2.1 Atomic Structure / Theory & Introduction to the Periodic Table
PRE-HEALTH PREP
Atomic Theory & Atomic Structure
General Chemistry 2.1
Atomic theory traces from Dalton's foundational postulates through Thomson's discovery of the electron, Millikan's measurement of its charge and mass, and Rutherford's gold foil experiment revealing the nuclear structure of the atom. The three subatomic particles—protons, neutrons, and electrons—define isotope notation, atomic number versus mass number, ion formation, and the calculation of atomic weight as a weighted average of isotopic masses. The periodic table organizes elements into groups (columns with similar properties) and periods (rows with repeating trends), classifying them as metals, non-metals, or metalloids across four named families: alkali metals, alkaline earth metals, halogens, and noble gases.
4 Scientists Who Shaped Modern Atomic Theory
The history of atomic theory begins with four pioneering scientists, each building on the discoveries of those before them. Their combined contributions gave us the modern understanding of what atoms are, what subatomic particles they contain, and how atomic structure works. You should know each scientist and the key discovery associated with their name.
Dalton's Atomic Theory
John Dalton is considered the father of modern atomic theory. His central proposal was elegantly simple: all matter is composed of tiny, indivisible particles called atoms. But Dalton went further than that. He recognized that atoms are not all identical—there are many different types, which we now call elements. Today, we know of approximately 120 different elements on the periodic table.
Dalton's atomic theory can be summarized in four key postulates:
- All matter is made up of atoms.
- The atoms of one element are different from the atoms of every other element.
- Compounds are formed by combining atoms of different elements.
- Atoms are neither created nor destroyed in chemical reactions—they are simply rearranged into new combinations to form different compounds.
Note: Dalton's fourth postulate is essentially a precursor to the law of conservation of mass. Atoms shuffle partners during chemical reactions, but the total number of each type of atom remains unchanged.
J.J. Thomson and the Discovery of the Electron
J.J. Thomson is credited with discovering the electron. Through his experiments with cathode rays, he determined the mass-to-charge ratio of this subatomic particle. He did not know the mass or the charge individually, but he established their ratio—a critical stepping stone for the scientists who followed. For most general chemistry courses, the essential fact is straightforward: Thomson discovered the electron.
Robert Millikan and the Oil Drop Experiment
Robert Millikan built directly on Thomson's work. Using his famous oil drop experiment, Millikan measured the charge of the electron. Once he knew the charge, he could combine it with Thomson's mass-to-charge ratio to calculate the mass of the electron as well. Millikan's contribution completed the picture that Thomson had started: together, they gave us both the mass and charge of the electron.
Ernest Rutherford and the Gold Foil Experiment
Ernest Rutherford fundamentally changed how we picture atomic structure. Before his work, the prevailing model was J.J. Thomson's plum pudding model, which imagined the atom as a uniform mixture of positive and negative charges—a soupy blend with no internal structure.
Rutherford tested this model by firing alpha particles (equivalent to helium nuclei) at an extremely thin sheet of gold foil. If atoms were truly solid, plum-pudding-style blobs, then every alpha particle should have bounced back or scattered to the side. Instead, Rutherford observed something astonishing: most alpha particles passed straight through the foil.
Example: The Volleyball Net Analogy
Imagine shooting a beam of sand at a volleyball net. Most grains pass right through the gaps without touching the net at all. Occasionally, a grain strikes a cord and bounces back. That is essentially what Rutherford observed in the gold foil experiment—most alpha particles sailed through, but a few were deflected or reflected, indicating they had struck something dense and compact.
From these results, Rutherford concluded that most of an atom is empty space, with nearly all its mass concentrated in a tiny, dense center he called the nucleus. This nuclear model of the atom remains the foundation of our understanding of atomic structure today.
Tip: Remember each scientist's key contribution—Dalton: atomic theory, Thomson: discovered the electron, Millikan: measured electron's mass and charge via the oil drop experiment, Rutherford: discovered the nucleus via the gold foil experiment.
| Scientist | Key Contribution |
|---|---|
| Dalton | Proposed his atomic theory |
| Thomson | Discovered electrons (studied cathode rays) |
| Millikan | Determined the mass & charge of an electron (oil drop experiment) |
| Rutherford | Discovered the nucleus (scattering of alpha particles in thin gold foil) |
Structure of the Atom
The modern atom consists of a central nucleus surrounded by a cloud of electrons. Inside the nucleus, you find two types of particles: protons, which carry a positive charge, and neutrons, which are electrically neutral (hence the name). The electrons occupying the space outside the nucleus carry a negative charge and are far smaller and lighter than either protons or neutrons.
Almost all of an atom's mass is concentrated in its nucleus. However, almost all of its volume belongs to the electron cloud. To appreciate the scale: if you could blow the nucleus up to the size of a basketball, the electron cloud would extend several miles in every direction, with electrons no bigger than tiny marbles zipping around within that vast space. Most of the atom, then, is empty space—a fact that mirrors Rutherford's gold foil experiment results.
Common Misconception: Diagrams showing electrons in neat circular orbits are oversimplifications. In reality, electrons exhibit complex wave-like behavior and do not follow fixed paths. You will explore this in greater detail when you study quantum mechanics and electron orbitals later in the course.
The three subatomic particles and their charges are the foundation for everything that follows in this lesson:
| Particle | Location | Charge | Relative Mass |
|---|---|---|---|
| Proton (p⁺) | Nucleus | +1 | ~1 amu |
| Neutron (n⁰) | Nucleus | 0 | ~1 amu |
| Electron (e⁻) | Electron cloud | −1 | ~0 amu (1/1841 amu) |
3 Subatomic Particles: Properties and Masses
Charge in Coulombs
When we say a proton has a "+1" charge and an electron has a "−1" charge, we are using simplified notation. These values represent the fundamental unit of charge. A "+2" charge simply means two excess protons (or two fewer electrons), and so on.
In scientific units, charge is measured in coulombs (C). The electron carries a charge of −1.602 × 10−19 C, while the proton carries +1.602 × 10−19 C. This value is significant in physics courses, but in general chemistry, you will almost always work with the simplified +1 and −1 notation.
Mass in Atomic Mass Units (amu)
Because subatomic particles are extraordinarily small, measuring their mass in grams would be inconvenient. Instead, chemists use the atomic mass unit (amu). The masses of the three subatomic particles in amu are:
- Proton: 1.00728 amu ≈ 1 amu
- Neutron: 1.00867 amu ≈ 1 amu (slightly heavier than the proton)
- Electron: 0.000549 amu ≈ 0 amu (about 1/1841 the mass of a proton)
Tip: You do not need to memorize the exact amu values. Just remember that protons and neutrons each weigh approximately 1 amu, while electrons are so light that their mass rounds to essentially zero when working with whole-number mass calculations.
| Particle | Charge | Mass (amu) |
|---|---|---|
| Proton | +1 | 1.00728 amu |
| Neutron | Neutral (0) | 1.00867 amu |
| Electron | −1 | 0.000549 amu |
Isotope Symbols, Atomic Number, and Mass Number
An isotope symbol conveys three critical pieces of information about an atom: the element's identity, its mass number, and (when applicable) its charge. Understanding how to read these symbols is an essential skill in general chemistry.
Atomic Number: What It Tells You
The atomic number (written as a subscript before the element symbol) tells you the number of protons in the nucleus. This number defines which element you have. Oxygen always has 8 protons, nitrogen always has 7, and carbon always has 6. If you change the number of protons, you change the element entirely.
On the periodic table, the atomic number is displayed above each element's symbol. It is the single most important identifier for any element.
Mass Number: Protons Plus Neutrons
The mass number (written as a superscript before the element symbol) represents the total count of protons plus neutrons in the nucleus. Because protons and neutrons each weigh approximately 1 amu, the mass number serves as a good approximation of the atom's mass in amu.
Example: Oxygen-16
Oxygen has an atomic number of 8 (so 8 protons). The most common isotope has a mass number of 16. To find the number of neutrons: 16 − 8 = 8 neutrons. The actual mass is approximately 15.999 amu, which rounds to the whole-number mass number of 16.
Isotopes: Same Element, Different Neutron Count
Isotopes are atoms of the same element that differ in their number of neutrons. All isotopes of an element have the same number of protons (and therefore the same atomic number), but they have different mass numbers.
Example: Isotopes of Oxygen
Oxygen-16 has 8 protons and 8 neutrons (mass number = 16). Oxygen-18 also has 8 protons but contains 10 neutrons (mass number = 18). Both are oxygen—they just differ in neutron count. Oxygen-16 is far more abundant in nature, accounting for over 99% of naturally occurring oxygen.
Charge and Ions
The third number in an isotope symbol is the charge, written as a superscript to the upper right of the element symbol. A charge indicates an imbalance between the number of protons and electrons in the atom.
Remember: the number of protons is fixed by the atomic number and defines the element. You cannot change the proton count without changing the element itself. Instead, atoms gain or lose electrons to acquire a charge.
Example: Finding Electron Count from Charge
Consider oxygen with a −2 charge (O2−). Oxygen has 8 protons (atomic number 8). A −2 charge means two more electrons than protons: 8 + 2 = 10 electrons. If oxygen had a +2 charge instead, it would have two fewer electrons: 8 − 2 = 6 electrons. A neutral oxygen atom (no charge written) has equal protons and electrons: 8 electrons.
2 Types of Ions: Cations and Anions
When an atom is electrically neutral (equal protons and electrons), it is simply called an atom. When it carries a charge, it becomes an ion. There are two types of ions:
| Ion Type | Charge | How It Forms | Example |
|---|---|---|---|
| Cation | Positive (+) | Atom loses electrons | Na⁺ (11 protons, 10 electrons) |
| Anion | Negative (−) | Atom gains electrons | Cl⁻ (17 protons, 18 electrons) |
Memory Aid: A cation is pawsitive. Or think of it this way: cations are positive, anions are negative. The cation vs. anion distinction becomes especially important when you study ionic bonding and nomenclature later in this chapter.
Atomic Weight and Isotopic Abundances
Atomic Weight vs. Mass Number
The atomic weight (also called atomic mass) published on the periodic table is not the same as the mass number of a single isotope. While the mass number is a whole-number count of protons plus neutrons for one specific isotope, the atomic weight is a weighted average of the masses of all naturally occurring isotopes of that element.
Example: Oxygen's Atomic Weight
The periodic table lists oxygen's atomic weight as 15.9994 amu. This value is very close to 16 because over 99% of naturally occurring oxygen is oxygen-16, with only a tiny fraction being oxygen-18. The weighted average skews heavily toward the dominant isotope.
How to Identify the Most Abundant Isotope
For most elements, you can identify the most abundant isotope by rounding the atomic weight to the nearest whole number. This works because one isotope typically dominates in nature:
- Carbon: atomic weight ≈ 12 → most abundant isotope is carbon-12 (~99%)
- Hydrogen: atomic weight ≈ 1 → most abundant isotope is hydrogen-1 (~99.98%)
- Helium: atomic weight ≈ 4 → most abundant isotope is helium-4 (~99.99%)
Common Mistake: When an element's atomic weight falls near the midpoint between two whole numbers, do not assume that the nearest whole number is the dominant isotope. This is a sign that the element has two major isotopes in roughly equal proportion. Bromine (79.9 amu) and chlorine (35.45 amu) are classic examples of this pattern.
The Bromine Exception
Bromine has an atomic weight of approximately 79.9 amu. Many students instinctively guess that bromine-80 is the most abundant isotope—but there is no bromine-80. Instead, bromine has two major isotopes: bromine-79 and bromine-81, present in nearly a 50/50 split (about 50.5% and 49.5%, respectively). Because the split is almost even, the average lands almost exactly in the middle at ~80.
How to Calculate Atomic Weight: The Weighted Average Formula
To calculate an element's atomic weight, multiply each isotope's mass (in amu) by its relative abundance (as a decimal), then add the products together.
Example: Calculating the Atomic Weight of Chlorine
Chlorine has two major isotopes: chlorine-35 (~75% abundant) and chlorine-37 (~25% abundant).
Calculation: (0.75 × 35 amu) + (0.25 × 37 amu) = 26.25 + 9.25 = 35.5 amu
This result is close to the published atomic weight of 35.4527 amu. The small difference arises because the exact masses are not perfect whole numbers and the exact abundances are not exactly 75% and 25%.
The general formula for any number of isotopes is:
Atomic Weight = (fraction1 × mass1) + (fraction2 × mass2) + ...
Exam Tip: On exams, you will often be given precise masses (e.g., 34.96885 amu) and exact abundances (e.g., 75.77%) rather than round numbers. The process is identical—convert each percentage to a decimal, multiply by the corresponding mass, and sum the results. Practice this calculation until it becomes second nature.
Introduction to the Periodic Table
The periodic table organizes approximately 120 known elements into a powerful framework based on their properties. Understanding its structure gives you immediate insight into how elements behave in general chemistry and how they relate to one another.
Periodic Table Groups and Periods
Groups are the vertical columns of the periodic table. Elements in the same group share similar physical and chemical properties. If you need to find an element with similar reactivity to fluorine, for instance, look within the same group: chlorine, bromine, and iodine are your best candidates.
Periods are the horizontal rows. The name "periodic" reflects a key pattern: as you move across a period from left to right, properties change in a regular, repeating fashion. Each period ends with a chemically inert noble gas. Then the next period begins, and you encounter elements with progressively similar properties at the same position—hence the periodicity.
Tip: Columns = Groups (similar properties grouped together). Rows = Periods (properties repeat periodically). This distinction comes up frequently on exams.
3 Categories: Metals, Non-Metals, and Metalloids
The periodic table is broadly divided into two major categories, with a small third group in between:
| Category | Location | Key Characteristics | Examples |
|---|---|---|---|
| Metals | Left of the staircase | Shiny, conductive, malleable; the majority of elements | Iron, nickel, copper, gold |
| Non-metals | Right of the staircase | Dull, poor conductors, brittle (if solid) | Oxygen, nitrogen, argon, sulfur |
| Metalloids (Semi-metals) | Along the staircase | Properties intermediate between metals and non-metals | Silicon, germanium, arsenic |
Notice that metals vastly outnumber non-metals and metalloids combined. There are only about six metalloids, positioned right along the staircase line that many periodic tables draw to separate metals from non-metals.
4 Named Periodic Table Groups You Should Know
Four groups on the periodic table have names that appear regularly in general chemistry:
| Group Name | Group Number | Elements | Key Property |
|---|---|---|---|
| Alkali Metals | Group 1 | Li, Na, K, Rb, Cs, Fr | Extremely reactive with water (explosive from Na onward) |
| Alkaline Earth Metals | Group 2 | Be, Mg, Ca, Sr, Ba, Ra | Reactive with water, but less so than alkali metals |
| Halogens | Group 17 | F, Cl, Br, I, At | Highly reactive non-metals; form salts with metals |
| Noble Gases | Group 18 | He, Ne, Ar, Kr, Xe, Rn | Chemically inert; rarely form compounds |
Note: Some groups have other names as well—oxygen's group is called the chalcogens, for instance—but those are not typically tested in introductory general chemistry.
Transition Metals, Lanthanides, and Actinides
The large block of elements in the center of the periodic table, between Groups 2 and 13, comprises the transition metals. These include familiar elements like iron, copper, and zinc. Below the main table, two additional rows contain the lanthanides and actinides, which are sometimes referred to as inner transition metals. These will come up occasionally throughout the course, but are not a major focus in introductory chapters.
6 Key Takeaways
- Dalton proposed modern atomic theory (all matter is atoms; different elements have different atoms; atoms rearrange but are not created or destroyed in reactions). Thomson discovered the electron. Millikan measured the electron's charge and mass via the oil drop experiment. Rutherford discovered the nucleus through the gold foil experiment.
- Atoms consist of a dense nucleus (containing protons and neutrons) surrounded by an electron cloud. Protons are positive (+1), neutrons are neutral (0), and electrons are negative (−1). Protons and neutrons each weigh ~1 amu; electrons are negligibly light.
- The atomic number (number of protons) defines the element. The mass number (protons + neutrons) identifies a specific isotope. Isotopes are atoms of the same element with different neutron counts.
- A neutral atom has equal protons and electrons. When the counts differ, the atom becomes an ion—a cation if positive (lost electrons), an anion if negative (gained electrons).
- Atomic weight on the periodic table is a weighted average of all naturally occurring isotopes' masses, calculated as: Σ(fractional abundance × isotopic mass). When the atomic weight falls between two whole numbers, the element likely has two major isotopes in roughly equal proportion.
- The periodic table is organized into groups (columns, similar properties) and periods (rows, repeating trends). Elements are classified as metals, non-metals, or metalloids. The 4 named groups to know: alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), and noble gases (Group 18).