The Periodic Table
A guided, interactive walkthrough of how the table is built and what it tells you about every element.
Chemistry is organized into ten major units. Each one builds on the last. This page covers Unit 2, The Periodic Table, which leans on the atomic structure ideas from Unit 1 and sets up the bonding work you will do in Unit 4.
By the end of this page you should be able to read a periodic table, explain why elements in a group behave similarly, predict how atoms change size and reactivity across the table, and draw a Lewis dot diagram for any main-group element or simple compound.
The periodic table is not random. It is arranged so that atoms with similar behavior line up. The rules are simple:
Rows are called Periods. There are 7 of them. Elements in the same period have the same number of principal energy levels (shells) that hold electrons.
Think of it like a library. The group tells you the "topic" (how many valence electrons), and the period tells you the "floor" (how many shells). Two elements in the same group behave like cousins because they have the same outer electron count, even though they live on different floors.
Run your finger down the bold zig-zag line on the right side of the table. That line is called the staircase. It divides the table into two big neighborhoods.
Nonmetals sit to the right of the staircase. They are dull, brittle, and poor conductors. Many of them are gases at room temperature.
Metalloids sit on the staircase itself. They have a mix of properties. The six metalloids you should know are B, Si, Ge, As, Sb, Te.
One quick note about hydrogen: it sits above the metals in Group 1, but it is a nonmetal. Its placement is about electron count, not behavior.
The last column, Group 18, is its own special category. These are the noble gases: He, Ne, Ar, Kr, Xe, Rn. They have completely filled valence shells, which makes them stable and unreactive. They do not normally form bonds. Every other element on the table is, in some sense, trying to become like them.
Here is the single most important idea in this unit. Atoms want a full outer shell. For most atoms that means eight valence electrons. That is where the word octet comes from.
This is why sodium (1 valence electron) gives up an electron to become Na$^+$, and why chlorine (7 valence electrons) grabs one to become Cl$^-$. Both end up with 8 valence electrons and the configuration of a noble gas.
Helium is the one exception worth remembering: its shell fills with 2 electrons, not 8. So when hydrogen gets stable, it aims for 2 electrons (duet rule), matching He.
An electron configuration is a shorthand for where an atom's electrons actually live. Every electron in an atom sits in a specific space around the nucleus called an orbital. Orbitals are grouped into subshells (named s, p, d, f), and subshells are grouped into shells (numbered 1, 2, 3, ...).
The number in front tells you the shell. The letter tells you the subshell. The superscript tells you how many electrons are in that subshell.
Add up all the superscripts and you get the total number of electrons in the atom, which equals the atomic number for a neutral atom. Sodium: $2 + 2 + 6 + 1 = 11$.
Electrons fill the lowest-energy subshell first and work their way up. The order is not a simple top-to-bottom list because subshells in different shells overlap in energy.
Notice the jump: after 3p you do 4s before 3d. That happens because 4s has slightly lower energy than 3d.
Writing out every subshell for a big atom gets tedious. Shortcut: replace the inner electrons with the nearest previous noble gas in brackets, then write the rest. Sodium becomes $[\text{Ne}] \, 3s^1$. Potassium becomes $[\text{Ar}] \, 4s^1$.
When an atom becomes an ion you add or subtract electrons from the outermost shell. Sodium becomes Na$^+$ by losing its $3s^1$, giving $1s^2 \, 2s^2 \, 2p^6$, which is the same as neon. Chlorine becomes Cl$^-$ by gaining one electron to fill its 3p, giving $[\text{Ar}]$.
An Aufbau diagram (sometimes called an orbital diagram) shows each orbital as a box and each electron as an arrow. It is a more detailed picture than the flat text configuration. "Aufbau" is German for "building up", because you build the atom by adding one electron at a time into the lowest available spot.
1. Aufbau Principle. Fill the lowest-energy orbital first. Same order as for configurations: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.
2. Pauli Exclusion Principle. An orbital can hold at most 2 electrons, and they must have opposite spins (one arrow up, one arrow down).
3. Hund's Rule. When there are multiple orbitals in the same subshell (like the three 2p orbitals), put one electron in each orbital first, all with parallel spins (all up), before pairing up.
In plain words: the bus seats fill one per row first, then doubles up.
Draw each subshell as a row of boxes: one box for s, three for p, five for d, seven for f. Label each row with the subshell name. Place arrows inside the boxes, first all up (Hund's rule), then pair up with down arrows.
The layout of the periodic table is not arbitrary, it mirrors the filling order of the orbitals.
STP stands for Standard Temperature and Pressure. The standard chemistry definition is:
At STP the large majority of elements are solids. Only a small set are not, so memorizing the list is fast:
Try the "State at STP" view on the periodic table above to see this visually. You will notice the gases cluster in the top right, with hydrogen being the exception way on the left.
An allotrope is one of two or more different forms of the same element, where the atoms are connected to each other in different arrangements. Same element, different structure, different properties.
The classic example is carbon. Diamond is a giant 3D lattice where each carbon bonds to four others. It is the hardest natural substance and does not conduct electricity. Graphite is stacked flat sheets of carbon. It is soft, slippery, and does conduct electricity. Both are pure carbon.
An element is a pure substance made of only one kind of atom. It cannot be broken down by chemical means. A compound is two or more different elements chemically bonded together in a fixed ratio. Compounds can be broken down into their elements, but only through chemical reactions, not by physical methods like filtering.
Here is the payoff for how the table is arranged. Three properties of atoms follow predictable patterns as you move across or down the table. Learn the patterns once and you can predict behavior for any element.
Down a group (top to bottom): Electronegativity decreases. Atomic radius increases. Ionization energy decreases.
Why? Two effects fight each other. Going across a period, the nucleus gets more positive charge but electrons stay in the same shell, so the atom pulls its electrons in tighter (smaller radius, harder to remove an electron, stronger pull on shared electrons). Going down a group, you add a whole new shell each time, so the outer electrons are farther from the nucleus and shielded by inner electrons (bigger radius, easier to lose an electron, weaker pull on shared electrons).
The simple bars above show one period and one group. The heat map below shows every element at once. Hover any cell to see its exact value. Switch trends to see how each pattern looks across the whole table.
Atoms become ions when they gain or lose electrons to reach a stable octet. The number of protons never changes in a chemical reaction, only the number of electrons. Protons carry the charge identity of the element, electrons cancel that charge out.
Anions (negative ions) form when atoms gain electrons. Nonmetals do this. Example: Cl gains 1 electron to become Cl$^-$.
This is the tricky part that gets tested a lot.
When an atom gains electrons, its radius gets larger. The same nucleus now has to hold more electrons, which spread out and repel each other.
Sodium has 1 valence electron. It loses that electron to reveal the stable Ne configuration underneath. The ion is noticeably smaller.
A common test question asks which noble gas an ion matches. The trick is to count electrons in the ion, not in the atom.
A Lewis electron-dot diagram is a simple way to draw an atom's valence electrons. You write the element symbol, then put dots around it for each valence electron. That is it.
So fluorine, with 7 valence electrons, ends up with one pair on three sides and a single dot on the fourth. Carbon, with 4, has one single dot on each side.
For an ion, put brackets around the symbol (with its dots) and write the charge outside as a superscript. A cation like Na$^+$ often shows no dots at all (it has given them away). An anion like Cl$^-$ shows 8 dots because it has gained the electron it needed for a full octet.
Potassium has 1 valence electron. Iodine has 7. K hands its one valence electron to I. Now K is K$^+$ (no dots) and I is I$^-$ (8 dots). Write each ion in its own brackets.
K gives up its valence electron. I now has 8. Each is stable.
When two nonmetals bond, neither one is willing to fully give up an electron. They share electrons instead. A shared pair of electrons is a covalent bond, often drawn as a single line between the two atoms. Each atom in the pair still needs a full octet (or duet, for hydrogen), so you keep adding electrons around each atom until everyone is happy.
1. Count total valence electrons by adding up each atom's valence count.
2. Put the least electronegative atom in the middle (hydrogen is never central).
3. Draw single bonds from the center to each outer atom. Each bond uses 2 electrons.
4. Distribute the remaining electrons as lone pairs, starting with the outer atoms, until they have octets (or duets).
5. If the center atom still lacks an octet, move lone pairs from an outer atom into double or triple bonds until everyone has 8.
If single bonds alone cannot give the center atom a full octet, share more. A double bond is two shared pairs (4 electrons) and is drawn as two lines. A triple bond is three shared pairs (6 electrons) and is drawn as three lines. Nitrogen gas, N$_2$, has a triple bond between its two atoms. Each N sees 6 electrons in the triple bond plus 2 in a lone pair, for a full octet.
A bonding pair is shared between two atoms and counts toward the octet of both.
A lone pair sits on one atom and counts only for that atom.
You should be able to go either direction: given a name, write the formula; given a formula, write the name. For this unit we are sticking with simple compounds. The key rule for ionic compounds is that positive and negative charges must cancel.
Click the card to flip. Use the arrows to move through the deck. Shuffle whenever you want a fresh order.
Try each one before opening the answer.
(A) K$^+$ with 8 dots around I, in brackets, charge minus
(B) K$^+$ alone, then [I with 8 dots]$^-$
(C) K with 8 dots, no brackets, paired with I with 8 dots
(D) I with 8 dots (no charge) next to K with 1 dot
(A) helium (B) neon (C) argon (D) krypton
Write H, draw a shared pair between H and Cl (two dots or a line), and then add Cl's three remaining lone pairs around Cl. Result: H shares 2 electrons (its duet), and Cl has 8 (2 shared + 6 in three lone pairs). Standard drawing:
H : Cl with three more pairs of dots around Cl (one on top, one on right, one on bottom)
Put N in the center with three H atoms around it. Draw a single bond from N to each H (3 bonds, 6 electrons). That leaves 2 electrons, which go as a lone pair on N. Result: three N-H bonds plus one lone pair on N. Nitrogen has 8 (6 in bonds + 2 in lone pair). Each H has 2.