Chlorine has an atomic mass of 35.45 amu on the periodic table. A student concludes that chlorine atoms have 17 protons and approximately 18.45 neutrons. What is wrong with this reasoning?
AThe atomic mass includes the mass of electrons, which shifts the value away from a whole number
B35.45 is a weighted average of two naturally occurring isotopes (Cl-35 and Cl-37), not the mass of any single chlorine atom
CNeutrons do not contribute to atomic mass — only protons and electrons are counted
DThe amu unit is defined relative to carbon-12, making all other values non-integers purely by mathematical convention
No chlorine atom has 18.45 neutrons — neutrons come in whole numbers. The 35.45 value is a weighted average of naturally occurring chlorine isotopes: about 75.8% Cl-35 (17 protons, 18 neutrons) and 24.2% Cl-37 (17 protons, 20 neutrons). The calculation: (0.758 × 35) + (0.242 × 37) ≈ 35.48 amu. The periodic table reports this mixture average, not a description of any individual atom. A real chlorine atom is always Cl-35 or Cl-37, never anything in between.
Question 2 Multiple Choice
Carbon-12 is used in mass spectrometry for structural analysis; carbon-14 is used in radiometric dating of biological material. This difference in application reflects:
ACarbon-14 having more electrons and therefore different chemical reactivity from carbon-12
BTheir different physical properties (radioactive instability in C-14) while both follow identical chemical pathways in biological systems
CCarbon-12 being more abundant and therefore cheaper to produce for routine laboratory use
DCarbon-14 having a different atomic number, giving it distinct metabolic behavior in living organisms
Carbon-12 and carbon-14 have the same number of protons and electrons — the same atomic number (6) — so they are chemically indistinguishable. Living organisms incorporate both into biological molecules without discrimination. The difference is physical: C-14 is radioactive and decays at a known half-life (5,730 years), while C-12 is stable. This radioactive decay makes C-14 useful for dating (the ratio of C-14 to C-12 decreases after death at a predictable rate), while C-12's stable mass makes it useful as a mass spectrometry standard. Chemistry is identical; physics differs.
Question 3 True / False
Two isotopes of the same element have different chemical properties because their mass difference affects how their electrons interact with surrounding atoms.
TTrue
FFalse
Answer: False
Chemical properties are determined by electron configuration — the number and arrangement of electrons — not by mass. Isotopes of the same element have the same atomic number, so they have the same number of protons and electrons in the same arrangement. They form the same chemical bonds, participate in the same reactions, and are incorporated into the same molecular structures. The mass difference produces different physical properties (density, diffusion rate, bond vibration frequency), but chemical behavior is essentially identical. This is why C-14 follows the same metabolic pathways as C-12.
Question 4 True / False
The atomic mass listed on the periodic table for any element with more than one stable naturally occurring isotope will never be a whole number.
TTrue
FFalse
Answer: True
Atomic mass is the weighted average of all naturally occurring isotopic masses weighted by natural abundance. For any element with multiple stable isotopes, the abundance proportions are fixed by natural processes and almost certainly do not produce an average that lands exactly on an integer. The non-integer value is direct evidence of isotope mixing in the natural sample. Elements with only one stable isotope (like fluorine, F-19) come closest to whole numbers, but even then small deviations occur due to nuclear binding energy effects (mass defect).
Question 5 Short Answer
A student measures the atomic mass of a pure sample of carbon-12 and gets exactly 12.000 amu. They then measure the atomic mass of natural carbon and get 12.011 amu. Explain the discrepancy.
Think about your answer, then reveal below.
Model answer: The 12.000 result is not a measurement — it is a definition. The amu is defined as exactly 1/12 the mass of a carbon-12 atom, so measuring pure C-12 gives 12.000 by definition. Natural carbon is not pure C-12: it is 98.9% carbon-12 and about 1.1% carbon-13 (mass ≈ 13.003 amu), with trace C-14. The 12.011 value is the weighted average: (0.989 × 12.000) + (0.011 × 13.003) ≈ 12.011. The heavier C-13 minority pulls the average above 12.000.
This example illustrates two things at once: that atomic mass units are defined relative to C-12 (making 12.000 a reference point, not a discovery), and that natural samples of elements are always isotopic mixtures. The periodic table's atomic mass values are not properties of individual atoms — they are properties of the naturally occurring mixture of isotopes. This distinction matters practically: in stoichiometry, when you use the molar mass of carbon (12.011 g/mol), you are implicitly using a weighted average that reflects the natural C-12/C-13 ratio in your sample.