Chapter 5 — Front Matter
Purpose and Overview
Every material you have ever touched — wood, glass, water, air, skin, metal, stone — is made of atoms. This is one of the most powerful ideas in the history of science. It is simple enough to state in a single sentence, yet it took more than two thousand years for scientists to gather enough evidence to accept it with confidence.
Atoms are unimaginably small. A single grain of table salt contains roughly ten million billion atoms. You cannot see individual atoms with your eyes, and for most of human history, no technology could detect them directly. And yet, chemistry is built on them. Every formula you will write, every reaction you will study, every calculation you will perform in this course depends on the idea that matter is made of atoms.
This chapter tells the story of how scientists figured that out. It begins with ancient speculation, moves through careful experiments, and arrives at the modern picture of atomic structure. Along the way, you will see how science works: how ideas get tested, revised, and sometimes replaced when new evidence demands it.
Opening Question
If you could cut a piece of gold in half, then cut one of those halves in half, and keep cutting — would you eventually reach a piece so small it could not be cut further? Or could you always cut again, forever? The answer to this ancient question shapes everything we believe about the physical world. By the end of this chapter, you will know what evidence convinced scientists which answer is correct — and what that means for every substance you will study in chemistry.
9Sections
10Objectives
22Key Terms
2Worked Examples
Learning Objectives
- 1Describe early Greek ideas about the nature of matter and explain why Democritus's proposal was significant.
- 2State the law of conservation of mass and the law of definite proportions and explain how they support atomic theory.
- 3List the five main points of Dalton's atomic theory, identifying what he got right and what has since been revised.
- 4Explain how Brownian motion and chemical evidence supported the existence of atoms.
- 5Describe and compare the solid sphere, plum pudding, nuclear, Bohr, and quantum mechanical models of the atom.
- 6Explain the results of Rutherford's gold foil experiment and describe the nuclear model it produced.
- 7Identify the three subatomic particles, their charges, relative masses, and locations in the atom.
- 8Define atomic number and explain how it determines element identity.
- 9Define mass number and calculate the number of protons, neutrons, and electrons in a neutral atom.
- 10Explain why atomic theory is the conceptual foundation of all major topics in chemistry.
Sections 5.1 – 5.2
From Ancient Philosophy to the First Atomic Theory
Early Ideas · Democritus · Aristotle · Dalton's Theory · The Laws Behind It
Section 5.1
Early Ideas About Matter
The Question Behind the Question
Long before there were chemistry labs, there were philosophers who asked a simple but profound question: What is matter made of, at its most basic level? If you could cut a piece of gold in half, then cut one of those halves in half, and keep cutting — would you eventually reach a piece so small it could not be cut further? Or could you always cut again, forever?
This is not a trivial question. The answer shapes everything else we believe about the physical world.
The Greek Tradition
Around 400 BCE, a Greek philosopher named Democritus proposed that all matter was made of tiny, indivisible particles. He called these particles atomos, which means "uncuttable" in Greek. Democritus imagined that different materials had different shaped atoms — rough materials had rough atoms, liquid materials had smooth slippery atoms, and sweet or bitter tastes came from the shape of the atoms touching the tongue.
Democritus did not conduct experiments. His atomic idea came from reasoning, not measurement. He could not test it, weigh it, or verify it with any instrument. Despite this, the core idea — that matter has a smallest particle — was remarkably on target.
Aristotle's Competing View
Aristotle, one of the most influential thinkers of the ancient world, disagreed with Democritus. Aristotle believed that matter was continuous, meaning it could be divided forever without ever reaching an uncuttable piece. He also believed that all matter was made of four elements — earth, water, air, and fire — combined in different proportions. A fifth element, aether, made up the heavens.
Because Aristotle's authority dominated European and Islamic scholarship for nearly two thousand years, his continuous model of matter outlasted the atomic model. The idea of atoms was largely set aside until the 1600s and 1700s, when scientists began studying matter experimentally rather than philosophically.
Science vs. Philosophy
Democritus proposed atoms through reasoning alone; he had no experiments to support the idea. This is why the atomic model was set aside for so long. Good scientific ideas need more than cleverness — they need evidence. The evidence for atoms came slowly, through careful measurement over two centuries.
Section 5.2
Dalton's Atomic Theory
In the early 1800s, an English schoolteacher and scientist named John Dalton developed the first modern atomic theory. Dalton was not the first to think about atoms, but he was the first to ground the atomic idea in quantitative chemical evidence. He based his theory on two laws that chemists had already established through careful measurement.
The Law of Conservation of Mass
In any chemical reaction, the total mass of the reactants equals the total mass of the products. Matter is neither created nor destroyed. When hydrogen burns in oxygen to form water, the mass of the water equals the combined mass of the hydrogen and oxygen that reacted.
The Law of Definite Proportions
A compound always contains the same elements in the same proportions by mass, regardless of the sample size or source. Water, whether from a river, a cloud, or a laboratory, always contains hydrogen and oxygen in a mass ratio of approximately 1 to 8.
The Five Main Points of Dalton's Atomic Theory
- 1All matter is made of very small particles called atoms.
- 2Atoms of the same element are identical to each other — same mass, same chemical properties. Atoms of different elements differ in mass and properties.
- 3Atoms cannot be created, destroyed, or divided in chemical reactions. This explains conservation of mass.
- 4Atoms of different elements combine in simple whole-number ratios to form compounds. This explains definite proportions.
- 5In a chemical reaction, atoms are rearranged. The same atoms are present before and after; they are just connected differently.
What Dalton Got Wrong
Dalton assumed atoms were indivisible solid spheres with no internal structure. We now know atoms contain protons, neutrons, and electrons. He also assumed all atoms of the same element are identical in mass — we now know that isotopes of the same element can have different masses. Despite these errors, his theory was a revolutionary and useful advance.
Sections 5.3 – 5.4
Evidence for Atoms and the Language of Elements
Multiple Proportions · Brownian Motion · Elements · Compounds · Chemical Formulas
Section 5.3
Evidence for Atoms
One of the most interesting challenges in science is building knowledge about things you cannot see. Atoms are far too small to observe directly. How, then, did scientists become confident that atoms were real and not just a convenient fiction? The answer is that evidence came from many directions at once. No single experiment proved atoms existed, but the combined weight of evidence — from chemistry, physics, and eventually microscopy — became overwhelming.
The Law of Multiple Proportions
The law of multiple proportions extends the law of definite proportions. When two elements can form more than one compound, the different compounds contain the elements in ratios that are simple whole-number multiples of each other. Carbon and oxygen form two compounds: carbon monoxide (CO) and carbon dioxide (CO₂). For every fixed mass of carbon, the amount of oxygen in CO₂ is exactly twice the amount in CO. This makes perfect sense if oxygen atoms are whole, countable units — but is very difficult to explain if matter is continuous.
Brownian Motion
In 1827, a botanist named Robert Brown was examining pollen grains suspended in water under a microscope. He noticed the pollen grains moved in an erratic, jittery, completely unpredictable pattern — not because of currents in the water, but constantly and randomly. This became known as Brownian motion.
For decades, no one could fully explain it. In 1905, Albert Einstein provided a mathematical explanation: Brownian motion is caused by the countless random collisions of water molecules against the larger pollen grains. Because collisions come from all directions at slightly different times, the pollen grain gets pushed in a constantly changing direction. This was powerful evidence that water was indeed made of moving particles — particles far too small to see, but real enough to push pollen grains around.
Why This Matters
Einstein's 1905 explanation of Brownian motion was one of the clearest demonstrations that molecules — and by extension, atoms — are physically real objects, not just mathematical conveniences. It marked a turning point in scientific acceptance of the atomic model.
Section 5.4
Atoms, Elements, and Compounds
An element is a pure substance made of only one kind of atom. Every atom of hydrogen has one proton. Every atom of carbon has six protons. Every atom of gold has seventy-nine protons. The number of protons in an atom determines which element it is. There are 118 known elements, each listed on the periodic table.
A compound is a pure substance made of two or more different elements chemically bonded together. The atoms in a compound are not just mixed — they are held together by chemical bonds, and the compound has properties entirely different from its component elements.
| Compound | Elements Present | Properties of Elements | Properties of Compound |
|---|
| Water (H₂O) | Hydrogen, Oxygen | H: flammable gas; O: supports combustion | Liquid at room temperature; neither burns nor supports combustion well |
| Table Salt (NaCl) | Sodium, Chlorine | Na: reactive metal; Cl: poisonous gas | White, stable, edible solid |
| Carbon Dioxide (CO₂) | Carbon, Oxygen | C: solid; O: gas that supports combustion | Gas; does not support combustion; used by plants in photosynthesis |
Compound vs. Mixture
In a mixture, components retain their individual properties and can be separated by physical means. In a compound, elements are chemically bonded in fixed ratios and can only be separated by a chemical reaction. This is why compounds obey the law of definite proportions: the ratio of atoms is fixed by the bonds between them.
Element Symbols and Chemical Formulas
Chemists represent atoms with one- or two-letter symbols. The first letter is always capitalized; the second, if present, is lowercase. These symbols are standardized worldwide. The symbol for hydrogen is H, oxygen is O, carbon is C, sodium is Na (from Latin natrium), iron is Fe (from ferrum), and gold is Au (from aurum).
Chemical formulas represent compounds using element symbols combined with subscript numbers. The formula H₂O means two hydrogen atoms bonded to one oxygen atom. NaCl means one sodium paired with one chlorine. CO₂ means one carbon bonded to two oxygens. Every formula in chemistry is a claim about atomic composition.
Sections 5.5 – 5.6
How the Model of the Atom Changed Over Time
Scientific Models · Solid Sphere · Plum Pudding · Gold Foil · Nuclear Model · Quantum Model
Section 5.5
Atomic Models and Scientific Revision
A model in science is a simplified representation of something we cannot directly observe or fully understand. A good model captures the most important features of a system accurately enough to make useful predictions. The history of atomic models is a perfect illustration of how models evolve — each was a genuine improvement on what came before, and each was eventually revised when new evidence revealed the picture was incomplete.
| Model | Proposed By | Key Feature | What Required Revision |
|---|
| Solid Sphere | Dalton (~1803) | Atoms are indivisible solid balls; different elements have different sizes and masses | Discovery of the electron (1897) showed atoms have internal structure |
| Plum Pudding | Thomson (1897) | Electrons embedded in a diffuse positive charge | Rutherford's gold foil experiment (1909–1911) showed positive charge is concentrated, not diffuse |
| Nuclear Model | Rutherford (1911) | Small, dense, positive nucleus; electrons occupy surrounding space | Could not explain why electrons don't fall into nucleus or the hydrogen emission spectrum |
| Bohr Model | Bohr (1913) | Electrons orbit nucleus in fixed energy levels | Only works accurately for hydrogen; electrons don't travel in defined circular paths |
| Quantum Mechanical | Heisenberg, Schrödinger (1920s) | Electrons exist in orbitals — regions of probability — around the nucleus | Current accepted model; continues to be refined |
Rutherford's Gold Foil Experiment
In 1909 and 1911, Ernest Rutherford, Hans Geiger, and Ernest Marsden fired a beam of small, positively charged alpha particles at an extremely thin sheet of gold foil. Behind the foil, a detector screen recorded where the alpha particles went.
Based on the plum pudding model, Rutherford expected all particles to pass through with only slight deflections — the way a bullet might pass through a thin sheet of fog. The results were shocking. Most alpha particles passed straight through, as expected if atoms were mostly empty space. But some were deflected at large angles. A very small number bounced almost straight back.
Rutherford's Famous Remark
Rutherford described his astonishment this way: "It was as if you fired artillery shells at tissue paper and they came back and hit you." There was simply no way to explain large-angle deflections using the plum pudding model. A diffuse positive cloud could not deflect a fast-moving alpha particle so sharply.
Rutherford's Conclusion: The Nuclear Model
Almost all of an atom's mass and all of its positive charge are concentrated in a tiny, dense region at the center — the nucleus. The rest of the atom is mostly empty space, with electrons occupying the surrounding region. This nuclear model replaced the plum pudding model and is the foundation of modern atomic structure.
Section 5.6
The Nuclear Atom
The nucleus is extraordinarily small compared to the total size of the atom. If an atom were enlarged to the size of a football stadium, the nucleus would be about the size of a marble at the center of the field. The electrons would be somewhere near the seats — but because electrons have essentially no volume, the seats themselves would mostly be empty space. Despite its tiny size, the nucleus contains virtually all of the atom's mass and is incredibly dense.
| Scale | Approximate Diameter | Comparison |
|---|
| Typical atom | 10⁻¹⁰ m (0.1 nm) | About 1 angstrom — invisible to any optical microscope |
| Atomic nucleus | 10⁻¹⁵ m | 100,000 times smaller than the atom itself |
| Stadium analogy | Nucleus = marble | If atom = stadium, nucleus = marble at center field; nearly everything else is empty space |
Why Atoms Are Stable
A natural question: if opposite charges attract, why don't electrons fall into the nucleus? The Bohr model partially addressed this with orbits; the quantum mechanical model provides a deeper answer through quantum physics. For our purposes in chemistry, the key point is simply that atoms are stable — electrons occupy well-defined regions around the nucleus without collapsing into it.
Sections 5.7 – 5.8
Protons, Neutrons, Electrons, and Atomic Numbers
Three Particles · Charges · Masses · Atomic Number Z · Mass Number A · Calculations
Section 5.7
Protons, Neutrons, and Electrons
Every atom contains three types of subatomic particles: protons, neutrons, and electrons. Each has a characteristic location, mass, and charge. Understanding these three particles is essential for everything that follows in this course.
| Particle | Symbol | Location | Relative Mass | Charge |
|---|
| Proton | p⁺ | Nucleus | 1 amu | +1 |
| Neutron | n⁰ | Nucleus | 1 amu | 0 |
| Electron | e⁻ | Outside nucleus | ~0 (1/1836 amu) | −1 |
Protons
A proton carries a charge of +1 and has a mass of approximately 1 atomic mass unit (amu). Protons are found in the nucleus. The number of protons in an atom is the single most important fact about it — it determines which element the atom is. Every atom of hydrogen has exactly 1 proton. Every atom of carbon has exactly 6. If you change the number of protons, you change the element.
Neutrons
A neutron carries no electrical charge and has a mass of approximately 1 amu — slightly more than the proton. Neutrons are found in the nucleus alongside protons. They contribute to atomic mass without affecting charge. The number of neutrons can vary for atoms of the same element; atoms of the same element with different neutron counts are called isotopes, which you will study in Chapter 6.
Electrons
An electron carries a charge of −1 and has a mass of approximately 1/1836 the mass of a proton — so small that it is treated as negligible in mass calculations. Electrons occupy the region outside the nucleus. In a neutral atom, the number of electrons equals the number of protons, balancing positive and negative charges. When an atom gains or loses electrons, it becomes a charged ion. You will study ions in Chapter 6.
Why Electron Mass Is Ignored
Because each electron is roughly 1/1836 the mass of a proton, even 100 electrons contribute less than 0.1 amu to atomic mass. For all practical chemistry calculations, the mass of an atom is simply the number of protons plus the number of neutrons — both located in the nucleus.
Section 5.8
Atomic Number and Mass Number
Atomic Number (Z)
The atomic number is the number of protons in the nucleus of an atom. It determines which element the atom is. Every atom of hydrogen has Z = 1. Every atom of carbon has Z = 6. In a neutral atom, Z also equals the number of electrons.
Mass Number (A)
The mass number is the total number of protons and neutrons in the nucleus.
A = Z + N where N = number of neutrons
Therefore: N = A − Z (number of neutrons = mass number − atomic number)
Worked Example 5.1
Finding subatomic particle counts from atomic number and mass number
- ProblemAn atom has an atomic number of 17 and a mass number of 35. How many protons, neutrons, and electrons does it have (neutral atom)?
- Step 1The atomic number equals the number of protons.
Protons = Z = 17
- Step 2In a neutral atom, electrons = protons.
Electrons = 17
- Step 3Use the mass number formula to find neutrons.
Neutrons = A − Z = 35 − 17 = 18
- IdentifyAtomic number 17 on the periodic table is chlorine (Cl).
Answer: 17 protons, 17 electrons, 18 neutrons. The element is chlorine.
Worked Example 5.2
Identifying an element from proton count
- ProblemAn atom has 11 protons. What element is it? How many electrons does it have if neutral?
- Step 1Look up atomic number 11 on the periodic table.
Z = 11 → Sodium (Na)
- Step 2A neutral atom has equal protons and electrons.
Electrons = 11
Answer: The element is sodium (Na). It has 11 electrons.
Nuclear Symbol Notation
Chemists write a shorthand notation to describe a specific atom completely. The nuclear symbol shows the element symbol with the mass number as a superscript and the atomic number as a subscript, both to the left of the symbol. For example, a common carbon atom is written as ¹²₆C — mass number 12 (superscript), atomic number 6 (subscript), element carbon (C). From this notation, you can immediately calculate all three subatomic particle counts. You will practice this notation extensively in Chapter 6.
For each atom, determine the number of protons, neutrons, and electrons (assume neutral atoms).
- Z = 6, A = 12
- Z = 8, A = 16
- Z = 11, A = 23
- Z = 17, A = 35
- Z = 26, A = 56
- Z = 79, A = 197
- Z = 20, A = 40
- Z = 47, A = 107
Section 5.9
Why Atomic Theory Matters in Chemistry
Every concept you will use in this course rests on the atomic model. Chemical formulas describe how many atoms of each element are present. Balanced equations ensure that the same atoms appear on both sides of a chemical reaction. Molar mass — which you will study in Chapter 15 — is the mass of one mole of atoms or molecules. Reaction mechanisms describe how bonds between atoms break and form. Without atomic theory, none of these ideas have meaning.
One of the deepest satisfactions of atomic theory is that it connects the invisible world to the visible one. When you dissolve salt in water, you can explain what is happening at the particle level: ionic bonds between sodium and chloride ions are disrupted by water molecules, and the ions spread through the solution. When iron rusts, iron atoms are reacting with oxygen molecules to form a new compound. When a fuel burns, you can trace the rearrangement of atoms from reactants to products. The visible world of objects and changes becomes understandable in terms of the invisible world of atoms, bonds, and electrons.
Atomic theory is also not just a chemistry idea. It connects chemistry to physics (which studies atomic forces and energy), biology (which studies molecules made of atoms), geology (which studies minerals made of atomic compounds), medicine (which uses atomic-level understanding of drugs and cells), and environmental science (which tracks atomic and molecular pollutants through Earth systems). The atom is the unit of connection between all the sciences that study matter.
The Pattern of Scientific Progress
The history of atomic models illustrates something important about science. No model was thrown out simply because it was imperfect. Each model was replaced when a better model — one supported by more evidence and capable of explaining more observations — became available. That openness to revision in the face of evidence is what makes science trustworthy. Dalton's model was revised. Thomson's model was replaced. Bohr's model was extended. Each revision made chemistry more accurate and more powerful. The model you are learning is the result of that cumulative process.
Chapter 5 — Back Matter
Chapter Summary
- Sec 5.1Around 400 BCE, Democritus proposed matter is made of tiny indivisible particles called atomos. Aristotle's competing continuous model dominated for nearly two thousand years. Experimental chemistry in the 1600s–1700s began to revive atomic thinking.
- Sec 5.2Dalton's first modern atomic theory (early 1800s) was grounded in the law of conservation of mass and the law of definite proportions. Its five main points correctly explained chemical combination and mass conservation, though it incorrectly assumed atoms were indivisible and uniform.
- Sec 5.3Chemical evidence (law of multiple proportions) and physical evidence (Brownian motion, explained by Einstein in 1905) supported the reality of atoms. Later experiments in the late 1800s and early 1900s confirmed atoms have internal structure.
- Sec 5.4An element is a pure substance made of one kind of atom; its identity is determined by proton count. A compound is a chemically bonded combination of elements with properties distinct from those of its components. Chemical formulas express atomic composition.
- Sec 5.5Atomic models evolved from Dalton's solid sphere through Thomson's plum pudding, Rutherford's nuclear model (from the gold foil experiment), Bohr's orbit model, to the modern quantum mechanical model. Each revision followed new experimental evidence — a demonstration of scientific progress.
- Sec 5.6The nucleus is a tiny, dense, positively charged center containing nearly all of the atom's mass. The surrounding electron region is vastly larger but essentially empty. If the atom were a stadium, the nucleus would be a marble at center field.
- Sec 5.7Protons (charge +1, mass ≈ 1 amu) and neutrons (charge 0, mass ≈ 1 amu) are in the nucleus. Electrons (charge −1, negligible mass) are outside the nucleus. In a neutral atom, electrons equal protons. Atomic mass ≈ protons + neutrons.
- Sec 5.8Atomic number Z = number of protons; it defines the element. Mass number A = protons + neutrons. Neutrons = A − Z. Nuclear symbol notation places the mass number as a superscript and atomic number as a subscript to the left of the element symbol.
- Sec 5.9Atomic theory is the foundation of every major chemistry topic: formulas, equations, the periodic table, bonding, stoichiometry, and chemical change. It connects the visible world of observable matter to the invisible world of atoms and electrons.
Atomic Theory Timeline
~400 BCE
Democritus proposes that matter is made of tiny, indivisible particles he calls atomos (uncuttable).
~350 BCE
Aristotle argues for a continuous model of matter composed of four elements. His view dominates for nearly two thousand years.
1789
Antoine Lavoisier establishes the law of conservation of mass through careful quantitative experiments.
~1794–1799
Joseph Proust establishes the law of definite proportions (constant composition).
1803–1808
John Dalton develops the first modern atomic theory, grounded in quantitative chemical evidence.
1827
Robert Brown observes the erratic random motion of pollen grains in water — Brownian motion.
1897
J. J. Thomson discovers the electron using cathode ray tubes; proposes the plum pudding model.
1905
Albert Einstein provides the mathematical explanation for Brownian motion as evidence for molecular reality.
1909–1911
Rutherford, Geiger, and Marsden perform the gold foil experiment. Rutherford proposes the nuclear model.
1913
Niels Bohr proposes fixed energy-level orbits for electrons; explains the hydrogen emission spectrum.
1920s–1930s
Heisenberg and Schrödinger develop the quantum mechanical model. Chadwick discovers the neutron (1932).
20th c. onward
Quantum mechanical model refined and confirmed; becomes the basis for modern chemistry and physics.
Key Terms
atom
The smallest particle of an element that retains the chemical properties of that element.
atomic number (Z)
The number of protons in the nucleus of an atom; determines which element the atom is.
mass number (A)
The total number of protons and neutrons in the nucleus of an atom. A = Z + N.
proton
A subatomic particle with charge +1 and mass ≈ 1 amu; found in the nucleus. The proton count defines the element.
neutron
A subatomic particle with no charge and mass ≈ 1 amu; found in the nucleus. Neutron count can vary among atoms of the same element.
electron
A subatomic particle with charge −1 and negligible mass; found outside the nucleus. Equal to proton count in neutral atoms.
nucleus
The small, dense, positively charged center of an atom, containing protons and neutrons and nearly all of the atom's mass.
element
A pure substance made of only one kind of atom, defined by its proton count. There are 118 known elements.
compound
A pure substance made of two or more elements chemically bonded in fixed ratios, with properties distinct from its component elements.
Dalton's atomic theory
The first modern atomic theory: matter is made of atoms; same-element atoms are identical; atoms rearrange in reactions; compounds form in fixed whole-number ratios.
law of conservation of mass
In a chemical reaction, total mass of reactants equals total mass of products; matter is neither created nor destroyed.
law of definite proportions
A compound always contains the same elements in the same proportions by mass, regardless of source or sample size.
law of multiple proportions
When two elements form more than one compound, the mass ratios of one element to a fixed mass of the other are simple whole-number multiples.
Brownian motion
The erratic, random movement of small particles suspended in a fluid, caused by collisions with invisible molecules; evidence for the particle nature of matter.
nuclear model
Rutherford's model in which the atom's positive charge and mass are concentrated in a small, dense nucleus, with electrons in the surrounding space.
plum pudding model
Thomson's model in which electrons are embedded throughout a diffuse positive charge; replaced by the nuclear model after the gold foil experiment.
Bohr model
Model in which electrons travel in fixed circular energy levels around the nucleus; explains hydrogen's emission spectrum but fails for larger atoms.
quantum mechanical model
The current atomic model in which electrons occupy three-dimensional orbitals — regions of probability — around the nucleus.
subatomic particles
Particles smaller than an atom; in chemistry, primarily protons, neutrons, and electrons.
ion
An atom or group of atoms that has gained or lost electrons and therefore carries an electric charge. Studied fully in Chapter 6.
scientific model
A simplified representation of a system used to explain observations and make predictions; models evolve as new evidence is gathered.
atomic mass unit (amu)
A unit of mass for subatomic particles; protons and neutrons each have a mass of approximately 1 amu.
Concept Check
Answer these questions to check your understanding before moving to the practice problems.
- What did Democritus mean by the word atomos?
- What two experimental laws did Dalton use as evidence for his atomic theory?
- State two things Dalton's atomic theory got right and one thing it got wrong.
- What is Brownian motion, and what does it suggest about the nature of matter?
- What was surprising about the results of Rutherford's gold foil experiment?
- What did Rutherford conclude from the gold foil experiment?
- Name the three main subatomic particles. Where is each located? What charge does each carry?
- Which subatomic particle determines the identity of an element?
- How do you calculate the number of neutrons in an atom if you know the atomic number and mass number?
- In a neutral atom, how does the number of electrons relate to the number of protons?
Atomic Structure Practice
For each atom below, determine the number of protons, neutrons, and electrons. Assume neutral atoms.
- Z = 1, A = 1
- Z = 2, A = 4
- Z = 6, A = 12
- Z = 7, A = 14
- Z = 8, A = 16
- Z = 11, A = 23
- Z = 13, A = 27
- Z = 17, A = 35
- Z = 19, A = 39
- Z = 26, A = 56
- Z = 29, A = 63
- Z = 47, A = 107
- Z = 79, A = 197
- Z = 92, A = 238
Use the periodic table to name the element for each atomic number.
- 3 protons
- 8 protons
- 12 protons
- 15 protons
- 20 protons
- 30 protons
- 35 protons
- 50 protons
- A sodium atom has 11 protons and 12 neutrons. What is its mass number?
- A phosphorus atom has 15 protons and 16 neutrons. What is its mass number?
- An iron atom has a mass number of 56 and atomic number 26. How many neutrons?
- A lead atom has a mass number of 207 and atomic number 82. How many neutrons?
- A calcium atom has 20 protons and mass number 40. How many neutrons?
- An aluminum atom has 13 protons and 14 neutrons. Write the mass number. How many electrons in the neutral atom?
Write the complete nuclear symbol (mass number superscript, atomic number subscript, element symbol) for each atom.
- Helium, mass number 4
- Carbon, mass number 12
- Nitrogen, mass number 14
- Oxygen, mass number 16
- Magnesium, mass number 24
- Sulfur, mass number 32
- Chlorine, mass number 37
Model Analysis
Answer in complete sentences.
- Dalton's atomic theory said atoms were indivisible solid spheres. What experimental evidence later showed this was incorrect?
- Thomson's plum pudding model placed positive charge spread throughout the atom. What specific observation in Rutherford's experiment required a different model?
- If the gold foil experiment had shown all alpha particles passing straight through without any deflection, what conclusion would Rutherford have been forced to draw?
- The Bohr model imagined electrons in fixed circular orbits. Describe one way the quantum mechanical model differs from this picture.
- Scientists sometimes say that models are tools, not truth. What does this mean? Use the history of atomic models to support your answer.
- In Rutherford's experiment, most alpha particles passed through the gold foil, but a small number were deflected at large angles. What does each observation tell us about atomic structure?
Common Misconceptions
Misconception 1
"Atoms are the smallest particles in existence."
This was Dalton's assumption. We now know that atoms are made of protons, neutrons, and electrons — and protons and neutrons are themselves made of smaller particles called quarks. For the purposes of chemistry, atoms are the relevant fundamental units, but they are not truly the smallest things in the universe.
Misconception 2
"Electrons orbit the nucleus like planets orbit the sun."
The Bohr model presented this picture, and it is still used in simplified diagrams. However, electrons do not travel in neat circular paths. In the quantum mechanical model, electrons exist in regions of probability called orbitals. We can describe where an electron is most likely to be found, but we cannot trace a defined path.
Misconception 3
"Atoms are mostly solid."
Atoms are almost entirely empty space. If the nucleus of a hydrogen atom were the size of a basketball, the electron would be roughly a mile away. The apparent solidity of everyday objects comes from electrical repulsion between electrons in neighboring atoms, not from atoms being densely packed with material.
Misconception 4
"Changing the number of neutrons changes the element."
Changing the number of protons changes the element. Changing the number of neutrons produces a different isotope of the same element. Isotopes of the same element have identical chemical behavior because they have the same number of protons and electrons.
Misconception 5
"Atoms look like the diagrams in textbooks."
Atoms are far too small to see directly. Textbook diagrams showing a nucleus surrounded by electron shells are simplified models. They are not photographs. Real atoms do not have sharply defined shells with electrons moving in visible paths.
Review Questions
- Who proposed the idea that matter is made of indivisible particles, and approximately when?
- What does the word atomos mean, and why was it chosen?
- Who proposed the competing continuous model of matter? What was this person's alternative explanation?
- State the law of conservation of mass in your own words. Give an everyday chemistry example.
- State the law of definite proportions in your own words. Give an example using a familiar compound.
- List the five main points of Dalton's atomic theory in your own words.
- Name at least two things we now know that contradict Dalton's original model.
- Explain the law of multiple proportions. Give an example using carbon and oxygen.
- What is Brownian motion? Describe what Brown observed, and explain what Einstein's 1905 explanation revealed about matter.
- Define element and give three examples.
- Define compound and give three examples.
- Explain the difference between a compound and a mixture. Why can table salt not be separated into sodium and chlorine by physical means?
- What does the chemical formula CO₂ tell you about the composition of carbon dioxide?
- Describe Dalton's solid sphere model. What did it get right? What did it leave unexplained?
- What did Thomson discover in 1897, and how did this require revision of the atomic model?
- Describe the setup and results of Rutherford's gold foil experiment. What happened that was unexpected?
- What conclusion did Rutherford draw? Describe the nuclear model he proposed.
- What did Bohr add to Rutherford's nuclear model? What phenomenon did the Bohr model explain?
- How does the quantum mechanical model differ from Bohr's model?
- If an atom were enlarged to the size of a football stadium, approximately how large would the nucleus be?
- What is a proton? Describe its charge, mass, and location in the atom.
- What is a neutron? Describe its charge, mass, and location in the atom.
- What is an electron? Describe its charge, mass, and location in the atom.
- Why is the mass of electrons usually ignored when calculating atomic mass?
- Define atomic number. What does it tell you about an atom?
- Define mass number. How is it calculated?
- An atom has atomic number 9 and mass number 19. How many protons, neutrons, and electrons? What element is it?
- An atom has 16 protons and 16 neutrons. What is its mass number? What element?
- Explain why atomic theory is considered the foundation of chemistry.
- Give one example of how atomic theory connects chemistry to another scientific field.
Chapter 5 Checkpoint
Chapter 5 Checkpoint
Complete without assistance — assesses all key skills and concepts from Chapter 5
Section 1 — Vocabulary and Concepts
- Define atom in your own words.
- Explain the difference between atomic number and mass number.
- What does it mean for two atoms to be isotopes of the same element?
Section 2 — Historical and Conceptual
- What two experimental laws supported Dalton's atomic theory? State each law in one sentence.
- Describe what Rutherford observed in the gold foil experiment and what conclusion he drew.
- Arrange these atomic models in chronological order: quantum mechanical, nuclear, plum pudding, solid sphere, Bohr.
Section 3 — Particle Identification
- An atom has Z = 15 and A = 31. How many protons, neutrons, and electrons does it have? What element is it?
- An atom has Z = 35 and A = 80. How many protons, neutrons, and electrons does it have? What element is it?
- An atom has 6 protons and 8 neutrons. Write its mass number, identify the element, and write its nuclear symbol.
- An atom has 20 protons and 20 neutrons. Identify the element and determine the number of electrons in a neutral atom of this element.
Section 4 — Application and Analysis
- A student claims that changing the number of neutrons in a carbon atom changes it into a different element. Is this correct? Explain why or why not.
- Why is the nucleus described as "tiny but massive"? Use specific scale comparisons in your answer.
- Explain in your own words why the history of atomic models shows that science is self-correcting.
MGU Library Connection
- Atomic Structure Reference — Appendix C of this textbook (Periodic Table Reference)
- Physical Science (preceding course) — Chapter on Matter and Its Properties
- Chapter 6: Atomic Structure and Isotopes — builds directly on proton, neutron, and electron counts from Chapter 5
- Chapter 7: Electrons and Atomic Structure — extends Chapter 5's subatomic picture to electron configurations
- Chapter 8: The Periodic Table — uses atomic number as the organizing principle introduced in Section 5.8
- MGU Dictionary entries: atom, element, compound, atomic number, mass number, nucleus, proton, neutron, electron
Looking Ahead
Chapter 6: Atomic Structure and Isotopes
Chapter 5 established the three subatomic particles and the basic structure of the atom. Chapter 6 builds on this foundation by exploring atomic structure in greater depth. You will learn about isotopes — atoms of the same element that have different numbers of neutrons and therefore different mass numbers. You will learn how to write nuclear symbols, calculate average atomic mass from isotope data, and discover how isotopes are used in medicine, archaeology, and scientific research. You will also study ions — atoms that have gained or lost electrons and carry an electrical charge. Understanding ions is essential for studying ionic bonding, chemical formulas, and reactions in aqueous solution. The atomic number and mass number calculations practiced in this chapter will be used constantly in Chapter 6.