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

% 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(absolute Δ)/N

• % Δ = 100(average Δ)/(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 |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

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

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

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

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

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

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

Experiments

1. Density Lab—Measure the mass and volume of a solid, liquid and gas, determine densities, and use the density to identify the substances.

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 = m/V |

|Δm | | |

| | | | | | | |

b. Highlight the type of 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.

| |

Liquid: Mass a clean, dry 10 mL graduated cylinder (m1). Add 10.0 mL liquid to the cylinder. Mass 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 = m/V |

| | | |10. mL | |

e. Highlight the type of 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.

| |

Gas: Add ½ scoop of baking soda (NaHCO3) to the flask and ½ fill the pipet with 6 M HCl (handle with care). Stopper the flask and mass the assembly (m1). Fill the gas collecting bottle with water. Measure the volume of water (V1). Refill the bottle, cover the mouth and place it inverted in the trough. 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. Mass the assembly (m2). Measure the volume of water remaining in the bottle (V2).

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

| m1 – m2 = | V1 – V2 = |d = m/V |

|Δm |ΔV | |

| | | |

|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.

| |

2. Penny Isotope Lab— Use the average mass of 50 pennies to determine the percentages of heavy pennies and light pennies, and compare to the actual percentages.

Count 50 pennies and mass the total.

a. Record the mass and calculate the average.

| | |

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

| |

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 |

| | | | |

|Total %: |Total %: |

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

| |

3. Radioactive Decay Lab—Construct a half-life graph.

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.

| |

b. Calculate the number of radioactive atoms that remain after each minute (Nt = Noe-kt).

|t |

4. Hydrogen Spectrum Lab—Use hydrogen spectrum data to determine the electron transition which generates each color.

a. Calculate Ephoton for each wavelength. Express all answers as with 3 sf x 10-19 J (E = 2.00 x 10-25 J•m/λ).

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

| | | | |

b. Calculate En for each value of n. Express all answers with 3 sf x 10-19 J (En = -2.18 x 10-18 J/n2).

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

| | | | | | |

c. Calculate ΔE for each transition listed below. Express all answers with 3 sf x 10-19 J (ΔE = En-final – En-initial).

|6 to 5 | |5 to 4 | |4 to 2 | |

|6 to 4 | |5 to 3 | |4 to 1 | |

|6 to 3 | |5 to 2 | |3 to 2 | |

|6 to 2 | |5 to 1 | |3 to 1 | |

|6 to 1 | |4 to 3 | |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 |

| | | | |

e. What do these transitions have in common?

| |

Practice Problems

A. Measurement

1. How many significant figures are there in?

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

| | | | |

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?

| |

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?

| |

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?

| |

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?

| |

9. Calculate the molar masses to 3 significant figures.

|NaCl |H2O |Cl2 |

| | | |

|(NH4)2SO4 |C4H7NO4 |CuSO4•5 H2O |

| | | |

10. Use dimensional analysis to determine the following

a. 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 | |

b. 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 | |

B. Atomic Nature of Matter

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

| |

| |

12. 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.

| |

13. Complete the chart below.

|Symbol |protons |neutrons |electrons |

|2713Al | | | |

|4019K+ | | | |

|3115P3- | | | |

| |26 |30 |23 |

| |17 |17 |18 |

| |79 |118 |79 |

14. 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% |

| |

15. 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

| |

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

| |

16. 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?

| |

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

| |

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

|Formula | | | | |

|Mass | | | | |

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

| | | | |

| ↓ |

|Is it uniform throughout? | |

|↓ No | |Yes ↓ | |

| | | | |

| | ↓ |

| |Does it have a variable composition? |

| |↓ No | |Yes ↓ |

| | | | |

| ↓ | |

|Can it be separated into simpler substances? | |

|↓No | |Yes↓ | |

| | | | |

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

| |

| |

C. Radioactivity

20. 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 | |

21. What is the most likely mode of decay for the isotopes?

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

| | | | |

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

|B-8 | |

|K-40 | |

|U-235 | |

|Co-60 | |

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

|_____ + 10n → 2411Na + 42He |

|147N + _____ → 178O + 11H |

|168O + 11H → _____ + 42He |

|5826Fe + 10n → 5927Co + _____ |

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

[pic]

| |

25. 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.

| |

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

| |

c. What is the rate constant for the decay?

| |

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

| |

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

| |

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

a. the rate constant.

| |

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

| |

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

| |

d. the percent that remain radioactive after 290 day.

| |

27. 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?

| |

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

| |

28. 238U (t½ = 4.5 x 109 yr) naturally decays to 206Pb.

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

| |

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

| |

D. Electron Structure—Bohr Model

29. 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 |

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

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

| |

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

| |

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

| |

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

| |

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

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

| |

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

| |

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

| |

d. Is energy absorbed or released during the transition?

| |

e. What is the wavelength of the light emission?

| |

Summary

Measurement in Chemistry

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

Measurements are made using the metric system, where the standard units are called SI units, which are based on the meter, kilogram, and second as the basic units of length, mass, and time, respectively. The SI temperature scale is the Kelvin scale, although the Celsius scale is frequently used in chemistry. The metric system employs a set of prefixes to indicate decimal fractions or multiples of the base units; k (10-3), c (10-2), m (10-3), μ (10-6) and n (10-9).

All measured quantities are inexact to some extent. The precision of a measurement indicates how closely different measurements of a quantity agree with one another. The accuracy of a measurement indicates how well a measurement agrees with the accepted value. Significant figures indicate the level of certainty in a measurement. Significant figures in a measured quantity include one estimated digit; the last digit of the measurement. Calculations involving measured quantities are reported with the appropriate number of significant figures. In multiplication and division, the number of significant figures is used. In addition and subtraction, the position of the least accurate significant figure is used. Relative difference between an experiment value (E) and a true value (T) is % difference:

% Δ = 100|E – T|/T. Relative spread of N number of trials is % deviation: % Δ = 100Σ|trial – mean|/N(mean).

Mass and volume measure amount of matter. Density relates mass to volume, d = m/V. Chemical processes involve interaction of particles, which are measured in moles. The number of particles in a mole is called Avogadro's number, which is 6.02 x 1023. This number is based on using periodic table masses to be equal to mass of formula unit labeled in grams. Molar mass (MM) is the sum of atomic masses in the chemical formula. For example, the mass of one H2O molecule is 18.0 u, so the molar mass of H2O is 18.0 g.

In the dimensional analysis approach, we keep track of units as we carry measurements through calculations. The given units are multiplied by a series of conversion factors, which are ratios of equivalent quantities. After canceling out units algebraically, what remain are the target units.

Atomic Nature of Matter

Atoms are the basic building blocks of matter; they are the smallest units of an element that can combine with other elements. Atoms are composed of even smaller subatomic particles. Experiments led to the discovery and characterization of subatomic particles. Thomson experimented with cathode rays in magnetic and electric fields, which led to the discovery of the electron and its charge-to-mass ratio. Millikan worked with oil-drops in a vacuum to determine the charge of the electron. Rutherford observed the scattering of α particles by gold metal foil and concluded that atoms have a dense, positive nucleus.

The atom's nucleus contains protons and neutrons, whereas electrons move in the space around the nucleus. The charges of subatomic particles in terms of the charge of an electron are: electron -1, proton +1 and neutron 0. Masses in terms of the mass of a proton are: proton and neutron 1, and electron 0.00055.

Elements are classified by their atomic number or Z value, which equals the number of protons. The mass number or A value is the sum of protons and neutrons. Atoms of the same element that differ in mass number are called isotopes.

In a neutral atom, the number of protons equals the number of electrons. An anion is formed when electrons exceed protons. A cation is formed when protons exceed electrons.

The unified atomic mass scale (u) is 1/12 the mass of a C-12 atom. The average atomic mass of an element is calculated using the formula: 100mav = %1m1 + %2m2 ...

The two kinds of pure substances are elements and compounds. Elements are identified by a chemical symbol. Compounds are composed of two or more elements joined chemically and identified by a chemical formula, which shows the composition. Molecular compounds have a defined size, whereas crystalline compounds are unbounded, where their formula shows the ratio of atoms in the compound.

Mixtures are composed of multiple pure substances in an object or container and have variable compositions. They can be homogeneous or heterogeneous. Homogeneous mixtures are also called solutions and are uniform throughout.

Radioactivity

There are four kinds of radioactive decay: emission of alpha particles (α or 42He), beta particle (β or 0-1e), positron particle (β+, 01e), and gamma radiation (00γ).

In nuclear equations, reactant and product nuclei are represented by AZX, which is its nuclear symbol. In a balanced equation the sum of reactant A and Z values equal the sum of product A and Z values.

Modes of decay can be predicted by comparing the number of neutrons with the average (A – Z). In general, neutron-rich nuclei emit beta particles; neutron-poor nuclei emit positron particles; and nuclei above Z = 83 emit alpha particles.

Nuclear transmutations, induced conversions of one nucleus into another, can be brought about by bombarding nuclei with either charged particles or neutrons.

The decay rate (radioactivity) is proportional to the number of radioactive atoms, rate = kNt. The time for half of the radioactive atoms to decay is constant, t½ = (ln2)/k. The time interval t for No number of radioactive atoms to reduce to Nt is determined by the formula, kt = ln(No/Nt).

Electron Structure—Bohr Model

The electronic structure of an atom describes the energies and arrangement of electrons around the atom. Much of what is known about the electronic structure of atoms was obtained by observing atomic spectra, which is the radiant energy emitted or absorbed by matter.

Equations for radiant energy, Ephoton = hf and speed of light, c = fλ are combined in Ephoton = hc/λ = 2 x 10-25 J•m/λ.

Bohr analyzed the wavelengths of light emitted by hydrogen atoms and proposed a model that explains its atomic spectrum. In this model the energy of the hydrogen atom depends on the value its quantum number n, where En = -2.18 x 1018 J/n2. The value of n is a positive integer (1, 2, 3 . . .). As n increases, the energy of the electron increases until it reaches a value of 0 J, where n equals infinity and the electron leaves the atom or ionizes. The lowest energy state where n = 1 is called the ground state. Other values of n correspond to excited states. Light is emitted when the electron drops from a higher energy state to a lower energy state and light is absorbed when excited from a lower energy state to a higher one. The energy of light emitted or absorbed equals the difference in energy between the two states, Ephoton = En-final – En-initial = 2.00 x 10-25 J•m/λ.

Practice Multiple Choice

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

1. Based on the data, the density of the solid in g/mL is

Mass of metal 19.611 g

Volume of water 12.4 mL

Volume of water + metal 14.9 mL

(A) 7.8444 (B) 7.844 (C) 7.84 (D) 7.8

| | |

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.

| | |

3. Which represents a pair of isotopes?

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

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

| | |

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%

| | |

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

| | |

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

| | |

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

| | |

8. 25198Cf → 2 10n + 13154Xe + . . .

What is the missing product in the nuclear reaction?

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

| | |

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

| | |

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

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

| | |

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

| | |

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

| | |

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

| | |

Practice Free Response

1. 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?

| |

b. How would you determine the volume collected?

| |

c. How would you determine the mass collected?

| |

d. How would you determine the density of butane?

| |

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

a. What is the molar mass of UF6?

| |

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

| |

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

| |

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

|Scientist |Contribution |

|J. J. Thomson | |

|Millikan | |

|Rutherford | |

|Bohr | |

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

| |

5. 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+ | | |

6. 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% |

| |

7. Consider the isotopes of copper.

a. Write the beta decay of Cu-64.

| |

b. When Cu-64 is bombarded with C-12, three neutrons and another particle are produced. Write the equation.

| |

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

| |

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

|NaCl |O2 |C6H12O6 |NH3 |

| | | | |

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

|Atom |Cation |Anion |Molecule |

| | | | |

10. 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 | | |

11. Complete and balance the following nuclear equations.

|3216S + 10n → 11H + _____ |

|74Be + 0-1e → _____ |

|_____ → 18776Os + 0-1e |

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

12. 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?

| |

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

| |

14. 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?

| |

15. Hydrogen atoms in the ground state are ionized by UV light.

a. Calculate the energy needed to ionize an electron from n = 1? (En = -2.18 x 10-18/n2 J).

| |

b. Calculate the wavelength of UV light.

(Ephoton = (2.00 x 10-25 J•m)/λ)

| |

16. Answer the following questions regarding light and its interactions with molecules, atoms, and ions.

a. The longest wavelength of light with enough energy to break the CI-CI bond in CI2(g) is 495 nm.

(1) Calculate the frequency in s-1 of the light.

| |

(2) Calculate the energy in J of a photon of the light.

| |

(3) Calculate the energy in kJ•mol-1 of the CI-CI bond.

| |

b. A certain line in the spectrum of atomic hydrogen is associated with the electronic transition in the H atom from the sixth energy level (n = 6) to the second energy level (n = 2).

(1) Indicate whether the H atom emits energy or whether it absorbs energy during the transition. Justify your answer.

| |

| |

| |

(2) Calculate the wavelength in nm of the radiation associated with the spectral line.

| |

| |

| |

| |

| |

| |

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download