Isotopes and atomic mass worksheet with answers
A worksheet on isotopes and atomic mass with answers should help students move from particle diagrams to calculations without making the periodic-table value seem like a mystery. This version combines definitions, notation, numerical examples and a clear answer guide for introductory secondary science.
Isotopes are atoms of the same element with identical proton counts but different neutron counts. The abundance of those isotopes explains why an element’s reported atomic mass is usually a weighted average rather than the mass number of its most common isotope.
Australian schools operate under state or territory curricula, so teachers in Adelaide, Brisbane, Melbourne or Perth often adapt free resources for mixed-ability classes. Printed worksheets and reusable files can help where departmental budgets are limited. Anyone redistributing an adapted MyFundRazor copy should first read the template usage terms.
How isotopes and atomic mass work
The atomic number, written as Z, equals the number of protons and identifies the element. Neutron number is N, while the mass number A is the total number of protons and neutrons: A = Z + N. Changing neutrons can therefore produce a different isotope of the same element, but it cannot change the element’s identity.
Isotope symbols put the mass number before the atomic number, as in ¹⁴₆C for carbon-14. Read this as carbon with six protons and eight neutrons. A mass number is always a whole number and belongs to one specific isotope. The atomic mass printed on a periodic table may contain decimals because it combines isotope masses according to their natural abundances.
Relative atomic mass is the weighted average of the masses of an element’s naturally occurring isotopes. The dalton, symbol u, is used for atomic mass, and one u is defined as one-twelfth of the mass of a carbon-12 atom. This standard makes precise comparisons between atoms, materials and laboratory samples possible.
Worksheet questions and answer guide
Complete the questions without notes, show each subtraction or addition, and then use the guide to check your reasoning. For isotope questions, begin with the number of protons before calculating neutrons.
- Define an isotope in one sentence.
- Hydrogen-2 has one proton. How many neutrons does it contain?
- Give the number of protons, neutrons and mass number for carbon-12.
- A chlorine atom has 17 protons and 20 neutrons. Identify the isotope.
- Why is a carbon-12 atom assigned a mass of exactly 12 u, while chlorine’s reported atomic mass is about 35.45?
- Two atoms each have eight protons. One has eight neutrons and the other has ten. Are they the same isotope?
- Uranium-238 contains 92 protons. Calculate its neutron number.
- If carbon-13 accounts for 1.07% of natural carbon, what does that percentage describe?
The answers are:
- Atoms of the same element with the same number of protons but different numbers of neutrons.
- One neutron, because 2 − 1 = 1.
- Six protons, six neutrons and a mass number of 12.
- Chlorine-37, because 17 + 20 = 37.
- Carbon-12 defines the atomic mass scale, while chlorine’s value is a weighted average of chlorine-35 and chlorine-37.
- No. They are oxygen-16 and oxygen-18, two different isotopes of oxygen.
- 146 neutrons, because 238 − 92 = 146.
- The percentage of carbon atoms in a natural sample that are carbon-13, rather than their share of the sample’s mass.
Students should receive full credit for correct working and sound explanations, even when their wording differs from the guide. For question 8, distinguish the count of atoms from the percentage of mass. A percentage close to 1 per cent still represents a very large number of atoms in a macroscopic sample.
Comparing familiar isotopes
Use these reference values to compare how isotopes differ. Abundances are approximate natural percentages and may vary slightly with the source or material.
| Isotope | Atomic number | Protons | Neutrons | Mass number | Natural abundance | Notes |
|---|---|---|---|---|---|---|
| Hydrogen-1 | 1 | 1 | 0 | 1 | 99.985% | Stable |
| Hydrogen-2 | 1 | 1 | 1 | 2 | 0.015% | Stable |
| Carbon-12 | 6 | 6 | 6 | 12 | 98.93% | Stable reference isotope |
| Carbon-13 | 6 | 6 | 7 | 13 | 1.07% | Stable |
| Carbon-14 | 6 | 6 | 8 | 14 | Trace | Radioactive |
| Oxygen-16 | 8 | 8 | 8 | 16 | 99.76% | Stable |
| Oxygen-17 | 8 | 8 | 9 | 17 | 0.04% | Stable |
| Oxygen-18 | 8 | 8 | 10 | 18 | 0.20% | Stable |
| Chlorine-35 | 17 | 17 | 18 | 35 | 75.78% | Stable |
| Chlorine-37 | 17 | 17 | 20 | 37 | 24.22% | Stable |
The oxygen entries show the central rule clearly: oxygen-16 and oxygen-18 both have eight protons, so both are oxygen. Their neutron counts differ, producing mass numbers 16 and 18. Carbon-14 has the same proton count as carbon-12 and carbon-13, but its extra neutrons make it radioactive.
The chlorine entries explain how fractional abundance can produce a non-whole atomic mass. A sample with roughly three-quarters chlorine-35 and one-quarter chlorine-37 must have a weighted average between 35 and 37, rather than either isotope’s mass alone.
Calculating atomic mass from abundances
The calculation follows one rule: average atomic mass equals the sum of each isotope mass multiplied by its fractional abundance. A percentage must be divided by 100 before it is used. Always label the result in u, and keep enough decimal places for the marks.
For a simple chlorine example, suppose 75% of atoms are chlorine-35 and 25% are chlorine-37. Using mass numbers gives (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.50 u. Using precise isotope masses and standard abundances gives the familiar periodic-table value of about 35.45 u, so the small difference is expected.
Boron provides another check: boron-10 has an abundance of 19.9%, and boron-11 has 80.1%. Its average is (10 × 0.199) + (11 × 0.801) = 1.990 + 8.811 = 10.801 u, which rounds to 10.81 u. The abundance-weighted result, not a simple midpoint, is the required atomic mass.
Australian classroom connections
Australian science lessons commonly place isotopes in Years 7 to 10 chemistry, although the exact placement varies across state and territory curricula. Australian spelling uses “practise” as a verb and “aluminium” in chemistry. In classroom talk, a teacher may encourage students to “give it a go” or say “no worries, check it again”, keeping the tone supportive while expectations remain clear.
Isotopes also connect to Australian science and industry. The Australian Nuclear Science and Technology Organisation at Lucas Heights in Sydney produces radioisotopes for medicine, research and industry. Uranium’s history in the Northern Territory and uranium-bearing zircon in Western Australia’s mineral sands provide local examples for discussing radioactive decay and geological dating.
School laboratories use local safety procedures and risk assessments, so worksheets should distinguish paper calculations from practical activities involving sources or equipment. Students extending the task into a sourced report can use the academic publishing guidance to check citation and academic presentation. Linking the chemistry to places such as Canberra’s universities or Sydney’s research facilities can make the topic feel relevant without changing the core calculations.
Use, adapt, and share the worksheet
Use the worksheet as a pre-topic check, then repeat the isotope calculation after teaching weighted averages. Let students mark their own answers, revise incorrect working and explain one result to a partner. The guide makes feedback quick for the teacher and useful for students who were absent.
Copy the questions into a free MyFundRazor spreadsheet or worksheet format, substitute another element if needed, and preserve the worked calculations for reference. Printable A4 copies suit students without reliable home internet, while reusable digital files support blended classrooms and regional learning hubs.
Print the sheet, attempt it without notes, check every response against the guide and share the completed resource with your department or a study group. That simple cycle gives Australian learners a free, repeatable way to master isotope notation and atomic mass.