Learn the ideas first, then practice: build ionic compounds, draw covalent ones, name them, and build acids. Every practice tab has Hint, Reset and Skip, 3 levels and a crown to win. The toolbox button (bottom right) is always there when you need to look something up.
A metal and a nonmetal. The metal gives electrons away and becomes a positive ion (cation). The nonmetal takes them and becomes a negative ion (anion). Opposite charges attract. Example: Na + Cl makes NaCl.
Only nonmetals. Neither atom wants to give electrons up, so they share them in bonds. The result is a molecule. Example: C + O makes CO2.
Watch out: ammonium salts such as NH4Cl are made only of nonmetals, but they are ionic because NH4+ is a polyatomic ion. Hydrogen is a nonmetal.
Every element pulls on electrons with a different strength, and that pull comes from where the element sits on the periodic table (chemists call it electronegativity). Think of it like people and where they come from.
A metal on the far left barely holds its outer electrons and a nonmetal on the far right pulls hard. Their backgrounds are so different that one side takes the electron completely. That makes ions, and the bond is just the electrostatic attraction between opposite charges, with no sharing.
Two nonmetals pull on electrons about equally, like two people from similar backgrounds. Neither can take from the other, so they share. That is a covalent bond. If one pulls a bit harder, the sharing is unequal (a polar covalent bond).
Is transferred weaker than shared? Not really, and this is where the analogy stops. An ionic crystal holds billions of ions with attractions in every direction, so salt melts at 801 °C. A covalent bond inside a molecule is strong too, but separate molecules only weakly attract each other, so many covalent substances (water, sugar, wax) melt easily. Both kinds of bond are strong in their own way. What differs is how the electrons are handled: transferred or shared.
Metals in group 1, 2 and 3 form 1+, 2+ and 3+ ions. Nonmetals in groups 15, 16 and 17 form 3−, 2− and 1− ions. Transition metals such as iron and copper can form more than one charge, so the name carries a Roman numeral: iron(III) is Fe3+.
The compound has no overall charge, so the positives must cancel the negatives. Crisscross: each ion's charge number becomes the other's subscript, then reduce. Mg2+ with Cl− gives MgCl2. Al3+ with O2− gives Al2O3.
In the Ionic compounds tab you can build them from ions, or start from atoms and strip or add electrons until every vacancy is paired with an added electron in an ionic bond.
Add up the valence electrons of every atom, draw the bonds, then give each atom a full set of electrons (an octet; hydrogen needs 2). Next count electron groups around the central atom. Bonds and lone pairs repel, so they spread as far apart as possible (VSEPR). Lone pairs change the molecular shape, but not the electron geometry.
| Type | Electron geometry | Molecular geometry | Angle |
|---|---|---|---|
| AX2 | linear | linear | 180° |
| AX3 | trigonal planar | trigonal planar | 120° |
| AX2E | trigonal planar | bent | < 120° |
| AX4 | tetrahedral | tetrahedral | 109.5° |
| AX3E | tetrahedral | trigonal pyramidal | < 109.5° |
| AX2E2 | tetrahedral | bent | < 109.5° |
| AX5 | trigonal bipyramidal | trigonal bipyramidal | 90° and 120° |
| AX6 | octahedral | octahedral | 90° |
A is the central atom, X is an atom bonded to it, E is a lone pair.
Cation first, then anion. Metals keep their element name. A single-atom nonmetal ion ends in -ide (chloride, oxide). Polyatomic ions keep their own names (nitrate, sulfate). Use a Roman numeral for metals with more than one charge. Never use prefixes. CaCl2 is calcium chloride.
Greek prefixes count the atoms: mono 1, di 2, tri 3, tetra 4, penta 5, hexa 6, hepta 7, octa 8, nona 9, deca 10. Leave off mono on the first element. The second element ends in -ide. CO is carbon monoxide, N2O4 is dinitrogen tetroxide.
Acids dissolved in water are named from their anion. The number of H is the charge of the anion, so the acid has no overall charge: sulfate (SO42−) needs 2 H and makes H2SO4.
| Anion ends in | Acid name | Example |
|---|---|---|
| -ide | hydro- + stem + -ic acid | HCl, hydrochloric acid |
| -ate | stem + -ic acid | HNO3, nitric acid |
| -ite | stem + -ous acid | HNO2, nitrous acid |
Chlorine's family adds per- (one more oxygen than -ate) and hypo- (one less oxygen than -ite): HClO4 is perchloric acid, HClO is hypochlorous acid.
Now use the ideas, not your memory. Work out charges from the periodic table, take on mixed challenges, and run experiments in the lab.
Goal: the total positive charge must equal the total negative charge, so the compound has no net charge.
Crisscross: write each ion's charge number, then swap them to become the other ion's subscript. Mg2+ and Cl− give Mg1Cl2, which is written MgCl2 (a subscript of 1 is never written).
Reduce: Ca2+ and O2− crisscross to Ca2O2, which reduces to CaO. A formula shows the simplest whole-number ratio.
Polyatomic ions: keep the ion together and use parentheses when you need more than one: Ca2+ and NO3− make Ca(NO3)2.
Tip: you can also click a bond line on the bench to break it, click one ion and then another to bond them without dragging, or select an ion and press Delete to remove it.
Step 1. Add up the valence electrons from every atom in the molecule. This is the total number of electrons your Lewis structure has to use.
Step 2. The atoms are already joined by single bonds. Click a bond to make it double or triple, and drag lone pairs onto atoms until all 0 electrons are placed and the atoms have full shells.
Step 3. Count the electron domains on the central atom (a double or triple bond is one domain), then drag a choice into each slot.
1. Find the anion from the name. hydro-…-ic comes from an -ide anion, -ic comes from -ate, -ous comes from -ite.
2. Add H+ until the charges cancel. The number of H+ equals the anion's charge. Sulfate (2−) needs 2 H+, phosphate (3−) needs 3.
Tip: click a hydrogen ion on the bench to remove it. Click the anion to clear the bench.
Choose two elements. The gap in how hard each one pulls on electrons decides whether they share them, trade them, or pool them.
Electronegativity values are on the Pauling scale. The table runs from hydrogen to xenon. The cut-offs of 0.4 and 1.7 are classroom rules of thumb; real bonds sit on a continuum.