Difficulty: Introductory | Prerequisites: Familiarity with elements and the periodic table layout.
This material covers how our understanding of the atom developed (Dalton through Rutherford) and the key numbers that define an atom: atomic number, mass number, and atomic mass. It also introduces ions and how the periodic table predicts the charges elements tend to form. You need to be comfortable with the layout of the periodic table (groups, periods, metals vs non-metals) before tackling this. Everything here feeds into electron configuration, bonding, and stoichiometry later in the course.
Atoms are made of protons (positive, in the nucleus), neutrons (neutral, in the nucleus), and electrons (negative, outside the nucleus). The atomic number (Z) is the number of protons and defines the element. The mass number (A) is protons plus neutrons. When an atom gains or loses electrons, it becomes an ion: cations are positive (lost electrons), anions are negative (gained electrons). The periodic table's group number predicts the ion charge for main-group elements.
Atom
The smallest unit of an element that retains the chemical properties of that element. Made up of protons, neutrons, and electrons.
Proton
A positively charged subatomic particle found in the nucleus. Mass: 1.673 x 10⁻²⁴ g. The number of protons defines which element an atom is.
Neutron
A neutral (no charge) subatomic particle found in the nucleus. Mass: 1.675 x 10⁻²⁴ g. Nearly identical in mass to the proton.
Electron
A negatively charged subatomic particle found outside the nucleus. Mass: 9.11 x 10⁻²⁸ g, roughly 1/1836 the mass of a proton. Electrons occupy nearly all the volume of the atom but contribute almost none of its mass.
Atomic number (Z)
The number of protons in the nucleus of an atom. In a neutral atom, Z also equals the number of electrons. Z defines the element.
Think of it as: the element's identity number on the periodic table.
Mass number (A)
The total number of protons and neutrons in the nucleus. A = Z + number of neutrons.
Atomic mass
The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units (amu). This is the number shown on the periodic table and is not a whole number because it accounts for isotope abundance.
Isotope
Atoms of the same element (same Z) with different numbers of neutrons (different A). Example: hydrogen has three isotopes: protium (0 neutrons), deuterium (1 neutron), tritium (2 neutrons).
Ion
An atom or group of atoms with a net electrical charge, formed when an atom gains or loses electrons.
Cation
A positively charged ion. Formed when an atom loses one or more electrons. Metals typically form cations.
In simple terms, "cat" has a "t" for "taken away" (electrons removed).
Anion
A negatively charged ion. Formed when an atom gains one or more electrons. Non-metals typically form anions.
John Dalton proposed five key ideas:
All matter consists of atoms.
All atoms of a given element are identical (we now know isotopes make this not quite true, but it was a useful starting model).
Atoms combine in whole-number ratios to form compounds.
Chemical reactions are rearrangements of atoms; atoms are neither created nor destroyed.
Atoms cannot be divided (also revised later with the discovery of subatomic particles).
Lavoisier: established the law of conservation of mass (mass is neither created nor destroyed in a chemical reaction).
Proust: established the law of constant composition (a given compound always contains the same elements in the same proportions by mass).
Dalton: law of multiple proportions (when two elements form more than one compound, the ratios of the masses are small whole numbers).
Thomson: characterised the electron using cathode ray tubes (CRT), discovering the first subatomic particle.
Rutherford: proposed the nuclear model of the atom (a small, dense, positive nucleus with electrons orbiting at a distance).
Faraday / Arrhenius: contributed to understanding the formation of ions.
Electron: charge = -1. Mass = 9.11 x 10⁻²⁸ g. Located outside the nucleus. Takes up nearly all the volume of the atom.
Proton: charge = +1. Mass = 1.673 x 10⁻²⁴ g. Located in the nucleus.
Neutron: charge = 0. Mass = 1.675 x 10⁻²⁴ g. Located in the nucleus.
Key point: almost the entire mass of the atom is in the nucleus (protons + neutrons), but the electrons occupy almost the entire volume.
Atomic number (Z): number of protons. Defines the element. In a neutral atom, Z = number of electrons.
Mass number (A): protons + neutrons. Always a whole number for a specific isotope.
Atomic mass: the weighted average across all naturally occurring isotopes. This is why the periodic table shows decimal values (e.g. chlorine is 35.45 amu, reflecting a mix of Cl-35 and Cl-37).
Protium: 1 proton, 0 neutrons (the most common).
Deuterium: 1 proton, 1 neutron.
Tritium: 1 proton, 2 neutrons.
When an atom gains or loses electrons, it becomes an ion. The protons and electrons are no longer equal, giving the atom a net charge.
Group 1 (alkali metals): form +1 ions (lose 1 electron).
Group 2 (alkaline earth metals): form +2 ions (lose 2 electrons).
Aluminium: forms a +3 ion.
Group 16 (e.g. oxygen, sulphur): form -2 ions (gain 2 electrons).
Group 17 (halogens): form -1 ions (gain 1 electron).
Nitrogen and phosphorus: form -3 ions (gain 3 electrons).
Isotopes are used in medicine (radioactive tracers, PET scans), archaeology (carbon-14 dating), and energy (uranium-235 in nuclear reactors). Understanding ion charges is essential for predicting the formulas of ionic compounds and for electrochemistry.
Students often confuse atomic mass with mass number. Atomic mass is a weighted average (a decimal), while mass number is a whole number for a specific isotope.
Students sometimes think that gaining electrons makes an atom a cation. It does not. Gaining electrons makes an atom an anion (negative charge). Losing electrons makes a cation (positive charge).
Students forget that in a neutral atom, the number of protons equals the number of electrons. Once you change the electron count, it is no longer neutral; it is an ion.
"Groups" on the periodic table are vertical columns, not horizontal rows. Rows are called "periods." The lecture notes contain an error stating "groups = horizontal"; this is incorrect.
⚠️ Be able to determine the number of protons, neutrons, and electrons for any atom or ion given its symbol, atomic number, and mass number.
⚠️ Know the charges that main-group elements form. Group 1 = +1, Group 2 = +2, Al = +3, halogens = -1, Group 16 = -2, N and P = -3.
⚠️ Distinguish between atomic number, mass number, and atomic mass. Expect a question asking which is the weighted average.
⚠️ Know the five postulates of Dalton's atomic theory and which ones have since been revised.
⚠️ Match each scientist to their contribution: Lavoisier (conservation of mass), Proust (constant composition), Thomson (electron), Rutherford (nuclear model).
Fill in the blank: the atomic number tells you the number of ______ in the nucleus.
True or false: a cation has more electrons than protons.
Fill in the blank: mass number = protons + ______.
True or false: Rutherford discovered the electron.
How many protons, neutrons, and electrons does a Na⁺ ion have? (Na: Z = 11, A = 23.)
Answers: 1. Protons. 2. False (a cation has fewer electrons than protons). 3. Neutrons. 4. False (Thomson discovered the electron; Rutherford proposed the nuclear model). 5. 11 protons, 12 neutrons (23 - 11), 10 electrons (11 - 1 from the +1 charge).
Q: An atom has 17 protons, 18 neutrons, and 18 electrons. What is the element, the mass number, and the charge?
A: The element is chlorine (Z = 17). Mass number = 17 + 18 = 35. Charge = 17 - 18 = -1, so it is Cl⁻.
Q: What is the difference between atomic mass and mass number?
A: Mass number is the count of protons + neutrons for a single isotope (always a whole number). Atomic mass is the weighted average of the mass numbers of all naturally occurring isotopes of that element (usually a decimal).
Q: Name two of Dalton's postulates that have since been modified.
A: (1) "All atoms of an element are identical" has been modified because isotopes show that atoms of the same element can have different masses. (2) "Atoms cannot be divided" has been modified because atoms are made of protons, neutrons, and electrons.
Q: Which scientist proposed the law of conservation of mass, and what does it state?
A: Lavoisier. Mass is neither created nor destroyed in a chemical reaction. The total mass of the reactants equals the total mass of the products.
Q: Why does sodium form a +1 ion while chlorine forms a -1 ion?
A: Sodium (Group 1) loses one electron to achieve a stable electron configuration, giving it a +1 charge. Chlorine (Group 17 / halogen) gains one electron to complete its outer shell, giving it a -1 charge.
Ion charges feed directly into writing formulas for ionic compounds (covered in chemical nomenclature). Isotopes and atomic mass connect to mass spectrometry, nuclear chemistry, and molar mass calculations in stoichiometry. Dalton's atomic theory is a historical foundation, but electron configuration (coming later) refines the model further.
Atomic theory, Dalton, Lavoisier, Proust, Thomson, Rutherford, Faraday, Arrhenius, proton, neutron, electron, atomic number, mass number, atomic mass, isotope, ion, cation, anion, protium, deuterium, tritium, periodic table groups, alkali metals, alkaline earth metals, halogens, CHEM 101, general chemistry, Michigan State University