Atoms and Periodic Trends for the MCAT: Everything You Need to Know
Learn essential MCAT chemistry topics on atoms and periodic trends, including atomic radius, ionization energy, electronegativity, and MCAT-style practice questions.
(Note: This guide is part of our MCAT General Chemistry series.)
Table of Contents
Part 1: Introduction to atoms and the periodic table
Part 2: Electronic structure of atoms
a) Atomic structure
b) Electron configuration
c) Electron spin
d) Quantum mechanical model
Part 3: The Periodic Table
a) Groups
b) Periodic trends
Part 4: Atomic phenomena
a) Electron emission
b) Radioactive decay
Part 5: High-Yield Terms
Part 6: Passage-Based Questions and Answers
Part 7: Standalone Questions and Answers
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Part 1: Introduction to atoms and the periodic table
Atoms are the fundamental unit of matter. As a result, atomic structure and the periodic table form the basis of many general chemistry and organic chemistry MCAT questions.
Several important terms in this guide are bolded. As you work through this guide, we encourage you to create definitions and examples that make the most sense to you. At the end of this guide, there are also MCAT-style practice questions that will test your knowledge on this subject.
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Part 2: Electronic structure of atoms
Atoms are the smallest unit of matter that can retain a unique identity. This is the result of the structure of individual atoms, which are composed of varying subatomic particles.
a) Atomic structure
Protons, neutrons, and electrons are the subatomic particles that make up atoms. A proton is a subatomic particle that contains a positive charge (+1e). “e” refers to the fundamental unit of charge and one atomic mass unit (AMU). A neutron is a subatomic particle that contains no charge and has a mass of one AMU. Groups of protons and neutrons—collectively referred to as nucleons—form the nucleus of an atom.
An electron is a subatomic particle that contains a negative charge (-1e) and has a mass that is 1823x smaller than that of a proton, or 1/1823 atomic mass units. In other words, when comparing the mass of a proton or neutron to an electron, the mass of an electron is negligible. Electrons can be arranged in shells around an atom. The electrons in the outer shell of an atom are called valence electrons. Valence electrons play an important role in chemical reactions and bonding.
Table 1 The key subatomic particles
| Subatomic particle | Charge | Mass |
|---|---|---|
0 AMU |
Note that the number of protons determines the element a specific atom is. For example, any atom containing 15 protons will be phosphorus, despite the number of neutrons that the atom has. This number of protons is referred to as the atomic number.
Atoms of the same atomic number but with differing numbers of neutrons are isotopes. Consider this: one isotope of phosphorus might consist of 16 neutrons and 15 protons, while another phosphorus isotope contains 17 neutrons and 15 protons. Both isotopes possess the same atomic number and are considered to be phosphorus but possess differing amounts of neutrons.
Figure 1 Two possible isotopes of phosphorus.
As a result, two isotopes of an element will possess differing atomic mass: a measure of mass that is calculated by summing the number of protons and neutrons within an atom. The atomic mass of many isotopes is used to calculate atomic weight: a weighted average of masses of all naturally occurring isotopes.
- Atomic number (Z): the number of protons in an atom, which determines the elemental identity of the isotope (e.g., any isotopes with 15 protons must be phosphorus)
- Mass number (A): the number of protons and neutrons in an atom (e.g., 30P contains 15 protons and 15 neutrons, resulting in an atomic mass of 30)
- Atomic weight: the weighted average of the masses of naturally occurring isotopes (e.g., if half of all naturally occurring isotopes are 30P and the remaining half is 31P, the atomic weight would be the average of the two mass numbers multiplied by their abundance. Since both isotopes are equally abundant, the atomic weight is 30.5 AMU.)
How is it possible that so many protons with positive charges are held together so tightly in the nucleus? When considering that like charges repel, shouldn’t the positive charges of the protons cause the nucleus to rip apart?
The strong nuclear force is a fundamental force that prevents this. In particular, it is an attractive force that holds protons and neutrons together—despite any alike positive charge.
b) Electron configuration
How are subatomic particles arranged to form an atom? While protons and neutrons are held together in the nucleus by the strong nuclear force, the arrangement of electrons outside the nucleus is much more complex.
An electron configuration is written by determining the number of electrons that are present in the species at hand. For most atoms, or neutrally charged particles, this number of electrons is equal to the atomic number of the element. (Why is this? Recall that the charge of one electron is equal to the charge of one proton. Thus, if a particle has zero charge, the number of electrons and number of protons must be equal). For ions, or particles with a net positive or negative charge, the number of electrons may not be equal to the atomic number of the element.
When “filling” in an atom, electrons are placed in concentric shells surrounding an atom. While these shells are not physical structures, they are useful in visualizing the spaces that electrons may occupy.
The number of concentric shells is determined by the group number of an atom. Thus, elements in the first row of the periodic table have one electron shell, elements in the second row have two electron shells, and so forth. As a general rule, electron shells of a higher number are located further from the nucleus. The outermost electron shell, or highest-numbered electron shell, is located the furthest. This outermost shell is also called the valence shell.
Valence electrons are electrons on the outermost shell of an atom. These electrons are responsible for the reactivity and individual properties of atoms. Each atom is at its most energetically stable when its valence shell is full of electrons. Some elements, like noble gases, already have a full shell.
Each shell is composed of orbitals: regions of certain shapes in which electrons are likely to be found. While there are many orbital shapes that are possible, the MCAT will focus on three particular orbitals denoted as s, p, and d. Elements can be categorized into s-block, p-block, and d-block elements based on their highest-energy orbital.
Figure 2 Elements categorized by outermost valence electron orbitals.
Categorizing elements based on their highest-energy orbital loosely correlates with the type of element each is. Most s-block elements and f-block elements are metals. The p-block elements are a mixture of metalloids and nonmetals.
There are two important considerations in determining the electron configuration of a neutral atom:
Determine the group number of the atom. Determine the group number of the atom by identifying the column of the periodic table the atom is present in. For instance, the halogens are in the seventeenth group, or column, or the periodic table.
Figure 3 Elements in the seventh group.
2. Understand the path of filling electrons in shells. Electrons are filled in shells based on a predetermined order.
All filling begins in the first “subshell,” or the first orbital (e.g., that of hydrogen).
Filling continues to the right of the periodic table. When one row, or family, is completed, filling continues with the leftmost element of the succeeding row.
The Aufbau principle states that each orbital in a shell must be sequentially filled before electrons can begin to fill the next shell.
Filling is concluded when the appropriate number of electrons is placed into orbitals. For neutrally charged atoms, this number of electrons is equal to the atomic number of the element.
Figure 4 Guidelines for filling electron configurations.
- Electrons are always placed in lower energy shells first, so we begin at the top left of the periodic table with the 1s orbital.
- For every increase in the group number, we have an additional electron that we need to account for. Thus, the electron configuration of hydrogen is 1s1, and helium is 1s2.
- After filling the first row, move to row 2. Lithium is 1s22s1.
Electron configurations can also be written using noble gas configuration. This is a form of shorthand for electron configuration, in which writing the electron configuration begins at the closest preceding noble gas. Thus, the electron configuration of lithium can also be written as: [He]2s1.
Filling of the s and p orbitals is fairly straightforward. The f-block elements, or metals, require slightly more consideration. Let’s consider the electron configuration of iron (Fe).
- Using noble gas configuration, we can start with the identity of the closest preceding noble gas: [Ar].
- Iron is located in group number 8. Thus, we must fill in 8 electrons.
- Using the path of electron filling, the 4s2 must be filled first using two electrons.
- Six electrons must be placed in the next orbital. While in the first three rows the s-block filling was followed by p-block filling, the fourth-row elements also contain a d-orbital. Perhaps confusingly, these are of a lower shell number.
- Thus, the six electrons are placed in the 3d6.
- The electron configuration for Fe is [Ar]4s23d6.
c) Electron spin
Electrons possess a characteristic called spin. An electron’s spin can be one of two values: spin up or spin down.
Within the 1s orbital, there are two electrons. Each electron in this pair must have opposite spin. Within larger orbitals—such as in p or d orbitals—Hund’s rule states that each orbital is occupied by a single electron before any orbital is occupied by two electrons. In addition, each of the single electrons will have the same spin.
Figure 5 An application of Hund's rule in filling orbitals with electrons. Note that each 2p orbital is filled with single electrons of the same spin before electrons are paired together.
The balance between spin up and spin down electrons will affect the magnetic properties of an element. If an atom’s electrons are all paired (e.g., in pairs of one spin up and one spin down electron), it is considered to be diamagnetic. Diamagnetic materials do not interact strongly with magnets.
On the other hand, paramagnetic atoms have unpaired electrons. This is a result of possessing more spin up electrons than spin down electrons. Paramagnetic materials are usually strongly affected by nearby magnetic fields, but do not possess a permanent magnetic charge.
Figure 6 Helium is diamagnetic, while lithium is paramagnetic.
Ferromagnetic materials are paramagnetic materials with an additional property: the electron spins of unpaired electrons can spontaneously align in the same direction. As a result, the ferromagnetic material creates its own magnetic field. This is the case for most magnetic materials you may encounter in everyday life, such as refrigerator magnets.
d) Quantum mechanical model
Imagine you are building a coffee table, with hundreds of screws available at your disposal. While the screws perform the same function (to hold the table together), no two screws can occupy the same position on the coffee table at once. (This is stated by the Pauli exclusion principle, which states that no two electrons can contain the same four quantum numbers). If someone asked you to describe a single screw in your table, it would be very difficult. You could discuss the length of the screw, the angle of orientation of the screw, the position in the 3D space of the screw, and much more, but that would take a lot of words to explain. What if there was an easy system to describe the placement of these screws?
While we don’t have a simplified system to describe the positions of the screws in a coffee table, the quantum mechanical model can describe the position and orbit of electrons in an atom through four quantum numbers:
Table 2 Quantum numbers used to describe individual electrons
| Quantum number | Range | Notes |
|---|---|---|
Determines the type of orbital ℓ = 0 → s orbital ℓ = 1 → p orbital ℓ = 2 → d orbital |
||
Determines the spin orientation of an electron +1/2 → spin up -1/2 → spin down |
Let’s apply this system to naming an electron: in this instance, the one and only electron that a neutral hydrogen atom possesses. In this case:
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Part 3: The Periodic Table
a) Groups
Understanding trends in the periodic table can be extraordinarily helpful on the MCAT. You will be provided with a periodic table on the exam, and you may be asked to use it as a tool to apply your knowledge of periodic trends to compare the properties of elements.
One of the most basic and important trends is that between groups. The periodic table is organized into 18 vertical columns, each known as a group. Each group in a periodic table contains elements with a different number of valence electrons. Group 2 elements possess 2 valence electrons, group 13 elements possess 3 valence electrons, and so on.
Elements in a group have similar atomic arrangements and thus properties. For instance, elements in group 18 (the noble gases) each possess a full outer shell of valence electrons. This means that the noble gases aren’t reactive, a trend we see as we go down the group.
Group 1 elements, the alkali metals, possess only one valence electron and are highly reactive. Group 2 elements, the alkaline earth metals, are also quite reactive. As a general rule of thumb, the elements at the far left of the periodic table are highly reactive compared to elements in the center of the periodic table.
The halogens are located in group 17 of the periodic table. All of these elements are found in a diatomic state, meaning that there are two atoms in a single molecule. Since these elements require only one additional valence electron to fulfill the octet rule, they are highly reactive and are often found in salt compounds.
The chalcogens are located in group 16 of the periodic table. They are also referred to as the oxygen group, since this group contains the element oxygen.
The transition metals are found in groups 3-12 of the periodic table. They may also be referred to as the d-block elements, since their valence shell is a d subshell.
The representative elements are found in all other groups of the periodic table. As a result, these include both the s- and p-block elements.
b) Periodic trends
The Periodic Table is cleverly organized such that it documents several trends, or changes in properties, of each element. It will be important to understand each of these periodic trends and be able to apply them well on Test Day.
Electronegativity refers to the attractive force the nucleus of an atom exerts to attract an electron. Recall that elements are at their most stable when the outermost electron shell is filled. Thus, elements with nearly full valence shells have a higher electronegativity compared to elements with only one or two electrons in the valence shell. Energetically speaking, atoms with nearly full electron shells can gain stability by attracting one last electron, while elements with only a handful of valence electrons can most easily gain stability by releasing these electrons. Thus, the electronegativity of elements increases upwards and across to the right of the periodic table. Fluorine (F), a halogen, is highly electronegative, whereas barium (Ba), an alkaline earth metal, is not very electronegative. The exception, of course, is the noble gases. Due to their high stability, the noble gases have nearly zero electronegativity.
Ionization energy refers to the amount of energy needed to remove an electron from the valence shell of an atom. Let’s take two atoms on the opposite sides of the periodic table: potassium (K) and chlorine (Cl). Between the two, chlorine has a higher ionization energy of its valence shell compared to potassium. Can you think of why that might be?
A neutral chlorine atom’s valence shell is filled with 7 electrons. The valence shell is almost full. From a stability standpoint, it is extraordinarily favorable to gain one last electron and complete the valence shell. On the other hand, potassium has one lone electron in the valence shell. From a stability standpoint, it is extraordinarily favorable to lose this electron—thus reverting the valence shell to the full electron shell of one lower period. Thus, ionization energy increases upwards and across to the right of the periodic table.
Note that the second ionization energy of an atom is significantly higher than that of the first ionization energy. This is because more energy is required to ionize a charged particle in comparison to ionizing a neutral particle.
Electron affinity refers to the amount of energy released by an electron when it attracts an electron. Electron affinity is closely related to ionization energy and electronegativity. Electron affinity increases upwards and across to the right of the periodic table. This is because attractive forces between the nucleus and valence electrons increase as distance to the valence shell decreases, and the addition of protons to a nucleus while holding this distance constant will increase the electrostatic attraction.
One final important trend to understand is that of atomic size. Atomic radius decreases upwards and across to the right of the periodic table. This is primarily due to a phenomenon known as effective nuclear charge.
The number of protons in the nucleus of the atom
The presence of shielding electrons
Figure 7 Effective nuclear charge is dependent on the presence of protons and electron shielding.
Ions, or charged atoms, may result in exceptions to the trends of atomic radius. Ions can either be positive or negative depending on whether an electron is gained or lost. Generally, negative ions are larger than the original atom as the attractive forces exerted by the nucleus are dispersed over a greater number of electrons. On the other hand, positive ions are smaller than the original atom because there are more positive nuclear forces to attract electrons toward the center of the atom.
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Part 4: Atomic phenomena
Valence electrons and orbitals are very important for interactions with other atoms, but single atoms can also participate in atomic phenomena. Some of these phenomena include the absorption and emission of light energy and radioactive decay.
For further discussion on these topics, be sure to refer to our guide on atomic and nuclear physics.
a) Electron emission
The Bohr model is a popular way of explaining the structure of an atom. The Bohr model states that electrons orbit around protons in an atom in discrete circular paths (orbitals) dependent on their energy levels. These energy levels are known as quanta, represented with the letter n.
Electrons are able to move between these quanta depending on the amount of energy they absorb or release. Although it’s now known that the Bohr model is not accurate for all atoms, it is a good approximation for the structure and behavior of the hydrogen atom.
When all the electrons of an atom are in the lowest orbitals possible, closest to the atom, the atom is known to be in its ground state. Here, electrons cannot move to any lower energy levels. When electrons of an atom “jump” to a higher orbital level than their ground state, the atom is in an excited state.
How exactly does this electron “jump” occur? When an electron absorbs enough energy to match the energy level of the next orbital, the electron will move to a higher energy level. This is a phenomenon known as the photoelectric effect. Often, electrons will fall back down to a ground state after being in their excited state. When they jump back down, they emit a similar level of energy as they’ve absorbed.
As electrons of an atom drop back from an excited state to their ground state, photons (light particles) can be released. The energy of the photon can be calculated using the following equation:
E = hc/λ
where E = energy released,
h = Planck’s constant,
c = 3.0 × 108 m/s,
λ = wavelength of emitted light
Depending on how much energy is released, the photon can have different wavelengths that we might see. Each atom has a characteristic absorption frequency at which it absorbs energy and a unique emission spectrum that contains the frequencies of the photons that are emitted. This emission spectrum can be used to identify the atom based on changes in electron excited and ground states.
The energy emitted by an electron as it travels down to its ground state can be calculated using the Rydberg equation:
E = hc/λ = RH(1/ni2 − 1/nf2)
where RH = the Rydberg unit of energy,
ni = initial orbital level,
nf = final orbital level
While it’s possible to calculate the energy of a photon, it is much more difficult to calculate other characteristics of the photon. The Heisenberg uncertainty principle states that both the position and velocity of a photon cannot be known at the same time.
b) Radioactive decay
Isotopes of an element vary in their stability. As a general rule of thumb, isotopes with an amount of neutrons approximately equal to the atomic number are more stable than their lighter variants.
Unstable isotopes can spontaneously undergo decay processes in which subatomic particles are ejected, releasing energy. There are three primary forms of radioactive decay: alpha, beta, and gamma radiation.
When solving radioactive decay problems, it is important to keep the atomic mass numbers and atomic numbers balanced across the equation. Let’s look at an example of 3015P undergoing alpha decay:
To solve this equation, we must satisfy the balance of both the top and bottom numbers. Once we have solved for atomic number, we can further use this information to identify the resulting element. As a result:
Beta decay occurs when an atom releases a beta particle, which takes two forms: 0-1e (an electron) or 0+1e(positron). We can also write 0-1e as 0-1β and write 0+1e as 0+1β. What happens when 3015P undergoes electron decay?
In addition to beta decay, there are also forms of beta particle capture, such as electron capture or positron capture. Here’s an example of electron capture:
Finally, gamma decay occurs when an atom releases a gamma particle, which looks like this: 00γ. A gamma particle is massless but contains high amounts of energy. As a result, it is the most destructive form of radioactive decay we have discussed thus far. Here is an example of gamma decay:
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Part 5: High-yield terms
Atomic number: the number of protons in an atom
Atomic radius: the distance from the center of an atom to the edge of its electron cloud
Atomic weight: the weighted average of the masses of naturally occurring isotopes
Bohr model: explains that electrons move around protons in an atom in discrete circular paths (orbitals) dependent on their energy levels
Diamagnetic: atoms have all paired electrons; do not interact strongly with magnets
Effective nuclear charge: a measure of the net positive charge that electrons in an atom feel
Electron: a subatomic particle that contains a negative charge (-1e) and has a mass that is 1823x smaller than that of a proton, or 1/1823 atomic mass units
Electron affinity: the amount of energy released by an electron when it attracts an electron
Electron capture/positron capture: when an atom captures a positron, resulting in one fewer proton
Electron configuration: conventional nomenclature for electronic structure that is used in general chemistry
Electronegativity: the attractive force an atom exerts to attract an electron
Emission spectrum: characteristic wavelengths emitted by an atom
Excited state: when electrons of an atom jump to a higher orbital level than normal
Gamma decay: when an atom releases a gamma particle
Ground state: when all the electrons of an atom are in the lowest orbitals possible, closest to the atom
Group: the vertical organization of elements in a periodic table; elements in a group have similar chemical properties
Half life: the time it takes for the atom to decay to half its original mass
Hund’s rule: states that each orbital is occupied by a single electron before any orbital is occupied by two electrons; each of the single electrons will have the same spin
Ionization energy: the amount of energy needed to remove an electron from the valence shell of an atom
Isotope: atoms of the same element that contain the same number of protons but a different number of neutrons
Mass number: the number of protons and neutrons in an atom
Neutron: a subatomic particle that contains no charge and has a mass of one atomic mass unit
Nucleus: the core of an atom consisting of protons and neutrons
Paramagnetism: refers to atoms that have unpaired electrons; strongly attracted to magnets
Pauli exclusion principle: states that no two electrons can contain the same four quantum numbers
Proton: a subatomic particle that contains a positive charge (+1e)
Quantum mechanical model: a system to describe an electron in an atom with four quantum numbers: n, ℓ, ml, and ms
Strong nuclear force: an attractive force that holds protons and neutrons together
Subatomic particles: particles such as neutrons, electrons, and protons that make up an atom
Valence electron: electrons on the outer shell of an atom
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Part 6: Passage-based questions and answers
While performing a synthesis experiment, chemists have come across an unknown element. To determine the identity of the element, the researchers perform a variety of tests to discern its chemical properties.
In the first experiment, the chemists bombard the element with radioactive particles to observe any possible decay patterns. The element was observed to undergo radioactive decay, which resulted in the production of an electron and neon gas.
In a second experiment, various reactivity tests are performed. It is discovered that the unknown element is highly reactive with alkali metals.
In a final experiment, the researchers discovered that the element has a high level of electronegativity and ionization energy.
Question 1: In what group of the periodic table is the element likely found?
A) Alkali metals
B) Alkaline earth metals
C) Halogens
D) Noble gases
Question 2: The type of radioactive decay described in the passage is an example of:
A) Alpha decay
B) Beta decay
C) Gamma decay
D) Positron emission
Question 3: Which of the following is a possible set of quantum numbers for the unidentified element?
A) 0, 1, 1, +1/2
B) 1, 1, 0, +1/2
C) 2, 0, 2, +1/2
D) 2, 1, 0, +1/2
Question 4: How does the atomic radius of this element compare to others in the same group and row of the periodic table?
A) Larger compared to others in its group; larger compared to others in its row
B) Larger compared to others in its group; smaller compared to others in its row
C) Smaller compared to others in its group; smaller compared to others in its row
D) Smaller compared to others in its group; larger compared to others in its row
Answers to passage-based questions
1. Answer choice C is correct. The unknown element readily reacts with alkali metals, which are characterized by one free valence electron that is readily given away. Group 17 elements, referred to as halogens, require one valence electron to complete the outer shell and thus readily react with alkali metals (choice C is correct). The alkaline earth metals are found in group 2 (choice B is incorrect). The noble gases are found in group 18 and are unreactive (choice D is incorrect).
2. Answer choice B is correct. Beta decay results in the emission of an electron and gaining one proton (choice B is correct). Alpha decay results in the release of an alpha particle composed of two protons and two neutrons (choice A is incorrect). Gamma decay results in the emission of a photon (choice C is incorrect). While positron emission is a form of beta decay, this scenario describes electron emission (choice D is incorrect).
3. Answer choice D is correct. Based on the information presented about the element’s radioactive decay, the unknown element is most likely fluorine (atomic number 9). The first quantum number, n, is equivalent to the periodic table row the element is located in. Thus, we would expect fluorine’s first quantum number to be 2 (choices A and B are incorrect). The second quantum number, l, ranges from 0 to n-1, meaning that the second quantum number for fluorine must be either 0 or 1 (choice C is incorrect). The third quantum number, ml, ranges from l to +l, meaning that it could be either -1 or +1 (choice D is correct).
4. Answer choice C is correct. Fluorine is located to the far right of the periodic table and in the first few rows. The atomic radius should be smaller compared to others in its group, as elements further down the periodic table possess more electron orbitals and thus are larger (choices A and B are incorrect). The atomic radius is also smaller compared to others in its row, as the effective nuclear charge (Z) increases as protons are added to the nucleus (choice C is correct).
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Part 7: Standalone questions and answers
Question 1: Which of the following represents a naturally-occurring isotope of oxygen?
A) 178O
B) 177O
C) 88O
D) 16.58O
A) Alpha decay
B) Electron emission
C) Positron emission
D) Gamma decay
Question 3: Which of the following best describes the difference between the Pauli Exclusion Principle and Hund’s Rule?
A) The Pauli Exclusion Principle states that all of the orbitals must be filled by a single electron before they pair up while Hund’s Rule describes the spin of the electrons.
B) The Pauli Exclusion Principle states that all of the orbitals must be filled by a single electron before they pair up while Hund’s Rule states that no electrons in an atom can have the same four quantum numbers.
C) The Pauli Exclusion Principle states that no electron in an atom can have the same four quantum numbers, while Hund’s Rule states that all of the orbitals must be filled by a single electron before they pair up.
D) The Pauli Exclusion Principle and Hund’s Rule are the same and describe the different ways of identifying electrons.
Question 4: Which of the following elements is the most electronegative?
A) O
B) Ne
C) F
D) Cl
Question 5: Which of the following elements has the largest atomic radius?
A) F
B) Li
C) K
D) Br
Answers to standalone questions
1. Answer choice A is correct. Oxygen has an atomic number of 8 (choice B is incorrect). The atomic mass number of an individual isotope must be an integer value (choice D is incorrect). It is highly unlikely that an isotope of oxygen could have an atomic mass of 8; this would mean the nucleus has 8 protons and zero neutrons! Such an imbalanced nucleus would be highly unstable (choice C is incorrect).
2. Answer choice B is correct. Recall that alpha decay results in the emission of two protons and two neutrons, resulting in a change of the atomic mass (choice A is incorrect). Gamma decay results in the release of a photon of light but does not change the identity of the atom (choice D is incorrect). Positron emission and electron emission are both forms of beta decay; however, only electron emission results in an increase of atomic number (choice B is correct).
3. Answer choice C is correct. The Pauli Exclusion Principle states that no two electrons can contain the same four quantum numbers (choices A and B are incorrect). Hund’s Rule states that each orbital is occupied by a single electron before any orbital is occupied by two electrons (choice C is correct).
4. Answer choice C is correct. Electronegativity increases upwards and to the right of the periodic table. Thus, the most electronegative element should be to the right of the periodic table (choice A is incorrect). The rightmost elements, the noble gases, are not very electronegative as they have full valence shells and are thus quite stable (choice B is incorrect). Fluorine must be more electronegative than chlorine, as fluorine is located in the row above chlorine (choice C is correct).
5. Answer choice C is correct. Atomic radius decreases upwards and to the right of the periodic table because attraction between protons and electrons increases. Thus, the largest element must be located closest to the bottom of the periodic table and furthest to the left (choice C is correct).
