The valence electrons are the electrons in an atom’s outer shell; they are the ones that take part in bonding. Pick an element, predict its number of valence electrons from its group, then press the four sides of the symbol to place one dot per valence electron (one dot on each side first, then pairs). Mission: 5 elements right (first prediction and dots) from at least 2 different groups.
Each press adds a dot to that side (a side holds up to 2; a third press clears it).
Rule for the main groups: groups 1 and 2 have 1 and 2 valence electrons; groups 13–18 have the group number minus 10 (3 to 8). Helium is the exception: 2. Elements in the same group have the same number of valence electrons, which is why they behave alike. See the electron shells in the Bohr atom model and the full periodic table.
A metal atom and a nonmetal atom meet. Predict the charge each ion will have, then move the metal’s outer electrons to the nonmetal one at a time and watch the charges appear. An ion is an atom that has lost electrons (cation, +) or gained electrons (anion, −). Mission: complete 4 pairs.
Press each yellow e⁻ button under the picture to move that electron to the nonmetal.
| Atoms | Your prediction | Ions formed | Formula |
|---|
Yellow = the metal’s valence electrons that move; blue = electrons that stay where they are (inner shells and the nonmetal’s own outer electrons); dashed circles = empty places that get filled. Real ions are not little shells with dots: this is a model. For the 3D crystal and the full energy story (Born–Haber cycle), see ionic bonding in 3D.
An ionic compound has no overall charge. Pick a positive ion and a negative ion, add tiles until the charges cancel, and check. The formula must also use the lowest whole-number ratio. Then pick the right name. Mission: 5 different formulas, including one with a 1 : 2 (or 2 : 1) ratio and one with a 2 : 3 (or 3 : 2) ratio.
Shortcut (criss-cross): the size of each ion’s charge becomes the other ion’s subscript. Al³⁺ and O²⁻ give Al₂O₃. Always reduce: Ca²⁺ and O²⁻ give Ca₂O₂ → CaO.
Salt is not made of NaCl molecules. It is an extended structure: a crystal lattice of Na⁺ and Cl⁻ ions repeating in every direction. Predict how many Cl⁻ ions touch one Na⁺ ion in this flat layer, then press an ion to test. Then answer 3 myth-buster questions. Mission: prediction tested and all 3 questions right on the first try.
Red = oxygen end of a water molecule (slightly negative, δ−), white = hydrogen ends (slightly positive, δ+).
Ion sizes: Na⁺ (102 pm) is much smaller than Cl⁻ (181 pm). In 3D every Na⁺ has 6 Cl⁻ neighbors and every Cl⁻ has 6 Na⁺ neighbors. Rotate the real 3D lattice on the ionic bonding page.
Nonmetal atoms share pairs of electrons: each shared pair is a covalent bond (a line). Pairs that are not shared are lone pairs (two dots). Set the number of shared pairs on each link and the lone pairs on each atom. A correct structure uses exactly the valence electrons available and gives every atom 8 electrons around it (an octet; hydrogen: 2, a duet). Mission: build 6 molecules, including CO₂ or O₂ (double bonds) and N₂ (triple bond).
Green ring = the atom has its octet (or duet); orange = not yet. A Lewis structure shows which atoms are joined and where the electrons are, not the 3D shape: water is really bent and ammonia is a pyramid. Explore shapes on the VSEPR molecular shapes page.
The type of bonding decides how a substance behaves. Choose a substance and a test, predict the result, then run it. When all 4 tests are done for a substance, classify its bonding in the data table. Mission: test 4 substances and classify each one right on the first choice.
| Substance | Melts by 300 °C? | Conducts as a solid? | Dissolves in water? | Water mixture (µS/cm) | Bond type |
|---|
Values: handbook melting points (NaCl 801 °C, KCl 770 °C, copper 1085 °C, quartz about 1710 °C; paraffin wax 46–68 °C; sucrose melts and decomposes near 186 °C). Conductivity at about 25 °C: distilled water about 2 µS/cm, 0.1 M sugar water about 3, 0.1 M NaCl about 10,700, 0.1 M KCl about 12,900. Tap water is usually 50–800, so use distilled water in a real test. Real-lab safety: goggles on; hot wax and melted sugar cause bad burns, so heating them is a teacher demonstration.
Eight mystery substances, eight data cards. Use the evidence to decide the type of bonding. Select a card, then press Place card here in a group. Mission: at least 7 of 8 right when you check. The names are revealed when you pass.
10 questions from a bank of 20. Score 7 or more to complete the mission. Each new round picks new questions in a new order.
Everything around you is made of atoms, and almost none of them are alone. Salt, water, sugar, sand and copper are all atoms or ions held together. The “glue” is the attraction between positive and negative charges, like tiny magnets: the positive nucleus of one atom attracts the negative electrons of another. Some substances form molecules, small separate groups of atoms (water, H₂O). Others form extended structures that repeat in every direction, such as a salt crystal, a diamond or a piece of copper.
Only the electrons in the outer shell take part in bonding. The group number tells you how many: oxygen in group 16 has 6. Noble gases (group 18) have a full outer shell of 8 and hardly ever react. When atoms bond, they very often end up with the same electron arrangement as a noble gas: 8 outer electrons (the octet rule; hydrogen ends up with 2). This is a pattern, not a wish: atoms do not “want” anything. Arrangements like a noble gas simply have low energy, and the bonded atoms have less energy than the separate atoms.
Sodium has 1 valence electron; chlorine has 7. When they react, sodium’s electron moves to chlorine. Sodium becomes Na⁺ (2, 8, like neon) and chlorine becomes Cl⁻ (2, 8, 8, like argon). Moving the electron actually costs a little energy, but a lot of energy is released when the oppositely charged ions attract and pack into a crystal lattice. A crystal of salt contains no NaCl molecules: each Na⁺ is surrounded by 6 Cl⁻. The formula NaCl only gives the ratio. Because the charges must cancel, Mg²⁺ needs two Cl⁻ (MgCl₂) and two Al³⁺ balance three O²⁻ (Al₂O₃).
Two nonmetal atoms both attract electrons strongly, so neither can take an electron from the other. Instead they share. In H₂O the oxygen shares one pair with each hydrogen: O ends up with 8 electrons around it (2 bonds + 2 lone pairs) and each H with 2. Atoms can share 2 pairs (a double bond, O=C=O) or 3 pairs (a triple bond, N≡N). In a Lewis structure, each line is a shared pair and each pair of dots is a lone pair. If the two atoms pull unequally (different electronegativity), the sharing is unequal and the bond is polar, like a tug-of-war that one side is winning.
Melting separates particles. In ionic crystals every ion is held by strong attractions in all directions, so NaCl melts at 801 °C. In wax or sugar only weak forces act between the molecules, so wax melts near 60 °C: the covalent bonds inside the molecules do not break. A covalent network such as quartz (SiO₂) or diamond is held entirely by covalent bonds, so it melts above 1700 °C. Electric current needs moving charged particles: ions in melted or dissolved salt, or free electrons in metals and in the layers of graphite. Dry salt, sugar and wax have none, so they do not conduct.
In Mission 1, compare O and S (both in group 16): same number of dots. In Mission 2, compare Na + Cl with Mg + O: why does MgO need two electrons? In Mission 5, build CO₂ with single bonds first and read the octet badges. In the Property Lab, test sugar and salt in water: both dissolve, but only one conducts.
| Type | Between | What the electrons do | Particles in the solid | Typical properties | Examples |
|---|---|---|---|---|---|
| Ionic | metal + nonmetal | transferred from metal to nonmetal | ions in a crystal lattice | high melting point, brittle, conducts only when melted or dissolved | NaCl, KCl, MgO, CaCl₂ |
| Covalent molecular | nonmetals | shared in pairs | separate molecules | low melting point, does not conduct | H₂O, CO₂, sugar, wax |
| Covalent network | nonmetals and metalloids (C; Si with O) | shared in a giant network | atoms all joined by covalent bonds | very high melting point, hard, insoluble, does not conduct (graphite is the exception) | diamond, quartz SiO₂, graphite |
| Metallic | metal atoms | shared by all atoms (“sea of electrons”) | metal ions in a sea of electrons | conducts as a solid, shiny, can be bent and hammered | copper, iron, aluminum |
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| Valence electrons | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 (He: 2) |
| Usual ion | +1 (Na⁺) | +2 (Mg²⁺) | +3 (Al³⁺) | shares (C) | −3 (N³⁻) | −2 (O²⁻) | −1 (Cl⁻) | none |
| Usual covalent bonds | H: 1 | — | B: 3 | 4 | 3 | 2 | 1 | 0 |
Naming ionic compounds: metal name + nonmetal root + -ide (sodium chloride, magnesium oxide, aluminum nitride); no prefixes. Covalent compounds of two nonmetals use prefixes: carbon dioxide, dinitrogen tetroxide. Related pages: Bohr atom, periodic table, periodic trends, ionic bonding 3D, VSEPR shapes.
Electronegativity (EN) measures how strongly an atom in a bond pulls on the shared electrons. Pauling values: Na 0.93, Mg 1.31, H 2.20, C 2.55, N 3.04, Cl 3.16, O 3.44, F 3.98. The larger the difference ΔEN, the more unequal the sharing. Bonds form a continuum: roughly, ΔEN < 0.4 is nonpolar covalent, 0.4–1.7 is polar covalent and above 1.7 is mostly ionic. These cutoffs are guidelines, not sharp lines. Choose a bond:
Partial charges are written δ+ and δ−. In H–Cl the chlorine end is δ−; in O–H the oxygen end is δ−. Polar bonds plus a bent shape make water a polar molecule, which is why it dissolves salt and sugar but not wax. EN rises across a period and falls down a group: see periodic trends.
The attraction between two ions follows Coulomb’s law: F ∝ q₁q₂ / r², and the energy released when they come together scales with q₁q₂ / r (q = ion charges, r = distance between the ion centers). Bigger charges and smaller ions give a stronger lattice. Lattice energy is the energy needed to separate one mole of an ionic solid into gaseous ions:
| Compound | Ion charges | r(cation) + r(anion) | Lattice energy (kJ/mol) | Melting point |
|---|---|---|---|---|
| LiF | +1, −1 | 76 + 133 = 209 pm | 1037 | 845 °C |
| NaCl | +1, −1 | 102 + 181 = 283 pm | 787 | 801 °C |
| KCl | +1, −1 | 138 + 181 = 319 pm | 715 | 770 °C |
| MgO | +2, −2 | 72 + 140 = 212 pm | ≈ 3800 | 2852 °C |
MgO has about 4 times the charge product of NaCl and smaller ions, so its lattice energy is almost 5 times larger. That is why MgO is used to line furnaces. The Born–Haber cycle that builds these numbers step by step is on the ionic bonding page.
Formal charge = (valence electrons) − (lone-pair electrons) − (number of bonds). In O=C=O every atom has formal charge 0. The structure O≡C–O also gives every atom an octet, but its formal charges are +1 on the triple-bonded O and −1 on the single-bonded O. The structure with formal charges closest to zero is the best one, so the Lewis builder only accepts O=C=O. (Try the other one: it tells you why.)
In a metal, each atom releases its valence electrons into a sea of electrons shared by the whole piece. The electrons move freely (electrical and thermal conduction, shine), and layers of ions can slide without breaking the bonding (metals bend instead of shattering). Graphite is a covalent network in which each carbon bonds to only 3 others in flat layers; the fourth valence electron is delocalized along the layer, so graphite conducts as a solid, while diamond (4 bonds per carbon) does not.
Hydrogen and helium follow a duet. Boron in BF₃ has only 6 electrons around it; atoms from period 3 can have more than 8 (PCl₅, SF₆). O₂ has two unpaired electrons (it is attracted to a magnet), which a simple Lewis structure cannot show. For molecular shapes from Lewis structures, use VSEPR; to revisit electron shells, see the Bohr atom and the periodic table.
Standards: NGSS MS-PS1-1 (models of molecules vs extended structures: Mission 4), HS-PS1-1 (outer-shell electrons and the periodic table: Missions 1, 2, 5), HS-PS1-2 (ion formation and formulas: Missions 2, 3), HS-PS1-3 (bulk properties and the strength of electrical forces: Missions 6, 7). In grades 7–8, Missions 4, 6 and 7 meet MS-PS1-1 (models of molecules and extended structures; note that its assessment boundary also excludes valence electrons and the ionic nature of subunits, so the ion talk in Mission 4 is enrichment there). Missions 1–3 and 5 go beyond the middle-school assessment boundary: enrichment for grade 8, core for HS-PS1-1 and HS-PS1-2. Objective: students predict ion charges and covalent bonds from valence electrons, write ionic formulas with zero total charge, draw Lewis structures with octets, and use melting-point and conductivity data to classify substances by bond type.
Misconceptions to target: NaCl molecules; atoms “want” 8 electrons; making bonds needs energy; melting sugar or wax breaks covalent bonds; salt water conducts because electrons flow through the water. Missions are saved on the student’s device; the certificate prints as proof of completion.
A chemical bond is an electrical attraction that holds atoms or ions together. It forms because the positive nuclei and the negative electrons of neighboring atoms attract. The three main types are ionic bonds (between ions), covalent bonds (shared electron pairs) and metallic bonds (metal ions in a sea of electrons).
Key takeaway: a bond is an electrical attraction between atoms or ions.In an ionic bond, electrons are transferred from a metal atom to a nonmetal atom, and the resulting positive and negative ions attract each other in a crystal lattice (table salt, NaCl). In a covalent bond, two nonmetal atoms share one or more pairs of electrons to form molecules (water, H₂O) or networks (diamond). Ionic compounds usually have high melting points and conduct when dissolved; covalent molecular substances usually melt easily and do not conduct.
Key takeaway: ionic = transfer and attraction of ions; covalent = sharing electron pairs.For main-group elements, use the group number of the periodic table. Groups 1 and 2 have 1 and 2 valence electrons. For groups 13 to 18, subtract 10 from the group number: carbon (group 14) has 4, oxygen (group 16) has 6, chlorine (group 17) has 7. Helium is the exception with 2. Elements in the same group have the same number of valence electrons.
Key takeaway: groups 1–2: the group number; groups 13–18: group number minus 10.Add up the valence electrons of all the atoms (H₂O: 1 + 6 + 1 = 8). Place the atoms, join neighboring atoms with single bonds (each line is 2 electrons), then add lone pairs until every atom has 8 electrons around it (hydrogen: 2). If you run out of electrons before every atom has an octet, turn lone pairs into double or triple bonds. Finally, check that the number of electrons drawn equals the number you started with.
Key takeaway: count valence electrons, connect with bonds, add lone pairs, use multiple bonds if needed.The octet rule is the pattern that atoms in compounds often end up with 8 electrons in their outer shell, the same arrangement as a noble gas. Atoms do not want anything; those arrangements are simply low in energy, so the bonded atoms or ions have less energy than the separate atoms. Hydrogen follows a duet (2 electrons), and some molecules, such as BF₃ or SF₆, do not follow the octet rule at all.
Key takeaway: the octet is a stable, low-energy pattern, not a desire.An electric current needs moving charged particles. When salt dissolves, the Na⁺ and Cl⁻ ions separate and can move through the water, so a 0.1 M salt solution conducts about 5,000 times better than distilled water. Sugar dissolves as whole, neutral molecules, so sugar water conducts about as poorly as distilled water. Dry salt does not conduct either, because its ions are locked in the crystal.
Key takeaway: dissolved ions move and carry charge; sugar molecules are neutral.Write the positive ion first and the negative ion second, and choose the numbers of ions so that the total charge is zero. Mg²⁺ and Cl⁻ need one Mg²⁺ for every two Cl⁻: MgCl₂. Al³⁺ and O²⁻ need two Al³⁺ (+6) for three O²⁻ (−6): Al₂O₃. Always use the lowest whole-number ratio, so calcium oxide is CaO, not Ca₂O₂.
Key takeaway: balance the charges to zero, then reduce to the lowest ratio.In metallic bonding, metal atoms release their valence electrons into a "sea" of electrons that is shared by the whole piece of metal. The positive metal ions are held together by their attraction to this sea. Because the electrons can move freely, metals conduct electricity and heat as solids, and because the bonding does not depend on fixed partners, metals can be bent and hammered.
Key takeaway: metal ions in a sea of shared, free-moving electrons.A polar covalent bond is a bond in which two atoms share electrons unequally because one atom has a higher electronegativity. The shared electrons spend more time near that atom, giving it a partial negative charge (δ−) and the other atom a partial positive charge (δ+). In H–Cl the difference in electronegativity is 0.96, so the bond is polar; in H–H it is 0, so the bond is nonpolar.
Key takeaway: unequal sharing creates partial charges δ+ and δ−.In an ionic crystal, every ion is attracted by all its oppositely charged neighbors (6 in table salt), so a lot of energy is needed to pull the ions apart. Table salt melts at 801 °C and magnesium oxide, with 2+ and 2− ions, at 2852 °C. Molecular substances such as wax melt at low temperatures because only the weak forces between molecules have to be overcome.
Key takeaway: strong attractions in every direction must be overcome.