PSI AP CHEMISTRY Atomic Theory and Models of the Atom Classwork:

PSI AP CHEMISTRY
Summer Assignment Review Unit
Free Response
Atomic Theory and Models of the Atom
Classwork:
1. An atom has 65 neutrons and a mass of 122 u. Write the proper nuclide symbol.
2. Atom “A” has 11 neutrons and Atom “B” also has 11 neutrons. Are these atoms
isotopes? Explain.
3. Which component of Dalton’s atomic theory reflects his understanding of the law of
conservation of mass for chemical processes? Why?
4. The atomic masses of elements are generally not whole numbers. Explain why.
5. Gallium has two stable isotopes: Ga - 69 (68.92 u) and Ga-71 (70.92 u). Using the
average atomic mass of gallium, what is the % abundance of the heavier of the two
isotopes?
6. Naturally occurring chlorine is 75.78% Cl - 35 (34.9689 u) and 24.22% Cl - 37
(36.9659 u). Calculate the average atomic mass.
7. What conclusions were made from Rutherford’s “Gold Foil Experiment”? How did this
experiment change the atomic model?
8. Explain why the Bohr model was insufficient and how the quantum model explained
this deficiency. How is the electron viewed differently in each model?
9. Consider the equations relating wavelength, frequency and energy of electromagnetic
radiation. How is the energy of a photon related to its frequency and wavelength?
10. Calculate the frequency of an X ray that has a wavelength of 8.21 nm.
11. What is the frequency of a photon that has an energy of 3.7 × 10-18 J
12. Provide a brief explanation for what each of the four quantum numbers describes.
Homework:
1. Radioactive americium - 241 is used in household smoke detectors and in bone
mineral analysis.
a) Give the number of electrons, protons, and neutrons in an atom of americium 241.
b) Write the proper nuclide symbol.
2. What characteristics do atoms of carbon-12, carbon-13, and carbon-14 have in
common? IN what ways are they different?
3. Identify the isotope that has atoms with
a) 117 neutrons, 77 protons, and 77 electrons
b) 30 neutrons, 28 protons, and 28 electrons
4. How did the discovery of isotopes conflict with Dalton’s atomic theory?
5. The average mass of any large number of atoms of a given element is always the
same for a given element. Explain.
6. . Naturally occurring boron is 19.9% B - 10 (mass = 10.01294 u) and 80.1% B - 11
(mass = 11.0093 u). Calculate the average atomic mass.
7. Uranium has an atomic mass equal to 238.0289. It consists of two isotopes: uranium 235 with an isotopic mass of 235.044 u and uranium-238 with an isotopic mass of
238.051 u. Calculate the % abundance of the uranium-235 isotope.
8. How did Rutherford interpret the deflection of α-particles in his gold foil experiment?
Did these findings support or disprove the plum pudding model of the atom? Explain.
9. How does Bohr's model of the atom explain the existence of line spectra? How are
spectral lines produced?
10. What is the energy of a photon that has a wavelength of 8.33 x 10-6 m?
11. What is the wavelength of a photon that has an energy of 5.25 × 10-19 J
12. How does the quantum model describe the location of an electron?
Periodic Table
Classwork:
1. An element is found to gain three electrons when it forms an ion.
a) What group number would this element be found in?
b) Is there enough information provided to determine what period it is in?
2. Look at the average atomic mass of Ar and K.
a) Explain why early scientists might have been tempted to have K follow Cl on
the periodic table.
b) Propose two reasons as to why they placed Ar after Cl instead of K.
3. Identify the following elements:
a) An alkali metal in the 5th period.
b) A transition metal
c) An atom in the 3rd period that forms a stable ion with a -1 charge.
4. Explain why atoms tend to gain or lose electrons relative to the number of valence
electrons.
5. What ions are the following elements likely to form?
a) Nitrogen
b) Calcium
c) Sulfur
6. Explain why the noble gases are inert (unreactive).
7. Why are the charges of transition metals (d-block) difficult to predict?
8. Write the formulas for the following binary ionic compounds.
a) magnesium oxide
b) manganese(II) chloride
c) calcium phosphide
d) copper(I) sulfide
9. Write the formulas for the following polyatomic ionic compounds.
a) potassium sulfate
b) aluminum phosphate
c) iron(III) carbonate
d) aluminum hydroxide
10. Aluminum reacts with a certain nonmetallic element to form a compound with the
general formula Al2X3. Element X must be from which group on the periodic table?
Homework:
1. What accounts for similarities of chemical properties for elements in the same
group(family)?
2. Provide the group names for the elements in Group 1, 2, 17 and 18. Provide and
example of an element in each of the above groups.
3. Identify the following elements:
a) A halogen in the 3rd period.
b) A metalloid
c) An atom in the 4th period that forms a stable ion with a +1 charge.
4. Locate the following elements on the periodic table and indicate which orbital type is
occupied by its valence electrons
a) Lithium
b) Silicon
c) Copper
5. What ions are the following elements likely to form?
a) Oxygen
b) Sodium
c) Bromine
6. A main group element in Period 4 forms the molecular compound H2E and the ionic
compound Na2E.
a) To which group does the element belong?
b) Write the name and symbol of the element.
7. Write the formulas for the following binary ionic compounds.
a) sodium sulfide
b) cobalt(II) chloride
c) lithium nitride
d) Tin(IV) oxide
8. Write the formulas for the following polyatomic ionic compounds.
a) barium nitrate
b) calcium phosphite
c) iron(II) chromate
d) potassium permanganate
9. How many total ions (cations and anions) are present in the following ionic
compounds?
a) sodium acetate
b) aluminum nitrate
c) Copper(II) chloride
10. The most common charge associated with silver in its compounds is +1. Indicate the
formulas you would expect for the ionic compounds formed between silver and the
following elements.
a) iodine
b) sulfur
c) phosphorous
Mole Concept
Classwork:
1. Answer the following questions for the compound aluminum sulfate.
a) What is the molar mass of this compound?
b) What is the mass of a 1.5 mole sample?
c) How many oxygen atoms are present in the 1.5 mol sample
2. What mass of rhodium contains as many atoms as there are in
a) gallium atoms in 36.0 g gallium
b) indium atoms in 36.0 g indium
3. a) Calculate the mass in grams, of 0.433 mol of calcium nitrate.
b) How many formula units of calcium nitrate are present?
c) How many nitrate ions are present?
4. Carbon has two isotopes C-12 (99%) and C-13 (1%).
a) How many atoms of C would be present in a 34 gram sample of pure diamond
(pure carbon)?
b) How many atoms of those are C-13 atoms?
5. A sample of Ni(CO)4, a toxic transition-metal complex, has 5.23 × 1024 atoms of carbon.
How many atoms of Ni does it contain?
6. How many grams of CO2 are in 7.50 liters of CO2 at STP?
7. Without doing any detailed calculations, rank the following samples in order of
increasing number of atoms: 0.50 mol H2O; 23 g Na; 6.0 x 1023 N2 molecules.
8. A reaction produces 0.0891 grams of ammonia gas (NH3).
a) How many grams of N2 must have reacted to produce this ammonia gas?
b) Assuming N2 gas was the entire source of N and all of it was converted to
ammonia, how many L of N2 gas reacted assuming the reaction was carried out
at STP conditions?
b) How many L of ammonia were produced assuming the reaction was carried
out at STP conditions?
9. Calculate the following quantities for 0.200L of a 0.400M Calcium Iodide (CaI2)
solution.
a) moles of CaI2
b) grams of CaI2 required to prepare the solution
b) moles of I10. What is the concentration (M) of a NaCl solution prepared by dissolving 9.3 g of
NaCl in sufficient water to give 350 mL of solution?
11. a) How many moles of sodium ions are present in 150mL of a 0.75 M sodium
phosphate solution?
b) What is the molarity of sodium ions?
Homework:
1. Answer the following questions for a 3.50g sample of C6H12O6.
a) What is the molar mass of this compound?
b) How many moles are in the sample?
c) How many hydrogen atoms are present in the sample?
2. a) Calculate the number of moles in 1.75 grams of sodium carbonate.
b) How many formula units of sodium carbonate are present?
c) How many sodium ions are present?
3. Boron has two isotopes B-10 (19.9%) and B-11 (80.1%).
a) How many atoms of B would be present in a 50 gram sample of pure boron?
b) How many atoms of those are B-10 atoms?
4. A sample of C12H22O11 contains 0.4662 moles of carbon atoms. How many moles of
hydrogen atoms are in the sample?
5. Without doing any detailed calculations, rank the following samples in order of
increasing number of atoms: 3.0 x 1023 molecules H2O2; 2 mol CH4; 32 g O2
6. One component of smog is nitrogen monoxide, NO. A car produces about 8.0 g of this
gas per day. What is the volume at STP?
7. A reaction produces 100 grams of water.
a) How many grams of H2 must have reacted to produce this amount of water if 1
mol of H2O is produced for every 1 mol of H2 that reacts?
b) Assuming H2 gas was the entire source of H and all of it was converted to
water, how many L of H2 gas reacted assuming the reaction was carried out at
STP conditions?
c) How many molecules of H2 reacted, assuming the reaction was carried out at
STP conditions?
8. Calculate the following quantities for 343 mL of a 1.27M Na2SO4 solution.
a) Moles of Na2SO4
b) grams of Na2SO4 required to prepare the solution
b) moles of Na+
9. What is the molarity of a 750 mL solution containing 50.0 g KCl?
10. a) How many moles of hydroxide ions are present in 300mL of a 2.50M Ca(OH) 2
solution?
b) What is the molarity of hydroxide ions?
TEACHER KEY
Atomic Theory and Models of the Atom
Classwork:
1. An atom has 65 neutrons and a mass of 122 u. Write the proper nuclide symbol.
122
La
57
2. Atom “A” has 11 neutrons and Atom “B” also has 11 neutrons. Are these atoms
isotopes? Explain.
No! Isotopes will have differing numbers of neutrons.
3. Which component of Dalton’s atomic theory reflects his understanding of the law of
conservation of mass for chemical processes? Why?
Atoms are rearranged in chemical processes, not created or destroyed.
4. The atomic masses of elements are generally not whole numbers. Explain why.
The reported masses on the periodic table are averages of all of the
isotopes of a particular element AND are standardized to C-12.
5. Gallium has two stable isotopes: Ga - 69 (68.92 u) and Ga-71 (70.92 u). Using the
average atomic mass of gallium, what is the % abundance of the heavier of the two
isotopes?
39.9%
5. Naturally occurring chlorine is 75.78% Cl - 35 (34.9689 u) and 24.22% Cl - 37
(36.9659 u). Calculate the average atomic mass.
35.45 u
7. What conclusions were made from Rutherford’s “Gold Foil Experiment”? How did this
experiment change the atomic model?
The atom is mostly empty space with the positive protons densely packed
together in the nucleus. Instead of having the electrons and protons mixed
together, this had the electrons in orbit around the positively charged
nucleus.
8. Explain why the Bohr model was insufficient and how the quantum model explained
this deficiency. How is the electron viewed differently in each model?
The Bohr model was not able to explain WHY the orbits were stable, the
PES spectra of period 2 elements, and hyperfine spectral lines. The quantum
model postulated stable orbitals IF we view the electron as a wave instead of as a
particle. The quantum orbitals represent areas of high probability of finding an
electron.
9. Consider the equations relating wavelength, frequency and energy of electromagnetic
radiation. How is the energy of a photon related to its frequency and wavelength?
The energy is inversely proportional to the wavelength and directly
proportional to the frequency.
10. Calculate the frequency of an X ray that has a wavelength of 8.21 nm.
3.65 x 1016 1/s
11. What is the frequency of a photon that has an energy of 3.7 × 10-18 J
5.6 x 1015 1/s
12. Provide a brief explanation for what each of the four quantum numbers describes.
Principal – Main energy level
Azimuthal – Orbital
Magnetic – Orientation of orbital
Spin – spin state of electron
Homework:
1. Radioactive americium - 241 is used in household smoke detectors and in bone
mineral analysis.
a) Give the number of electrons, protons, and neutrons in a neutral atom of
americium - 241.
Protons = 95 Neutrons = 146 electrons = 95
b) Write the proper nuclide symbol.
241
Am
95
2. What characteristics do atoms of carbon-12, carbon-13, and carbon-14 have in
common? In what ways are they different?
Same atomic number or number of protons. They each have
different numbers of neutrons and atomic mass
3. Identify the isotope that has atoms with
a) 117 neutrons, 77 protons, and 77 electrons
Ir -177
b) 30 neutrons and an atomic mass of 58
Nickel-58
4. How did the discovery of isotopes conflict with Dalton’s atomic theory?
Dalton postulated that atoms of the same element were all identical.
Isotopes are atoms of the same element but differ in mass or # of neutrons.
5. The average mass of any large number of atoms of a given element is always the
same for a given element. Explain.
In a large sample, the number of atoms of each isotope will match the
natural distribution or abundances of each.
6. . Naturally occurring boron is 19.9% B - 10 (mass = 10.01294 u) and 80.1% B - 11
(mass = 11.0093 u). Calculate the average atomic mass.
10.811 u
7. Uranium has an atomic mass equal to 238.0289. It consists of two stable isotopes:
uranium -235 with an isotopic mass of 235.044 u and uranium-238 with an isotopic mass
of 238.051 u. Calculate the % abundance of the uranium-235 isotope.
0.72%
8. How did Rutherford interpret the deflection of α-particles in his gold foil experiment?
Did these findings support or disprove the plum pudding model of the atom? Explain.
Rutherford believed the deflections were best explained by envisioning the
protons being densely packed in a nucleus causing the deflection of the alpha
particles.
These findings did not support the plum pudding model as it predicted all
alpha particles would move through the gold foil.
9. How does Bohr's model of the atom explain the existence of line spectra? How are
spectral lines produced?
The Bohr model postulates discrete orbits where electrons can exist. The
spectral lines are created when light is emitted when an electron moves from
one orbit to a less energetic one.
10. What is the energy of a photon that has a wavelength of 8.33 x 10-6 m?
2.39 X 10-20 J
11. What is the wavelength (in nm) of a photon that has an energy of 5.25 × 10-19 J?
379 nm
12. How does the quantum model describe the location of an electron?
The quantum model uses four numbers to represent the possible quantum
states of an electron in the atom. These represent areas of high probability of
locating the electron or in other words – area of relatively high electron
density.
Periodic Table
Classwork:
1. An element is found to gain three electrons when it forms an ion.
a) What group number would this element be found in?
Group 15 or 5A
b) Is there enough information provided to determine what period it is in?
No
2. Look at the average atomic mass of Ar and K.
a) Explain why early scientists might have been tempted to have K follow Cl on
the periodic table.
K is next if grouped by increasing atomic mass
b) Propose two reasons as to why they placed Ar after Cl instead of K.
K is not inert like the rest of the noble gases - Ar is. Secondly, Ar
has the next highest atomic number, 1 less than K.
3. Identify the following elements:
a) An alkali metal in the 5th period.
Rb
b) A transition metal
Cr
c) An atom in the 3rd period that forms a stable ion with a -1 charge.
Cl
4. Explain why atoms tend to gain or lose electrons relative to the number of valence
electrons.
Atoms gain or lose electrons to obtain a full valence shell - which for most
atoms is 8 (full s and p).
5. What ions are the following elements likely to form?
a) Nitrogen
N3b) Calcium
Ca2+
c) Sulfur
S26. Explain why the noble gases are inert (unreactive).
They have a full valence shell
7. Why are the charges of transition metals (d-block) difficult to predict?
They can lose 2 electrons from their outermost s but then
occasionally some of the “d” electrons are lost as well.
8. Write the formulas for the following binary ionic compounds.
a) magnesium oxide
MgO
b) manganese(II) chloride
MnCl2
c) calcium phosphide
Ca3P2
d) copper(I) sulfide
Cu2S
9. Write the formulas for the following polyatomic ionic compounds.
a) potassium sulfate
K2SO4
b) aluminum phosphate
AlPO4
c) iron(III) carbonate
Fe2(CO3)3
d) aluminum hydroxide
Al(OH)3
10. Aluminum reacts with a certain nonmetallic element to form a compound with the
general formula Al2X3. Element X must be from which group on the periodic table?
Group 16 or 6A
Homework:
1. What accounts for similarities of chemical properties for elements in the same
group(family)?
Same number of valence electrons
2. Provide the group names for the elements in Group 1, 2, 17 and 18. Provide and
example of an element in each of the above groups.
Group 1 = alkali metals
Group 2 = alkaline earth metals
Group 17 = Halogens
Group 18 = noble gases
3. Identify the following elements:
a) A halogen in the 3rd period.
Cl
b) A metalloid
Si
c) An atom in the 4th period that forms a stable ion with a +1 charge.
K
4. Locate the following elements on the periodic table and indicate which orbital type(s)
is occupied by its valence electrons
a) Lithium
2s
b) Silicon
3s and 3p
c) Copper
4s
5. What ions are the following elements likely to form?
a) Oxygen
O2b) Sodium
Na+
c) Bromine
Br6. A main group element in Period 4 forms the molecular compound H2E and the ionic
compound Na2E.
a) To which group does the element belong?
Group 16 or 6A
b) Write the name and symbol of the element.
Se
7. Write the formulas for the following binary ionic compounds.
a) sodium sulfide
Na2S
b) cobalt(II) chloride
CoCl2
c) lithium nitride
Li3N
d) Tin(IV) oxide
SnO2
8. Write the formulas for the following polyatomic ionic compounds.
a) barium nitrate
Ba(NO3)2
b) calcium phosphite
Ca3(PO3)2
c) iron(II) chromate
FeCrO4
d) potassium permanganate
KMnO4
9. How many total ions (cations and anions) are present in the following ionic
compounds?
a) sodium acetate
NaC2H3O2 = 2
b) aluminum nitrate
Al(NO3)3 = 4
c) Copper(II) chloride
CuCl2 = 3
10. The most common charge associated with silver in its compounds is +1. Indicate the
formulas you would expect for the ionic compounds formed between silver and ions of
the following elements.
a) iodine
AgI
b) sulfur
Ag2S
c) phosphorous
Ag3P
Mole Concept
Classwork:
1. Answer the following questions for the compound aluminum sulfate.
a) What is the molar mass of this compound?
Al2(SO4)3 = 342.15 g/mol
b) What is the mass of a 1.5 mole sample?
513 g
c) How many oxygen atoms are present in the 1.5 mol sample
1.08 x 1025 O atoms
2. What mass of rhodium contains as many atoms as there are in
a) gallium atoms in 36.0 g gallium
53.1 g
b) indium atoms in 36.0 g indium
32.3 g
3. a) Calculate the mass in grams, of 0.433 mol of calcium nitrate.
71.1 g
b) How many formula units of calcium nitrate are present?
2.61 x 1023 formula units
c) How many nitrate ions are present?
5.21 x 1023 nitrate ions
4. Carbon has two isotopes C-12 (99%) and C-13 (1%).
a) How many atoms of C would be present in a 34 gram sample of pure diamond
(pure carbon)?
1.71 x 1024 atoms C
b) How many atoms of those are C-13 atoms?
1.71 x 1022 atoms C-13
5. A sample of Ni(CO)4, a toxic transition-metal complex, has 5.23 × 1024 atoms of carbon.
How many atoms of Ni does it contain?
1.31 x 1024 atoms Ni
6. How many grams of CO2 are in 7.50 liters of CO2 at STP?
14.7 grams
7. Without doing any detailed calculations, rank the following samples in order of
increasing number of atoms: 0.50 mol H2O; 23 g Na; 6.0 x 1023 N2 molecules.
23 g Na < 0.50 mol H2O < 6.0 x 1023 N2 molecules
8. A reaction produces 0.0891 grams of ammonia gas (NH3).
a) How many grams of N2 must have reacted to produce this ammonia gas?
0.0734 grams
b) Assuming N2 gas was the entire source of N and all of it was converted to
ammonia, how many L of N2 gas reacted assuming the reaction was carried out
at STP conditions?
0.059 L
b) How many L of ammonia were produced assuming the reaction was carried
out at STP conditions?
0.12 L ammonia
9. Calculate the following quantities for 0.200L of a 0.400M Calcium Iodide (CaI2)
solution.
a) moles of CaI2
0.0800 moles
b) grams of CaI2 required to prepare the solution
23.5 grams
b) moles of I0.160 moles I
10. What is the concentration (M) of a NaCl solution prepared by dissolving 9.3 g of
NaCl in sufficient water to give 350 mL of solution?
0.458 M
11. a) How many moles of sodium ions are present in 150mL of a 0.75 M sodium
phosphate solution?
0.338 moles Na
b) What is the molarity of sodium ions?
2.25 M
Homework:
1. Answer the following questions for a 3.50g sample of C6H12O6.
a) What is the molar mass of this compound?
180 g/mol
b) How many moles are in the sample?
0.0194 moles
c) How many hydrogen atoms are present in the sample?
1.4 x 1023 atoms H
2. a) Calculate the number of moles in 1.75 grams of sodium carbonate.
0.0165 moles
b) How many formula units of sodium carbonate are present?
9.9 x 1021 formula units
c) How many sodium ions are present?
1.98 x 1022 sodium ions
3. Boron has two isotopes B-10 (19.9%) and B-11 (80.1%).
a) How many atoms of B would be present in a 50 gram sample of pure boron?
2.78 x 1024 atoms B
b) How many atoms of those are B-10 atoms?
5.54 x 1023 atoms B-10
4. A sample of C12H22O11 contains 0.4662 moles of carbon atoms. How many moles of
hydrogen atoms are in the sample?
0.8547 moles H
5. Without doing any detailed calculations, rank the following samples in order of
increasing number of atoms: 6.0 x 1023 molecules H2O2; 2 mol CH4; 32 g O2
32 g O2 < 6.0 x 1023 molecules H2O2 < 2 mol CH4
6. One component of smog is nitrogen monoxide, NO. A car produces about 8.0 g of this
gas per day. What is the volume at STP?
5.97 L
7. A reaction produces 100 grams of water.
a) How many grams of H2 must have reacted to produce this amount of water if 1
mol of H2O is produced for every 1 mol of H2 that reacts?
11.1 grams H2
b) Assuming H2 gas was the entire source of H and all of it was converted to
water, how many L of H2 gas reacted assuming the reaction was carried out at
STP conditions?
124.4 L H2
c) How many molecules of H2 reacted, assuming the reaction was carried out at
STP conditions?
3.3 x 1024 molecules H2
8. Calculate the following quantities for 343 mL of a 1.27M Na2SO4 solution.
a) Moles of Na2SO4
0.436 moles
b) grams of Na2SO4 required to prepare the solution
61.86 grams
b) moles of Na+
0.872 moles
9. What is the molarity of a 750 mL solution containing 50.0 g KCl?
0.894 M
10. a) How many moles of hydroxide ions are present in 300mL of a 2.50M Ca(OH)2
solution?
1.5 moles
b) What is the molarity of hydroxide ions?
5M