Atomic Structure and the Periodic Table

What is Matter Made Of?

If you keep cutting a solid block of iron into smaller and smaller pieces, eventually you reach a point where it can no longer be divided. All substances in the universe are made of tiny particles called atoms.

An atom is defined as the smallest part of an element that can exist. Matter is classified into elements, compounds, and mixtures. An element is a substance that is made of only one type of atom, and there are approximately 100 different elements displayed on the periodic table.

Examples include Iron (Fe), which is a metallic element arranged in a regular, giant lattice structure. Oxygen (O) is a non-metal that naturally exists as diatomic molecules ($O_2$), meaning two oxygen atoms are chemically joined. It is strictly classified as an element, not a compound, because it contains only one type of atom.

Inside the Atom

Atoms are incredibly small, with a radius of about $0.1$ nanometres ($1 \times 10^{-10}$ m). They consist of a central nucleus containing protons and neutrons, surrounded by negatively charged electrons in energy levels (shells).

The radius of the nucleus is about $1/10,000$ of that of the atom (approximately $1 \times 10^{-14}$ m). Despite its tiny size, virtually all the mass of an atom is concentrated in the central nucleus.

The number of protons, known as the atomic number, is unique to each element and dictates its identity. The mass number is the sum of the protons and neutrons in the nucleus.

The Periodic Table and Atomic Structure

Before the discovery of protons, scientists like Dmitri Mendeleev ordered elements by atomic weight. However, this caused problems because elements were sometimes placed in groups with completely different chemical properties.

Mendeleev overcame this by leaving gaps for undiscovered elements and swapping the order of some elements to match their chemical properties. The later discovery of isotopes explained why ordering by atomic weight was incorrect, as isotopes have different masses but identical chemical properties.

The modern periodic table arranges elements in order of increasing atomic (proton) number. This sequential arrangement ensures that elements with similar properties fall naturally into the same vertical columns without needing to be manually swapped.

Electronic Configuration and Table Placement

The arrangement of electrons in shells is called the electronic configuration. For the first 20 elements, the shells follow a strict sequence of maximum capacities: 2 electrons in the first shell, 8 in the second, 8 in the third, and 2 in the fourth ($2, 8, 8, 2$).

The periodic table is directly linked to this atomic structure. Periods (horizontal rows) represent the number of occupied electron shells an atom has.

Groups (vertical columns) represent the number of valence electrons in the highest occupied energy level (outer shell). Elements in the same group have similar chemical properties exactly because they have the same number of electrons in their outer shell.

Worked Example: Deducing Table Position

You can determine an element's position on the periodic table purely from its atomic number.

Deduce the group and period of Sodium (atomic number 11).

Step 1: Write the electronic configuration using the $2, 8, 8, 2$ rule.

• Sodium has 11 electrons: 2 in the first shell, 8 in the second, and 1 in the third ($2, 8, 1$).

Step 2: Use the number of shells to find the period.

• There are 3 numbers, meaning 3 occupied shells, so Sodium is in Period 3.

Step 3: Use the final number to find the group.

• The last number is 1, meaning 1 electron in the outer shell, so Sodium is in Group 1.

Group Trends and Chemical Properties

Because elements in the same group share the same number of outer electrons, they react in similar ways. However, their reactivity changes as you move down the group due to shielding, where inner electron shells reduce the electrostatic attraction between the positive nucleus and the negative outer electrons.

Group 1 metals all have one outer electron and react by losing it to form $1+$ ions. Reactivity increases down the group because the outer electron is further from the nucleus and more shielded, meaning the electrostatic attraction is weaker and the electron is lost more easily.

Group 7 halogens have seven outer electrons and react by gaining one to form $1-$ ions. Reactivity decreases down the group because a larger atomic radius and more shielding mean the nucleus has a weaker attraction for an incoming electron.

Group 0 elements (noble gases) do NOT react easily. They are inert because they already have a stable, full outer shell of 8 electrons (except Helium, which has a full shell of 2).