Look up atomic masses for all 118 elements. Calculate formula masses, convert between amu and grams, and explore mass defect and nuclear binding energy.
Atomic mass is the mass of an atom measured in atomic mass units (amu or u), where 1 amu is defined as exactly 1/12 the mass of a carbon-12 atom. The standard atomic weight listed on the periodic table is a weighted average of the masses of all naturally occurring isotopes of an element, which is why it is rarely an integer.
Atomic mass is fundamental to quantitative chemistry. It connects the atomic scale to the macroscopic world through Avogadro's number: the atomic mass in amu numerically equals the molar mass in g/mol. This means 12.011 amu of carbon weighs 12.011 g per mole, which is 6.022 × 10²³ atoms.
This calculator provides precise atomic masses for all elements, converts between amu and grams/kilograms, calculates the mass of a specified number of atoms, and explores the mass defect — the difference between the sum of individual proton and neutron masses and the actual nuclear mass, which accounts for nuclear binding energy via E = mc².
Quickly look up precise atomic masses and perform conversions between amu, grams, moles, and number of atoms. The mass defect calculator connects atomic mass to nuclear physics. This atomic mass calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
Standard Atomic Mass = Σ (fractional abundance × isotopic mass) Mass of N atoms = N × atomic mass (amu) Mass in grams = moles × molar mass 1 amu = 1.66054 × 10⁻²⁴ g Mass defect: Δm = Z×m_p + N×m_n - m_atom Binding energy: E = Δm × c² = Δm × 931.5 MeV/amu
Result: 10.55 g
Copper has atomic mass 63.546 amu. For 10²³ atoms: mass = 10²³ × 63.546 amu × 1.66054×10⁻²⁴ g/amu = 10.55 g. Equivalently, 10²³ atoms = 0.1661 mol, and 0.1661 × 63.546 g/mol = 10.55 g.
The modern atomic mass scale is based on carbon-12. Before 1961, two different scales existed: chemists used oxygen-16 and physicists used oxygen as a mixture of isotopes. Unification on carbon-12 was a compromise that pleased both communities (the maximum shift in any element's mass was only 0.004%).
Einstein's famous equation E = mc² has a direct chemical application in the mass defect. When nucleons bind together, they release energy and the resulting nucleus weighs slightly less than the sum of its parts. For ⁵⁶Fe, the mass defect is ~0.53 amu or ~492 MeV — enough to accelerate a proton to 99.99% the speed of light.
Mass spectrometers separate isotopes by their mass-to-charge ratio and can measure atomic masses to seven decimal places. Applications include geological dating (U-Pb, K-Ar methods), forensic analysis (stable isotope ratios), medical diagnostics (PET scans using ¹⁸F), and nuclear energy (²³⁵U enrichment).
Mass number (A) is the integer count of protons + neutrons. Atomic mass is the precise mass in amu, which accounts for binding energy and is not exactly an integer.
Two reasons: (1) standard atomic mass is a weighted average of isotopes, and (2) even individual isotope masses aren't exactly integers due to mass defect (nuclear binding energy). This keeps planning practical and lowers the chance of preventable errors.
1 amu = 1.66054 × 10⁻²⁴ g. This equals 1/(6.022 × 10²³), the reciprocal of Avogadro's number divided by the gram.
Mass defect is the difference between the sum of individual particle masses and the actual nuclear mass. This "missing mass" has been converted to binding energy (E = mc²) that holds the nucleus together.
Carbon-12 (¹²C) is the standard: 1 amu = exactly 1/12 of the mass of a ¹²C atom. This keeps planning practical and lowers the chance of preventable errors.
Chlorine has two stable isotopes: ³⁵Cl (75.77%) and ³⁷Cl (24.23%). The weighted average is 0.7577 × 34.969 + 0.2423 × 36.966 = 35.45 amu.