CHEM 111, Chapter 2 – Lecture 3 | Source: Lecture notes / textbook Ch. 2
Difficulty: Introductory | Prerequisites: Basic understanding of elements and the periodic table (Chapter 1 notes).
This chapter sits at the foundation of general chemistry. Before you can talk about reactions, stoichiometry, or bonding, you need to understand what atoms are made of, how they differ from one another, and how they combine into molecules and ions. The material here builds directly on Chapter 1's introduction to matter and measurement. If you are comfortable with elements and the periodic table, you are ready for this. Everything that follows in the course (chemical equations, mole calculations, bonding theory) depends on the vocabulary and concepts introduced here.
Atoms of the same element can have different numbers of neutrons, making them isotopes with different mass numbers. We write atomic structure using the atomic number (Z, protons), mass number (A, protons + neutrons), and element symbol. Seven elements naturally exist as diatomic molecules, and atoms become ions by gaining or losing electrons.
Isotopes
Atoms of the same element that have different numbers of neutrons and therefore different mass numbers. Think of it as: same element, same number of protons, but a heavier or lighter version because of extra or fewer neutrons.
Atomic number (Z)
The number of protons in the nucleus of an atom. This is what defines which element an atom is. In simple terms, Z is the element's identity card.
Mass number (A)
The total count of protons plus neutrons in an atom's nucleus. In simple terms, it tells you how heavy that particular atom is (ignoring electrons, which have negligible mass).
Neutron number (N)
The number of neutrons in the nucleus. Calculated as A minus Z. Think of it as: total particles in the nucleus minus the protons.
Diatomic elements
Seven elements that naturally exist as two-atom molecules rather than single atoms: H₂, O₂, F₂, Br₂, I₂, N₂, Cl₂. Think of it as: these elements are "social" and always travel in pairs.
Ion
An atom or group of atoms that carries a net electrical charge because it has gained or lost electrons. In simple terms, a charged version of an atom.
Cation
A positively charged ion, formed when a neutral atom loses one or more electrons. There are now more protons than electrons. Think of it as: the "cat" in cation has paws (plus, positive).
Anion
A negatively charged ion, formed when a neutral atom gains one or more electrons. There are now more electrons than protons. Think of it as: "a negative ion."
All isotopes of an element share the same atomic number (same number of protons) but differ in neutron count.
Hydrogen has three isotopes, which makes it the clearest example:
Protium (ordinary hydrogen): mass number 1, atomic number 1 (0 neutrons)
Deuterium (heavy hydrogen): mass number 2, atomic number 1 (1 neutron)
Tritium (radioactive hydrogen): mass number 3, atomic number 1 (2 neutrons)
All three are hydrogen (Z = 1), but their physical properties differ slightly because of the extra neutrons.
Z = atomic number = number of protons. This defines the element.
N = number of neutrons.
A = mass number = Z + N = protons + neutrons.
The standard notation places A as a superscript and Z as a subscript to the left of the element symbol X.
Three accepted formats (all equivalent):
Element name followed by mass number: Carbon-12
Element symbol followed by mass number: C-12
Superscript notation: ¹²C
All three tell you the same information: the element and its mass number.
Seven elements that exist naturally as two-atom molecules, not as lone atoms:
H₂, O₂, F₂, Br₂, I₂, N₂, Cl₂
A useful mnemonic: HOFBrINCl (pronounced "Hoff-brinkle").
You will need these when writing balanced chemical equations, because these elements never appear as single atoms in their standard state.
A neutral atom has equal numbers of protons and electrons.
Remove electrons from a neutral atom, and a cation (positive ion) forms. More protons than electrons means a net positive charge.
Add electrons to a neutral atom, and an anion (negative ion) forms. More electrons than protons means a net negative charge.
The number of protons never changes when an ion forms. Only the electron count changes.
Mass number formula:
A = Z + N
(mass number = atomic number + neutron number)
Rearranged forms:
N = A − Z (to find neutrons when you know the mass number and atomic number)
Z = A − N (to find protons when you know the mass number and neutron count)
Isotope notation:
ᴬ_Z X, where A is the mass number (superscript), Z is the atomic number (subscript), and X is the element symbol.
Isotopes have direct practical uses. Deuterium (heavy hydrogen) is used in nuclear reactors as a moderator, and tritium is used in self-luminous exit signs and as a tracer in environmental science. Carbon-14 (another isotope you will encounter later) is the basis of radiocarbon dating in archaeology. Understanding ions is essential for making sense of how batteries work, how nerves transmit signals, and why table salt dissolves in water.
Students often think that isotopes of an element are different elements. They are not. Isotopes share the same atomic number and the same chemical behaviour; only the neutron count (and therefore mass) differs.
Students frequently confuse atomic number with mass number. The atomic number (Z) is protons only. The mass number (A) includes protons and neutrons. If someone asks "how many particles are in the nucleus," the answer is A, not Z.
A common error is thinking that forming an ion changes the number of protons. It does not. Ion formation changes only the number of electrons.
Students sometimes forget that cations are positive and anions are negative, or mix up which is which. Remember: cation has a "t" that looks like a plus sign.
⚠️ You will almost certainly be asked to calculate the number of protons, neutrons, or electrons in a given isotope or ion. Practise A = Z + N in all three rearranged forms.
⚠️ Expect a question requiring you to list or identify the seven diatomic elements. Memorise HOFBrINCl.
⚠️ Ion problems are common: "If sodium (Z = 11) loses one electron, what is the charge?" You need to know that removing an electron from a neutral atom gives a +1 cation.
⚠️ Isotope notation questions will ask you to interpret or write ᴬ_Z X symbols. Be comfortable reading these in all three formats (name-mass, symbol-mass, superscript notation).
True or false: Isotopes of the same element have different numbers of protons. (False: they have different numbers of neutrons.)
Fill in the blank: A = ___ + N. (Z)
True or false: A cation has more electrons than protons. (False: a cation has fewer electrons than protons.)
Name four of the seven diatomic elements. (Any four of: H₂, O₂, F₂, Br₂, I₂, N₂, Cl₂.)
True or false: Tritium and protium are different elements. (False: both are hydrogen, with different mass numbers.)
Q: An atom has 6 protons and 8 neutrons. What is its mass number, and what element is it?
A: Mass number = 6 + 8 = 14. The element is carbon (Z = 6). This isotope is carbon-14.
Q: How many neutrons are in an atom of ²³Na (sodium-23, Z = 11)?
A: N = A − Z = 23 − 11 = 12 neutrons.
Q: A neutral chlorine atom (Z = 17) gains one electron. What type of ion forms, and what is its charge?
A: An anion forms with a charge of −1. It now has 18 electrons and 17 protons.
Q: Write the isotope notation for an atom with 26 protons and 30 neutrons.
A: Z = 26 (iron, Fe), A = 26 + 30 = 56. The notation is ⁵⁶Fe or Fe-56.
Q: List all seven diatomic elements.
A: Hydrogen (H₂), oxygen (O₂), fluorine (F₂), bromine (Br₂), iodine (I₂), nitrogen (N₂), chlorine (Cl₂).
Q: Deuterium has a mass number of 2 and an atomic number of 1. How many neutrons does it have?
A: N = 2 − 1 = 1 neutron.
This connects to atomic mass and the mole (Chapter 3) because you need to understand isotopes before you can calculate weighted average atomic masses from the periodic table. It also connects to electron configurations and periodic trends (Chapters 6 and 7), where the number of protons (Z) determines an element's position on the periodic table and its chemical properties. The concept of ions introduced here is the foundation for ionic bonding and compound naming (later in Chapter 2 and Chapter 3).
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