Difficulty: Introductory | Prerequisites: Matter, Measurement and Properties study notes (or equivalent comfort with elements vs. compounds).
This topic moves from "what matter is" to "what matter is made of." You will learn how atoms are structured, how the periodic table organises the elements, and how electron arrangement drives chemical behaviour. The periodic trends you learn here explain why elements react the way they do, which is the bridge to bonding and reactions in the next unit.
Atoms consist of protons, neutrons and electrons. The periodic table arranges elements by atomic number, grouping those with similar properties into columns. The number of valence electrons (outermost shell) determines how an element bonds. Atomic size and ionisation energy follow predictable trends across periods and down groups.
Atom
The basic unit of matter. An atom is the smallest particle of an element that retains the chemical identity of that element.
Proton
A positively charged subatomic particle found in the nucleus. The number of protons defines which element an atom is.
Neutron
An electrically neutral subatomic particle found in the nucleus alongside protons. Neutrons contribute to mass but not to charge.
Electron
A negatively charged subatomic particle that orbits the nucleus in energy levels (shells). Electrons are responsible for chemical bonding.
Atomic number (Z)
The number of protons in an atom's nucleus. It uniquely identifies the element: every atom of carbon has Z = 6.
Mass number (A)
The total number of protons plus neutrons in the nucleus. In simple terms, it tells you the atom's "weight" in whole-number terms.
Isotope
Atoms of the same element (same number of protons) that have different numbers of neutrons, and therefore different mass numbers.
Atomic weight (atomic mass)
The weighted average of the masses of all naturally occurring isotopes of an element, reported in atomic mass units (amu).
Valence electrons
The electrons in an atom's outermost shell. They determine how an atom bonds and reacts. Think of them as the electrons that "do the chemistry."
Electron configuration
The distribution of an element's electrons among its shells and subshells (s, p, d, f). It is written as a sequence like 1s2 2s2 2p6.
Ionisation energy
The energy required to remove the outermost electron from a gaseous atom. Higher ionisation energy means the atom holds its electrons more tightly.
An element is a pure substance that cannot be broken down chemically. Each element is represented by a one- or two-letter symbol (H for hydrogen, Ca for calcium, Fe for iron). The first letter is always capitalised; the second, if present, is always lower case.
The periodic table arranges all known elements by increasing atomic number. It is divided into metals (left side, good conductors, shiny), nonmetals (right side, poor conductors, dull) and metalloids (along the "stair-step" line, intermediate properties).
Four elements make up roughly 96% of the human body by mass: oxygen, carbon, hydrogen and nitrogen. Magnesium and sulfur are classed as major minerals, while iodine and others are trace elements needed in very small amounts.
Every atom has a dense, positively charged nucleus surrounded by a cloud of electrons.
Protons carry a +1 charge and sit in the nucleus. The number of protons (atomic number, Z) defines the element.
Neutrons carry no charge and also sit in the nucleus. They add mass and help stabilise the nucleus.
Electrons carry a -1 charge and occupy energy levels (shells) around the nucleus. In a neutral atom, the number of electrons equals the number of protons.
A chemical formula such as H2O uses element symbols with subscripts to show the ratio of atoms. The subscript tells you how many atoms of that element are present in one molecule or formula unit.
Isotopes are atoms of the same element that differ in neutron count. Carbon-12 and carbon-14 are both carbon (6 protons), but carbon-12 has 6 neutrons while carbon-14 has 8.
The atomic weight (or atomic mass) listed on the periodic table is a weighted average of all naturally occurring isotopes, based on their relative abundance. For example, lead (Pb) is listed at 207.2 amu because it is a mixture of several stable isotopes.
To calculate atomic weight from isotope data, multiply each isotope's mass by its fractional abundance and add the results.
Periods are the horizontal rows (1 through 7). Moving across a period, atomic number increases by one with each element, and properties change gradually.
Groups (families) are the vertical columns. Main group elements are numbered 1A through 8A; transition metals fill groups 1B through 8B. Elements in the same group share similar chemical properties because they have the same number of valence electrons.
Key group names worth memorising:
Group 1A: alkali metals (Li, Na, K, etc.)
Group 2A: alkaline earth metals (Be, Mg, Ca, etc.)
Group 7A: halogens (F, Cl, Br, I)
Group 8A: noble gases (He, Ne, Ar, Kr)
Metals sit on the left and centre, nonmetals on the upper right, and metalloids along the diagonal stair-step boundary. Transition metals occupy the middle block and often form ions with variable charges, which is why their compounds are named with Roman numerals (e.g. iron(III) oxide).
Valence electrons are the electrons in the outermost energy level. For main group elements, the group number (1A through 8A) tells you how many valence electrons the atom has. Carbon is in group 4A, so it has 4 valence electrons.
Electrons fill shells and subshells in a predictable order. The four subshell types (s, p, d, f) hold different maximum numbers of electrons: s holds 2, p holds 6, d holds 10, f holds 14.
Maximum electrons per shell:
n = 1: 2 electrons (1s only)
n = 2: 8 electrons (2s + 2p)
n = 3: 18 electrons (3s + 3p + 3d)
n = 4: 32 electrons (4s + 4p + 4d + 4f)
The general formula is 2n2, where n is the shell number.
Electron configurations are written by filling subshells in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. For example, oxygen (Z = 8) is 1s2 2s2 2p4.
Atomic size (atomic radius)
Increases going down a group, because each period adds another electron shell, pushing the outermost electrons further from the nucleus.
Decreases going across a period from left to right, because additional protons pull the electron cloud inward more tightly while the shell number stays the same.
Ionisation energy
Decreases going down a group. Outer electrons are further from the nucleus and easier to remove.
Increases going across a period from left to right. More protons create a stronger pull on the outermost electron.
These two trends are essentially mirror images of each other: where atoms are large they lose electrons easily (low ionisation energy), and where atoms are small they hold electrons tightly (high ionisation energy). Understanding this relationship helps predict reactivity and bonding behaviour.
Carbon-14 dating uses isotope ratios to determine the age of archaeological samples. The concept of isotopes is what makes this possible.
Semiconductor manufacturing depends on metalloids (silicon, germanium) whose intermediate conductivity can be precisely tuned.
Medical imaging (PET scans) uses radioactive isotopes of elements like fluorine to detect cancer.
Students often confuse atomic number with mass number. Atomic number (Z) = protons only. Mass number (A) = protons + neutrons.
Students sometimes think electrons orbit the nucleus in neat circular paths like planets. Electrons occupy probability regions called orbitals, not fixed tracks.
Students forget that the atomic weight on the periodic table is an average, not the mass of any single atom. No individual chlorine atom weighs 35.45 amu.
Students assume that a larger atom must have more electrons than a smaller one. Atomic size depends on shell count and nuclear charge, not electron count alone.
⚠️ Know how to find protons, neutrons and electrons from the atomic number and mass number. This is tested constantly.
⚠️ Be able to write electron configurations for the first 20 elements.
⚠️ Expect questions asking you to predict periodic trends: which atom is larger, which has a higher ionisation energy.
⚠️ Understand how to calculate atomic weight from isotope masses and abundances.
⚠️ Know the group names (alkali metals, halogens, noble gases) and what properties each group shares.
True or false: An atom of sodium (Na, Z = 11) has 11 neutrons.
False. Z = 11 means 11 protons. The number of neutrons depends on the isotope (most common Na has 12 neutrons).
How many valence electrons does sulfur (group 6A) have?
Six.
Fill in the blank: Atomic size ______ across a period from left to right.
Decreases.
True or false: Isotopes of the same element have identical chemical properties.
Mostly true. Chemical properties depend on electron configuration, which is the same across isotopes of the same element.
What is the electron configuration of nitrogen (Z = 7)?
1s2 2s2 2p3.
Q: An atom has 26 protons, 30 neutrons and 26 electrons. Identify the element, its atomic number and its mass number.
A: The element is iron (Fe). Atomic number = 26. Mass number = 26 + 30 = 56.
Q: Chlorine has two stable isotopes: Cl-35 (75.77% abundance, mass 34.97 amu) and Cl-37 (24.23%, mass 36.97 amu). Calculate its atomic weight.
A: (0.7577 x 34.97) + (0.2423 x 36.97) = 26.50 + 8.96 = 35.46 amu. This matches the periodic table value.
Q: Arrange these elements in order of increasing atomic size: Na, Mg, K.
A: Mg < Na < K. Mg is to the right of Na in the same period (smaller), and K is below Na in the same group (larger).
Q: Which has a higher ionisation energy, oxygen or sulfur? Explain.
A: Oxygen. It is above sulfur in group 6A, so its valence electrons are closer to the nucleus and harder to remove.
Q: Write the electron configuration for calcium (Z = 20).
A: 1s2 2s2 2p6 3s2 3p6 4s2.
Electron configuration and valence electrons lead directly into the bonding unit, where you will use them to draw Lewis structures and predict molecular shapes. Periodic trends (ionisation energy, atomic size) explain why metals lose electrons easily while nonmetals tend to gain them, which is the foundation of ionic bonding. Isotopes come back in nuclear chemistry and radiochemistry later in the course.
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