 kg x 103 g x (10-2)3 m3 = 0.455 g

m3 1 kg 1 cm3 cm3

4. mass and volume measure amount of matter

a. density: d = m/V

1. units depend on units for m and V

2. dH2O = 1.00 g/mL = 1.00 g/cm3 = 1000 kg/m3

b. number of particles: mole = 6.022 x 1023 particles

1. periodic table mass equals formula mass in g

2. molar mass (MM)—sum of mass of atoms in chemical formula (use 3 significant figures)

a. Al: 27.0 g/mol

b. H2O: 18.0 g/mol

c. conversions (dimensional analysis)

1. mass Δ moles (given formula or MM)

__ g x 1 mole/(MM) g = __ mole

2. volume Δ mass (given density–d)

__ mL x (d) g/1 mL = __ g

3. volume Δ mass Δ moles (given d and MM)

__ mL x (d) g/1 mL x 1 mole/(MM) g = __ mole

B. Atomic Nature of Matter (2.1 to 2.7)

1. historical perspective

a. Dalton's atomic theory (1805)

1. unique, indestructible atoms for each element

2. atoms are rearranging, not created during chemical change

3. compounds are groups of atoms in fixed ratio

b. subatomic structure

1. J. J. Thomson (1897): measure charge-to-mass ratio of electrons with cathode rays

2. Millikan (1909): measure electron charge with oil drops in a vacuum chamber

3. Rutherford (1910): characterized dense, + nucleus with alpha (α) radiation and gold foil

2. components of the atom

a. subatomic particles

|Particle |Location |Charge |Mass |Symbol |

|Proton |nucleus |+ 1 |1.0 |11p or 11H |

|Neutron |nucleus |0 |1.0 |10n |

|Electron |outside |- 1 |.00055 |o-1e |

b. atomic number (Z)

1. number of protons

2. defines type of atom

c. mass number (A)

1. protons + neutrons

2. isotopes (same Z, different A)

3. nuclear symbol: AZX

d. ions are atoms where # electrons ≠ # protons

1. e > p: (–) charged (anion): XN-

2. e < p: (+) charged (cations): XN+

e. unified atomic mass unit (u)

1. 1 u = 1/12 the mass of a C-12 atom

2. average atomic mass (periodic table mass)

a. isotopes have fixed % in natural sample

b. 100mav = %1m1 + %2m2 + ...

3. forms of matter

a. pure substance has a unique composition of atoms ∴ unique formula and set of properties

1. elements—one type of atom

(diatomic: H2, N2, O2, F2, Cl2, Br2, I2)

2. compounds—two or more types of atoms

a. molecular—formula defines size

b. crystalline—formula shows ratio of atoms

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b. mixture of pure substances in an object or container

1. variable composition (no set formula)

2. uniform: homogenous mixture = solution

3. non-uniform: heterogeneous

c. summary

[pic]

Monatomic Molecular Molecular Homogeneous

Element Element Compound Mixture

[pic]

Crystalline Compound

C. Radioactivity (21.1 to 21.4)

1. forms of natural radiation

|Type |Symbol |Mass # (A) |Charge # (Z) |Stopping |

| | | | |Shield |

|alpha |α |42He |4 |+2 |paper |

|beta |β− |0-1e |0 |-1 |Al |

|positron |β+ |01e |0 |+1 |destroyed |

|gamma |γ |00γ |0 |0 |Pb |

2. balancing nuclear reactions using nuclear symbols: AZX

• balance A and Z values

• determine symbol by Z number

• 23892U → 42He + 23490Th

3. nuclear instability

a. isotopes that are outside the "belt of stability" tend to be radioactive

b. modes of decay

1. atomic number > 83—α (alpha)

22688Ra → 22286Rn + 42He

2. Aisotope > Aaverage: 10n → 11p + 0-1β (beta)

146C → 147N + 0-1e

3. Aisotope < Aaverage: 11p → 10n + 01β (positron)

116C → 115B + 01e

alpha decay

beta decay positron decay

4. transmutations

a. induced nuclear reactions by bombardment

b. 147N + 42He → 178O + 11H

c. produce trans-uranium elements

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5. radioactive decay

a. rate of decay ∝ number of radioactive atoms (Nt)

1. rate = kNt (k: rate constant)

2. time for half of remaining atoms to decay (t½) is constant: k = (ln2)/t½

[pic]

b. ln(No/Nt) = kt or Nt = Noe-kt

1. No = original amount

2. t and k must have same time units

D. Electron Structure—Bohr Model(6.3 to 6.4)

1. atomic spectrum

a. colors emitted by energized atoms (unique for each element)

[pic]

b. calculations: Ephoton = 2.00 x 10-25 J•m/λ

1. λ = wavelength (m)

2. f = frequency (s-1) = c/λ (ν on AP test)

3. Ephoton = hf = hc/λ

(hc = (6.63 x 10-34 J•s)(3.0 x 108 m/s) = 2.00 x 10-25 J•m)

2. Bohr model—atoms with one electron only

a. energy levels (n)

1. Eelectron = -B/n2

2. for H: En = -2.18 x 10-18 J/n2

3. ground state (n = 1) electron has lowest (most negative) energy

4. excited state (n > 1), electron energy increases until ionized (E = 0 J)

5. ΔEelectron = En-final – En-initial

a. ΔEelectron > 0 when increasing n

b. ΔEelectron < 0 when decreasing n

b. |ΔEelectron| = Ephoton

[pic]

434 nm 486 nm 656 nm

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Measurement

1. How many significant figures are there in?

|0.008090 mL |1300.40 atm |13400 m |one liter |

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2. Express the answers to the correct number of sf.

|(3.016)(4.23) | |12.0 + 1.01 + 6 | |

|0.0031 | |101.4 | |

3. How much do you have when you double 12.28 g?

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4. A student measures the mass of an object to be 12.045 g. The true mass is 12.000 g. What is the percent error?

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5. Determine the % deviation for the following massings.

|Mass |48.307 g |49.886 g |50.911 g |49.524 g |

|mean | |

|Δ | | | | |

|Average | |

|Δ | |

|% Δ | |

6. Convert the following:

|345 nm → m | |

|3640 cm2 → m2 | |

|350 mL → L | |

|155 cm3 → L | |

7. A student adds 7.76 g of pellets to a graduated cylinder containing 5.00 mL. The total volume of the pellets and water is 7.87 mL. What is the density of the pellets?

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8. A student measure the mass of an empty graduated cylinder (10.076 g), then fills it with 10.0 mL of liquid. The total mass of cylinder and liquid is 18.799 g. What is the density?

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9. Calculate the molar masses to 3 significant figures.

|NaCl |H2O |Cl2 |

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|(NH4)2SO4 |C4H7NO4 |CuSO4•5 H2O |

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10. Use dimensional analysis to determine the following for aluminum (MM = 27.0 g/mol, d = 2.70 g/cm3)

|2.48 g Al → mol | |

|5.00 cm3 Al → g | |

|155 cm3 Al → mol | |

11. Use dimensional analysis to determine the following for carbon dioxide (MM = 44.0 g/mol, d = 1.82 g/L)

|85.0 g CO2 → L | |

|3.15 mol CO2 → g | |

|3.22 L CO2 → mol | |

Density Lab

12. Measure the mass and volume of a solid, liquid and gas, determine densities, and use the density to identify the substance.

Solid: Add 5.0 mL (V1) water to a 10 mL graduated cylinder. Mass the cylinder + water (m1). Add solid. Record the volume to the nearest 0.1 mL (V2). Mass (m2).

a. Record the collected data. Calculate the change in mass (Δm) and change in volume (ΔV) and density (d).

| m2 – m1 = | V2 – V1 = ΔV |d |

|Δm | | |

| | | | |5.0 mL | | |

b. Check the identity of the solid based on its density.

|Al |Zn |Pb |

|d = 2.7 g/mL |d = 7.1 g/mL |d = 11 g/mL |

c. Calculate the number of moles of solid.

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Liquid: Mass a clean, dry 10 mL graduated cylinder (m1). Add 10.0 mL liquid to the cylinder. Mass the cylinder + liquid (m2).

d. Record the collected data. Calculate the change in mass (Δm), change in volume (ΔV) and density (d).

| m2 – m1 = Δm |V |d |

| | | |10. mL | |

e. Circle the identity of the liquid based on its density.

|C2H6O |H2O |C3H8O3 |

|d = 0.79 g/mL |d = 1.0 g/mL |d = 1.1 g/mL |

f. Calculate the number of moles of liquid.

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Gas: Add ½ scoop of baking soda (NaHCO3) to the flask. Mass the flask (m1). ½ fill the pipet-tube assembly with 6 M HCl (handle with care). Mass the assembly (m2). Fill the gas collecting bottle with water. Measure the volume of water (V1). Stopper the flask with the assembly and insert the open tube into the gas collecting bottle. Add the HCl to the flask, one drop at a time, until the gas collecting bottle is nearly full. Disconnect the stopper. Mass the flask (m3) and assembly (m4). Measure the volume of water remaining (V2).

g. Record the collected data. Calculate the change in mass (Δm) and change in volume (ΔV) and density (d).

| m1 + m2 – (m3 + m4) = Δm | V1 – V2 = ΔV|d |

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|d = 8.4 x 10-5 g/mL |d = 1.3 x 10-3 g/mL |d = 2 x 10-3 g/mL |

i. Calculate the number of moles of gas.

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Atomic Nature of Matter

13. How did Rutherford show that atoms have a nucleus?

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14. Below is a modern view of an isotope of a sulfur atom.

16 protons

16 neutrons

18 electrons

Write the nuclear symbol for this ionized isotope.

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15. Complete the chart below.

|Symbol |protons |neutrons |electrons |

|2713Al | | | |

|4019K+ | | | |

|3115P3- | | | |

| |26 |30 |23 |

| |17 |17 |18 |

| |79 |118 |79 |

16. Calculate the average atomic mass of Si, which consists of three isotopes listed below.

|Isotope |Si-28 |Si-29 |Si-30 |

|Atomic Mass |27.98 |28.98 |29.97 |

|Abundance |92.20% |4.70% |3.10% |

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17. Chlorine is primarily two isotopes Cl-35 and Cl-36 and has an average atomic mass of 35.45 u.

a. Which isotope is more abundant? Explain

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b. Estimate the approximate abundances for the two isotopes, Cl-35 and Cl-36 without using your calculator.

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18. Antimony has two isotopes: Sb-123 and Sb-121. Sb-121 has a mass of 120.9 u and an abundance of 57.25 %. Antimony has an average atomic mass of 121.75.

a. What is the abundance of Sb-123?

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b. What is the atomic mass of Sb-123?

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19. Consider the following molecules. Based on the model below, write the unique formula and formula mass.

|Formula | | | | |

|Mass | | | | |

20. How is a molecular compound different from a non-molecular compound?

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21. Fill in the flow chart from the word bank (compound, element, heterogeneous, homogeneous, matter, pure substance, solution).

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| ↓ |

|Is it uniform throughout? | |

|↓ No | |Yes ↓ | |

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| |Does it have a variable composition? |

| |↓ No | |Yes ↓ |

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| ↓ | |

|Can it be separated into simpler substances? | |

|↓No | |Yes↓ | |

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Penny Isotope Lab

22. Mass 50 pennies, determine the average mass, use the average to determine the percentages of heavy pennies and light pennies, and compare to the actual numbers.

Count 50 pennies and mass the total.

a. Record the mass and calculate the average.

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b. Calculate the percentage of 2.5-g and 3.1-g penny.

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Separate the pennies into pre-1982, 1982 and post-1982 piles. Mass each 1982 penny.

c. Record the number of pennies in each group and the percentage of pennies that are 3.1 g and 2.5 g.

|pre-1982 |1982 |post-1982 |

|3.1 g |3.1 g |2.5 g |2.5 g |

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|Total %: |Total %: |

d. Calculate the % difference between the actual % of 2.5 g penny (c) and the calculated % (b).

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Radioactivity

23. Write nuclear equation for the radioactive process.

|Alpha emission of Ra-226 | |

|Beta emission of I-131 | |

|Positron emission of C-11 | |

|Th-231 decays to Pa-231 | |

|Th-232 decays to Ra-228 | |

24. What is the most likely mode of decay for the isotopes.

|H-3 |N-11 |Co-60 |Rn-222 |

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25. Write a nuclear equation for the most likely mode of decay.

|B-8 | |

|K-40 | |

|U-235 | |

|Co-60 | |

26. Fill in the missing part of the nuclear transmutations.

|___ + 10n → 2411Na + 42He |

|147N + ___ → 178O + 11H |

|168O + 11H → ___ + 42He |

|5826Fe + 10n → 5927Co + ___ |

27. Pu-239 undergoes nuclear fission when bombarded by a neutron. Determine the missing (?) product.

[pic]

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28. The graph shown below illustrates the decay of 8842Mo, which decays via positron emission.

[pic]

a. Write a nuclear equation for the decay of 8842Mo.

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b. What is the half-life of the decay?

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c. What is the rate constant for the decay?

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d. What percent of the original sample remains after 12 minutes?

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e. How many minutes does it take the sample to go from 0.8 g to 0.5 g?

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29. The half-life of radioactive S-35 is 88 days. Determine

a. the rate constant.

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b. the number of days for the sample to be ¼ as radioactive (without a calculator).

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c. the number of days for the sample to be ¼ as radioactive (with a calculator).

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d. the percent that remain radioactive after 290 day.

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30. C-14 (t½ = 5715 yrs) decays to N-14. As a result, the 14C/12C ratio in organic material decreases upon death.

a. What is the rate constant k for C-14?

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b. How old is an organic artifact whose 14C/12C ratio is 65.4 % of a living plant?

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31. 238U (t½ = 4.5 x 109 yr) naturally decays to 206Pb.

a. What is the rate constant k for U-238?

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b. What is the age of the rock whose 206Pb/238U is 1.32/1?

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Radioactive Decay Lab

32. Calculate the number of radioactive atoms that remain after given periods of time, graph the data and compare the graph to a ½-life graph.

a. Calculate the rate constant k given t½ = 2.77 minutes.

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b. Calculate the number of radioactive atoms that remain after each minute.

|t |

Electron Structure—Bohr Model

33. Calculate the missing value.

|Photon energy |Wavelength |

|3.25 x 10-18 J | |

| |1.216 x 10-7 m |

|6.60 x 10-19 J | |

| |345 nm |

34. A hydrogen electron transitions from n = 1 to n = 4.

a. What is the electron's energy at n = 1?

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b. What is the electron's energy at n = 4?

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c. Does the electron gain or lose energy in the transition?

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d. What is the change in energy of the electron?

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35. A hydrogen electron transitions from n = 9 to n = 7.

a. What is the energy of the electron when n = 9.

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b. What is the energy of the electron when n = 7

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c. What is the change in energy for the transition?

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d. Is energy absorbed or released during the transition?

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e. What is the wavelength of the light emission?

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Hydrogen Spectrum Lab

36. Use hydrogen spectrum data to determine the electron transition which generates each color.

View the hydrogen spectrum through the spectroscope and compare this with diagram in your D-1 lecture notes.

a. Calculate Ephoton for each wavelength (_._ _ x 10-19 J).

|4.10 x 10-7 m |4.34 x 10-7 m |4.86 x 10-7 m |6.56 x 10-7 m |

|4.88 | | | |

b. Calculate En for each value of n (_._ _ x 10-19 J).

|1 |2 |3 |4 |5 |6 |

|-21.8 | | | | | |

c. Calculate ΔE for each transition (_._ _ x 10-19 J).

|From |n = 6 |n = 5 |n = 4 |n = 3 |n = 2 |

|To 5 |0.26 | | | | |

|To 4 | | | | | |

|To 3 | | | | | |

|To 2 | | | | | |

|To 1 | | | | | |

d. Match the results from (a) and (c) and record the transition that produced each spectral line.

|4.10 x 10-7 m |4.34 x 10-7 m |4.86 x 10-7 m |6.56 x 10-7 m |

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e. What do these transitions have in common?

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f. Which transition listed in (c) would produce the shortest wavelength?

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g. Is this wavelength in the infrared or ultraviolet region?

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Practice Quiz

Multiple Choice (No calculator)

Briefly explain why the answer is correct in the space provided.

|1 |

2. Which scientist is correctly matched with the discovery?

(A) Millikan discovered the electron charge-to-mass ratio.

(B) Thomson discovered the charge of an electron.

(C) Bohr discovered the four quantum numbers.

(D) Rutherford discovered the nucleus.

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3. Which represents a pair of isotopes?

(A) 146C and 147N (B) 189F and 3517Cl

(C) 5626Fe2+ and 5626Fe3+ (D) 3517Cl and 3617Cl

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4. Copper has two isotopes, 63Cu and 65Cu. What is the abundance of 63Cu if the average atomic mass is 63.5?

(A) 90% (B) 75% (C) 50% (D) 20%

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5. Which of the following is correct about beta particles?

I. mass number of 4 and a charge of +2

II. more penetrating than alpha particles

III. electron

(A) I only (B) III only (C) I and II (D) II and III

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6. For the types of radiation given, which is the correct order of increasing ability to penetrate a piece of lead?

(A) α < γ < β (B) α < β < γ

(C) β < α < γ (D) β < γ < α

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7. 24996Cm is radioactive and decays by the loss of one beta particle. The other product is

(A) 24594Pu (B) 24997Bk (C) 24896Cm (D) 25096Cm

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8. 25198Cf → 2 10n + 13154Xe + . . .

What is the missing product in the nuclear reaction?

(A) 11842Mo (B) 11844Ru (C) 12042Mo (D) 12044Ru

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9. The radioactive decay of C-14 to N-14 occurs by

(A) beta particle emission (B) alpha particle emission

(C) positron emission (D) electron capture

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10. What is the resulting nucleus after 21484Po emits 2 α and 2 β particles?

(A) 20683Bi (B) 21083Bi (C) 20682Pb (D) 20882Pb

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11. 23592U + 10n → 14155Cs + 3 10n + X

Neutron bombardment of uranium can induce the reaction represented above. Nuclide X is which of the following?

(A) 9235Br (B) 9435Br (C) 9137Rb (D) 9237Rb

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12. If 87.5 % of a sample of pure Pb-210 decays in 36 days, what is the half-life of Pb-210?

(A) 6 days (B) 8 days (C) 12 days (D) 14 days

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13. The half-life of isotope Y is 12 minutes. What mass of Y was originally present if 1 g is left after 60 minutes?

(A) 8 g (B) 16 g (C) 24 g (D) 32 g

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Free Response (Calculator)

1. A student measures the mass of an object to be 75.011 g. The true mass is 77.500 g. What is the percent error?

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2. Determine the % deviation for the following massings.

|Mass |17.215 g |17.135 g |16.988 g |17.455 g |

|mean | |

|Δ | | | | |

|Average | |

|Δ | |

|% Δ | |

3. A student adds 33.552 g of pellets to a graduated cylinder containing 5.00 mL. The total volume of the pellets and water is 8.05 mL. What is the density of the pellets?

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4. An empty graduated cylinder (25.044 g) is filled with 50.0 mL of liquid. The total mass of cylinder and liquid is 69.886 g. What is the density of the liquid?

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5. You have a butane lighter for the purpose of measuring butane's density at room temperature and pressure.

a. How would you collect a sample of butane?

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b. How would you determine the volume collected?

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c. How would you determine the mass collected?

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d. How would you determine the density of butane?

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6. Consider the gas, UF6, which has a density of 15.7 g/L.

a. What is the molar mass of UF6?

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b. How many moles of UF6 have a volume of 1.00 L?

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c. What is the volume of 25.0 g of UF6?

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7. Write a brief description of the scientists' contribution.

|Scientist |Contribution |

|J. J. Thomson | |

|Millikan | |

|Rutherford | |

|Bohr | |

8. Write the symbol for an atom that contains 24 protons, 28 neutrons and 21 electrons.

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9. Consider two variations of 2311 Na; 2411Na and 2311Na+. How is each different from Na-23 and what is it called?

| |How is it different? |What is it called? |

|2411Na | | |

|2311Na+ | | |

10. Calculate the average atomic mass of Pb given the atomic masses and abundances of its stable isotopes.

|Isotope |Pb-204 |Pb-206 |Pb-207 |Pb-208 |

|Atomic Mass |204 u |206 u |207 u |208 u |

|Abundance |1.4% |24.1% |22.1% |52.4% |

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11. Copper has two isotopes, Cu-63 and Cu-65. What is the % Cu-65 if the average mass of copper is 63.5?

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12. Determine the molar mass of the following pure substances.

|NaCl |O2 |C6H12O6 |NH3 |

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13. Place the forms of H (H, H+, H-, H2) in the table.

|Atom |Cation |Anion |Molecule |

| | | | |

14. Predict the mode of decay for the following nuclei and then write a balanced nuclear equation for the process.

|Nuclei |Decay Mode |Balanced Nuclear Equation |

|N-13 | | |

|Cu-68 | | |

|Np-241 | | |

15. Complete and balance the following nuclear equations.

|3216S + 10n → 11H + ___ |

|74Be + 0-1e → ___ |

|___→ 18776Os + 0-1e |

|23592U + 10n → 13554Xe + 2 10n + ___ |

16. What is produced after 21484Po produces 2 α and 2 β?

| |

17. The rate constant for tritium (H-3) is 0.0564 yr-1.

a. What is the half-life?

| |

b. How long will it take 75% of the sample to decay?

| |

18. The half-life of Cr-51 is 27.8 days.

a. Calculate the rate constant of Cr-51.

| |

b. What percent of the sample will remain after 100 days?

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19. If 87.5 % of a sample of pure 131I decays in 24 days, what is the half-life of 131I? (without calculator)

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20. The half-life of Y is 30 s. What mass of Y was originally present if 1 g is left after 60 s?

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21. How old is a wooden artifact, which has a C-14 (t½ = 5715 yr) activity of 10.4 counts/min•g compared to living wood that has a C-14 activity of 13.6 counts/min•g?

| |

22. A line in the spectrum of hydrogen is associated with the electronic transition from n = 6 to n = 2.

a. Find the energy that is given off in the transition from

n = 6 to n = 2 for hydrogen (En = -2.18 x 10-18/n2 J).

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b. Calculate the wavelength of the radiation associated with the spectral line. (Ephoton = (2.00 x 10-25 J•m)/λ)

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