1. Why atoms bond
Atoms bond to lower their potential energy. Almost every atom does this by ending up with a full outer (valence) shell, which for most elements means eight valence electrons. This is the octet rule. Hydrogen and helium are the exceptions: they aim for two electrons (a full first shell), not eight.
The core principle
Bonded atoms have less energy than free atoms. Energy is released when a bond forms and absorbed when a bond breaks. This single fact drives every chemical reaction in the universe.
Noble gases (Group 18) already have full valence shells, which is why they almost never form compounds. Every other element is trying to look like a noble gas. The three ways to get there are: give electrons away, take them, or share them. Those three strategies are the three bond types.
2. Electronegativity decides the bond type
Electronegativity is how strongly an atom pulls bonding electrons toward itself. Values run roughly from 0.7 (francium) to 4.0 (fluorine). Find them on Reference Table S. The difference in electronegativity between two bonded atoms predicts the bond type.
| Electronegativity difference | Bond type | What happens |
|---|---|---|
| 0.0 to ~0.4 | Nonpolar covalent | Electrons shared equally |
| ~0.4 to ~1.7 | Polar covalent | Electrons shared unequally |
| ≥ ~1.7 | Ionic | Electrons transferred |
Regents rule of thumb
The cutoff numbers above are guidelines, not gospel. The Regents tests this conceptually: big difference (metal + nonmetal) → ionic; small difference (two nonmetals) → covalent; no difference (same element, like O₂) → pure nonpolar covalent.
3. Ionic bonds
An ionic bond is the electrostatic attraction between a positive ion (cation) and a negative ion (anion). It forms by transfer of electrons, typically from a metal to a nonmetal.
Example: Sodium chloride (NaCl)
Sodium has one valence electron it wants to lose. Chlorine has seven valence electrons and wants one more. Sodium gives its electron to chlorine. Now Na is Na⁺ (electron config of neon) and Cl is Cl⁻ (electron config of argon). Opposite charges attract. That attraction is the ionic bond.
Properties of ionic compounds
Crystal lattice
Ionic compounds form rigid 3D networks of alternating + and − ions. There are no individual "molecules" of NaCl, just one giant lattice.
High melting points
The lattice is held together by strong electrostatic forces in every direction. Lots of energy is needed to break it apart.
Hard but brittle
Shift a layer and like charges line up. They repel and the crystal cracks.
Conduct when molten or dissolved
In the solid lattice, ions can't move. Melt the salt or dissolve it in water and the ions move freely, carrying current.
Regents trap
Solid NaCl does not conduct electricity. Aqueous NaCl(aq) does. Melted (molten) NaCl(ℓ) also does. The exam loves this distinction.
4. Covalent bonds
A covalent bond is a shared pair of electrons between two atoms, usually two nonmetals. Each atom contributes one electron and both atoms "feel" the shared pair as part of their valence shell.
Single, double, and triple bonds
| Bond | Electron pairs shared | Example | Strength | Length |
|---|---|---|---|---|
| Single | 1 (2 electrons) | H–H, Cl–Cl | Weakest | Longest |
| Double | 2 (4 electrons) | O=O, CO₂ | Stronger | Shorter |
| Triple | 3 (6 electrons) | N≡N | Strongest | Shortest |
More shared pairs = more electron density between the nuclei = stronger pull = shorter, stronger bond. That is why N₂ (triple bond) is so unreactive: breaking it is a huge energy cost.
The diatomic seven
Memorize these
Seven elements exist naturally as diatomic molecules (two atoms covalently bonded to each other): H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. The mnemonic is "Have No Fear Of Ice Cold Beer". If you see one of these elements alone in an equation, it must be written as the diatomic.
Molecular vs network covalent
Molecular covalent
Discrete molecules (H₂O, CO₂, CH₄). Low melting points, often gases or liquids at room temperature. Don't conduct electricity.
Network covalent
Giant 3D networks of atoms (diamond, silicon dioxide). Extremely hard, very high melting points, don't conduct.
5. Metallic bonds
A metallic bond is the attraction between metal cations and a "sea" of mobile, delocalized valence electrons that flow freely through the lattice. Picture positive ions sitting in a pool of shared electrons.
Why metals are the way they are
- Conduct electricity (solid or molten): mobile electrons carry charge.
- Conduct heat: mobile electrons transfer kinetic energy fast.
- Malleable and ductile: shift layers and the electron sea reshapes around the new positions. No like charges line up to repel.
- Lustrous: the electron sea reflects light.
- High melting points (usually): the attraction between cations and the electron sea is strong.
Regents trap
Metals conduct electricity in both solid and liquid (molten) states. Ionic compounds only conduct when molten or dissolved. This is a clean distinguishing question.
6. Lewis dot structures
Lewis dot structures show valence electrons as dots around the element symbol. They help you predict how atoms bond.
Drawing rules for single atoms
- Write the element symbol.
- Count valence electrons (Group 1 = 1, Group 2 = 2, Group 13 = 3, Group 14 = 4, Group 15 = 5, Group 16 = 6, Group 17 = 7, Group 18 = 8).
- Place dots one at a time on the four sides of the symbol (top, right, bottom, left) before pairing any up.
Drawing rules for molecules
- Sum all valence electrons in the molecule.
- Put the least electronegative atom in the center (never H, never F).
- Draw single bonds from center to each surrounding atom (each bond = 2 electrons).
- Distribute remaining electrons to give every atom an octet (or H a duet).
- If you run short, convert lone pairs into double or triple bonds.
Worked example: CO₂
Total valence electrons: C contributes 4, each O contributes 6 → total = 16. Put C in the center: O–C–O. Give each O three lone pairs (6 each). That uses 4 + 12 = 16 electrons but leaves C with only 4 electrons (not an octet). Convert one lone pair on each O into a second bond between C and O. Result: O=C=O. Carbon has 8 electrons, each oxygen has 8 electrons. Done.
7. Bond polarity and molecular polarity
This is two separate questions, and you have to answer both.
Bond polarity
Look at the two atoms in the bond and compare electronegativities (Table S).
- Same atoms or very small difference (< 0.4): nonpolar covalent
- Different atoms, moderate difference (~0.4 to ~1.7): polar covalent
- Large difference (> ~1.7): ionic
Molecular polarity
A molecule with polar bonds may itself be polar or nonpolar. It depends on shape. If the polar bonds point symmetrically in opposite directions, the pulls cancel and the molecule is nonpolar. If they don't cancel, the molecule is polar.
| Molecule | Shape | Bond polarity | Molecule polarity | Why |
|---|---|---|---|---|
| H₂ | Linear | Nonpolar | Nonpolar | Same atoms |
| HCl | Linear | Polar | Polar | Only one bond, can't cancel |
| CO₂ | Linear | Polar | Nonpolar | Two polar bonds, symmetric, cancel |
| H₂O | Bent | Polar | Polar | Bent shape, dipoles don't cancel |
| NH₃ | Pyramidal | Polar | Polar | Asymmetric, dipoles don't cancel |
| CH₄ | Tetrahedral | Slightly polar | Nonpolar | Symmetric, four pulls cancel |
The water test
Water (H₂O) is bent because the oxygen has two lone pairs that push the hydrogens together. Both O–H bonds are polar and they don't point in opposite directions, so the dipoles add instead of canceling. That bent shape and resulting polarity is responsible for most of water's strange and life-supporting properties.
8. Intermolecular forces (IMFs)
Intermolecular forces are attractions between molecules, not within them. They determine boiling point, melting point, surface tension, viscosity, and solubility. They are always weaker than the bonds inside the molecule.
Critical distinction
Boiling water means breaking the intermolecular forces (the hydrogen bonds between H₂O molecules). It does not break the H–O bonds inside each water molecule. The steam coming off your kettle is still H₂O.
The three types, weakest to strongest
| IMF | Found in | Strength | Cause |
|---|---|---|---|
| London dispersion forces | All molecules (only force in nonpolar molecules) | Weakest | Temporary, random electron movement creates tiny instantaneous dipoles |
| Dipole-dipole | Polar molecules | Moderate | Permanent positive end of one molecule attracts permanent negative end of another |
| Hydrogen bonds | Molecules with H bonded directly to N, O, or F | Strongest IMF | Especially strong dipole-dipole where H is very exposed |
Mnemonic: "FON" the hydrogen bonder
Hydrogen bonding only happens when H is directly bonded to F, O, or N. Remember the trio "FON." H–F, H–O (in water and alcohols), H–N (in ammonia and amines) all do hydrogen bonding. H–C does not.
Why this matters for boiling points
Stronger IMFs mean more energy needed to pull molecules apart, which means higher boiling point. Compare three molecules of similar mass:
| Molecule | Molar mass | IMFs present | Boiling point |
|---|---|---|---|
| CH₄ (methane) | 16 g/mol | London only | −162 °C |
| NH₃ (ammonia) | 17 g/mol | London + dipole + H-bonding | −33 °C |
| H₂O (water) | 18 g/mol | London + dipole + extensive H-bonding | 100 °C |
Same approximate mass, wildly different boiling points. That is IMFs talking.
The dissolving rule: like dissolves like
Polar substances dissolve in polar solvents. Nonpolar substances dissolve in nonpolar solvents. Salt (ionic, very polar) dissolves in water (polar). Oil (nonpolar) does not. Oil dissolves in hexane (nonpolar).
9. Naming compounds
Ionic compounds (metal + nonmetal)
- Write the metal name (cation) first, unchanged: sodium, calcium, aluminum.
- Write the nonmetal name (anion), changing the ending to -ide: chloride, oxide, sulfide.
- If the metal has more than one possible charge (most transition metals), put its charge in Roman numerals: iron(II), iron(III), copper(I), copper(II).
- Polyatomic ions keep their name unchanged. Use Reference Table E.
Examples
- NaCl → sodium chloride
- CaO → calcium oxide
- Fe₂O₃ → iron(III) oxide
- NaNO₃ → sodium nitrate (NO₃⁻ is a polyatomic ion from Table E)
- (NH₄)₂SO₄ → ammonium sulfate
Molecular (covalent) compounds (two nonmetals)
Use Greek prefixes to count atoms of each element. Drop the "mono" on the first element.
| Prefix | Number |
|---|---|
| mono- | 1 |
| di- | 2 |
| tri- | 3 |
| tetra- | 4 |
| penta- | 5 |
| hexa- | 6 |
Examples
- CO → carbon monoxide
- CO₂ → carbon dioxide
- N₂O₄ → dinitrogen tetroxide
- P₂O₅ → diphosphorus pentoxide
Key terms
| Octet rule | Atoms tend to bond to achieve 8 valence electrons (2 for H, He). |
| Electronegativity | Tendency of an atom to attract bonding electrons. Found on Table S. |
| Ionic bond | Electrostatic attraction between oppositely charged ions, formed by electron transfer. |
| Covalent bond | Bond formed by sharing of electron pairs between two atoms. |
| Metallic bond | Attraction between metal cations and a sea of mobile electrons. |
| Polar covalent bond | Covalent bond where electrons are shared unequally. |
| Nonpolar covalent bond | Covalent bond where electrons are shared equally. |
| Dipole | A molecule or bond with separated regions of positive and negative charge. |
| Intermolecular forces | Attractions between separate molecules. Determine phase, boiling point, solubility. |
| Hydrogen bond | Strong dipole-dipole attraction when H is directly bonded to F, O, or N. |
| Lewis dot structure | Diagram showing valence electrons as dots around element symbols. |
| Network solid | Solid in which atoms are covalently bonded in a continuous 3D lattice (diamond, SiO₂). |
Practice questions
Q1. Which type of bond is formed when electrons are transferred from one atom to another?
Answer: (3) ionic. Transfer of electrons creates a cation and an anion that attract each other electrostatically. That is the definition of an ionic bond. Covalent involves sharing, metallic involves a delocalized sea, and hydrogen bonding is an intermolecular force, not a bond between atoms.
Q2. Which substance has the strongest intermolecular forces at STP?
Answer: (4) H₂O. Water has extensive hydrogen bonding (H bonded directly to O), which is the strongest IMF. HCl has dipole-dipole but no H bonding (Cl is not F, O, or N). CO₂ is nonpolar overall (symmetric) so only London. CH₄ is nonpolar, only London.
Q3. Which property is characteristic of metallic bonding?
Answer: (3) mobile, delocalized valence electrons. The "sea of electrons" model is the defining feature of metallic bonding and explains conductivity, malleability, and luster. Ionic crystals are brittle (option 1). Metals are good conductors (not 2). Most metals have high, not low, melting points (not 4).
Q4. (Part B-2) A molecule of CO₂ contains polar covalent bonds, yet the molecule itself is nonpolar. Explain this observation in terms of molecular shape and dipole direction.
Sample full-credit response: Each C=O bond is polar because oxygen is significantly more electronegative than carbon, so the bonding electrons are pulled toward oxygen, creating a dipole in each bond. The CO₂ molecule has a linear shape (O=C=O). The two C=O dipoles point in exactly opposite directions and have equal magnitude, so they cancel each other out. The net dipole moment of the molecule is zero, making CO₂ a nonpolar molecule overall despite having polar bonds.
Q5. (Part C) Compare and contrast the properties of ionic compounds and metallic substances. In your response, address (a) the type of particles present, (b) electrical conductivity in solid and liquid states, and (c) what happens at the particle level when each substance is struck with a hammer.
Sample full-credit response:
(a) Particles: Ionic compounds contain positive metal cations and negative
nonmetal anions arranged in a rigid crystal lattice. Metallic substances contain positive
metal cations surrounded by a sea of mobile, delocalized valence electrons.
(b) Conductivity: Solid ionic compounds do not conduct electricity because
their ions are locked in place in the lattice. Molten or aqueous ionic compounds do conduct
because the ions are free to move and carry charge. Metallic substances conduct electricity
in both solid and liquid states because the delocalized valence electrons are always free
to move.
(c) Hitting with a hammer: When an ionic crystal is struck, the layers of
ions shift. Like charges suddenly line up next to each other and repel, splitting the
crystal. Ionic compounds are therefore brittle. When a metal is struck, the cation layers
slide past one another, but the electron sea simply flows around the new positions and
keeps holding the cations together. The metal deforms but does not break. Metals are
therefore malleable and ductile.