MGU Science Series — High School Chemistry
Part II: Atoms, Elements, and the Periodic Table
Protons · Neutrons · Isotopes · Ions · Average Atomic Mass
Chapter 6 of 28
MGU Science Series · High School Chemistry · Library Textbook Edition
Chapter 6:
Atomic Structure
and Isotopes
Protons, neutrons, electrons, nuclear symbols, isotopes, average atomic mass, ions, and the atomic mass unit — the full language of atomic identity
■ Part II: Atoms, Elements, and the Periodic Table ■
■ 9 Sections • Chapter 6 of 28 ■
Follows Chapter 5: Atoms and Atomic Theory · Prepares students for Chapter 7: Electrons and Atomic Structure
Chapter 6 — Front Matter
Purpose and Overview
In Chapter 5, you learned that atoms are real particles with a nucleus, and that protons, neutrons, and electrons are their three main components. You learned that the atomic number — the number of protons — is what makes one element different from every other element.
But atomic structure goes deeper than that introduction suggested. Atoms of the same element do not always have the same number of neutrons. Atoms can gain or lose electrons, changing their electrical charge. The mass of an atom can be expressed in special units that make calculations practical. And atomic mass values on the periodic table are not whole numbers — they are weighted averages that reflect the natural mixture of atom types found in any real sample of an element.
This chapter builds that deeper understanding. You will learn to use nuclear symbols to describe any atom precisely. You will learn what isotopes are, why they exist, and why they matter. You will calculate average atomic mass the way chemists do. You will learn what ions are and how charged atoms form. And you will see how isotopes connect to real work in medicine, archaeology, environmental science, and research.
Opening Question
The periodic table lists carbon's atomic mass as 12.011 — not 12. Why not a clean whole number? And why do some carbon atoms behave identically to others in every chemical reaction, yet have different masses? By the end of this chapter, you will understand exactly why, and you will be able to calculate that 12.011 value yourself from scratch.
9Sections
10Objectives
15Key Terms
9+Worked Examples
Learning Objectives
- 1Identify protons, neutrons, and electrons by location, charge, and mass.
- 2Explain how the atomic number defines element identity.
- 3Calculate the number of neutrons in any atom given its atomic number and mass number.
- 4Define isotopes and explain why atoms of the same element can have different mass numbers.
- 5Read and write nuclear symbols for specific isotopes.
- 6Calculate the average atomic mass of an element from isotope masses and natural abundances.
- 7Determine the charge on an ion given its proton and electron counts.
- 8Distinguish cations from anions and explain how each forms.
- 9Define the atomic mass unit and explain its reference standard.
- 10Describe real-world applications of isotopes in medicine, dating, and research.
Chapter 6 — Contents
Sections at a Glance
- 6.1The Parts of the AtomProtons, neutrons, electrons; nucleus location; charges; relative masses; neutral atom summary
- 6.2Atomic Number and Element IdentityWhy proton count defines the element; atomic number on the periodic table; electrons in neutral atoms
- 6.3Mass Number and IsotopesMass number defined; isotopes defined; naming isotopes; finding neutrons from mass number; worked examples
- 6.4Nuclear SymbolsNuclear symbol format; reading and writing nuclear symbols; why both numbers appear; worked examples
- 6.5Average Atomic MassNatural abundance; weighted average calculation; what the value means; reasonableness checks; worked examples
- 6.6Ions and Electric ChargeGaining and losing electrons; calculating ion charge; ions vs. element identity; worked examples
- 6.7Cations and AnionsCations defined; anions defined; common examples; naming conventions; memory strategies
- 6.8Atomic Mass UnitsWhy amu exists; definition of one amu; particle masses in amu; atomic mass on the periodic table
- 6.9Isotopes in Medicine, Dating, and ResearchStable vs. radioactive isotopes; medical imaging; carbon-14 dating; rock dating; ecological tracing; cancer treatment
Sections 6.1 – 6.2
Inside the Atom: Particles and Identity
Parts of the Atom · Atomic Number · Element Identity
Section 6.1
The Parts of the Atom
Every atom in the universe is made of three kinds of particles: protons, neutrons, and electrons. These three particles have different locations, different masses, and different electrical charges. Understanding how they work together is the foundation of everything else in chemistry.
The Nucleus
At the center of every atom is the nucleus. The nucleus is extraordinarily small compared to the total size of the atom, but it contains almost all of the atom's mass. If an atom were the size of a football stadium, the nucleus would be roughly the size of a marble on the fifty-yard line.
The nucleus contains two kinds of particles: protons and neutrons. These particles are called nucleons because they live in the nucleus. Protons carry a positive electrical charge of +1. Neutrons carry no electrical charge at all — they are electrically neutral, which is how they got their name.
Electrons
Electrons occupy the space outside the nucleus. They move through a large region surrounding the nucleus called the electron cloud. Each electron carries a charge of −1. That negative charge exactly balances a proton's +1 charge, which means a neutral atom has equal numbers of protons and electrons.
| Particle | Location | Charge | Relative Mass |
|---|
| Proton | Nucleus | +1 | 1 amu |
| Neutron | Nucleus | 0 | 1 amu |
| Electron | Electron cloud | −1 | 1/1836 amu (≈ 0) |
Key Insight — Where the Mass Lives
Electrons weigh almost nothing compared to protons and neutrons. For all practical purposes, the mass of an atom comes entirely from the protons and neutrons in its nucleus. Electrons contribute so little that they are ignored in almost every mass calculation.
A Neutral Atom — Three Defining Features
Protons — determine what element the atom is. Never changes without a nuclear reaction.
Neutrons — contribute to mass. Can vary among atoms of the same element.
Electrons — equal to protons in a neutral atom, giving zero net charge.
Section 6.2
Atomic Number and Element Identity
The single most important number associated with any atom is its atomic number. The atomic number is the number of protons in the nucleus. No two elements share the same atomic number. Every atom with exactly 6 protons is carbon. Every atom with exactly 79 protons is gold. If you change the number of protons, you change the element — completely.
The periodic table is arranged in order of increasing atomic number. When you look at carbon's entry, the number 6 near its symbol is the atomic number. In a neutral atom, the number of electrons always equals the atomic number, because the positive and negative charges must balance.
| Element | Atomic Number | Protons (neutral) | Electrons (neutral) |
|---|
| Carbon | 6 | 6 | 6 |
| Sodium | 11 | 11 | 11 |
| Chlorine | 17 | 17 | 17 |
| Gold | 79 | 79 | 79 |
Common Misconception
Gaining or losing electrons does not change what element an atom is — that is set by the proton count alone. Only a nuclear reaction that adds or removes protons can change one element into another.
Sections 6.3 – 6.5
Isotopes, Nuclear Symbols, and Average Atomic Mass
Mass Number · Isotopes · Nuclear Symbols · Weighted Averages
Section 6.3
Mass Number and Isotopes
The mass number is the sum of the protons and neutrons in the nucleus of a specific atom. Because protons and neutrons each have a mass of approximately 1 amu, the mass number is also a close approximation of the atom's mass in atomic mass units.
Key Formulas
Mass number = Number of protons + Number of neutrons
Number of neutrons = Mass number − Atomic number
Atoms of the same element that have different numbers of neutrons are called isotopes. Isotopes are chemically nearly identical because chemical behavior depends on electrons, and isotopes of the same element have the same number of electrons. They differ only in mass. Isotopes are named by their element name followed by their mass number: carbon-12, carbon-13, carbon-14.
Worked Example 6.1
Finding the number of neutrons
- ProblemChlorine has atomic number 17. A particular chlorine atom has a mass number of 35. How many neutrons does it have?
- Solve
Neutrons = Mass number − Atomic number
Neutrons = 35 − 17 = 18
- Check17 protons + 18 neutrons = 35 total nucleons = mass number 35. ✓
Answer: This chlorine-35 atom has 17 protons, 18 neutrons, and (if neutral) 17 electrons.
Worked Example 6.2
Identifying isotopes
- ProblemAtom A: 7 protons, 7 neutrons. Atom B: 7 protons, 8 neutrons. Are they isotopes?
- ThinkBoth have 7 protons → both are nitrogen. Their neutron counts differ (7 vs. 8), giving mass numbers 14 and 15.
- SolveSame element, different neutron count → same element, different mass number = isotopes.
Answer: Yes — nitrogen-14 and nitrogen-15 are isotopes of each other.
Section 6.4
Nuclear Symbols
The nuclear symbol is the standardized notation that communicates an isotope's element, mass number, and atomic number in one compact form. The mass number (A) appears as a superscript on the upper left of the element symbol; the atomic number (Z) appears as a subscript on the lower left.
Nuclear Symbol Format
ᴬX → mass number upper-left, atomic number lower-left, element symbol center
ᴢ
Reading the symbol gives you: element identity (symbol), proton count (Z), total nucleons (A), and by subtraction, the neutron count (A − Z).
Worked Example 6.3
Reading a nuclear symbol
- ProblemA nuclear symbol shows 238 on the upper left, 92 on the lower left, and the symbol U. Identify all atomic information.
- Solve
Element: Uranium (U)
Mass number: 238
Atomic number: 92 (protons)
Neutrons: 238 − 92 = 146
Electrons: 92 (if neutral)
Answer: Uranium-238. 92 protons, 146 neutrons, 92 electrons (neutral).
Worked Example 6.4
Writing a nuclear symbol
- ProblemAn oxygen atom has 8 protons and 10 neutrons. Write its nuclear symbol.
- Solve
Element symbol: O
Atomic number: 8
Mass number: 8 + 10 = 18
Isotope name: Oxygen-18
- ResultWrite 18 as the upper-left superscript, 8 as the lower-left subscript, and O as the symbol.
Answer: Oxygen-18 (¹⁸₈O). Mass number 18, atomic number 8.
Section 6.5
Average Atomic Mass
Most elements exist in nature as a mixture of isotopes. The percentages in which each isotope occurs in a natural sample are called its natural abundance. Because a natural sample of any element contains multiple isotopes in specific proportions, the periodic table reports a weighted average mass — not the mass of any single isotope, but the typical mass of an atom in a naturally occurring sample.
Average Atomic Mass Formula
Average atomic mass = (mass₁ × fraction₁) + (mass₂ × fraction₂) + ...
Fractional abundance = percentage ÷ 100. Multiply each isotope's mass by its fractional abundance, then sum all products.
Worked Example 6.5
Calculating average atomic mass — carbon
- GivenCarbon-12: mass = 12.000 amu, abundance = 98.89%. Carbon-13: mass = 13.003 amu, abundance = 1.11%.
- Step 1Convert to decimal fractions: 0.9889 and 0.0111.
- Step 2
Carbon-12: 12.000 × 0.9889 = 11.867 amu
Carbon-13: 13.003 × 0.0111 = 0.144 amu
- Step 3
11.867 + 0.144 = 12.011 amu
Answer: 12.011 amu — matches the periodic table value for carbon.
Worked Example 6.6
Calculating average atomic mass — chlorine
- GivenChlorine-35: 34.969 amu, 75.77%. Chlorine-37: 36.966 amu, 24.23%.
- Solve
Cl-35: 34.969 × 0.7577 = 26.496 amu
Cl-37: 36.966 × 0.2423 = 8.957 amu
Total: 26.496 + 8.957 = 35.453 amu
Answer: 35.45 amu. Because Cl-35 is ~75% abundant, the average is pulled close to 35 — as expected.
Reasonableness Check
The average atomic mass must always fall between the lightest and heaviest isotopes, and must be pulled toward the more abundant one. If your answer falls outside the range of isotope masses, recheck your arithmetic.
Sections 6.6 – 6.8
Ions, Charge, and Atomic Mass Units
Ions and Charge · Cations and Anions · The Atomic Mass Unit
Section 6.6
Ions and Electric Charge
A neutral atom has equal numbers of protons and electrons. When an atom gains or loses electrons, this balance breaks, and the atom acquires a net electrical charge. An atom that has gained or lost electrons is called an ion.
Ion Charge Formula
Charge = Number of protons − Number of electrons
Positive result → positive ion (more protons than electrons)
Negative result → negative ion (more electrons than protons)
Worked Example 6.7
Finding ion charge
- ProblemA magnesium atom (atomic number 12) loses 2 electrons. What is its charge?
- Solve
Protons: 12
Electrons: 12 − 2 = 10
Charge: 12 − 10 = +2
Answer: The magnesium ion has a charge of 2+. Written Mg²⁺.
Worked Example 6.8
Finding electrons in an ion
- ProblemA sulfur ion (atomic number 16) has a charge of 2−. How many electrons does it have?
- Solve
Charge = Protons − Electrons
−2 = 16 − Electrons
Electrons = 16 + 2 = 18
Answer: The sulfide ion has 18 electrons.
Section 6.7
Cations and Anions
Ions are divided into two categories: cations (positively charged, formed when atoms lose electrons) and anions (negatively charged, formed when atoms gain electrons). Metals tend to lose electrons and form cations. Nonmetals tend to gain electrons and form anions.
| Ion | Symbol | Protons | Electrons | Type |
|---|
| Sodium ion | Na⁺ | 11 | 10 | Cation |
| Magnesium ion | Mg²⁺ | 12 | 10 | Cation |
| Aluminum ion | Al³⁺ | 13 | 10 | Cation |
| Chloride ion | Cl⁻ | 17 | 18 | Anion |
| Oxide ion | O²⁻ | 8 | 10 | Anion |
| Sulfide ion | S²⁻ | 16 | 18 | Anion |
Naming Convention
Cations keep the element name: the sodium ion, the calcium ion. Anions formed from a single element change their ending to -ide: chloride, fluoride, oxide, sulfide, nitride. You will use this throughout Chapters 9 and 11.
Worked Example 6.9
Identifying cations and anions
- Problem(a) A fluorine atom (atomic number 9) gains 1 electron. (b) A calcium atom (atomic number 20) loses 2 electrons. Identify each.
- Solve (a)Gained electron → extra negative charge → anion. Fluoride ion, F⁻.
- Solve (b)Lost electrons → more protons than electrons → cation. Calcium ion, Ca²⁺.
Answer: F⁻ is an anion. Ca²⁺ is a cation.
Section 6.8
Atomic Mass Units
A proton has a mass of approximately 1.67 × 10⁻²⁴ grams — an inconveniently small number for everyday chemistry. To avoid working with such tiny values, chemists use the atomic mass unit (amu). One atomic mass unit is defined as exactly one-twelfth the mass of a carbon-12 atom. Carbon-12 was chosen as the reference because it is stable, abundant, and precisely measurable.
| Particle | Mass in amu | Note |
|---|
| Proton | 1.0073 amu | ≈ 1 amu for most calculations |
| Neutron | 1.0087 amu | ≈ 1 amu for most calculations |
| Electron | 0.000549 amu | Negligible; ignored in mass calculations |
| Carbon-12 | 12.000 amu (exact) | The reference standard |
Why Mass Numbers Are Good Approximations
Because each proton and neutron contributes almost exactly 1 amu, the mass number gives a very close estimate of any atom's mass in amu. The precise atomic masses on the periodic table reflect small corrections and isotope averaging, but mass numbers are excellent working approximations.
Section 6.9
Isotopes in the Real World
Medicine · Archaeology · Geology · Ecology · Cancer Treatment
Section 6.9
Isotopes in Medicine, Dating, and Research
Isotopes are not just a theoretical concept. They have powerful real-world applications in medicine, archaeology, environmental science, geology, and fundamental research. Some isotopes are stable — they do not change over time. Others are unstable and eventually release energy and particles as their nuclei break apart. These unstable isotopes are called radioactive isotopes, or radioisotopes, and they are especially useful in medicine and dating.
Medical Imaging with Isotopes
In nuclear medicine, radioactive isotopes called tracers can be followed as they move through the body. Technetium-99m is given to patients (usually by injection) and detected by a camera that senses its radiation. Doctors can watch where it travels to identify problems in organs, bones, the heart, and the thyroid gland — millions of scans every year rely on this isotope. Iodine-131 concentrates specifically in the thyroid, making it useful for diagnosing thyroid disorders and treating thyroid cancer. PET scans use fluorine-18 to image metabolic activity in the brain, heart, and other organs.
Carbon-14 Dating
Carbon-14 is a radioactive isotope that forms naturally in the upper atmosphere. As long as a plant or animal is alive, it continuously takes in carbon — including a small, predictable fraction of carbon-14. When the organism dies, it stops taking in new carbon, and the carbon-14 already present begins to decay at a known rate. By measuring how much carbon-14 remains in an ancient sample, scientists can calculate when the organism died. This technique works for organic materials up to about 50,000 years old and has been used to date ancient wooden artifacts, charcoal, bone, fabric, and plant remains.
Radiometric Dating of Rocks
For materials millions or billions of years old, geologists use isotopes with much longer decay times. Uranium-238 decays to lead-206 over billions of years and is used to date ancient rocks and meteorites. Potassium-40 decays to argon-40 and is used to date volcanic rocks. These techniques have established that Earth is approximately 4.5 billion years old.
Stable Isotopes in Ecological Research
Stable isotopes that do not decay are also research tools. Different isotopes of the same element behave slightly differently in physical and chemical processes, allowing scientists to track them through ecosystems, food webs, water cycles, and climate records. The ratio of oxygen-18 to oxygen-16 in ancient ice cores reveals past temperatures. The ratio of nitrogen isotopes in animal tissues reveals what the animal ate. Carbon isotope ratios in ancient human bones have helped archaeologists reconstruct prehistoric diets.
| Application | Isotope Example | How It Works |
|---|
| Medical imaging | Technetium-99m | Detected as it moves through the body |
| Thyroid treatment | Iodine-131 | Concentrates in thyroid; destroys abnormal cells |
| PET scans | Fluorine-18 | Images metabolic activity in organs |
| Archaeological dating | Carbon-14 | Decay rate reveals age of organic material |
| Rock dating | Uranium-238 | Decay rate reveals age of geological formations |
| Ecological tracing | Oxygen-18, Carbon-13 | Stable ratios reveal climate and diet history |
| Cancer treatment | Cobalt-60, Iodine-131 | Radiation damages and destroys cancer cells |
Chapter 6 — Back Matter
Chapter Summary
- Sec 6.1Atoms consist of protons (+1, ≈1 amu) and neutrons (0, ≈1 amu) in the nucleus, and electrons (−1, ≈0 amu) in the electron cloud. Mass is almost entirely in the nucleus.
- Sec 6.2The atomic number — number of protons — defines the element. No two elements share an atomic number. In a neutral atom, electrons equal protons.
- Sec 6.3The mass number is protons + neutrons. Isotopes are atoms of the same element with different neutron counts and therefore different mass numbers. Neutrons = mass number − atomic number.
- Sec 6.4The nuclear symbol places the mass number (upper left) and atomic number (lower left) next to the element symbol, encoding full atomic identity in one compact form.
- Sec 6.5Average atomic mass is a weighted average of all naturally occurring isotopes, each multiplied by its fractional abundance. It represents the typical mass in a natural sample, not the mass of any single atom.
- Sec 6.6Ions form when atoms gain or lose electrons. Charge = protons − electrons. Gaining electrons gives a negative charge; losing electrons gives a positive charge. The element identity (proton count) is unchanged.
- Sec 6.7Cations are positively charged ions (lost electrons). Anions are negatively charged ions (gained electrons). Metals tend to form cations; nonmetals tend to form anions. Anion names end in -ide.
- Sec 6.8One atomic mass unit is defined as exactly one-twelfth the mass of a carbon-12 atom. Protons and neutrons each mass ≈1 amu; electrons are negligible. Periodic table masses are in amu.
- Sec 6.9Isotopes have wide real-world use: medical imaging (technetium-99m), cancer treatment (iodine-131), archaeological dating (carbon-14), geological dating (uranium-238), and ecological tracing (stable isotope ratios).
Key Terms Review
atomic number
The number of protons in the nucleus of an atom. Defines the element; never changes for a given element.
mass number
The total number of protons and neutrons in the nucleus of a specific atom. Mass number = protons + neutrons.
isotopes
Atoms of the same element with the same number of protons but different numbers of neutrons, giving different mass numbers.
nuclear symbol
Standardized notation showing the element symbol with mass number as upper-left superscript and atomic number as lower-left subscript.
natural abundance
The percentage of each isotope found in a naturally occurring sample of an element.
average atomic mass
The weighted average of the masses of all naturally occurring isotopes, accounting for their relative abundances. Listed on the periodic table.
ion
An atom that has gained or lost electrons and therefore carries a net electrical charge.
cation
A positively charged ion formed when an atom loses one or more electrons. Charge = protons − electrons (positive).
anion
A negatively charged ion formed when an atom gains one or more electrons. Charge = protons − electrons (negative).
atomic mass unit (amu)
Standard unit of atomic mass, defined as exactly one-twelfth the mass of a carbon-12 atom. ≈ 1.66 × 10⁻²⁴ grams.
nucleus
The small, dense, positively charged center of an atom containing protons and neutrons. Contains essentially all of the atom's mass.
proton
A subatomic particle with a charge of +1 and a mass of ≈1 amu, located in the nucleus. The proton count defines the element.
neutron
A subatomic particle with no electrical charge and a mass of ≈1 amu, located in the nucleus. Varies among isotopes.
electron
A subatomic particle with a charge of −1 and negligible mass, found in the electron cloud. Determines chemical behavior.
radioisotope
An unstable isotope whose nucleus decays over time, releasing energy and particles. Used in medicine, dating, and research.
Formula Review
Atomic Number and Mass Practice
Complete the table for each neutral atom.
| Element | Atomic Number | Mass Number | Protons | Neutrons | Electrons |
|---|
| Helium | 2 | 4 | | | |
| Carbon | 6 | 13 | | | |
| Nitrogen | 7 | 15 | | | |
| Aluminum | 13 | 27 | | | |
| Phosphorus | 15 | 31 | | | |
| Iron | 26 | 56 | | | |
| Silver | 47 | 108 | | | |
| Gold | 79 | 197 | | | |
- An atom has 20 protons and 20 neutrons. Identify the element and state its mass number.
- An atom has 15 protons and 16 neutrons. Identify the element and state its mass number.
- An atom has an atomic number of 11 and a mass number of 23. How many neutrons does it have?
- An atom has 53 protons and 74 neutrons. What is its mass number? What element is this?
- Two atoms: Atom A has 8 protons and 8 neutrons. Atom B has 8 protons and 10 neutrons. (a) What element are both? (b) What are their mass numbers? (c) Are they the same substance or different? Explain.
Isotope Practice
- Write the isotope name for each: (a) 6 protons, 6 neutrons (b) 6 protons, 7 neutrons (c) 6 protons, 8 neutrons (d) 8 protons, 8 neutrons (e) 8 protons, 9 neutrons (f) 8 protons, 10 neutrons.
- For each isotope, identify element, mass number, atomic number, protons, neutrons, and electrons (neutral): (a) Uranium-235 (Z = 92) (b) Iodine-131 (Z = 53) (c) Technetium-99 (Z = 43) (d) Cobalt-60 (Z = 27) (e) Potassium-40 (Z = 19).
- Hydrogen has three isotopes: hydrogen-1 (ordinary), deuterium (hydrogen-2), and tritium (hydrogen-3). All have 1 proton. (a) How many neutrons does each have? (b) Why are all three called hydrogen? (c) Are they isotopes of each other?
- Explain in your own words why isotopes of the same element are chemically almost identical, even though they have different masses.
- Neon has three naturally occurring isotopes: neon-20, neon-21, and neon-22, all with 10 protons. How many neutrons does each have?
Ion Charge Practice
| Particle | Protons | Electrons | Charge | Cation or Anion? |
|---|
| Na⁺ | 11 | | | |
| Ca²⁺ | 20 | | | |
| Al³⁺ | 13 | | | |
| Cl⁻ | 17 | | | |
| O²⁻ | 8 | | | |
| S²⁻ | 16 | | | |
| K⁺ | 19 | | | |
| N³⁻ | 7 | | | |
- A magnesium atom (Z = 12) loses 2 electrons. (a) How many protons? (b) How many electrons? (c) What is the charge? (d) Cation or anion?
- A sulfur atom (Z = 16) gains 2 electrons. (a) How many electrons does the ion have? (b) What is the charge? (c) Write the ion symbol.
- An ion has 26 protons and 23 electrons. (a) What element is this? (b) What is the charge? (c) Cation or anion?
- An oxide ion (O²⁻) has a charge of 2−. Oxygen has Z = 8. (a) How many protons? (b) How many electrons? (c) How many electrons did the oxygen atom gain?
- Explain why gaining or losing electrons does not change what element an atom is, but gaining or losing protons would change it completely.
Average Atomic Mass Practice
- Boron has two naturally occurring isotopes: boron-10 (10.013 amu, 19.9%) and boron-11 (11.009 amu, 80.1%). Calculate the average atomic mass. Does your answer agree with the periodic table?
- Magnesium has three isotopes: Mg-24 (23.985 amu, 78.99%), Mg-25 (24.986 amu, 10.00%), Mg-26 (25.982 amu, 11.01%). Calculate the average atomic mass.
- Silicon has three isotopes: Si-28 (27.977 amu, 92.23%), Si-29 (28.976 amu, 4.67%), Si-30 (29.974 amu, 3.10%). Calculate the average atomic mass.
- An element has two isotopes: Isotope A (62.930 amu, 69.17%) and Isotope B (64.928 amu, 30.83%). Calculate the average atomic mass and identify the element using your periodic table.
- Bromine's average atomic mass is approximately 79.90 amu. Isotopes: Br-79 (78.918 amu) and Br-81 (80.916 amu). (a) Without calculating, predict which isotope is more abundant and explain. (b) If Br-79 abundance is 50.69%, what is Br-81's abundance? (c) Verify by calculating average atomic mass.
- An element's two isotopes: Isotope 1 (106.905 amu, 51.84%) and Isotope 2 (108.905 amu, 48.16%). (a) Calculate the average atomic mass. (b) Identify the element.
Real-World Applications
- A doctor uses technetium-99m to image a patient's thyroid. Technetium has Z = 43. (a) How many protons does technetium-99m have? (b) How many neutrons? (c) Why would a doctor prefer a radioactive isotope over a stable one for imaging?
- An archaeologist finds charcoal containing about 50% of the carbon-14 found in living wood. Carbon-14 has a half-life of about 5,730 years. (a) How many half-lives have passed? (b) Approximately how old is the charcoal?
- Heavy water is made with deuterium (hydrogen-2, 1 proton and 1 neutron) instead of ordinary hydrogen (hydrogen-1, 1 proton and 0 neutrons). (a) Are ordinary hydrogen and deuterium isotopes? Explain. (b) If heavy water behaves chemically like regular water, what does this reveal about the role of electrons vs. neutrons in chemical behavior?
- Iodine has Z = 53. (a) How many protons are in every iodine atom, regardless of isotope? (b) How many neutrons are in iodine-131? (c) Why does iodine concentrate in the thyroid gland specifically?
- Uranium-238 (Z = 92) decays to lead-206 (Z = 82). (a) How many neutrons are in uranium-238? (b) How many neutrons are in lead-206? (c) Why would uranium-238 be more useful than carbon-14 for dating a rock that is 500 million years old?
Error Analysis
Each problem below contains a mistake. Identify the error and correct it.
Problem 1
"Carbon-14 has 14 protons, because the mass number equals the number of protons."
What is wrong with this statement? What is the correct proton count for carbon-14, and how many neutrons does it have?
Problem 2
Average atomic mass of two isotopes (10.000 amu, 20%) and (11.000 amu, 80%) calculated as: (10.000 + 11.000) ÷ 2 = 10.500 amu
What error did the student make? Show the correct calculation.
Problem 3
A sulfur ion has 16 protons and a charge of 2−. A student says it has 14 electrons.
Is the student correct? Show the correct calculation for the number of electrons.
Problem 4
"Copper-63 and zinc-63 are isotopes of each other because they have the same mass number."
Is this correct? Explain what isotopes actually are and why this statement is wrong.
Problem 5
"If an oxygen atom gains 2 electrons to become an oxide ion, it has changed into a different element because it now has more particles."
Correct the student's reasoning. What determines element identity? What actually changed?
Review Questions
- What three particles make up an atom? Where is each particle located, and what charge does each carry?
- What is the difference between atomic number and mass number?
- Two atoms have the same atomic number but different mass numbers. What can you conclude?
- A chlorine-37 atom has Z = 17. (a) How many protons? (b) How many neutrons? (c) How many electrons if neutral?
- Why does the periodic table list atomic masses as decimal numbers rather than whole numbers?
- Explain the difference between a cation and an anion.
- A neutral calcium atom has Z = 20. When it reacts with chlorine, it forms a Ca²⁺ ion. How many electrons does the calcium ion have?
- Describe in your own words how average atomic mass is calculated. Why is it a weighted average rather than a simple average?
- An element has Z = 26. It forms two different ions: one with 2+ charge and one with 3+ charge. (a) What element is this? (b) How many electrons does the 2+ ion have? (c) How many electrons does the 3+ ion have?
- Why do isotopes of the same element react chemically in almost identical ways, even though they have different masses?
Looking Ahead
Chapter 7: Electrons and Atomic Structure
Chapter 6 focused on the nucleus — protons, neutrons, mass, isotopes, and ions. Now the focus shifts decisively to electrons. Electrons are the architects of chemical behavior. They determine how atoms bond, what reactions occur, and why the periodic table has the pattern it does. Chapter 7 introduces electron energy levels, the electron cloud model, orbital regions, electron configurations, valence electrons, and Lewis dot symbols. These ideas form the foundation for everything in Part III: Bonding, Compounds, and Molecular Structure.
MGU Library Connection
- Periodic Table Reference — Appendix C of this textbook
- Atomic Mass Unit and Measurement Reference — Appendix B of this textbook
- Chapter 5: Atoms and Atomic Theory — proton discovery, nuclear model, Dalton's theory
- Chapter 15: The Mole and Molar Mass — average atomic mass connects to molar mass calculations
- Chapter 26: Nuclear Chemistry — radioactivity, half-life, fission, fusion, and radiation safety
- MGU Dictionary entries: isotope, ion, cation, anion, atomic mass unit, average atomic mass, nuclear symbol