Struggling to calculate average atomic mass? You are not alone — it trips up students at every level. Our Average Atomic Mass Calculator handles the math instantly. Just enter your isotope masses and natural abundances, and it delivers the exact weighted average — with full step-by-step working so you actually understand the result.

This guide goes far beyond what any other calculator page offers. You will get the exact formula, five fully worked examples, a ready-to-use isotope reference table, an explanation of monoisotopic mass (a concept most pages ignore), the science behind mass spectrometry, real-world applications, common mistakes, and 14 FAQ answers. Everything you need — in one place.

1. What Is Average Atomic Mass?

Average atomic mass is the weighted average of the masses of all naturally occurring isotopes of an element, where each isotope's mass is weighted by how abundant it is in nature.

Here is why this matters: most elements do not exist as a single type of atom. They exist as a mixture of isotopes — atoms with the same number of protons but different numbers of neutrons, and therefore different masses. A single fixed mass cannot represent the element accurately. The solution is the weighted average.

This weighted average is exactly what you see on the periodic table — which is why the values are almost never whole numbers.

Quick Examples

• Hydrogen: Three isotopes — Protium (¹H), Deuterium (²H), and Tritium (³H). The average atomic mass is 1.008 u.
• Carbon: Two main isotopes — C-12 (98.93%) and C-13 (1.07%). The average atomic mass is 12.011 u.
• Chlorine: Two major isotopes — Cl-35 (75.78%) and Cl-37 (24.22%). The average atomic mass is 35.45 u.
• Bromine: Two almost equally abundant isotopes — Br-79 (50.69%) and Br-81 (49.31%). The average atomic mass is 79.904 u.

Key insight: Average atomic mass is sometimes called atomic weight — the terms are used interchangeably in most chemistry contexts, though technically "weight" implies a gravitational force while "mass" does not.

2. Atomic Mass vs. Average Atomic Mass — The Key Difference

These two terms are frequently confused. Here is the clear distinction:

Bottom line: When you pick up a chemistry textbook and read "the atomic mass of carbon is 12.011," that is the average atomic mass — the weighted blend of all carbon isotopes found in nature. No single carbon atom actually has a mass of 12.011 u.

3. Average Atomic Mass vs. Molar Mass

This is one of the most commonly confused pairs in chemistry. Here is the clean answer:

The numbers are always identical — the unit simply changes. This is why average atomic mass is so important: it directly gives you the molar mass you need for stoichiometry, solution preparation, and every quantitative chemistry calculation.

4. The Average Atomic Mass Formula — Fully Explained

The Core Formula

AM = (f₁ × m₁) + (f₂ × m₂) + (f₃ × m₃) + … + (fₙ × mₙ)

Where:

• AM = Average atomic mass (result in amu / u / Da)
• fₙ = Fractional abundance of the nth isotope (a decimal between 0 and 1)
• mₙ = Atomic mass of the nth isotope in amu
• n = Total number of isotopes

The Golden Rule

The sum of all fractional abundances must equal exactly 1.0:

f₁ + f₂ + f₃ + … + fₙ = 1.0000

If using percent abundances, they must sum to 100%. If your numbers do not sum correctly, your answer will be wrong — this is the most common error students make (see Section 15).

Expanded Form for Percent Abundance

If you are working with percentages rather than decimals, use this equivalent form:

AM = [(m₁ × %₁) + (m₂ × %₂) + … + (mₙ × %ₙ)] ÷ 100

Why It Is a Weighted Average — Not a Simple Average

A simple average of Cl-35 and Cl-37 would give (35 + 37) ÷ 2 = 36. But chlorine's average atomic mass is 35.45 — much closer to 35 than to 37. Why? Because Cl-35 is roughly 3 times more abundant than Cl-37 in nature. The weighted average correctly reflects that imbalance.

5. Percent Abundance vs. Fractional Abundance — Clearly Explained

Both formats appear in textbooks and data tables. You need to know which one you are working with before you calculate — mixing them is one of the biggest sources of errors.

Quick conversion: Fractional abundance = Percent abundance ÷ 100

Our calculator accepts both formats. Just select the correct unit before you enter your values, and the calculator handles the conversion automatically.

6. How to Use the Average Atomic Mass Calculator

1. Select the number of isotopes your element has. Choose from 2 to 10. If you are not sure, check the isotope reference table in Section 8.
2. Choose your abundance format — percent (%) or decimal (fractional). The calculator accepts both.
3. Enter the isotopic mass of the first isotope in amu. Use precise values from the reference table for accurate results — avoid rounded whole numbers.
4. Enter the abundance of the first isotope in your chosen format.
5. Repeat for each remaining isotope.
6. Click Calculate. The tool verifies that your abundances sum correctly and then shows the weighted average atomic mass with full step-by-step working.

Pro tip: Always use precise isotope masses (e.g., 34.968853 u for Cl-35) rather than rounded values (e.g., 35 u). Rounded values introduce errors — for example, using 35 and 37 for chlorine gives 35.48 amu, while precise values give the more accurate 35.452 amu shown on the periodic table.

7. Five Fully Worked Examples

Example 1: Chlorine (Cl) — Two Isotopes

Given data:

• Cl-35: mass = 34.968853 u, natural abundance = 75.78%
• Cl-37: mass = 36.965903 u, natural abundance = 24.22%

Step 1 — Convert % to fractional (decimal) abundance:

• f₁ = 75.78 ÷ 100 = 0.7578
• f₂ = 24.22 ÷ 100 = 0.2422

Step 2 — Verify they sum to 1:

0.7578 + 0.2422 = 1.0000 ✅

Step 3 — Apply the formula:

AM = (34.968853 × 0.7578) + (36.965903 × 0.2422)

AM = 26.502 + 8.950

AM = 35.452 u ✅ (matches periodic table: 35.45 u)

⚠️ Note on precision: Using rounded masses of 35 and 37 (as some calculators do) gives 35.48 amu — a less accurate result. Always use precise isotope masses.

Example 2: Carbon (C) — Two Isotopes

• C-12: mass = 12.000000 u, natural abundance = 98.93%
• C-13: mass = 13.003355 u, natural abundance = 1.07%

Calculation:

AM = (12.000000 × 0.9893) + (13.003355 × 0.0107)

AM = 11.8716 + 0.13914

AM = 12.011 u ✅

Carbon-12 dominates so heavily (98.93%) that the average is pulled very close to 12 — the small fraction of C-13 nudges it just slightly higher to 12.011.

Example 3: Magnesium (Mg) — Three Isotopes

• Mg-24: mass = 23.985042 u, abundance = 78.99%
• Mg-25: mass = 24.985837 u, abundance = 10.00%
• Mg-26: mass = 25.982593 u, abundance = 11.01%

Verify: 78.99 + 10.00 + 11.01 = 100.00% ✅

Calculation:

AM = (23.985042 × 0.7899) + (24.985837 × 0.1000) + (25.982593 × 0.1101)

AM = 18.947 + 2.499 + 2.861

AM = 24.305 u ✅

Example 4: Bromine (Br) — Nearly Equal Isotopes

• Br-79: mass = 78.918338 u, abundance = 50.69%
• Br-81: mass = 80.916291 u, abundance = 49.31%

Calculation:

AM = (78.918338 × 0.5069) + (80.916291 × 0.4931)

AM = 40.012 + 39.900

AM = 79.904 u ✅

Bromine is a perfect example of two nearly equal-abundance isotopes pulling the average almost exactly between them — which is why the periodic table value of 79.904 sits right between 79 and 81.

Example 5: Iron (Fe) — Four Isotopes

• Fe-54: mass = 53.939609 u, abundance = 5.845%
• Fe-56: mass = 55.934938 u, abundance = 91.754%
• Fe-57: mass = 56.935394 u, abundance = 2.119%
• Fe-58: mass = 57.933276 u, abundance = 0.282%

Verify: 5.845 + 91.754 + 2.119 + 0.282 = 100.000% ✅

Calculation:

AM = (53.939609 × 0.05845) + (55.934938 × 0.91754) + (56.935394 × 0.02119) + (57.933276 × 0.00282)

AM = 3.153 + 51.333 + 1.207 + 0.163

AM = 55.845 u ✅

Iron's average is dominated almost entirely by Fe-56 (91.754% abundant) — the other three isotopes barely nudge the result.

8. Isotope Reference Table — Masses & Abundances

This table gives you the precise IUPAC isotope masses and natural abundances you need to feed directly into the average atomic mass calculator — no separate lookup required. Use these exact values for accurate results.

Source: IUPAC Commission on Isotopic Abundances and Atomic Weights (CIAAW) and NIST Atomic Weights database.

9. Average Atomic Mass of the First 40 Elements

Use this quick-reference table whenever you need the average atomic mass of a common element without running a full calculation.

Interesting pattern: Elements with atomic number 9 (Fluorine), 11 (Sodium), 13 (Aluminium), 15 (Phosphorus), 27 (Cobalt), and 33 (Arsenic) have only one stable isotope — so their average atomic mass equals the mass of that single isotope. These are called monoisotopic elements.

10. Monoisotopic Mass — The Concept No One Else Explains

Here is an important concept that neither Omni Calculator nor Calculator-Online explains on their average atomic mass pages — yet it is critical for anyone doing advanced chemistry, pharmacology, or mass spectrometry.

What Is Monoisotopic Mass?

The monoisotopic mass is the mass calculated by using only the most abundant lightest stable isotope of each element — rather than a weighted average of all isotopes.

Average Atomic Mass vs. Monoisotopic Mass: Side by Side

When Do You Use Which?

• Use average atomic mass for: stoichiometry, mole calculations, solution preparation, general chemistry coursework, and any bulk-quantity measurement. This is the periodic table value.
• Use monoisotopic mass for: mass spectrometry data interpretation, drug development and pharmacokinetics, protein and peptide analysis, high-resolution analytical chemistry.

Real-world consequence: For small molecules the difference is tiny and usually ignorable. For large biomolecules like proteins — which may contain thousands of atoms — the cumulative difference between monoisotopic and average masses can be several Daltons. Using the wrong value in mass spec analysis causes you to misidentify compounds.

11. How Scientists Measure Isotope Abundance: Mass Spectrometry

When you enter an isotope abundance into the average atomic mass calculator, that number did not come from thin air. It was measured in a laboratory using a technique called mass spectrometry. Here is how it works — in plain English.

The Five Steps of Mass Spectrometry

1. Ionization: A sample of the element is vaporized and hit with high-energy electrons. This knocks electrons off the atoms, creating positively charged ions.
2. Acceleration: The ions are pulled through an electric field. Every ion gets the same kinetic energy, which means lighter ions move faster than heavier ones.
3. Deflection: The moving ions enter a curved magnetic field. Lighter ions deflect more sharply; heavier ions deflect less. This separates them by mass-to-charge ratio (m/z).
4. Detection: A detector at the end records how many ions land at each position. This creates a mass spectrum — a graph of signal intensity vs. m/z value.
5. Calculation: The position of each peak gives the isotope mass. The height (intensity) of each peak gives the relative abundance. Scientists convert relative peak heights to percentages, then plug those into the average atomic mass formula.

Reading a Mass Spectrum for Chlorine

Chlorine's mass spectrum shows exactly two peaks:

• A tall peak at m/z = 35 — representing ³⁵Cl (about 75.78% of all chlorine)
• A shorter peak at m/z = 37 — representing ³⁷Cl (about 24.22% of all chlorine)

The peak height ratio of approximately 3:1 reflects the abundance ratio directly. Divide each height by the total to get percent abundance. Then plug into the formula and you get 35.452 u.

The takeaway: Every number in the isotope reference table above was determined by this process. When you use the average atomic mass calculator, you are completing the final step of a process that took scientists decades of experimental refinement to make precise.

12. IUPAC Standard Atomic Weights — What They Mean

The numbers on your periodic table are official figures published by IUPAC (the International Union of Pure and Applied Chemistry), updated periodically as measurement technology improves.

How IUPAC Determines Standard Atomic Weights

• They collect isotopic composition data from many natural sources worldwide — rocks, seawater, atmosphere, biological samples, and more.
• They calculate a weighted global average across all these measurements.
• Values are reviewed and revised on a rolling basis. The 2021 update is the most recent comprehensive revision.

Why Some Elements Now Have Interval Weights

Here is something most calculator pages never mention: for 10 elements — hydrogen, lithium, boron, carbon, nitrogen, oxygen, silicon, sulfur, chlorine, and argon — IUPAC now publishes an interval rather than a single value. This is because their isotopic composition genuinely varies depending on where or what you measure.

For most classroom and lab purposes, using the conventional single value (like 35.45 for chlorine) is perfectly fine. The interval matters only in high-precision isotope ratio studies such as geochemistry, climate science, and forensic analysis.

13. Elements With No Stable Isotopes — A Special Case

Some elements on the periodic table do not have a standard average atomic mass at all — because they have no stable isotopes. Every form of these elements is radioactive and eventually decays.

For these elements, the periodic table shows the mass number of the most stable known isotope in square brackets instead of a decimal average.

You cannot meaningfully calculate an average atomic mass for these elements using the standard formula because no stable naturally occurring isotope mixture exists. The average atomic mass calculator only applies to elements with stable or very long-lived isotopes.

14. Real-World Applications of Average Atomic Mass

Average atomic mass is not just a textbook number. It has direct, consequential uses across science and industry:

• 🧪 Stoichiometry & Chemical Reactions: Every mole calculation depends on average atomic mass. When you weigh out 12 grams of carbon in a lab, you are using 1 mole of carbon — a fact that only makes sense because of the defined relationship between average atomic mass (12.011 u) and molar mass (12.011 g/mol).
• 🌍 Radiocarbon Dating: Carbon-14 is a radioactive isotope with mass 14.003242 u and a half-life of 5,730 years. Scientists compare the C-14 to C-12 ratio in a sample against the known average, then use the decay rate to calculate when the organism died — accurately dating samples up to 50,000 years old.
• ⚛️ Nuclear Power: Uranium-235 (235.044 u) and Uranium-238 (238.051 u) have very different fission properties. Nuclear engineers use precise atomic mass data to calculate fuel enrichment levels, chain reaction rates, and energy output. The tiny mass difference between fission products and starting material (mass defect) directly gives the energy released via E = mc².
• 💊 Drug Development (Isotope Labeling): Pharmaceutical researchers replace hydrogen atoms with deuterium (²H, mass 2.014 u) in drug molecules to slow metabolism and improve effectiveness. The FDA has approved deuterium-labeled drugs — and every design calculation relies on precise average atomic mass and isotope mass values.
• 🔬 Food Authenticity & Fraud Detection: The carbon, oxygen, and strontium isotopic ratios in food vary based on geography and growing conditions. Forensic food scientists use these ratios (derived from average atomic mass calculations) to verify the origin of olive oil, honey, wine, and other products — detecting fraud and mislabeling.
• 🏭 Semiconductor Manufacturing: High-purity silicon chips are increasingly made with isotopically enriched Si-28 (monoisotopic silicon). The elimination of Si-29 and Si-30 reduces lattice strain and phonon scattering, improving chip performance at the nanoscale. This application requires precise knowledge of each isotope's mass and the natural average.
• 🚀 Cosmochemistry & Planetary Science: Scientists analyze isotope ratios in meteorites, moon rocks, and Mars samples to understand solar system formation. The specific average atomic masses of iron, silicon, and nickel isotopes act as chemical fingerprints that date cosmic events billions of years ago.

15. Seven Common Mistakes to Avoid

1. Using rounded mass numbers instead of precise isotope masses.
   Entering 35 for Cl-35 instead of 34.968853 u introduces an error of ~0.03 u per isotope. This adds up quickly in multi-isotope calculations. Always use values from a precise reference table.
2. Forgetting to convert percent abundance to decimal form.
   If your formula uses fractional abundance (0 to 1), entering 75.78 instead of 0.7578 makes your result 100× too large. Always check which format the formula or calculator requires before inputting values.
3. Abundances that do not sum to 100% (or 1.0).
   If your isotope abundances sum to 99% or 101%, your result is proportionally wrong. Always check the sum before clicking Calculate. Our calculator flags this error automatically.
4. Confusing average atomic mass with the mass number.
   The mass number of carbon-12 is exactly 12 (a whole number). The average atomic mass of carbon is 12.011 (a decimal). These are different quantities used in different contexts. Using 12 in stoichiometry where 12.011 is needed introduces a consistent rounding error.
5. Calculating a simple average instead of a weighted average.
   A simple average of Cl-35 and Cl-37 gives (35 + 37) / 2 = 36 — which is completely wrong. The correct weighted average is 35.45 because Cl-35 is roughly 3× more abundant than Cl-37. Always weight by natural abundance.
6. Using average atomic mass when monoisotopic mass is needed.
   In mass spectrometry, you need monoisotopic mass. Using the periodic table average value shifts your expected m/z peaks, causing compound misidentification. Know which value your application demands before you calculate.
7. Ignoring the IUPAC interval for variable-composition elements.
   For hydrogen, carbon, chlorine, and 7 other elements, IUPAC now assigns a range rather than a single value. For high-precision work (geochemistry, isotope ratio analysis), using a fixed single value introduces a real uncertainty. Know when the interval matters for your application.

16. Brief History of Atomic Mass

• 1803 — John Dalton (England): First proposed a systematic table of relative atomic masses, using hydrogen (H = 1) as the reference. His values were rough but revolutionary — the first quantitative atomic model in history.
• 1860 — Stanislao Cannizzaro (Italy): Resolved the great confusion between atomic and molecular masses using Avogadro's hypothesis. His landmark lecture at the Karlsruhe Congress gave chemists a coherent, agreed-upon set of atomic weights for the first time.
• 1869 — Dmitri Mendeleev (Russia): Arranged all known elements by increasing atomic mass and discovered that chemical properties repeat periodically — creating the periodic table. Atomic mass was the very spine of his organizational system.
• 1913 — Frederick Soddy (UK): Discovered isotopes — atoms of the same element with different masses. This instantly explained why so many atomic masses were not whole numbers: they were averages of isotopic mixtures. Soddy won the Nobel Prize in Chemistry in 1921 for this discovery.
• 1919 — Francis Aston (UK): Built the first high-precision mass spectrograph and directly measured isotope masses and abundances for over 200 isotopes. He showed that individual isotope masses deviate slightly from whole numbers due to nuclear binding energy — the "mass defect." He won the Nobel Prize in Chemistry in 1922.
• 1961 — IUPAC International Standard: Carbon-12 was officially adopted as the universal atomic mass standard, replacing the earlier oxygen-16 reference. The definition: 1 u = exactly 1/12 the mass of ¹²C. This remains the international standard today.
• 2016 — IUPAC Interval Weights: IUPAC introduced interval atomic weights for 10 elements, acknowledging that their isotopic composition varies measurably across different Earth sources. A single fixed value is no longer considered scientifically precise enough for these elements in high-accuracy work.

17. Frequently Asked Questions (14 FAQs)

Q1: What is average atomic mass?

Average atomic mass is the weighted average of the masses of all naturally occurring isotopes of an element, calculated by multiplying each isotope's mass by its fractional abundance and summing the results. It is the value shown on the periodic table and numerically equals an element's molar mass in g/mol.

Q2: How do you calculate average atomic mass step by step?

Follow four steps: (1) Write down each isotope's precise mass in amu. (2) Convert each isotope's percent abundance to a decimal by dividing by 100. (3) Multiply each isotope's mass by its fractional abundance. (4) Add all the products together. The result is the average atomic mass.

Q3: What is the formula for average atomic mass?

AM = (f₁ × m₁) + (f₂ × m₂) + … + (fₙ × mₙ), where f is the fractional (decimal) abundance of each isotope and m is its mass in amu. All f values must sum to exactly 1.0 for an accurate result.

Q4: Why is atomic mass not a whole number?

Because the value on the periodic table is a weighted average of multiple isotopes, not the mass of a single atom. Even if each individual isotope had a mass very close to a whole number, the weighted blend of several isotopes at different abundances almost always produces a decimal result.

Q5: What is the difference between average atomic mass and molar mass?

They are numerically identical but differ in scale and unit. Average atomic mass describes one atom, measured in amu (u). Molar mass describes one mole (6.022 × 10²³ atoms) of that element, measured in g/mol. For carbon: average atomic mass = 12.011 u/atom; molar mass = 12.011 g/mol. Same number, different context.

Q6: What is percent abundance?

Percent abundance is the percentage of atoms of a given isotope found naturally on Earth. For example, Cl-35 has 75.78% percent abundance — meaning 75.78 out of every 100 chlorine atoms in any natural sample are the Cl-35 isotope. All isotope percent abundances for one element must sum to exactly 100%.

Q7: What is fractional abundance?

Fractional abundance is percent abundance expressed as a decimal. Simply divide the percentage by 100: 75.78% becomes 0.7578. All fractional abundances for an element must sum to exactly 1.0. Both formats are valid in the average atomic mass formula — just never mix them in the same calculation.

Q8: Why does my calculated average atomic mass not match the periodic table?

The most common causes are: (1) using rounded mass numbers (e.g., 35) instead of precise isotope masses (e.g., 34.968853 u); (2) abundances that do not sum to 100%; (3) mixing percent and decimal formats in the same calculation; or (4) missing a minor isotope. Check each of these against the isotope reference table in Section 8.

Q9: What element has the most stable isotopes?

Tin (Sn) holds the record with 10 stable isotopes: Sn-112, Sn-114, Sn-115, Sn-116, Sn-117, Sn-118, Sn-119, Sn-120, Sn-122, and Sn-124. Its average atomic mass is 118.710 u. This is why a full-featured average atomic mass calculator must support at least 10 isotopes.

Q10: What is the average atomic mass of chlorine?

Chlorine's average atomic mass is 35.45 u. It has two stable isotopes: Cl-35 (mass 34.968853 u, abundance 75.78%) and Cl-37 (mass 36.965903 u, abundance 24.22%). Using the formula: AM = (34.968853 × 0.7578) + (36.965903 × 0.2422) = 35.452 u, which rounds to 35.45 on the periodic table.

Q11: What is the difference between monoisotopic mass and average atomic mass?

Average atomic mass is a weighted average of all naturally occurring isotopes — it is the periodic table value used in everyday chemistry. Monoisotopic mass uses only the mass of the lightest most-abundant stable isotope of each element. For small molecules the difference is tiny. For large biomolecules or in mass spectrometry, using the wrong one can cause significant calculation errors.

Q12: How is isotope abundance measured experimentally?

Through mass spectrometry. The element is ionized, accelerated, and deflected by a magnetic field. Ions separate by mass-to-charge ratio — lighter isotopes deflect more, heavier ones less. A detector records the signal intensity at each m/z position, giving peak heights that reflect relative abundance. These are converted to percent abundances and fed into the average atomic mass formula.

Q13: Do all elements have an average atomic mass?

No. Elements with no stable isotopes — like technetium (Tc), promethium (Pm), and all artificial superheavy elements — do not have a meaningful average atomic mass because no stable, naturally occurring mixture of isotopes exists. The periodic table shows the mass number of the most stable known isotope in square brackets for these elements (e.g., [98] for technetium).

Q14: Why do some elements have interval atomic weights on the periodic table?

For 10 elements — including hydrogen, carbon, nitrogen, oxygen, chlorine, and sulfur — the isotopic composition varies measurably across different natural samples (different waters, soils, rock types, organisms). A single fixed value cannot represent this natural variability accurately. IUPAC now publishes a range (interval) for these elements, such as [1.00784; 1.00811] for hydrogen. For typical classroom use, a single representative value (like 1.008) is sufficient.

18. Summary & Key Takeaways

• Average atomic mass is the weighted average of all naturally occurring isotopes of an element, based on their percent abundances. It is the value shown on the periodic table.
• The formula is: AM = (f₁ × m₁) + (f₂ × m₂) + … + (fₙ × mₙ), where f values are fractional (decimal) abundances that must sum to 1.0.
• Average atomic mass ≠ mass number. The mass number is a whole-number count of protons + neutrons for one specific isotope. Average atomic mass is a decimal weighted average across all isotopes.
• Average atomic mass = molar mass numerically — same number, different unit (u vs. g/mol). This makes it the foundation of all stoichiometry.
• Always use precise IUPAC isotope masses (e.g., 34.968853 u for Cl-35), not rounded values (35 u). Rounded values produce less accurate results.
• Monoisotopic mass uses only the lightest stable isotope of each element — essential for mass spectrometry and advanced analytical chemistry, but not for everyday stoichiometry.
• Isotope abundances are measured experimentally using mass spectrometry — a technique that separates ions by mass-to-charge ratio and measures peak intensities as relative abundances.
• IUPAC publishes standard atomic weights. For 10 elements (H, C, N, O, Si, S, Cl, Ar, and others), they now publish a weight interval rather than a fixed value because isotopic composition varies in nature.
• Elements with no stable isotopes (like Tc, Pm, and superheavy elements) do not have a standard average atomic mass. The periodic table shows the mass number of their most stable isotope in brackets instead.
• Real-world applications span stoichiometry, radiocarbon dating, nuclear power, drug development, food forensics, semiconductor manufacturing, and cosmochemistry.
• The top mistakes are: using rounded masses, mixing percent and decimal formats, abundances not summing to 100%, confusing mass number with average atomic mass, and using average atomic mass when monoisotopic mass is needed.
• Tin (Sn) has the most stable isotopes of any element at 10 — the reason a quality average atomic mass calculator must support up to 10 isotopes.