Electronic Configuration & Trends in the Periodic Table
The Modern Periodic Table
In the modern periodic table, elements are arranged in order of increasing atomic number (the number of protons). This ordering was established after the work of Henry Moseley in 1913, replacing Dmitri Mendeleev’s original arrangement by atomic mass.
The table is organised into horizontal rows called periods and vertical columns called groups. There are seven periods and eight main groups (Groups 1–7 and Group 0). Elements in the same group share similar chemical properties because they have the same number of electrons in their outer shell.
The group number (for Groups 1–7) equals the number of electrons in the outer shell. The period number equals the number of occupied electron shells. For example, potassium (2, 8, 8, 1) is in Group 1 (one outer electron) and Period 4 (four shells).
Electronic Configuration
Electronic configuration describes how electrons are distributed across the energy levels of an atom. Electrons always fill the lowest available energy level first. The first shell can hold a maximum of 2 electrons, while the second and third shells each hold up to 8.
To work out the electronic configuration of any element in the first 20, start with its atomic number (which equals its total number of electrons) and fill the shells in order. For example, magnesium has atomic number 12, so its configuration is 2, 8, 2 — two electrons in the first shell, eight in the second, and two in the third.
Hydrogen (1): 1. Helium (2): 2. Lithium (3): 2, 1. Beryllium (4): 2, 2. Boron (5): 2, 3. Carbon (6): 2, 4. Nitrogen (7): 2, 5. Oxygen (8): 2, 6. Fluorine (9): 2, 7. Neon (10): 2, 8. Sodium (11): 2, 8, 1. Magnesium (12): 2, 8, 2. Aluminium (13): 2, 8, 3. Silicon (14): 2, 8, 4. Phosphorus (15): 2, 8, 5. Sulfur (16): 2, 8, 6. Chlorine (17): 2, 8, 7. Argon (18): 2, 8, 8. Potassium (19): 2, 8, 8, 1. Calcium (20): 2, 8, 8, 2.
Linking Electronic Configuration to Position
The electronic configuration of an element tells you exactly where it sits in the periodic table. Look at sodium (2, 8, 1): the outer shell has 1 electron, so sodium is in Group 1; there are 3 occupied shells, so sodium is in Period 3. This pattern holds for every element in the first 20.
This link also works in reverse — if you know an element is in Group 6, Period 3, you can immediately write its electronic configuration: 2, 8, 6 (which is sulfur, atomic number 16). The exam often asks you to deduce the group and period from a configuration, or to predict a configuration from a position.
Elements in the same group have the same number of outer-shell electrons, which is why they react in similar ways. For example, all Group 1 metals (lithium, sodium, potassium) have one outer electron and react by losing it to form a +1 ion.
Metals and Non-Metals
The periodic table is divided into metals and non-metals. Metals are found on the left-hand side and in the centre of the table; non-metals sit on the right. The dividing line runs in a rough staircase shape from boron down to astatine.
This division is linked to electronic configuration. Metals typically have 1, 2 or 3 electrons in their outer shell, which they can easily lose to form positive ions. Non-metals have 4 or more outer-shell electrons (or a full outer shell for the noble gases) and tend to gain or share electrons when they react.
Most metals are shiny solids at room temperature, have high melting points, high density, and conduct electricity well. Most non-metals have low boiling points and are poor conductors of electricity. There are exceptions — mercury is a liquid metal, and carbon (as graphite) conducts electricity despite being a non-metal — but these general trends hold for the majority of elements.
Trends Across a Period
Moving across a period from left to right, the number of outer-shell electrons increases by one for each element, while the number of shells stays the same. This means elements change from reactive metals on the left (Group 1), through less reactive metals, to non-metals, and finally to very unreactive noble gases on the right (Group 0).
In Period 3, for example, sodium (2, 8, 1) is a highly reactive metal, silicon (2, 8, 4) is a metalloid, chlorine (2, 8, 7) is a reactive non-metal, and argon (2, 8, 8) is an unreactive noble gas. The pattern is driven entirely by the electronic configuration.
Trends Down a Group
Moving down a group, each element has one more occupied electron shell than the element above it. The outer-shell electrons are further from the nucleus and are shielded by the inner shells. This affects how easily the element reacts. For metals (like Group 1), reactivity increases down the group because it becomes easier to lose the outer electron. For non-metals (like Group 7), reactivity decreases because it becomes harder to attract an extra electron into the outer shell.
Down a group, each element has more electron shells, so the outer electrons are further from the nucleus and more shielded. For Group 1 metals this makes it easier to lose the outer electron, so reactivity increases. For Group 7 non-metals, the greater distance makes it harder to gain an extra electron, so reactivity decreases.
The noble gases (helium, neon, argon) sit in Group 0 on the far right of the periodic table. They have full outer shells — helium has 2, neon has 2, 8, and argon has 2, 8, 8. Because their outer shells are already complete, noble gases are very stable and unreactive. They exist as single atoms (monatomic gases) and do not easily form bonds with other elements.