Health Education Systems Inc (HESI) A2 Exam Quiz: Isotopes And Atomic Mass
18 questions · exam conditions
0:00
Isotopes And Atomic MassQuestion 1 of 18

An unknown element X has an isotope with a mass number of 58 and 30 neutrons. A common ion of this element has a +2 charge. How many electrons are in this ion?

26
28
30
32
← Back to quizzes

Health Education Systems Inc (HESI) A2 Exam Quiz

Health Education Systems Inc (HESI) A2 Exam Quiz: Isotopes And Atomic Mass

Practice Isotopes And Atomic Mass in Health Education Systems Inc (HESI) A2 Exam 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 Isotopes And Atomic Mass, giving you a quick way to practice the rules, question types, and explanations that matter most for Health Education Systems Inc (HESI) A2 Exam.

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

An unknown element X has an isotope with a mass number of 58 and 30 neutrons. A common ion of this element has a +2 charge. How many electrons are in this ion?

  1. 26 (correct answer)
  2. 28
  3. 30
  4. 32
Explanation: When you encounter questions about atomic structure and ions, you need to work systematically through three key relationships: mass number equals protons plus neutrons, atomic number equals the number of protons, and ion charge tells you how electrons differ from protons. Start by finding the number of protons. Since mass number = protons + neutrons, and you know the mass number is 58 with 30 neutrons: 58=protons+3058 = \text{protons} + 30, so there are 28 protons. This means element X has an atomic number of 28 (which is nickel, though you don't need to identify it). In a neutral atom, electrons equal protons, so a neutral atom would have 28 electrons. However, this ion has a +2 charge, meaning it has lost 2 electrons. Therefore: 28−2=2628 - 2 = 26 electrons in the ion. Looking at the wrong answers: B) 28 represents the number of electrons in a neutral atom, but ignores the +2 charge. C) 30 confuses neutrons with electrons—a common mistake when students mix up the atomic particles. D) 32 might result from incorrectly adding electrons instead of subtracting them for a positive charge, or from calculation errors. Remember this pattern: positive ions have fewer electrons than protons (they've lost electrons), while negative ions have more electrons than protons (they've gained electrons). Always identify the number of protons first using mass number minus neutrons, then adjust the electron count based on the charge.

Question 2

If a neutral atom of Carbon-14 (14C{}^{14}\text{C}) gains a neutron, what is the resulting particle?

  1. An ion of Carbon-15.
  2. A neutral atom of Carbon-15. (correct answer)
  3. A neutral atom of Nitrogen-15.
  4. A neutral atom of Carbon-14.
Explanation: When you encounter questions about atomic structure changes, focus on what happens to the proton and neutron counts, as these determine the element's identity and mass number. Let's trace what happens when a neutral Carbon-14 atom gains a neutron. Carbon-14 has 6 protons (which defines it as carbon) and 8 neutrons (14 - 6 = 8). When it gains one neutron, you now have 6 protons and 9 neutrons, giving a mass number of 15. Since the number of protons hasn't changed, this is still carbon, just a different isotope. Because no electrons were gained or lost, the atom remains electrically neutral. Looking at the wrong answers: Choice A incorrectly suggests the result would be an ion. Ions form when atoms gain or lose electrons, not neutrons, so the charge remains neutral. Choice C falls into a common trap - while gaining a neutron does increase the mass number to 15, it doesn't change the element from carbon to nitrogen. That would require gaining a proton, not a neutron. Choice D ignores the change entirely, but adding a neutron must increase the mass number from 14 to 15. The correct answer is B: a neutral atom of Carbon-15. Study tip: Remember the atomic identity hierarchy: protons determine the element (carbon vs. nitrogen), neutrons determine the isotope (Carbon-14 vs. Carbon-15), and electrons determine the charge (neutral vs. ion). When analyzing nuclear changes, track each particle type separately to avoid mixing up these distinct effects.

Question 3

Silicon (Si) has an average atomic mass of 28.0855 amu. Its three main isotopes are Si-28, Si-29, and Si-30. If the abundance of Si-28 (mass ≈ 27.98 amu) is 92.23%, which statement is most likely true about the remaining two isotopes?

  1. Si-29 and Si-30 have a combined abundance of about 7.77% and are both more massive than Si-28. (correct answer)
  2. Si-29 is the most abundant isotope of silicon overall.
  3. The combined abundance of Si-29 and Si-30 is less than 5%, and they are less massive than Si-28.
  4. Si-28 must be radioactive because it is so much more abundant than the other isotopes.
Explanation: When you encounter isotope abundance problems, you're working with weighted averages where the atomic mass reflects the contribution of each isotope based on both its mass and abundance. Given that Si-28 (27.98 amu) has 92.23% abundance, the remaining isotopes Si-29 and Si-30 must account for 100% - 92.23% = 7.77% combined abundance. Since silicon's average atomic mass (28.0855 amu) is higher than Si-28's mass, the other isotopes must be more massive to pull the average upward. This mathematical relationship confirms that both Si-29 and Si-30 have greater masses than Si-28, making answer A correct. Answer B is wrong because Si-28 is clearly stated as having 92.23% abundance, making it the most abundant isotope by far. Answer C incorrectly claims the combined abundance is less than 5% (it's 7.77%) and wrongly states that Si-29 and Si-30 are less massive than Si-28. If they were less massive, the average atomic mass would be lower than 27.98 amu, not higher at 28.0855 amu. Answer D makes an unfounded connection between abundance and radioactivity—isotope stability has no direct relationship to natural abundance percentages. For isotope problems on the HESI, remember that the average atomic mass is always a weighted average. If the average is higher than the most abundant isotope's mass, the minor isotopes must be heavier to shift the average upward. This mathematical logic will guide you to the right answer every time.

Question 4

A sample contains two particles: one is a neutral atom with 8 protons and 8 neutrons, and the other is a particle with 8 protons and 10 neutrons. Which term accurately describes the relationship between these two particles?

  1. They are ions of the same element.
  2. They are isotopes of the same element. (correct answer)
  3. They are different elements.
  4. They are allotropes of the same element.
Explanation: When you encounter questions about atomic structure, focus on what determines an element's identity versus what creates variations within that element. The key is understanding that the number of protons defines the element, while variations in neutrons create different forms of the same element. Both particles have 8 protons, which means they're both oxygen atoms (atomic number 8). The first particle has 8 neutrons, while the second has 10 neutrons. Since they share the same number of protons but differ in neutron count, they are isotopes of oxygen. Isotopes are atoms of the same element with different mass numbers due to varying neutron counts. Let's examine why the other options are incorrect: A) These aren't ions because both particles are described in terms of protons and neutrons only—no mention of electron gain or loss, which would be necessary to create charged ions. C) These can't be different elements because they have the same number of protons (8). The proton number (atomic number) is what distinguishes one element from another on the periodic table. D) Allotropes refer to different structural forms of the same element in the same physical state, like diamond versus graphite for carbon. This term applies to molecular arrangements, not atomic composition differences. Study tip: Remember "P.I.N." - Protons determine Identity, Ions involve electron changes, Neutrons create isotopes. When you see particles with the same proton count but different neutron counts, you're always dealing with isotopes, regardless of their charge state.

Question 5

An element, X, has two naturally occurring isotopes. X-10 has an abundance of 20% and a mass of 10.01 amu. X-11 has an abundance of 80% and a mass of 11.01 amu. What is the average atomic mass of element X?

  1. 10.21 amu
  2. 10.51 amu
  3. 10.81 amu (correct answer)
  4. 11.00 amu
Explanation: When you encounter isotope problems, you're dealing with weighted averages. Since isotopes of the same element have different masses but occur in different proportions in nature, the average atomic mass reflects both the mass of each isotope and how commonly it's found. To calculate average atomic mass, multiply each isotope's mass by its decimal abundance, then sum the results. For element X, convert the percentages to decimals: X-10 has 0.20 abundance, X-11 has 0.80 abundance. Average atomic mass=(0.20×10.01)+(0.80×11.01)\text{Average atomic mass} = (0.20 \times 10.01) + (0.80 \times 11.01) =2.002+8.808=10.808 amu= 2.002 + 8.808 = 10.808 \text{ amu} Rounding to appropriate significant figures gives 10.81 amu. Answer A (10.21 amu) represents a common error where students accidentally flip the abundances, calculating (0.80 × 10.01) + (0.20 × 11.01). Answer B (10.51 amu) results from incorrectly using the percentages as whole numbers rather than decimals, or from calculation errors in the weighted average. Answer D (11.00 amu) occurs when students simply average the two masses without considering abundance, treating each isotope as equally common. Notice that the correct answer (C) is much closer to 11.01 amu than to 10.01 amu, which makes sense since X-11 is four times more abundant than X-10. Always check that your calculated average is weighted toward the more abundant isotope—this serves as a quick reasonableness check for isotope problems.

Question 6

A student observes that the atomic mass of most elements on the periodic table is not a whole number (e.g., copper is 63.546 amu). Which of the following provides the best explanation for this observation?

  1. The mass of individual protons and neutrons is not exactly 1 amu, and this small difference accumulates.
  2. The presence of multiple naturally occurring isotopes with different masses is averaged to give the final atomic mass. (correct answer)
  3. Experimental errors in measuring atomic masses prevent the values from being precise whole numbers.
  4. Electrons have a small but significant mass that, when added to the proton and neutron mass, results in a non-integer value.
Explanation: When you encounter questions about atomic masses on the periodic table, you're dealing with the concept of weighted averages based on isotopic abundance. The key insight is that elements exist as mixtures of isotopes in nature. The correct answer is B because atomic masses represent weighted averages of all naturally occurring isotopes of an element. For example, copper has two main isotopes: copper-63 (about 69% abundance) and copper-65 (about 31% abundance). The atomic mass of 63.546 amu reflects this mixture: (0.69×63)+(0.31×65)=63.546(0.69 × 63) + (0.31 × 65) = 63.546 amu. This weighted average accounts for both the mass and relative abundance of each isotope. Let's examine why the other options are incorrect. Option A suggests that the non-integer masses of protons and neutrons cause this effect. While it's true that these particles don't have exactly 1 amu mass, this factor alone wouldn't create the specific decimal patterns we observe. Option C attributes the decimals to experimental error, but modern mass spectrometry is extremely precise—these values are accurate, not imprecise. Option D incorrectly emphasizes electron mass contribution. Electrons are about 1/1836 the mass of a proton, so their contribution is negligible compared to the isotopic averaging effect. Remember this pattern: when you see non-whole atomic masses on the HESI, think "isotopic mixture." The periodic table values always represent what you'd measure from a natural sample containing multiple isotopes, not individual atoms.

Question 7

The periodic table lists the atomic mass of chlorine (Cl) as 35.45 amu. The two major isotopes of chlorine are Cl-35 and Cl-37. Based on the average atomic mass, what can be concluded about their relative abundances?

  1. Cl-35 and Cl-37 are present in nearly equal amounts.
  2. Cl-37 is significantly more abundant than Cl-35.
  3. Cl-35 is significantly more abundant than Cl-37. (correct answer)
  4. The abundance cannot be determined from the information given.
Explanation: When you encounter questions about atomic mass and isotopes, remember that the periodic table's atomic mass represents a weighted average of all naturally occurring isotopes. This average reflects not just the masses of the isotopes, but how abundant each one is in nature. To determine relative abundance, look at where the average atomic mass falls relative to the individual isotope masses. Chlorine's average atomic mass is 35.45 amu, which falls between the two isotopes Cl-35 (35 amu) and Cl-37 (37 amu). Crucially, notice that 35.45 is much closer to 35 than to 37. Since the average is weighted by abundance, this tells you that Cl-35 must be significantly more abundant than Cl-37. If you calculated the exact percentages, you'd find that Cl-35 makes up about 76% of natural chlorine, while Cl-37 is only about 24%. Looking at the wrong answers: Choice A suggests equal abundances, but if that were true, the average would be right in the middle at 36 amu, not 35.45. Choice B claims Cl-37 is more abundant, which would push the average closer to 37, not 35.45. Choice D suggests the abundance can't be determined, but the position of the weighted average clearly indicates which isotope dominates. For HESI chemistry questions, always remember this pattern: the weighted average atomic mass will be closer to whichever isotope is more abundant. Use the position of the average as your guide to abundance ratios.

Question 8

An ion of magnesium, 25Mg2+{}^{25}\text{Mg}^{2+}, is analyzed. How many protons, neutrons, and electrons does this ion contain?

  1. 12 protons, 13 neutrons, 10 electrons (correct answer)
  2. 12 protons, 13 neutrons, 12 electrons
  3. 12 protons, 13 neutrons, 14 electrons
  4. 13 protons, 12 neutrons, 11 electrons
Explanation: When you encounter atomic notation like 25Mg2+^{25}\text{Mg}^{2+}, you're being asked to decode three key pieces of information: the number of protons, neutrons, and electrons in this specific ion. Start with what the symbol tells you. Magnesium (Mg) always has 12 protons—this is its atomic number and never changes. The superscript 25 is the mass number, which equals protons plus neutrons. Since we have 12 protons, the neutrons must be 25−12=1325 - 12 = 13. The 2+^{2+} charge means this ion has lost 2 electrons, so from the neutral atom's 12 electrons, we now have 12−2=1012 - 2 = 10 electrons. Choice A correctly identifies 12 protons, 13 neutrons, and 10 electrons. Choice B makes the common error of forgetting that ions have different numbers of electrons than protons—it gives 12 electrons, which would be true for a neutral magnesium atom, not the 2+^{2+} ion. Choice C incorrectly suggests 14 electrons, which would require the ion to have gained electrons and carry a negative charge. Choice D completely misinterprets the notation, incorrectly assigning 13 as the number of protons instead of recognizing that 25 is the mass number. Remember this pattern: atomic number = protons (never changes), mass number = protons + neutrons, and for ions, electrons = protons minus the positive charge (or plus the negative charge). Always pay careful attention to charge indicators—they're crucial for determining electron count.

Question 9

A medical researcher is using two different isotopes of oxygen, oxygen-16 and oxygen-18, to trace metabolic pathways. Which statement best explains why this technique is effective?

  1. The two oxygen isotopes have different numbers of electrons, leading to distinct reaction rates.
  2. The two oxygen isotopes have nearly identical chemical properties but different masses, allowing them to be distinguished by mass spectrometry. (correct answer)
  3. Oxygen-18 is significantly more reactive than oxygen-16 due to its two extra neutrons, which destabilize the nucleus.
  4. The different number of neutrons changes the element's atomic number, which alters how it bonds with other atoms.
Explanation: When you encounter questions about isotopes in biological research, focus on the fundamental principle that isotopes of the same element have identical chemical properties but different physical properties due to varying neutron numbers. Isotope tracing works because oxygen-16 and oxygen-18 behave chemically identically in metabolic reactions—they bond the same way, participate in the same pathways, and don't alter normal cellular processes. However, their different masses (16 vs. 18 atomic mass units) make them distinguishable using mass spectrometry. Researchers can inject oxygen-18 and then track where it appears in metabolic products, revealing pathway details without disrupting normal metabolism. Choice A incorrectly states that isotopes have different numbers of electrons. Isotopes of the same element always have the same number of protons and electrons—only neutron numbers differ. This means their chemical behavior is virtually identical. Choice C falsely claims oxygen-18 is more reactive due to nuclear instability. Both oxygen-16 and oxygen-18 are stable isotopes. The extra neutrons in oxygen-18 don't significantly affect chemical reactivity since chemical properties depend on electron configuration, not neutron number. Choice D contains a fundamental error about atomic structure. The atomic number (number of protons) defines the element and never changes among isotopes. Oxygen-16 and oxygen-18 both have 8 protons, so both have the same atomic number and bonding characteristics. Remember: isotopes = same chemistry, different mass. This principle makes them perfect biological tracers since they don't interfere with normal processes but remain detectable.

Question 10

The modern standard for defining the atomic mass unit (amu) is based on a specific isotope. Which statement accurately identifies this standard?

  1. One amu is exactly equal to the mass of a single proton.
  2. One amu is defined as 1/16th the mass of a neutral oxygen-16 atom.
  3. One amu is defined as 1/12th the mass of a neutral carbon-12 atom. (correct answer)
  4. One amu is exactly equal to the mass of a single hydrogen-1 atom.
Explanation: When you encounter questions about atomic mass units on the HESI, you're being tested on fundamental chemistry concepts that underpin dosage calculations and molecular understanding in healthcare. The atomic mass unit (amu) serves as a standardized reference point for measuring atomic masses. The modern definition establishes one amu as exactly 1/12th the mass of a neutral carbon-12 atom. This means if you could isolate a single carbon-12 atom and divide its mass by 12, you'd get one amu. This carbon-12 standard was adopted internationally in 1961 because carbon-12 is stable, abundant, and provides a convenient reference point for all other atomic masses. Looking at why the other options miss the mark: Option A incorrectly suggests that one amu equals a proton's mass, but a proton actually has a mass of approximately 1.007 amu—close, but not exact. Option B references an outdated system that used oxygen-16 as the standard before 1961, when the scientific community switched to the more precise carbon-12 reference. Option D proposes that one amu equals the mass of a hydrogen-1 atom, but hydrogen-1 has a mass of about 1.008 amu since it includes both a proton and an electron. Remember this pattern: when the HESI asks about scientific standards or definitions, the correct answer will reflect the current, internationally accepted measurement system. The carbon-12 standard is fundamental to understanding molecular weights, which you'll need for medication calculations and chemical processes in healthcare settings.

Question 11

Particle A has 17 protons, 18 neutrons, and 18 electrons. Particle B has 17 protons, 20 neutrons, and 17 electrons. What is the relationship between Particle A and Particle B?

  1. Particle A is a negative ion and Particle B is a neutral atom of a different element.
  2. They are both neutral isotopes of the same element.
  3. Particle A is an anion and Particle B is a neutral atom; they are isotopes of the same element. (correct answer)
  4. Particle A is a cation and Particle B is a neutral atom; they are isotopes of the same element.
Explanation: When you encounter questions about atomic particles, focus on three key numbers: protons (which determine the element), electrons (which determine charge), and neutrons (which create isotopes). Let's analyze each particle systematically. The number of protons identifies the element - both particles have 17 protons, making them both chlorine atoms. Now examine the charges: Particle A has 17 protons (+17 charge) and 18 electrons (-18 charge), giving it a net charge of -1, making it an anion (negative ion). Particle B has 17 protons and 17 electrons, so it's electrically neutral. Since they have the same number of protons but different numbers of neutrons (18 vs 20), they are isotopes of the same element. Looking at the wrong answers: Choice A incorrectly states they're different elements - impossible since both have 17 protons. Choice B claims both are neutral, but Particle A has an extra electron making it negatively charged. Choice D incorrectly identifies Particle A as a cation (positive ion) when it actually has more electrons than protons, making it an anion. Choice C correctly identifies that Particle A is an anion (negative ion due to the extra electron), Particle B is a neutral atom (equal protons and electrons), and they are isotopes of the same element (same atomic number, different mass numbers). Study tip: Remember the pattern - protons = element identity, electron difference = ion type, neutron difference = isotopes. Always count protons first to confirm you're dealing with the same element.

Question 12

Consider the two primary isotopes of carbon: Carbon-12 and Carbon-13. A neutral atom of Carbon-13 contains:

  1. one more proton and one less neutron than Carbon-12.
  2. the same number of protons and one more neutron than Carbon-12. (correct answer)
  3. one more electron and the same number of neutrons as Carbon-12.
  4. one more proton and the same number of electrons as Carbon-12.
Explanation: When you encounter questions about isotopes, remember that isotopes are atoms of the same element with different numbers of neutrons. The key insight is that all atoms of the same element must have the same number of protons (this defines the element), but can have varying numbers of neutrons. Carbon always has 6 protons, regardless of which isotope we're discussing. The numbers in Carbon-12 and Carbon-13 refer to their mass numbers (protons + neutrons). Since Carbon-12 has a mass number of 12 and 6 protons, it must have 6 neutrons (12 - 6 = 6). Carbon-13, with a mass number of 13 and the same 6 protons, must have 7 neutrons (13 - 6 = 7). Since both are neutral atoms, they each have 6 electrons to balance their 6 protons. Therefore, Carbon-13 has the same number of protons and one more neutron than Carbon-12, making answer B correct. Looking at the incorrect options: Answer A suggests different numbers of protons, which would make these different elements entirely, not isotopes of carbon. Answer C proposes different numbers of electrons, but neutral atoms have equal numbers of protons and electrons. Answer D also suggests different numbers of protons, which again would mean different elements. Study tip: For isotope questions on the HESI, always remember that isotopes differ only in neutron number. Use the formula: Mass number = Protons + Neutrons. The atomic number (number of protons) stays constant for any given element, and neutral atoms always have equal numbers of protons and electrons.

Question 13

An atom has a mass number of 31 and contains 16 neutrons. This atom is electrically neutral. What is the identity of the element and how many electrons does it have?

  1. Sulfur (S), with 16 electrons
  2. Phosphorus (P), with 15 electrons (correct answer)
  3. Phosphorus (P), with 16 electrons
  4. Silicon (Si), with 14 electrons
Explanation: When you encounter atomic structure questions, you need to use the fundamental relationships between mass number, atomic number, protons, neutrons, and electrons to identify elements and their properties. Start with what you know: mass number = 31, neutrons = 16, and the atom is electrically neutral. Since mass number equals protons plus neutrons, you can find the number of protons: 31−16=1531 - 16 = 15 protons. The number of protons defines the element's identity—this is the atomic number. An atom with 15 protons is phosphorus (P). Since the atom is electrically neutral, it must have equal numbers of protons and electrons, so it has 15 electrons. Looking at the wrong answers: Choice A incorrectly identifies the element as sulfur, which would need 16 protons, not 15. Choice C correctly identifies phosphorus but claims 16 electrons—this would make the atom negatively charged (an anion), contradicting the "electrically neutral" condition. Choice D suggests silicon with 14 electrons, but silicon has 14 protons, which doesn't match our calculation of 15 protons. For HESI atomic structure questions, remember this sequence: use mass number and neutrons to find protons, use protons to identify the element via the periodic table, then use the charge state to determine electrons. Electrically neutral atoms always have equal protons and electrons, while ions will have different numbers. Practice this systematic approach to avoid mixing up element identity with electron count.

Question 14

A neutral atom of potassium-39 (K-39, with 19 protons) is considered. A change in which of its subatomic particles would result in the formation of an isotope of potassium?

  1. A change in the number of protons.
  2. A change in the number of neutrons. (correct answer)
  3. A change in the number of electrons.
  4. A change in both protons and neutrons.
Explanation: When you encounter questions about atomic structure and isotopes, focus on understanding what defines an element versus what creates variants of that element. An element's identity is determined solely by its number of protons (atomic number), while isotopes are different versions of the same element. Potassium-39 has 19 protons, which makes it potassium. To create an isotope of potassium, you need to change the mass while keeping the element identity intact. Since neutrons contribute to atomic mass but not to the element's chemical identity, changing the number of neutrons creates isotopes. For example, potassium-40 and potassium-41 are isotopes of potassium-39 because they have different numbers of neutrons (21 and 22 respectively) but still have 19 protons. Choice A is incorrect because changing the number of protons would create a completely different element, not an isotope of potassium. If you added or removed protons, you'd get calcium (20 protons) or argon (18 protons), not a potassium isotope. Choice C is wrong because changing electrons creates ions (charged particles), not isotopes. The number of electrons affects charge but doesn't change mass significantly or create isotopes. Choice D is incorrect because changing both protons and neutrons would create a different element entirely, just like option A, rather than maintaining the potassium identity needed for an isotope. Remember this key distinction: protons determine the element, neutrons create isotopes, and electrons create ions. On the HESI, isotope questions test whether you understand that only neutron changes preserve elemental identity.

Question 15

Regardless of the number of neutrons or electrons, all atoms and ions of the element silver (Ag) have the same:

  1. mass number.
  2. atomic mass.
  3. number of protons. (correct answer)
  4. charge.
Explanation: This question tests your understanding of atomic structure and what defines an element's identity. When you encounter questions about atoms and elements, focus on the fundamental characteristics that distinguish one element from another. The defining characteristic of any element is its atomic number, which equals the number of protons in the nucleus. Silver (Ag) has an atomic number of 47, meaning every silver atom contains exactly 47 protons. This number never changes - if an atom has 47 protons, it's silver; if it has a different number of protons, it's a different element entirely. Even when silver forms ions (gaining or losing electrons) or exists as different isotopes (having different numbers of neutrons), the proton count remains constant at 47. Let's examine why the other options are incorrect. Choice A (mass number) is wrong because mass number equals protons plus neutrons, and since neutrons can vary in isotopes, mass number changes. For example, silver-107 and silver-109 are both silver but have different mass numbers. Choice B (atomic mass) is incorrect for the same reason - atomic mass reflects the weighted average of all isotopes and varies based on neutron content. Choice D (charge) is wrong because atoms become ions by gaining or losing electrons, which changes their charge while keeping the same number of protons. Remember this key principle: the number of protons defines the element's identity and never changes. On the HESI, when you see questions about "all atoms of an element," think immediately about protons as the unchanging identifier.

Question 16

The isotope commonly used in medical imaging, Technetium-99m (Tc-99), has an atomic number of 43. How many protons and neutrons are in the nucleus of this isotope?

  1. 43 protons and 56 neutrons (correct answer)
  2. 43 protons and 99 neutrons
  3. 56 protons and 43 neutrons
  4. 99 protons and 43 neutrons
Explanation: Understanding atomic structure is essential for healthcare professionals, especially when working with medical isotopes used in imaging and treatment. When you see an isotope notation like Technetium-99m, you need to decode what the numbers tell you about the atom's nucleus. The atomic number (43 for Technetium) always tells you the number of protons in the nucleus. This never changes for a given element. The mass number (99 in Tc-99) represents the total number of protons plus neutrons. To find the number of neutrons, you subtract: neutrons = mass number - atomic number = 99 - 43 = 56 neutrons. Looking at the answer choices: Option A correctly identifies 43 protons and 56 neutrons. Option B incorrectly assumes the mass number (99) equals the number of neutrons, ignoring that this number includes both protons and neutrons. Option C reverses the numbers, mistakenly using 56 (the calculated neutrons) as protons and 43 (the atomic number) as neutrons. Option D incorrectly treats the mass number (99) as the number of protons, which would make this an entirely different element. Remember this simple formula: Mass number - Atomic number = Number of neutrons. The atomic number is always the number of protons and defines what element you're dealing with. On the HESI, isotope questions often appear in contexts involving medical procedures, so understanding this basic nuclear chemistry will help you tackle both chemistry and healthcare application questions.

Question 17

Which of the following pairs represents isotopes of the same element?

  1. A particle with 19 protons and 20 neutrons, and a particle with 20 protons and 20 neutrons.
  2. A particle with 6 protons and 6 neutrons, and a particle with 6 protons and 7 neutrons. (correct answer)
  3. A particle with 17 protons and 18 neutrons, and a particle with 17 protons and 18 neutrons that has gained one electron.
  4. A particle with 1 proton and 1 neutron, and a particle with 2 protons and 2 neutrons.
Explanation: When you encounter questions about isotopes, remember that isotopes are atoms of the same element that have identical numbers of protons but different numbers of neutrons. The number of protons defines the element's identity, while varying neutron counts create different isotopes. To identify isotopes, you need to find particles with the same proton number but different neutron numbers. Looking at option B, both particles have 6 protons, making them both carbon atoms. However, one has 6 neutrons while the other has 7 neutrons, creating carbon-12 and carbon-13 isotopes respectively. This perfectly demonstrates the isotope relationship. Option A is incorrect because the particles have different proton numbers (19 vs. 20), making them completely different elements - potassium and calcium. Option C represents the same exact atom (17 protons, 18 neutrons) in different charge states. The particle that "gained one electron" is simply an ion of the first particle, not a different isotope. Both are chlorine-35, just with different electron counts. Option D shows particles with different proton numbers (1 vs. 2), representing hydrogen and helium - again, different elements entirely. For HESI chemistry questions, remember this key pattern: isotopes = same protons, different neutrons. Don't be fooled by electron differences (which create ions) or proton differences (which create different elements). Focus solely on the proton-neutron relationship to identify true isotope pairs.

Question 18

Two neutral atoms are isotopes of uranium. Which property must be different for these two atoms?

  1. The number of protons in the nucleus.
  2. The number of electrons in the atomic orbitals.
  3. The mass number. (correct answer)
  4. The chemical reactivity in a reaction.
Explanation: When you encounter questions about isotopes, focus on the fundamental definition: isotopes are atoms of the same element that have identical numbers of protons but different numbers of neutrons. Since isotopes are atoms of the same element, they must have the same number of protons in their nuclei - this defines what element they are. Two uranium isotopes will both have 92 protons. The key difference lies in their neutron count, which directly affects their mass number (the sum of protons and neutrons). For example, uranium-235 has 143 neutrons while uranium-238 has 146 neutrons, giving them mass numbers of 235 and 238 respectively. Looking at the wrong answers: Choice A is incorrect because isotopes of the same element always have identical proton numbers - that's what makes them uranium atoms. Choice B is wrong because neutral atoms have equal numbers of protons and electrons, so if the proton count is the same, the electron count must also be the same. Choice D is incorrect because isotopes have nearly identical chemical properties since chemical behavior depends primarily on electron configuration, which remains the same between isotopes. Choice C is correct because the mass number must differ between isotopes. Since they have the same number of protons but different numbers of neutrons, their mass numbers (protons + neutrons) will inevitably be different. Remember this pattern: isotopes = same element, same protons, different neutrons, different mass numbers. The mass number is the distinguishing feature that separates one isotope from another.