Difficulty: Introductory | Prerequisites: None (first-week material)
This covers the building blocks of chemistry: what atoms are made of, how electrons are arranged, and how to read the periodic table for clues about an element's behaviour. If you can nail electron configurations and the difference between core and valence electrons, most of the bonding and reactivity material later in the course clicks into place.
Atoms are defined by their protons, and their mass number is protons plus neutrons. Electrons arrange themselves in orbitals following the Aufbau principle, and the outermost (valence) electrons determine how an element bonds and reacts. The periodic table organises all of this: group number tells you valence electrons, the d-block houses the transition metals, and seven elements naturally pair up as diatomic molecules.
Mass number
The total count of protons and neutrons in an atom's nucleus. In simple terms, it tells you how heavy a particular atom is relative to others of the same element.
Isotope
Atoms of the same element that have different numbers of neutrons, and therefore different mass numbers, but identical chemical properties. Think of it as: same element, same proton count, different weight.
Ion
A charged particle formed when a neutral atom gains or loses electrons. In simple terms, if an atom picks up extra electrons it becomes a negative ion (anion); if it loses electrons it becomes a positive ion (cation).
Electron configuration
The notation describing how electrons are distributed across an atom's orbitals and energy levels (e.g. 1s² 2s² 2p⁶ 3s² 3p² for silicon). Think of it as an address system for every electron in the atom.
Valence electrons
The electrons in an atom's outermost shell, responsible for chemical bonding and reactivity. These are the electrons that "do the work" in reactions.
Core electrons
Inner-shell electrons not involved in bonding. They sit closer to the nucleus and shield the valence electrons from the full nuclear charge.
Diatomic molecule
An element that naturally exists as a pair of two bonded atoms under standard conditions (e.g. O₂, N₂, Cl₂). The mnemonic is: H, N, O, F, Cl, Br, I (sometimes remembered as "Have No Fear Of Ice Cold Beer").
Transition metal
An element in the d-block of the periodic table, characterised by variable oxidation states and the ability to form coloured compounds. Think of it as: the large middle section of the periodic table, between the s-block and p-block.
Aufbau principle
Electrons fill orbitals starting from the lowest available energy level before occupying higher ones. In simple terms, electrons are lazy and always take the lowest-energy seat first.
Noble gas configuration
The stable, fully filled outer-shell electron arrangement that ions tend to adopt. When sodium loses an electron to become Na⁺, its configuration matches neon, which is why it is stable.
The mass number equals protons + neutrons.
An atom with 6 protons and 6 neutrons has a mass number of 12 (carbon-12).
Protons define which element it is; neutrons change only the mass number.
Isotopes share the same atomic number but differ in neutron count.
Cl-35: 17 protons, 18 neutrons.
Cl-37: 17 protons, 20 neutrons.
Both are chlorine, both react the same way, but Cl-37 is slightly heavier.
Protons define the element. Electrons determine the charge.
Gaining electrons produces a negative ion (anion); losing electrons produces a positive ion (cation).
The nitride ion (N³⁻) has 7 protons and 10 electrons (7 + 3 extra), giving it a 3– charge.
Na⁺ has 11 protons and 10 electrons (lost 1), giving it a 1+ charge.
Electrons fill orbitals in order of increasing energy (Aufbau principle).
Notation lists each sublevel and its electron count: 1s² 2s² 2p⁶ 3s² 3p² (silicon, 14 electrons).
The pattern follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, and so on.
Valence electrons sit in the highest principal energy level and participate in bonding.
Core electrons sit in lower energy levels and do not bond.
Silicon (1s² 2s² 2p⁶ 3s² 3p²): 4 valence electrons (in the n=3 shell), 10 core electrons.
The group number on the periodic table tells you the valence electron count for main-group elements.
Group 2A (alkaline earth metals): 2 valence electrons.
Group 7A (halogens): 7 valence electrons.
Group 8A (noble gases): 8 valence electrons (helium is the exception with 2).
Valence electron count drives chemical properties: elements in the same group react similarly.
Ions form to achieve a noble gas (filled outer shell) configuration.
Na (1s² 2s² 2p⁶ 3s¹) loses one electron to become Na⁺ (1s² 2s² 2p⁶), matching neon.
Cl gains one electron to become Cl⁻, matching argon.
Stability comes from having a complete outer shell, which is why these ions form so readily.
Found in the d-block (groups 3 through 12).
Characterised by variable oxidation states and coloured compounds.
Examples: manganese (Mn), molybdenum (Mo), tungsten (W).
Magnesium (Mg) is in the s-block and is not a transition metal, despite being a metal.
Seven elements exist naturally as pairs of atoms: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.
You must write these as diatomic formulas in balanced equations (e.g. O₂, not O).
Lithium and boron exist as single atoms in elemental form and are not diatomic.
Mass number
Mass number (A) = number of protons (Z) + number of neutrons (N)
Electron configuration order of filling
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p
Each sublevel holds a fixed number of electrons: s = 2, p = 6, d = 10, f = 14.
Isotopes are the basis of carbon dating in archaeology and of medical imaging tracers (e.g. technetium-99m in bone scans). Electron configurations explain why copper conducts electricity so well and why neon glows in signs: the arrangement of electrons determines every physical and chemical property you observe in daily life.
Students often think isotopes have different chemical properties. They do not. Isotopes differ only in mass (neutron count), not in how they react.
Students often confuse atomic mass with mass number. Atomic mass is the weighted average of all isotopes on the periodic table; mass number is the integer count for one specific atom.
Students often think that losing electrons makes an ion bigger. The opposite is true: cations are smaller than their parent atom because the same nuclear charge now pulls on fewer electrons.
Students sometimes place magnesium among the transition metals because it is a metal. Magnesium sits in the s-block (Group 2A) and is an alkaline earth metal, not a transition metal.
⚠️ Expect to write electron configurations from memory. Know the Aufbau filling order cold.
⚠️ Calculating protons, neutrons, and electrons in ions is a near-guaranteed exam question. Remember: protons = atomic number, neutrons = mass number minus atomic number, electrons = protons minus the charge.
⚠️ Be ready to identify which elements are diatomic. Forgetting to write O₂ instead of O in a balanced equation is a common mark lost.
⚠️ Distinguishing core from valence electrons: if given a configuration, you should be able to state both counts instantly.
True or false: Isotopes of the same element have different numbers of protons. (False: they have different numbers of neutrons.)
Fill in the blank: The mass number equals ______ + ______. (Protons + neutrons.)
True or false: A cation is smaller than its parent neutral atom. (True.)
Fill in the blank: Silicon has ______ valence electrons and ______ core electrons. (4 valence, 10 core.)
True or false: Magnesium is a transition metal. (False: it is in the s-block, Group 2A.)
Q: An atom has 17 protons and 20 neutrons. What is its mass number, and which element is it?
A: Mass number = 17 + 20 = 37. The element is chlorine (Cl-37).
Q: Write the full electron configuration for silicon (Z = 14).
A: 1s² 2s² 2p⁶ 3s² 3p².
Q: How many valence electrons and core electrons does silicon have?
A: 4 valence electrons (3s² 3p²) and 10 core electrons (1s² 2s² 2p⁶).
Q: Write the electron configuration of Na⁺. Which noble gas does it match?
A: 1s² 2s² 2p⁶. It matches neon.
Q: The nitride ion (N³⁻) has how many protons and how many electrons?
A: 7 protons and 10 electrons.
Q: List all seven diatomic elements.
A: Hydrogen (H₂), nitrogen (N₂), oxygen (O₂), fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂).
Q: Among Al³⁺, Cl⁻, Rb⁺, and Br⁻, which ion has the largest radius and why?
A: Rb⁺. It is a large alkali-metal cation with a high principal quantum number (period 5), so even after losing one electron it remains the largest of the four.
This connects to chemical nomenclature because you need to know ion charges (driven by electron configuration) to write correct ionic formulas. It also connects to periodic trends: electronegativity, ionic radius, and ionisation energy all follow from how tightly the nucleus holds its electrons. Later in the course, bonding theory (Lewis structures, VSEPR) builds directly on valence electron counts.
atomic structure, mass number, isotopes, neutron count, ions, cations, anions, electron configuration, Aufbau principle, orbital filling order, valence electrons, core electrons, inner-shell electrons, noble gas configuration, diatomic molecules, BrINClHOF, transition metals, d-block elements, s-block, periodic table groups, alkaline earth metals, halogens, noble gases, proton number, atomic number, CHEM 101, general chemistry