Portage Learning CHEM 103 General Chemistry I Module 2 Exam Questions
with Correct Verified Answers - 2026/2027
1. What are the three main subatomic particles that make up an atom? Protons, neutrons, and electrons.
2. What is the charge and approximate mass (in amu) of a proton? Charge of +1, mass of approximately 1 amu.
3. What is the charge and approximate mass of a neutron? No charge (neutral), mass of approximately 1 amu.
4. What is the charge and approximate mass of an electron? Charge of -1, mass of approximately 1/1836 amu
(negligible compared to protons/neutrons).
5. Where are protons and neutrons located within an atom? In the nucleus, at the center of the atom.
6. Where are electrons located within an atom? In the electron cloud/orbitals surrounding the nucleus.
7. What determines the identity of an element? The number of protons (atomic number) in the nucleus.
8. What is the atomic number of an element? The number of protons in the nucleus of an atom of that
element.
9. What is the mass number of an atom? The sum of the number of protons and neutrons in the nucleus.
10. How do you calculate the number of neutrons in an atom, given the mass number and atomic number?
Number of neutrons = mass number - atomic number (number of protons).
11. What are isotopes? Atoms of the same element (same number of protons) that have different numbers of
neutrons, and therefore different mass numbers.
12. How is atomic mass (as shown on the periodic table) related to isotopes of an element? Atomic mass is the
weighted average of the masses of all naturally occurring isotopes of that element, based on their relative
abundance.
13. What is the charge of an atom that has gained one electron? A charge of -1 (an anion).
14. What is the charge of an atom that has lost two electrons? A charge of +2 (a cation).
15. What term describes a positively charged ion? A cation.
16. What term describes a negatively charged ion? An anion.
17. How many electrons does a neutral atom of an element have? The same number as its atomic number
(equal to the number of protons).
18. What is the atomic symbol notation, including mass number and atomic number, used to represent a specific
isotope? Mass number written as a superscript and atomic number as a subscript to the left of the element
symbol (e.g., carbon-14 written with mass number 14 and atomic number 6).
19. What historical model of the atom described it as a 'plum pudding' with electrons embedded in a positive
sphere? J.J. Thomson's plum pudding model.
20. What experiment led Ernest Rutherford to propose that atoms have a small, dense, positively charged nucleus?
The gold foil experiment, in which most alpha particles passed through the foil but some were deflected at
large angles.
21. What is the relationship between the wavelength and frequency of electromagnetic radiation? Wavelength
and frequency are inversely proportional; as wavelength increases, frequency decreases, related by the
equation c = λν (speed of light equals wavelength times frequency).
22. What is the equation relating the energy of a photon to its frequency? E = hν, where E is energy, h is
Planck's constant, and ν is frequency.
23. As the frequency of light increases, what happens to the energy of its photons? The energy increases, since
energy is directly proportional to frequency.
24. What color of visible light has the shortest wavelength? Violet.
, 25. What color of visible light has the longest wavelength? Red.
26. Which type of electromagnetic radiation has a higher energy: ultraviolet light or infrared light? Ultraviolet
light, because it has a shorter wavelength and higher frequency than infrared light.
27. What does the Bohr model propose about the location and energy of electrons in an atom? Electrons orbit the
nucleus in fixed, quantized energy levels (orbits), and can only exist at these specific discrete energy levels.
28. What happens when an electron absorbs energy and moves to a higher energy level? The electron becomes
'excited' and moves from its ground state to a higher (excited) energy level.
29. What happens when an excited electron returns to a lower energy level? It releases energy in the form of a
photon of light, corresponding to the energy difference between the two levels.
30. What produces the distinct line spectrum observed for hydrogen and other elements? Electrons transitioning
between specific quantized energy levels, each transition releasing a photon of a specific wavelength/energy.
31. What is the ground state of an electron? The lowest possible energy level/state that an electron can occupy
in an atom.
32. What is an excited state of an electron? Any energy level higher than the ground state that an electron
temporarily occupies after absorbing energy.
33. Why is the Bohr model considered insufficient to fully describe atoms with more than one electron? It does
not account for electron-electron interactions and fails to accurately predict spectra for multi-electron
atoms; it was later replaced by the quantum mechanical model.
34. What does the quantum mechanical model describe regarding the position of an electron in an atom? It
describes the probability of finding an electron in a particular region of space (an orbital), rather than a
precise, fixed path as in the Bohr model.
35. What is an atomic orbital? A three-dimensional region of space around the nucleus where there is a high
probability of finding an electron.
36. What shape is characteristic of an s orbital? Spherical.
37. What shape is characteristic of a p orbital? Dumbbell-shaped (two lobes).
38. How many orbitals are found within a p sublevel? Three (px, py, and pz).
39. How many orbitals are found within a d sublevel? Five.
40. How many electrons can a single atomic orbital hold at most? Two electrons, with opposite spins.
41. What principal quantum number (n) describes? The main energy level (shell) of an electron, with higher
numbers indicating higher energy and greater average distance from the nucleus.
42. What does the angular momentum (azimuthal) quantum number (l) describe? The shape of the
orbital/sublevel (e.g., l = 0 for s, l = 1 for p, l = 2 for d, l = 3 for f).
43. What does the magnetic quantum number (ml) describe? The specific orientation of an orbital in space
within a given sublevel.
44. What does the spin quantum number (ms) describe? The direction of the electron's intrinsic spin, either
+1/2 or -1/2.
45. What does the Pauli exclusion principle state? No two electrons in the same atom can have the same set of
all four quantum numbers; therefore, two electrons in the same orbital must have opposite spins.
46. What does Hund's rule state regarding filling degenerate orbitals (orbitals of equal energy)? Electrons fill
degenerate orbitals singly (one electron per orbital) before any orbital is doubly occupied, and unpaired
electrons have parallel spins.
47. What does the Aufbau principle state? Electrons fill atomic orbitals starting from the lowest available
energy level before filling higher energy levels.
with Correct Verified Answers - 2026/2027
1. What are the three main subatomic particles that make up an atom? Protons, neutrons, and electrons.
2. What is the charge and approximate mass (in amu) of a proton? Charge of +1, mass of approximately 1 amu.
3. What is the charge and approximate mass of a neutron? No charge (neutral), mass of approximately 1 amu.
4. What is the charge and approximate mass of an electron? Charge of -1, mass of approximately 1/1836 amu
(negligible compared to protons/neutrons).
5. Where are protons and neutrons located within an atom? In the nucleus, at the center of the atom.
6. Where are electrons located within an atom? In the electron cloud/orbitals surrounding the nucleus.
7. What determines the identity of an element? The number of protons (atomic number) in the nucleus.
8. What is the atomic number of an element? The number of protons in the nucleus of an atom of that
element.
9. What is the mass number of an atom? The sum of the number of protons and neutrons in the nucleus.
10. How do you calculate the number of neutrons in an atom, given the mass number and atomic number?
Number of neutrons = mass number - atomic number (number of protons).
11. What are isotopes? Atoms of the same element (same number of protons) that have different numbers of
neutrons, and therefore different mass numbers.
12. How is atomic mass (as shown on the periodic table) related to isotopes of an element? Atomic mass is the
weighted average of the masses of all naturally occurring isotopes of that element, based on their relative
abundance.
13. What is the charge of an atom that has gained one electron? A charge of -1 (an anion).
14. What is the charge of an atom that has lost two electrons? A charge of +2 (a cation).
15. What term describes a positively charged ion? A cation.
16. What term describes a negatively charged ion? An anion.
17. How many electrons does a neutral atom of an element have? The same number as its atomic number
(equal to the number of protons).
18. What is the atomic symbol notation, including mass number and atomic number, used to represent a specific
isotope? Mass number written as a superscript and atomic number as a subscript to the left of the element
symbol (e.g., carbon-14 written with mass number 14 and atomic number 6).
19. What historical model of the atom described it as a 'plum pudding' with electrons embedded in a positive
sphere? J.J. Thomson's plum pudding model.
20. What experiment led Ernest Rutherford to propose that atoms have a small, dense, positively charged nucleus?
The gold foil experiment, in which most alpha particles passed through the foil but some were deflected at
large angles.
21. What is the relationship between the wavelength and frequency of electromagnetic radiation? Wavelength
and frequency are inversely proportional; as wavelength increases, frequency decreases, related by the
equation c = λν (speed of light equals wavelength times frequency).
22. What is the equation relating the energy of a photon to its frequency? E = hν, where E is energy, h is
Planck's constant, and ν is frequency.
23. As the frequency of light increases, what happens to the energy of its photons? The energy increases, since
energy is directly proportional to frequency.
24. What color of visible light has the shortest wavelength? Violet.
, 25. What color of visible light has the longest wavelength? Red.
26. Which type of electromagnetic radiation has a higher energy: ultraviolet light or infrared light? Ultraviolet
light, because it has a shorter wavelength and higher frequency than infrared light.
27. What does the Bohr model propose about the location and energy of electrons in an atom? Electrons orbit the
nucleus in fixed, quantized energy levels (orbits), and can only exist at these specific discrete energy levels.
28. What happens when an electron absorbs energy and moves to a higher energy level? The electron becomes
'excited' and moves from its ground state to a higher (excited) energy level.
29. What happens when an excited electron returns to a lower energy level? It releases energy in the form of a
photon of light, corresponding to the energy difference between the two levels.
30. What produces the distinct line spectrum observed for hydrogen and other elements? Electrons transitioning
between specific quantized energy levels, each transition releasing a photon of a specific wavelength/energy.
31. What is the ground state of an electron? The lowest possible energy level/state that an electron can occupy
in an atom.
32. What is an excited state of an electron? Any energy level higher than the ground state that an electron
temporarily occupies after absorbing energy.
33. Why is the Bohr model considered insufficient to fully describe atoms with more than one electron? It does
not account for electron-electron interactions and fails to accurately predict spectra for multi-electron
atoms; it was later replaced by the quantum mechanical model.
34. What does the quantum mechanical model describe regarding the position of an electron in an atom? It
describes the probability of finding an electron in a particular region of space (an orbital), rather than a
precise, fixed path as in the Bohr model.
35. What is an atomic orbital? A three-dimensional region of space around the nucleus where there is a high
probability of finding an electron.
36. What shape is characteristic of an s orbital? Spherical.
37. What shape is characteristic of a p orbital? Dumbbell-shaped (two lobes).
38. How many orbitals are found within a p sublevel? Three (px, py, and pz).
39. How many orbitals are found within a d sublevel? Five.
40. How many electrons can a single atomic orbital hold at most? Two electrons, with opposite spins.
41. What principal quantum number (n) describes? The main energy level (shell) of an electron, with higher
numbers indicating higher energy and greater average distance from the nucleus.
42. What does the angular momentum (azimuthal) quantum number (l) describe? The shape of the
orbital/sublevel (e.g., l = 0 for s, l = 1 for p, l = 2 for d, l = 3 for f).
43. What does the magnetic quantum number (ml) describe? The specific orientation of an orbital in space
within a given sublevel.
44. What does the spin quantum number (ms) describe? The direction of the electron's intrinsic spin, either
+1/2 or -1/2.
45. What does the Pauli exclusion principle state? No two electrons in the same atom can have the same set of
all four quantum numbers; therefore, two electrons in the same orbital must have opposite spins.
46. What does Hund's rule state regarding filling degenerate orbitals (orbitals of equal energy)? Electrons fill
degenerate orbitals singly (one electron per orbital) before any orbital is doubly occupied, and unpaired
electrons have parallel spins.
47. What does the Aufbau principle state? Electrons fill atomic orbitals starting from the lowest available
energy level before filling higher energy levels.
