TEAS: Science Quiz: Identify Atomic Structure
20 questions · exam conditions
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Identify Atomic StructureQuestion 1 of 20

Which subatomic particle has a mass of approximately 1 atomic mass unit (amu) and no charge?

Proton
Electron
Neutron
Photon
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TEAS: Science Quiz

TEAS: Science Quiz: Identify Atomic Structure

Practice Identify Atomic Structure in TEAS: Science with focused quiz questions that help you check what you know, review explanations, and build confidence with test-style prompts.

What this quiz covers

This quiz focuses on Identify Atomic Structure, giving you a quick way to practice the rules, question types, and explanations that matter most for TEAS: Science.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

All questions

Question 1

Which subatomic particle has a mass of approximately 1 atomic mass unit (amu) and no charge?

  1. Proton
  2. Electron
  3. Neutron (correct answer)
  4. Photon
Explanation: This question tests your knowledge of the fundamental properties of subatomic particles—the building blocks of atoms. When you encounter questions about atomic structure, focus on the key characteristics: mass, charge, and location within the atom. The neutron is the subatomic particle that fits both criteria perfectly. Neutrons have a mass of approximately 1 amu (actually 1.008 amu, very close to 1) and carry no electrical charge, making them electrically neutral. They reside in the atom's nucleus alongside protons and contribute significantly to an atom's total mass. Let's examine why the other options don't work: Choice A, the proton, does have a mass of approximately 1 amu (1.007 amu), but it carries a positive charge (+1), so it fails the "no charge" requirement. Choice B, the electron, is electrically neutral with no charge, but its mass is extremely small—only about 1/1836 of an amu, nowhere near 1 amu. Choice D, the photon, has no charge but also has no rest mass at all; photons are massless particles of light. For TEAS success, memorize this simple pattern: protons and neutrons both have masses around 1 amu and live in the nucleus, but protons are positive while neutrons are neutral. Electrons are much lighter and orbit the nucleus. This fundamental distinction appears frequently on the exam, so make sure you can quickly recall each particle's mass, charge, and location.

Question 2

Which of the following electron configurations represents an atom in its ground state?

  1. 1s22s22p53s11s^2 2s^2 2p^5 3s^1 for an atom with 10 electrons
  2. 1s22s22p63s11s^2 2s^2 2p^6 3s^1 for an atom with 11 electrons (correct answer)
  3. 1s22s12p63s21s^2 2s^1 2p^6 3s^2 for an atom with 11 electrons
  4. 1s12s22p63s21s^1 2s^2 2p^6 3s^2 for an atom with 11 electrons
Explanation: Ground state configurations fill orbitals in order of increasing energy. For 11 electrons, the correct configuration is 1s22s22p63s11s^2 2s^2 2p^6 3s^1, filling 1s, then 2s, then 2p completely before 3s. Choice A shows an excited state with 2p not filled before 3s. Choices C and D violate the aufbau principle by not filling lower energy orbitals first.

Question 3

An element in Group 2A (alkaline earth metals) would most likely form an ion with which charge?

  1. +1+1 because it needs to lose one electron to achieve noble gas configuration
  2. +2+2 because it needs to lose two electrons to achieve noble gas configuration (correct answer)
  3. −2-2 because it needs to gain two electrons to achieve stable octet configuration
  4. −1-1 because it has one electron more than the previous noble gas
Explanation: Group 2A elements have 2 valence electrons and achieve noble gas configuration by losing both electrons, forming +2 ions. Choice A describes Group 1A behavior. Choice C would require gaining 6 electrons to complete the p subshell, which is energetically unfavorable. Choice D incorrectly describes the electron relationship to noble gases.

Question 4

What is the primary factor that determines the chemical properties of an element?

  1. The total number of electrons in all energy levels
  2. The number of electrons in the outermost energy level (correct answer)
  3. The total mass of protons and neutrons in the nucleus
  4. The ratio of neutrons to protons in the nucleus
Explanation: Chemical properties are primarily determined by valence electrons (electrons in the outermost energy level) because these participate in chemical bonding. Choice A includes inner electrons that don't participate in bonding. Choice C describes atomic mass, which doesn't determine chemical properties. Choice D relates to nuclear stability, not chemical behavior.

Question 5

Which of the following atoms would require the least energy to remove one electron?

  1. Lithium (Li) because it has the smallest nuclear charge among the alkali metals
  2. Sodium (Na) because it has optimal size and nuclear charge for easy electron removal
  3. Potassium (K) because its valence electron is farthest from the nuclear attractive force
  4. Rubidium (Rb) because it has the largest atomic radius and weakest nuclear attraction (correct answer)
Explanation: Ionization energy decreases down a group, so rubidium (lowest in Group 1A among the choices) has the lowest ionization energy due to its large size and significant electron shielding. Choice A is incorrect as lithium has the highest ionization energy among alkali metals. Choices B and C have higher ionization energies than rubidium.

Question 6

Based on periodic trends, which of the following elements has the highest electronegativity?

  1. Sodium (Na)
  2. Oxygen (O) (correct answer)
  3. Aluminum (Al)
  4. Sulfur (S)
Explanation: When you encounter electronegativity questions, you're being tested on periodic trends—specifically how an atom's ability to attract electrons in a bond changes across the periodic table. Electronegativity increases as you move from left to right across a period and decreases as you move down a group. Looking at the elements provided, you need to consider their positions on the periodic table. Oxygen sits in period 2, group 16, making it both high up and far to the right. This positioning gives oxygen the highest electronegativity among these choices, with a value of 3.44 on the Pauling scale. Choice A, sodium, sits in period 3, group 1, placing it far to the left where metals have very low electronegativity (0.93). Sodium readily gives up electrons rather than attracting them. Choice C, aluminum, is also a metal in period 3, group 13, with low electronegativity (1.61) since it's still on the left side of the periodic table. Choice D, sulfur, is closer to oxygen in period 3, group 16, giving it relatively high electronegativity (2.58), but it's still lower than oxygen because it's one period down. The key pattern to remember: fluorine has the highest electronegativity of all elements, followed by oxygen, nitrogen, and chlorine. When comparing electronegativity on the TEAS, look for elements that are both high up and far right on the periodic table—these will typically be your answer.

Question 7

Which of the following atoms has the lowest electronegativity?

  1. Chlorine (Cl)
  2. Fluorine (F)
  3. Cesium (Cs) (correct answer)
  4. Oxygen (O)
Explanation: Electronegativity measures an atom's ability to attract electrons when forming chemical bonds. To determine which atom has the lowest electronegativity, you need to understand how this property varies across the periodic table. Electronegativity increases as you move from left to right across a period and decreases as you move down a group. This happens because atoms on the right side of the periodic table have more protons pulling on electrons, while atoms lower on the table have their outer electrons farther from the nucleus due to additional electron shells. Cesium (Cs) has the lowest electronegativity among these options. Located in the bottom-left corner of the periodic table (Period 6, Group 1), cesium has its valence electron very far from the nucleus and experiences significant electron shielding from inner electron shells. This makes it extremely willing to give up electrons rather than attract them. Option A, chlorine (Cl), is incorrect because it's located in Group 17 and has high electronegativity, making it very electron-attracting. Option B, fluorine (F), is wrong because it actually has the highest electronegativity of all elements—it's in the top-right corner of the periodic table. Option D, oxygen (O), is incorrect because it's also highly electronegative, being in Period 2 and Group 16. For TEAS questions about periodic trends, remember that electronegativity follows the pattern: increases left-to-right, decreases top-to-bottom. The most electronegative elements are in the upper right (excluding noble gases), while the least are in the lower left.

Question 8

An element is located in Group 2 of the periodic table. How many valence electrons does a neutral atom of this element possess?

  1. 1
  2. 2 (correct answer)
  3. 8
  4. 10
Explanation: When you encounter questions about valence electrons and periodic table groups, you're being tested on the fundamental relationship between an element's position and its electron configuration. The key insight is that for main group elements (Groups 1, 2, and 13-18), the group number directly tells you the number of valence electrons. Group 2 elements include beryllium, magnesium, calcium, strontium, barium, and radium. Each of these elements has exactly 2 electrons in its outermost shell, which makes them highly reactive metals that readily lose these 2 electrons to form +2 ions. Looking at the incorrect options: Choice A (1 valence electron) describes Group 1 elements like sodium and potassium, not Group 2. Choice C (8 valence electrons) represents the noble gases in Group 18, which have complete outer shells and are chemically inert. Choice D (10 electrons) might seem tempting if you're thinking about total electrons in inner shells, but the question specifically asks about valence electrons—only those in the outermost shell count. The correct answer is B (2 valence electrons) because Group 2 elements always have 2 electrons in their outermost shell. Study tip: Memorize this pattern for TEAS success: Groups 1 and 2 have 1 and 2 valence electrons respectively, while Groups 13-18 have 3-8 valence electrons. This simple rule will help you quickly answer periodic table questions without getting bogged down in complex electron configurations.

Question 9

An element's position in a specific period on the periodic table is determined by which of the following?

  1. The number of principal energy levels (correct answer)
  2. The number of valence electrons
  3. The total mass of the nucleus
  4. The chemical reactivity of the element
Explanation: When you encounter questions about the periodic table's organization, focus on understanding what determines an element's placement both horizontally (periods) and vertically (groups). The periodic table is organized so that elements in the same horizontal row, called a period, all have the same number of principal energy levels (electron shells). As you move from left to right across any period, each element adds one more proton and typically one more electron, but they're filling the same outermost energy level. For example, all elements in Period 2 (lithium through neon) have exactly two principal energy levels, while all Period 3 elements (sodium through argon) have three principal energy levels. Let's examine why the other options don't determine period placement. Option B, the number of valence electrons, actually determines an element's group (vertical column) placement, not its period. Elements in the same group have similar valence electron configurations. Option C, the total mass of the nucleus, doesn't follow the periodic table's organization pattern—mass generally increases across periods but isn't what defines them. Option D, chemical reactivity, varies significantly within the same period; for instance, Period 3 contains highly reactive sodium and unreactive argon. For TEAS questions about periodic trends, remember this key distinction: periods (horizontal rows) are determined by the number of electron shells, while groups (vertical columns) are determined by valence electrons. This fundamental organization principle explains most periodic properties and trends you'll encounter.

Question 10

Which subatomic particles are located within the nucleus and have a positive and a neutral charge, respectively?

  1. Proton and neutron (correct answer)
  2. Electron and neutron
  3. Proton and electron
  4. Neutron and electron
Explanation: When you encounter atomic structure questions, focus on the three key subatomic particles and their locations within the atom. The nucleus sits at the center and contains only two types of particles, while the third type orbits around it. Protons are positively charged particles found in the nucleus, and neutrons are neutral particles (no charge) also located in the nucleus. These two particles make up nearly all of an atom's mass. The question asks specifically for nuclear particles with positive and neutral charges respectively, which perfectly describes protons and neutrons. Looking at the answer choices: A) Proton and neutron correctly identifies both nuclear particles with their respective positive and neutral charges. B) Electron and neutron is wrong because electrons are negatively charged and orbit outside the nucleus in electron shells—they're not nuclear particles. C) Proton and electron incorrectly includes the electron, which again is not in the nucleus and carries a negative charge, not neutral. D) Neutron and electron reverses the charge order and incorrectly includes the electron as a nuclear particle. The key trap here is including electrons, which students sometimes confuse as being "part of the atom's core" when they're actually in the outer regions. Remember that only protons and neutrons live in the nucleus. For TEAS atomic structure questions, memorize this simple framework: nucleus contains protons (positive) and neutrons (neutral), while electrons (negative) orbit outside. This distinction between nuclear and non-nuclear particles appears frequently on the exam.

Question 11

Which of the following is equal to the number of protons in an atom's nucleus and uniquely defines the element?

  1. The mass number
  2. The isotope number
  3. The atomic number (correct answer)
  4. The charge of the ion
Explanation: When you encounter questions about atomic structure, focus on the fundamental properties that define atoms and elements. The key is understanding what makes each element unique on the periodic table. The atomic number is equal to the number of protons in an atom's nucleus and serves as the unique identifier for each element. Every atom of hydrogen has exactly 1 proton (atomic number 1), every carbon atom has 6 protons (atomic number 6), and every oxygen atom has 8 protons (atomic number 8). This number never changes for a given element—it's what makes carbon "carbon" and oxygen "oxygen." Let's examine why the other options don't work. Choice A, the mass number, represents the total number of protons plus neutrons in the nucleus. While this tells you about the atom's mass, it doesn't uniquely identify the element because different isotopes of the same element have different mass numbers. Choice B, "isotope number," isn't a standard term in chemistry—isotopes are identified by their mass numbers, not a separate numbering system. Choice D, the charge of an ion, tells you whether an atom has gained or lost electrons, but ions of different elements can have the same charge, so this doesn't uniquely identify the element. Remember this pattern for TEAS science questions: when asked what uniquely defines or identifies an element, always think "atomic number equals protons." The number of protons is the unchanging signature of each element, regardless of how many neutrons or electrons it might have.

Question 12

Which of the following correctly describes the metallic character trend in the periodic table?

  1. Metallic character increases across a period and decreases down a group
  2. Metallic character decreases across a period and increases down a group (correct answer)
  3. Metallic character remains constant across periods but increases down groups
  4. Metallic character increases both across periods and down groups
Explanation: Metallic character decreases across a period (atoms get smaller, hold electrons tighter) and increases down a group (atoms get larger, lose electrons more easily). Choice A has both trends backwards. Choice C incorrectly states metallic character is constant across periods. Choice D incorrectly suggests metallic character increases in both directions.

Question 13

As you move from left to right across a period in the periodic table, which trend is observed for atomic radius?

  1. Atomic radius increases due to additional electron shells being added to the atoms
  2. Atomic radius decreases due to increasing nuclear charge pulling electrons closer to the nucleus (correct answer)
  3. Atomic radius remains constant because the same number of electron shells are present
  4. Atomic radius fluctuates randomly due to varying electron-electron repulsion forces within the atoms
Explanation: Across a period, atomic radius decreases because the nuclear charge increases while electrons are added to the same shell, causing stronger attraction. Choice A describes the trend down a group, not across a period. Choice C is incorrect because radius does change significantly across periods. Choice D is wrong because the trend is predictable, not random.

Question 14

Which of the following correctly describes the number of protons, neutrons, and electrons in a neutral atom of 39K^{39}K?

  1. 19 protons, 20 neutrons, 19 electrons (correct answer)
  2. 20 protons, 19 neutrons, 20 electrons
  3. 39 protons, 19 neutrons, 39 electrons
  4. 19 protons, 39 neutrons, 19 electrons
Explanation: For 39K^{39}K, the atomic number (19) gives the number of protons. The mass number (39) minus the atomic number gives neutrons: 39 - 19 = 20. In a neutral atom, electrons equal protons, so 19 electrons. Choice B reverses protons and neutrons. Choice C incorrectly uses the mass number for protons and electrons. Choice D incorrectly uses the mass number for neutrons.

Question 15

What is the primary reason that noble gases have very high ionization energies?

  1. They have complete outer electron shells, making their electrons very stable and difficult to remove (correct answer)
  2. They have very small atomic radii, which concentrates their nuclear charge in a tiny space
  3. They have the highest nuclear charges in their respective periods, creating strong attractive forces
  4. They have no d-orbital electrons, which eliminates electron-electron repulsion and stabilizes the atom
Explanation: Noble gases have complete valence shells (octet configuration), making them extremely stable and requiring high energy to remove electrons. Choice B is incorrect as noble gases actually have larger radii than earlier period elements. Choice C is wrong since noble gases don't necessarily have the highest nuclear charges. Choice D incorrectly focuses on d-orbitals rather than the key factor of complete valence shells.

Question 16

An atom has the electron configuration 1s22s22p63s23p64s11s^2 2s^2 2p^6 3s^2 3p^6 4s^1. What period and group is this element in?

  1. Period 3, Group 1A because the highest energy electrons are in the third energy level
  2. Period 4, Group 1A because the highest energy electrons are in 4s and there's one valence electron (correct answer)
  3. Period 4, Group 7A because there are four energy levels and seven total valence electrons
  4. Period 1, Group 4A because the 4s orbital contains one electron in the first subshell
Explanation: The period is determined by the highest principal energy level with electrons (4), and the group is determined by valence electrons (1 electron in 4s makes this Group 1A). Choice A uses the wrong energy level for period determination. Choice C miscounts valence electrons. Choice D completely misinterprets the notation and electron configuration principles.

Question 17

What is the maximum number of electrons that can be held in all the p orbitals of the second energy level?

  1. 2 electrons because there is only one p orbital in the second level
  2. 6 electrons because there are three p orbitals, each holding two electrons (correct answer)
  3. 10 electrons because p orbitals can accommodate up to ten electrons total
  4. 8 electrons because the second level follows the octet rule for electron capacity
Explanation: The second energy level (n=2) contains three 2p orbitals (2px2p_x, 2py2p_y, 2pz2p_z), and each orbital can hold a maximum of 2 electrons, giving 6 total. Choice A incorrectly states there's only one p orbital. Choice C confuses p orbital capacity with d orbital capacity. Choice D misapplies the octet rule to orbital capacity.

Question 18

Which of the following best describes the trend in electronegativity across a period from left to right?

  1. Electronegativity decreases because atoms become larger and less able to attract electrons effectively
  2. Electronegativity increases because nuclear charge increases while atomic size decreases (correct answer)
  3. Electronegativity remains constant because the number of valence electrons increases proportionally
  4. Electronegativity fluctuates irregularly because electron-electron repulsion varies unpredictably
Explanation: Across a period, electronegativity increases because nuclear charge increases while electrons are added to the same shell, making atoms smaller and better at attracting electrons. Choice A describes the opposite trend. Choice C is incorrect as electronegativity changes significantly across periods. Choice D wrongly suggests the trend is unpredictable.

Question 19

An element has the electron configuration 1s22s22p63s23p41s^2 2s^2 2p^6 3s^2 3p^4. What is the atomic number of this element?

  1. 12
  2. 14
  3. 16 (correct answer)
  4. 18
Explanation: The atomic number equals the total number of electrons in a neutral atom. Adding the superscripts: 2 + 2 + 6 + 2 + 4 = 16 electrons, so atomic number is 16 (sulfur). Choice A (12) would be 1s22s22p63s21s^2 2s^2 2p^6 3s^2. Choice B (14) would be 1s22s22p63s23p21s^2 2s^2 2p^6 3s^2 3p^2. Choice D (18) would be 1s22s22p63s23p61s^2 2s^2 2p^6 3s^2 3p^6.

Question 20

Which trend correctly describes ionization energy as you move down a group in the periodic table?

  1. Ionization energy increases because atoms become larger and electrons are easier to remove
  2. Ionization energy decreases because atoms become larger and electrons are farther from the nucleus (correct answer)
  3. Ionization energy remains constant because the nuclear charge increases proportionally with atomic size
  4. Ionization energy increases because the nuclear charge increases and holds electrons more tightly
Explanation: Down a group, ionization energy decreases because atomic radius increases, placing valence electrons farther from the nucleus and making them easier to remove despite increased nuclear charge. Choice A has the wrong direction for the trend. Choice C is incorrect as ionization energy changes significantly down groups. Choice D ignores the shielding effect and increased distance.