This unit explains how atoms hold together and how those connections decide a molecule's shape, polarity, and physical properties. By the end, you should be able to classify a bond, draw a Lewis structure, predict a molecule's shape and polarity, and use intermolecular forces to explain boiling points and solubility.
Read the overview first to see how the four topics connect, then work through each review note and try the checkpoint question before reading the answer.
Unit 4 connects the atom-level ideas from earlier units, like valence electrons and electronegativity trends, to the properties you can actually measure in lab. Each topic builds on the one before it, so a mistake in a Lewis structure carries forward into the wrong shape, the wrong polarity, and the wrong boiling point prediction.
Honors Chemistry Unit 4 covers chemical bonding. You learn the three main bond types (ionic, covalent, and metallic) and how electronegativity difference helps you tell polar covalent bonds from nonpolar ones. You draw Lewis structures, including multiple bonds, resonance, formal charge, and octet exceptions, then use VSEPR theory to predict shapes like linear, trigonal planar, tetrahedral, trigonal pyramidal, and bent. Finally, you decide whether molecules are polar and use intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole) to explain trends in boiling point, melting point, and solubility.
Electronegativity difference helps predict how electrons are distributed, but it does not determine bond type by itself. NaCl forms an ionic lattice of Na+ and Cl- ions, while atoms in covalent substances share electrons. Metals have delocalized valence electrons, which helps explain why they conduct electricity and are malleable.
VSEPR theory says bonding pairs and lone pairs around a central atom spread out as far as possible. Count the electron domains to get the electron arrangement, then look only at where the atoms sit to name the shape. That is why CH4, NH3, and H2O all have four domains but three different shapes.
A polar molecule has a net dipole, so its molecules attract through dipole-dipole forces, plus hydrogen bonds if H is bonded to N, O, or F. Every molecule has London dispersion forces, which grow with polarizability (how easily the electron cloud distorts, which rises with more electrons) and with more surface contact. Stronger attractions mean higher boiling points, and a substance dissolves well when solute-solvent attractions can compete with the attractions inside each pure substance.
Where electrons sit, in a bond, a lone pair, or a temporary shift in the electron cloud, affects bond type, shape, and polarity. These features, along with polarizability and molecular surface contact, affect how strongly particles attract. That reasoning helps explain why the normal boiling point of water is 100°C while that of methane is about -162°C at 1 atm, even though both are small molecules.
Atoms bond because the bonded arrangement has lower energy than separate atoms.
A Lewis structure shows valence electrons as bond lines (2 electrons each) and dots, usually in lone pairs.
A polar molecule has an uneven charge distribution, with a partially negative end (δ-) and a partially positive end (δ+).
Intermolecular forces are attractions between separate molecules.
Atoms bond because the bonded arrangement has lower energy than separate atoms. Most main-group atoms end up with eight valence electrons (the octet rule), while hydrogen needs only two.
Ionic bonds form when a metal transfers electrons to a nonmetal, creating cations and anions held together by electrostatic attraction. Ionic solids have high melting points and conduct electricity only when melted or dissolved, because only then can the ions move. Covalent bonds form when nonmetals share electron pairs. A single bond shares one pair, a double bond two, and a triple bond three. For bonds between the same two elements in similar environments, bond strength generally rises from single to double to triple, while bond length decreases. For example, the N≡N triple bond takes about 945 kJ/mol to break. Metallic bonds are metal cations in a sea of delocalized electrons, which is why metals conduct and are malleable.
Electronegativity difference () gives a quick guide. Many textbooks use under 0.4 for nonpolar covalent, 0.4 to 1.7 for polar covalent, and above 1.7 for ionic, though some use 0.5 and 2.0. Treat these as guidelines, not hard walls.
Worked example. Classify the bonds in NaCl, HCl, and Cl2 using Pauling values (Na 0.93, H 2.20, Cl 3.16). For NaCl, , which is above 1.7, so ionic. For HCl, , so polar covalent with the partial negative charge on Cl. For Cl2, , so nonpolar covalent.
Check yourself. Using Mg 1.31 and O 3.44, what type of bond forms in MgO, and would solid MgO conduct electricity? The difference is 3.44 - 1.31 = 2.13, so MgO is ionic. Solid MgO does not conduct because its ions are locked in the lattice, but molten MgO conducts because its ions can move.
A Lewis structure shows valence electrons as bond lines (2 electrons each) and dots, usually in lone pairs. To draw one, count the total valence electrons, adding one for each negative charge and subtracting one for each positive charge. Choose the likely central atom, which is usually the least electronegative atom. Hydrogen is never central, and halogens are usually terminal. Connect the atoms with single bonds, give each outer hydrogen 2 electrons and most other outer atoms 8 electrons, then place any remaining electrons on the central atom. If the central atom still lacks an octet, check whether forming double or triple bonds gives a better structure with more favorable formal charges. Do not force an octet for valid electron-deficient exceptions such as BF3.
Formal charge equals valence electrons minus nonbonding electrons minus half the bonding electrons. Aim for formal charges near zero when picking the best structure. Some species, like O3 and NO3-, need resonance structures. Boron in BF3 has only 6 electrons, and sulfur in SF6 holds 12.
VSEPR theory turns the Lewis structure into a 3D shape. Each lone pair or bond on the central atom counts as one domain, even a double or triple bond.
Worked example. Draw CO2 and find its shape. Valence electrons are . The skeleton O-C-O uses 4, leaving 12, which go on the oxygens as 3 lone pairs each. Carbon has only 4 electrons, so move one lone pair from each oxygen into a bond, giving O=C=O with all formal charges 0. Carbon has 2 domains and no lone pairs, so CO2 is linear at 180°.
Check yourself. What is the shape of NH3, and why is it not tetrahedral? Nitrogen has four electron domains, three bonds and one lone pair, so its electron arrangement is tetrahedral. Molecular shape describes only atom positions, so NH3 is trigonal pyramidal with bond angles of about 107°.
| Electron domains | Lone pairs on center | Molecular shape | Approximate bond angle | Example |
|---|---|---|---|---|
| 2 | 0 | Linear | 180° | CO2 |
| 3 | 0 | Trigonal planar | 120° | BF3 |
| 4 | 0 | Tetrahedral | 109.5° | CH4 |
| 4 | 1 | Trigonal pyramidal | about 107° | NH3 |
| 4 | 2 | Bent | about 104.5° | H2O |
| 5 | 0 | Trigonal bipyramidal | 90° and 120° | PCl5 |
| 6 | 0 | Octahedral | 90° | SF6 |
A polar molecule has an uneven charge distribution, with a partially negative end (δ-) and a partially positive end (δ+). Deciding polarity takes two checks, and you need both.
First, check the bonds. If bonded atoms differ in electronegativity, the bond has a dipole, drawn as an arrow with a plus sign at its tail pointing toward the more electronegative atom. Second, check the shape. If the bond dipoles are arranged symmetrically, they cancel like equal tugs in opposite directions, and the molecule is nonpolar. If the shape is lopsided, usually because of lone pairs on the central atom or different outer atoms, the dipoles add up to a net dipole moment and the molecule is polar.
That is why CO2 (linear) and CCl4 (tetrahedral with four identical outer atoms) are nonpolar even though their bonds are polar, while H2O (bent) and CHCl3 (one outer atom differs) are polar.
Worked example. Compare NH3 and BF3 (N 3.04, H 2.20, B 2.04, F 3.98). In NH3, each N-H bond has , so the bonds are polar. The lone pair makes NH3 trigonal pyramidal, so the dipoles all lean toward nitrogen and do not cancel. NH3 is polar. In BF3, each B-F bond has . Despite the rough cutoff chart, BF3 is a molecular compound with strongly polar covalent B-F bonds. Its trigonal planar shape places three identical dipoles 120° apart, so they cancel and BF3 is nonpolar.
Check yourself. CH4 and CH3Cl are both tetrahedral. Which one is polar? CH3Cl is polar. In CH4, the four C-H bond dipoles cancel because of the molecule's symmetry. In CH3Cl, replacing one H with Cl breaks that symmetry, so the bond dipoles do not cancel.
Intermolecular forces are attractions between separate molecules. This section also includes ion-dipole attractions between ions and polar molecules. Attractions between neutral molecules are generally weaker than covalent bonds within a molecule. When water boils, H2O molecules separate from one another, but their O-H bonds remain intact.
London dispersion forces exist in every substance and come from temporary dipoles as electron clouds shift. They grow with more electrons and more surface contact, so I2 is a solid at room temperature while F2 is a gas. Dipole-dipole forces act between polar molecules. Hydrogen bonding is an extra-strong dipole-dipole attraction that happens only when H is bonded directly to N, O, or F. Ion-dipole forces, like the pull between Na+ and water, help many salts dissolve.
For comparable molecular substances, stronger intermolecular forces generally mean higher boiling points and lower vapor pressures when the vapor pressures are compared at the same temperature. For solubility, "like dissolves like" is a useful guideline. Polar substances tend to dissolve in polar solvents, and nonpolar substances in nonpolar ones.
Worked example. Rank CH4, HCl, and H2O by boiling point. CH4 is nonpolar and has only dispersion forces. HCl is polar and adds dipole-dipole forces. H2O has hydrogen bonds because H is bonded to O. The prediction is CH4 < HCl < H2O, and their normal boiling points at 1 atm agree, at about -162°C, -85°C, and 100°C, respectively.
Size can outweigh polarity, though. Nonpolar octane (C8H18) boils at about 126°C, above water, because its many electrons give large dispersion forces.
Check yourself. Does CH3OCH3, dimethyl ether, form hydrogen bonds with other CH3OCH3 molecules? It does not. The molecule contains oxygen, but every hydrogen is bonded to carbon rather than oxygen. Pure dimethyl ether therefore has dipole-dipole and dispersion forces but no hydrogen bonding between its molecules.
| Force | Found between | Example | Relative strength |
|---|---|---|---|
| London dispersion | All molecules and atoms | CH4, I2, noble gases | It is often weak in small particles and grows as electron clouds become larger and have more surface contact. |
| Dipole-dipole | Polar molecules | HCl, CH3Cl | For similarly sized molecules, it is often stronger than dispersion alone. |
| Hydrogen bonding | Molecules with H bonded to N, O, or F | H2O, NH3, HF | For similarly sized neutral molecules, it is usually among the strongest intermolecular attractions. |
| Ion-dipole | An ion and a polar molecule | Na+ in water | It is often stronger than the listed molecule-molecule attractions, but its strength depends on ion charge, molecular dipole, and distance. |
| Term | Definition |
|---|---|
| Valence electrons | Electrons available to participate in bonding. For main-group elements, these are usually the electrons in the outermost occupied energy level. |
| Octet rule | The tendency of atoms to gain, lose, or share electrons until they have eight valence electrons, or two for hydrogen. |
| Electronegativity | How strongly an atom attracts shared electrons in a bond, with fluorine highest at about 4.0 on the Pauling scale. |
| Lewis structure | A diagram showing valence electrons as bond lines and dots for lone pairs or unpaired electrons. |
| VSEPR theory | A model predicting molecular shape by assuming electron domains around a central atom spread as far apart as possible. |
| Electron domain | Any lone pair or bond around a central atom, counted once regardless of whether the bond is single, double, or triple. |
| Polar molecule | A molecule with a net dipole moment because its bond dipoles do not cancel. |
| Intermolecular forces | Attractions between separate molecules. Related ion-dipole attractions occur between ions and polar molecules. These attractions help determine properties such as boiling point, melting point, and solubility. |
| Hydrogen bond | A strong intermolecular attraction involving H bonded to N, O, or F and a lone pair on a neighboring molecule. |
| Delocalized electrons | Electrons not tied to a single atom or bond, such as the shared electron sea in a metal. |
Bond polarity and molecular polarity are different questions. CO2 and CCl4 both have polar bonds, but their symmetrical shapes make the dipoles cancel. Always draw the shape before deciding, and look for lone pairs on the central atom or mismatched outer atoms that break the symmetry.
For a molecular substance such as water, a state change mainly overcomes attractions between molecules. Boiling water produces H2O vapor, not hydrogen and oxygen gas. Ionic, metallic, and network-covalent solids require different explanations because they do not consist of separate molecules.
Hydrogen bonding needs H attached directly to N, O, or F. CH4 has plenty of hydrogen but only dispersion forces, and CH3OCH3 has oxygen but its hydrogens sit on carbon. Check what each H is actually bonded to.
In VSEPR, any bond between two atoms counts as one domain, so CO2 has 2 domains, not 4. Lone pairs do count as domains, but the shape name describes only where the atoms are. Water has a tetrahedral electron arrangement but a bent molecular shape.