Atomic Mass 113

Mass Number vs Atomic Number

Two different numbers describe an atom and they are easy to confuse:

• Atomic number (Z) is the number of protons in the nucleus. It defines what element the atom is. Hydrogen has Z = 1, oxygen has Z = 8, and nihonium has Z = 113.
• Mass number (A) is the total count of protons plus neutrons. It is always a whole number for a specific isotope. Carbon-12 has A = 12; uranium-238 has A = 238.
• Atomic mass is the actual measured mass of an atom in atomic mass units (amu, also called daltons, Da). It is close to but not identical to A, because the protons and neutrons together weigh slightly less than their separate masses (the mass defect, related to nuclear binding energy).

So "113" can label either an element (Z = 113, nihonium) or an isotope (A = 113, of any element with around 49–50 protons and the rest neutrons). They are distinct.

Cadmium-113

Cadmium has atomic number 48. Cadmium-113 has 48 protons and 65 neutrons, summing to mass number 113.

• Symbol: ¹¹³Cd or Cd-113
• Natural abundance: about 12% of natural cadmium.
• Stability: technically observationally stable, with an extremely long half-life (~10¹⁵ years for one of its decay modes) — for practical purposes treated as stable.
• Notable property: very large neutron capture cross-section. ¹¹³Cd absorbs thermal neutrons exceptionally well, which is why cadmium is used in nuclear reactor control rods.
• Atomic mass (precise): ≈ 112.9044 amu — slightly less than 113 because of nuclear binding energy.

Indium-113

Indium has atomic number 49. Indium-113 has 49 protons and 64 neutrons, summing to mass number 113.

• Symbol: ¹¹³In or In-113
• Natural abundance: about 4.3% of natural indium (the more common natural isotope is ¹¹⁵In, at about 95.7%).
• Stability: stable.
• Atomic mass (precise): ≈ 112.9041 amu.
• Notable use: ¹¹³ᵐIn (a metastable nuclear isomer at the same mass number) is used in nuclear medicine for some diagnostic imaging.

Other Isotopes With Mass 113

Several other isotopes share the mass number 113, though most are unstable:

• Tin-113 (¹¹³Sn): radioactive, half-life about 115 days. Used as the parent for ¹¹³ᵐIn generators in some nuclear medicine workflows.
• Silver-113 (¹¹³Ag): radioactive, half-life about 5.4 hours.
• Antimony-113 (¹¹³Sb): radioactive, half-life about 6.7 minutes.
• Tellurium-113 (¹¹³Te): radioactive, very short half-life.
• Various others across iodine, xenon, and lower-Z elements, with progressively shorter half-lives.

This is what nuclear physicists call an isobar chain — a sequence of isotopes that all share the same mass number A = 113 but differ in proton count Z. The most stable members are ¹¹³In and ¹¹³Cd; the others decay toward those.

113 amu in Practical Mass Spectrometry

In mass spectrometry, peaks at mass-to-charge ratio m/z = 113 are common. They can come from:

• Singly charged ions of indium or cadmium (where their natural isotopes hit 113).
• Molecular fragments whose nominal mass adds up to 113 — for example, certain six- or seven-atom organic fragments.
• Doubly charged ions of much heavier species (m/z = 113 from a 226-amu doubly charged ion).

Identifying the source of an m/z 113 peak depends on isotope pattern, instrument resolution, and the chemistry of the sample. High-resolution instruments can distinguish 113.000 from 113.040 from 113.084, for example, helping to pin down the elemental composition.

Why Atomic Mass Is Close to But Not Equal to A

The mass number A is a count: 113 nucleons (protons + neutrons). Each free proton has mass ~1.0073 amu and each free neutron has mass ~1.0087 amu. Naively, 113 nucleons would weigh ~113.99 amu.

Real nuclei weigh less than this because some of the mass is converted to binding energy when the nucleons bond together — Einstein's E = mc² in the nuclear context. For ¹¹³Cd, the actual atomic mass is about 112.9044 amu, around 1.09 amu less than the naive sum. That mass defect corresponds to the energy holding the nucleus together.

The mass defect is why nuclear reactions release so much energy compared to chemical reactions — even small differences in atomic mass between reactants and products correspond to enormous energies.

Element 113 (Nihonium) Is Different

Element 113 has atomic number 113 — meaning 113 protons in its nucleus. Its known isotopes are heavy: the first synthesized isotope, ²⁷⁸Nh, has mass number 278, not 113. So element 113 has nothing in common, mass-wise, with the isotopes discussed above.

For the full story of element 113's discovery, naming as nihonium, and place in the periodic table, see element 113: nihonium and element 113 in the periodic table.

Common Mistakes

• Confusing mass number with atomic number. A = 113 (an isotope mass) is not the same as Z = 113 (the element nihonium).
• Treating atomic mass as exact. The mass number is exact (a count of nucleons), but the atomic mass is a measured quantity in amu and includes mass defect.
• Assuming all 113-mass isotopes are stable. Only ¹¹³In is firmly stable; ¹¹³Cd is observationally stable; the rest are radioactive and decay toward the stable members.
• Misreading mass spec peaks. An m/z 113 peak does not always mean a 113-amu species — multiply-charged ions and adducts can land at the same nominal mass.

Quick-Reference Card

• Mass number 113: 113 nucleons total (protons + neutrons).
• Most common stable isotopes at A = 113: ¹¹³In, ¹¹³Cd.
• Cadmium-113 atomic mass: ≈ 112.9044 amu.
• Indium-113 atomic mass: ≈ 112.9041 amu.
• Element with Z = 113: nihonium (Nh), unrelated to mass number 113.
• Isobar chain: ¹¹³Sn → ¹¹³In and ¹¹³Cd are stable endpoints.

For more on 113 in chemistry and physics, see 113 on the periodic table, 113 chemical compounds, and element 113 (nihonium).