The periodic table organizes elements by atomic number, revealing patterns in their properties. Understanding how the table is structured and why trends exist gives you a foundation for predicting how elements behave and bond with each other.

Structure of the Periodic Table

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Organization of Elements

The periodic table arranges all known elements in order of increasing atomic number (the number of protons). This arrangement isn't random; it groups elements so that patterns in their chemical behavior become visible.

Groups are the vertical columns. Elements in the same group have similar chemical properties because they share the same number of valence electrons.
Periods are the horizontal rows. Elements in the same period have the same number of electron shells (energy levels).
Metals occupy the left and center of the table. They tend to be shiny, conduct heat and electricity well, and are malleable (can be hammered into shapes).
Nonmetals sit on the right side. They're typically gases or brittle solids at room temperature, and they're poor conductors.
Metalloids fall along the staircase boundary between metals and nonmetals. Elements like silicon and germanium share properties of both, which is why silicon is so useful in semiconductors.

Element Categories and Properties

Certain groups have special names because their members behave in distinctive ways:

Alkali metals (Group 1): Highly reactive metals that readily lose one electron. Reactivity increases as you go down the group. Lithium reacts gently with water, while potassium reacts violently enough to ignite.
Alkaline earth metals (Group 2): Reactive metals that lose two electrons, though they're less reactive than alkali metals. Calcium and magnesium are common examples.
Halogens (Group 17): Highly reactive nonmetals that readily gain one electron to complete their outer shell. Fluorine and chlorine are common examples.
Noble gases (Group 18): Extremely stable because they already have full outer electron shells. They rarely react with other elements.
Transition metals (d-block): Found in the middle of the table. They often form colored compounds (think of the blue in copper sulfate) and frequently serve as catalysts in chemical reactions.
Lanthanides and actinides (f-block): These two rows sit below the main table. They have unique magnetic and optical properties. Many actinides are radioactive.

Atomic Properties

Fundamental Atomic Characteristics

Atomic number is the number of protons in an atom's nucleus. It defines which element you're dealing with. Carbon always has 6 protons; oxygen always has 8.
Atomic mass is the weighted average mass of an element's naturally occurring isotopes, measured in atomic mass units (amu). It's a weighted average because some isotopes are more common in nature than others. For example, chlorine's atomic mass is about 35.5 amu because it's roughly 75% chlorine-35 and 25% chlorine-37.
Atomic radius is the distance from the nucleus to the outermost electrons. Two trends to know:
- Increases going down a group because each new period adds another electron shell, making the atom physically larger.
- Decreases going left to right across a period because protons are being added to the nucleus without adding a new shell. The stronger nuclear charge pulls the existing electrons in tighter.
Electronegativity measures how strongly an atom attracts electrons in a chemical bond.
- Increases from left to right across a period.
- Decreases from top to bottom within a group.

Notice that atomic radius and electronegativity trend in opposite directions. That makes sense: smaller atoms hold their electrons more tightly, so they're better at pulling in shared electrons from a bond.

Organization of Elements, Elements and The Periodic Table | Introductory Chemistry – Lecture & Lab

Electron Configuration and Valence Electrons

Electron configuration describes how electrons are distributed across an atom's orbitals (energy sublevels). The Aufbau principle tells you to fill orbitals in order of increasing energy: 1s, then 2s, then 2p, then 3s, and so on.

Here's how to write an electron configuration step by step:

Find the element's atomic number. That tells you the total number of electrons in a neutral atom.
Fill orbitals in energy order: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p$ and so on.
Each s sublevel holds up to 2 electrons, each p holds up to 6, and each d holds up to 10.
Write the result with superscripts showing how many electrons are in each sublevel. For example, sodium (atomic number 11) is $1s^2 \, 2s^2 \, 2p^6 \, 3s^1$ .

Valence electrons are the electrons in the outermost shell. They're the ones that participate in chemical bonding, which is why they matter so much. Elements in the same group have the same number of valence electrons, and that's exactly why they behave similarly. Sodium and potassium both have 1 valence electron, so both are reactive metals that form +1 ions.

The octet rule says atoms tend to gain, lose, or share electrons until they have eight valence electrons (a full outer shell), mimicking the stable configuration of noble gases. Hydrogen is an exception since it only needs two electrons to fill its single shell.

Periodic Trends

Ionization energy is the energy required to remove an electron from a neutral atom in the gas phase. Think of it as how tightly an atom holds onto its electrons.

Increases left to right across a period. More protons means a stronger pull on the electrons, so it takes more energy to remove one.
Decreases going down a group. Valence electrons are farther from the nucleus and shielded by inner electron shells, so they're easier to remove.

This is why alkali metals (far left, low ionization energy) lose electrons so easily, while noble gases (far right, high ionization energy) hold onto theirs.

Electron affinity is the energy change when a neutral atom gains an electron. A more negative (larger magnitude) electron affinity means the atom releases more energy when it gains an electron, meaning it "wants" that extra electron more.

Generally increases in magnitude from left to right across a period.
Trends within groups are less consistent because of factors like electron-electron repulsion in smaller atoms.

Electronegativity indicates how strongly an atom attracts shared electrons in a bond. Fluorine has the highest electronegativity of any element (3.98 on the Pauling scale). Noble gases typically don't have assigned electronegativity values because they rarely form bonds.

Atomic radius is typically measured as half the distance between the nuclei of two bonded atoms of the same element.

Increases going down a group as new electron shells are added.
Decreases across a period as increasing nuclear charge pulls electrons closer.

Ionic radius differs from atomic radius because the atom has gained or lost electrons:

Cations (positive ions, formed by losing electrons) are smaller than their parent atoms. Fewer electrons means less electron-electron repulsion, and the same nuclear charge pulls the remaining electrons in more tightly. For example, $Na$ has an atomic radius of about 186 pm, but $Na^+$ shrinks to about 102 pm.
Anions (negative ions, formed by gaining electrons) are larger than their parent atoms. Extra electrons increase repulsion, and the same nuclear charge can't hold the larger electron cloud as tightly. For example, $Cl$ has an atomic radius of about 99 pm, but $Cl^-$ expands to about 181 pm.

A helpful way to remember: losing electrons shrinks the atom (cation = smaller), gaining electrons expands it (anion = larger).

Putting the Trends Together

All of these trends share a common cause: the balance between nuclear charge (protons pulling electrons in) and electron shielding (inner electrons blocking that pull). Across a period, nuclear charge wins because you're adding protons without adding new shells. Down a group, shielding wins because each new shell puts more distance between the nucleus and the valence electrons.

If you can remember that one idea, you can reason through any trend on the periodic table without memorizing each one separately.

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