Three-dimensional molecular shape, bond polarity, molecular polarity, and the forces that act between molecules — the foundation for understanding physical properties from boiling points to the structure of DNA
When you draw a Lewis structure on paper, it looks flat. Dots and lines stretch left and right across the page. But molecules are not flat. They exist in three dimensions, and their shapes are just as real as the shapes of anything else you can hold in your hand.
The shape of a molecule is not just an interesting fact. Shape determines how molecules interact with each other, how they dissolve in water, whether they have high or low boiling points, and whether they can fit into a biological receptor the way a key fits a lock. The shape of water explains why ice floats. The shape of DNA's building blocks allows genetic information to be copied. The shape of a drug molecule determines whether it will work in the human body.
In this chapter, you will learn how to predict the three-dimensional shape of a molecule from its Lewis structure, how to determine whether a molecule is polar or nonpolar, and how molecular polarity and shape together determine the forces that act between molecules. These forces — called intermolecular forces — control physical properties such as melting point, boiling point, viscosity, and solubility. By the end of this chapter, you will see molecules not just as flat diagrams but as three-dimensional objects whose geometry shapes the physical world around you.
Water (H₂O) and carbon dioxide (CO₂) are both small molecules made of nonmetals. Yet water is a liquid at room temperature while CO₂ is a gas. Water is the universal solvent of life; CO₂ does not dissolve freely in your cells. A major reason for this difference is not what the molecules are made of — it is what shape they take in three dimensions. By the end of this chapter, you will be able to explain this difference precisely.
Consider two molecules: carbon dioxide (CO₂) and water (H₂O). Both are small, simple molecules. Both are compounds of nonmetals. Both appear constantly in Earth's atmosphere and in living things. Yet CO₂ is a gas at room temperature that does not dissolve well in biological fluids, while water is a liquid at room temperature and the universal solvent of life. A major reason for this difference is molecular shape.
Carbon dioxide is linear. Its two oxygen atoms sit on opposite ends of a straight line with carbon in the center. Water is bent. Its two hydrogen atoms are angled away from each other — not opposite, but at about 104.5°. This difference in shape produces dramatic differences in the properties of the two substances.
A bond angle is the angle formed between two bonds that share a central atom. In water, the bond angle is about 104.5°. In carbon dioxide, the bond angle is exactly 180°. In methane (CH₄), all four bond angles are 109.5°. These are not coincidences — they are predictable consequences of how electron pairs arrange themselves around a central atom.
The shape of a molecule is not random or arbitrary. It follows from the arrangement of electrons around the central atom. Understanding this connection — from Lewis structure to three-dimensional shape to physical property — is the central skill of this chapter.
Electron pairs in the valence shell of a central atom repel each other and arrange themselves to be as far apart as possible. This arrangement determines the shape of the molecule.
VSEPR stands for Valence Shell Electron Pair Repulsion. Pronounced "vesper."
An electron group is any region of electron density around a central atom — including single bonds, double bonds, triple bonds, and lone pairs. Multiple bonds (double and triple) count as one electron group because all their electrons occupy the same region of space between the bonded atoms.
To use VSEPR, count the total number of electron groups around the central atom. The arrangement that keeps them farthest apart depends on how many groups there are: two groups achieve 180° separation; three groups spread into a flat triangle at 120°; four groups spread into a three-dimensional tetrahedron at 109.5°.
| Electron Groups | Electron Arrangement | Ideal Angle | Example Central Atom |
|---|---|---|---|
| 2 | Linear | 180° | BeCl₂ (Be) |
| 3 | Trigonal planar | 120° | BF₃ (B) |
| 4 | Tetrahedral | 109.5° | CH₄ (C), NH₃ (N), H₂O (O) |
Not all electron pairs repel equally. Lone pairs are closer to the central atom than bonding pairs, because bonding pairs are pulled outward — they are shared between two atoms. Lone pairs belong entirely to one atom and spread out more in space. As a result, lone pairs exert stronger repulsive forces than bonding pairs and compress the angles between bonding pairs.
In methane (CH₄), four identical bonding pairs produce bond angles of exactly 109.5°. In ammonia (NH₃), one lone pair pushes the three bonding pairs closer together — angles compress to about 107°. In water (H₂O), two lone pairs compress the bond angle further to about 104.5°.
| Molecule | Central Atom | Bonding Pairs | Lone Pairs | Bond Angle |
|---|---|---|---|---|
| CH₄ (methane) | C | 4 | 0 | 109.5° |
| NH₃ (ammonia) | N | 3 | 1 | ~107° |
| H₂O (water) | O | 2 | 2 | ~104.5° |
Do not forget to count lone pairs when determining the total number of electron groups. Lone pairs are invisible in the final molecular shape, but they must be included in the electron group count because they push the bonding pairs into their actual positions.
For each molecule, draw the Lewis structure, count the electron groups on the central atom, and list the number of bonding pairs and lone pairs.
Electron geometry describes the arrangement of all electron groups around a central atom, including lone pairs. Molecular geometry describes the arrangement of atoms only — lone pairs are not atoms, so they do not appear in the final shape. When lone pairs are present, the electron geometry and molecular geometry are different.
| Electron Groups | Lone Pairs | Molecular Geometry | Bond Angle | Example |
|---|---|---|---|---|
| 2 | 0 | Linear | 180° | CO₂, BeCl₂ |
| 3 | 0 | Trigonal planar | 120° | BF₃, SO₃ |
| 3 | 1 | Bent | ~117° | SO₂ |
| 4 | 0 | Tetrahedral | 109.5° | CH₄, CCl₄ |
| 4 | 1 | Trigonal pyramidal | ~107° | NH₃, PH₃ |
| 4 | 2 | Bent | ~104.5° | H₂O, H₂S |
A molecule is linear when two bonding groups and no lone pairs surround the central atom. The electron pairs point in exactly opposite directions, producing a straight molecule with a 180° bond angle. Carbon dioxide (CO₂) is the classic example — two double bonds, no lone pairs on carbon, oxygen atoms on opposite sides. Hydrogen cyanide (HCN), with a triple bond to nitrogen and a single bond to hydrogen, is also linear.
When three electron groups surround a central atom with no lone pairs, the groups arrange into a flat triangle with 120° angles. All atoms lie in the same plane. Boron trifluoride (BF₃) is trigonal planar. Formaldehyde (H₂C=O) is also trigonal planar — carbon has a double bond to oxygen (one electron group) and two single bonds to hydrogen, giving three groups total. When one lone pair is present among three electron groups, the molecular shape becomes bent, with bond angles slightly less than 120°.
Four electron groups with no lone pairs point toward the corners of a tetrahedron — a three-dimensional shape with four triangular faces. Bond angles are all 109.5°. Methane (CH₄) is the classic tetrahedral molecule: four identical C–H bonds pointing in four different directions in space, not flat. Carbon tetrachloride (CCl₄) and silicon tetrafluoride (SiF₄) are also tetrahedral.
When four electron groups are present but one is a lone pair, the molecular geometry is trigonal pyramidal. The three bonding pairs point to three corners of a tetrahedron; the lone pair occupies the fourth corner but is not an atom, so it is invisible in the shape. The result looks like a triangular pyramid. Ammonia (NH₃) is the best-known example, with a bond angle of approximately 107°.
When four electron groups surround a central atom but two are lone pairs, only two atoms are bonded to the central atom. The molecule is bent (sometimes called V-shaped), resembling a boomerang. Water (H₂O) is the most important example, with a bond angle of approximately 104.5°. Two lone pairs compress the bond angle even more than one does.
Electronegativity is a measure of an atom's ability to attract shared electrons toward itself in a covalent bond. Electronegativity increases across a period (left to right) and decreases down a group. Fluorine is the most electronegative element. Oxygen, nitrogen, and chlorine are also highly electronegative. Carbon, hydrogen, and most metals have lower electronegativity values.
When two atoms with different electronegativities share electrons, the electrons are not shared equally. The more electronegative atom pulls the electrons closer to itself. This creates a polar covalent bond — one end is slightly negative, the other slightly positive. Chemists use the Greek letter delta (δ) to show this: the more electronegative atom is labeled δ−, the other is labeled δ+.
When two identical atoms share a bond — as in H₂, O₂, or Cl₂ — the electronegativity values are exactly equal. The electrons are shared perfectly. This is a nonpolar covalent bond. Bonds between atoms with very similar electronegativity (such as C–H) are also effectively nonpolar.
| Electronegativity Difference | Bond Type | Example |
|---|---|---|
| ≈ 0 (identical atoms) | Nonpolar covalent | H–H, Cl–Cl, C–C |
| 0 – ~0.4 | Slightly polar covalent (often treated as nonpolar) | C–H |
| ~0.4 – ~1.7 | Polar covalent | O–H, N–H, C–O, H–Cl |
| > ~1.7 | Ionic | Na–Cl, Mg–O |
O–H (very polar), N–H (polar), H–F (very polar), C–O (moderately polar), C–N (moderately polar), H–Cl (polar), C–F (very polar). These patterns appear throughout organic chemistry and biochemistry.
The boundary between polar covalent and ionic is not a sharp line — it is a gradual transition. Electronegativity differences give approximate guidance, not absolute categories. A difference of 1.7 is a rough threshold, not a chemical law.
For each bond, identify which atom is more electronegative, label atoms with δ+ and δ−, and classify the bond as nonpolar covalent, polar covalent, or ionic.
A molecule can have polar bonds but still be nonpolar overall. This happens when the individual bond polarities cancel because of molecular geometry. Think of it this way: if two people pull equally on opposite ends of a rope, there is no net motion in either direction — the forces cancel. A similar cancellation happens in symmetric molecules.
To determine whether a molecule is polar, you must consider two things: (1) whether the individual bonds are polar, and (2) whether the bond polarities cancel or add together based on the shape of the molecule.
Step 1: Draw the Lewis structure.
Step 2: Determine the molecular geometry.
Step 3: Identify all polar bonds.
Step 4: Determine whether the bond polarities cancel (symmetric) or add together (asymmetric).
Step 5: If polarities cancel → nonpolar molecule. If they do not cancel → polar molecule (has a dipole moment).
Carbon dioxide (CO₂) has very polar C=O bonds, but is linear — the two bond polarities point in exactly opposite directions and cancel perfectly. CO₂ is nonpolar despite having polar bonds. Carbon tetrachloride (CCl₄) has four polar C–Cl bonds, but its perfect tetrahedral symmetry causes them to cancel. Boron trifluoride (BF₃) has three polar B–F bonds in a symmetric trigonal planar geometry — they cancel. All three are nonpolar molecules.
Water (H₂O) is bent, not linear. The two O–H bond polarities point at an angle — they do not cancel. The oxygen end has a net partial negative charge (δ−) and the hydrogen end has a net partial positive charge (δ+). Water is a polar molecule with a significant dipole moment. Ammonia (NH₃) is also polar — three N–H bond polarities all point toward nitrogen, adding together, and the lone pair on nitrogen adds to the asymmetry. Hydrogen chloride (HCl), a simple diatomic molecule, is polar because chlorine is more electronegative than hydrogen.
Determine whether each molecule is polar or nonpolar. Explain your reasoning by considering both bond polarity and molecular geometry.
In Chapter 10, you studied intramolecular forces — the covalent and ionic bonds that hold atoms together within a compound. Those forces are strong, and breaking them requires large amounts of energy and produces chemical change.
Intermolecular forces are different. They are the attractions between separate molecules. These forces are much weaker than covalent or ionic bonds, but they are enormously important because they determine the physical properties of substances: whether a substance is a solid, liquid, or gas at a given temperature; how high its boiling point is; how easily it dissolves in water; how viscous (thick) it is.
When a liquid boils, covalent bonds within the molecules do not break. The molecules simply gain enough kinetic energy to overcome the intermolecular attractions holding them together, and they fly off as gas-phase molecules. Boiling overcomes intermolecular forces — it does not break intramolecular bonds.
| Force Type | Present In | Relative Strength | Origin |
|---|---|---|---|
| London dispersion | All molecules | Weakest (increases with size) | Temporary dipoles from electron motion |
| Dipole-dipole | Polar molecules | Moderate | Attraction between permanent dipoles |
| Hydrogen bonding | Molecules with H bonded to F, O, or N | Strongest intermolecular force | Attraction between bare H and lone pair on F, O, or N |
A hydrogen atom must be covalently bonded directly to a highly electronegative atom — F, O, or N.
That hydrogen is then attracted to a lone pair on another F, O, or N in a neighboring molecule.
Hydrogen bonded to carbon, sulfur, or other elements does not form true hydrogen bonds — those atoms are not electronegative enough to strip enough electron density away from hydrogen.
Water is the most familiar example of hydrogen bonding. Each water molecule can form hydrogen bonds with up to four neighboring water molecules: two through the two hydrogen atoms (each H can interact with an O on a neighbor) and two through the two lone pairs on the oxygen (which can attract H atoms from neighbors). This extensive network of hydrogen bonds is responsible for water's remarkable properties.
Hydrogen bonds are significantly stronger than most other intermolecular forces — typically 5 to 30 kJ/mol — though still much weaker than covalent bonds (hundreds of kJ/mol). This difference is enough to dramatically raise boiling points. Compare water (H₂O, molar mass 18 g/mol, boiling point 100°C) to hydrogen sulfide (H₂S, molar mass 34 g/mol, boiling point −60°C). Despite being smaller and lighter, water's hydrogen-bonding network gives it a much higher boiling point.
Substances that form hydrogen bonds include water (H₂O), ammonia (NH₃), hydrogen fluoride (HF), alcohols (contain O–H groups), carboxylic acids (contain O–H groups), amines (contain N–H groups), DNA base pairs (N–H···N and N–H···O), and proteins (N–H···O interactions along the backbone). Hydrogen bonding is essential to life — the double helix structure of DNA is held together between strands by hydrogen bonds, and the three-dimensional shapes of proteins are stabilized partly by internal hydrogen bonds.
F, O, and N are so electronegative that they pull nearly all the electron density away from the bonded hydrogen, leaving a hydrogen nucleus that is almost bare — a proton with very little electron shielding. This nearly exposed proton is strongly attracted to lone pairs on nearby electronegative atoms, creating a stronger attraction than ordinary dipole-dipole forces.
London dispersion forces (also called van der Waals forces) are the weakest type of intermolecular attraction, but they are present in all molecules — polar and nonpolar alike. They arise because electrons are in constant motion. At any instant, the electron distribution in a molecule may be slightly uneven, creating a temporary dipole. This temporary dipole induces a dipole in a neighboring molecule, causing a fleeting attraction. Over time these average to zero, but at any moment they create measurable attraction.
London dispersion forces are stronger in molecules with more electrons. A molecule with many electrons has a larger, more polarizable electron cloud — its electrons can shift more easily to create temporary dipoles. This is why larger molecules generally have higher boiling points than smaller ones, even when both are nonpolar.
Among the halogens: F₂ is a gas, Cl₂ is a gas, Br₂ is a liquid, and I₂ is a solid at room temperature. This progression reflects increasing London dispersion forces as the number of electrons increases.
Polar molecules have permanent dipoles — one end is consistently partially negative and the other consistently partially positive. When polar molecules are near each other, they align so positive ends face neighboring negative ends. This is a dipole-dipole force. These forces are stronger than London dispersion forces between similarly sized molecules. Polar molecules tend to have higher boiling points than nonpolar molecules of similar size because they experience both dipole-dipole forces and London dispersion forces, while nonpolar molecules experience only London dispersion forces.
Solubility is closely connected to intermolecular forces. Polar solvents — especially water — dissolve polar solutes and ionic compounds well because they can engage in dipole-dipole interactions or hydrogen bonds with those solutes. Nonpolar solvents dissolve nonpolar solutes well through London dispersion forces.
Water and oil do not mix. Oil consists of nonpolar hydrocarbon molecules with only London dispersion forces. Water has strong hydrogen bonds. Water molecules are much more strongly attracted to each other than to the oil molecules, so they exclude the oil. Soap works because its molecules have a polar end (interacts with water) and a nonpolar end (interacts with grease), allowing it to bridge the two.
| Stronger Intermolecular Forces Produce… | Weaker Intermolecular Forces Produce… |
|---|---|
| Higher melting point | Lower melting point |
| Higher boiling point | Lower boiling point |
| Higher viscosity (thicker liquid) | Lower viscosity (thinner liquid) |
| Lower vapor pressure | Higher vapor pressure |
CH₄ (−161°C) < HCl (−85°C) < H₂O (+100°C)Identify the type(s) of intermolecular force present between molecules of each substance and rank where asked.
Everything you have studied in this chapter comes together to explain something fundamental: why substances have the physical properties they do, and how molecular structure determines function in the real world.
Water's high boiling point (100°C) is due to its extensive hydrogen bonding. When water freezes, the hydrogen bonds lock into a crystal lattice slightly less dense than liquid water — which is why ice floats. If ice sank, lakes would freeze from the bottom up, destroying aquatic ecosystems. Water's polarity also makes it an excellent solvent for ionic and polar substances, earning it the name "universal solvent of life."
The two strands of the DNA double helix are held together between complementary base pairs by hydrogen bonds. These bonds are strong enough to maintain the double helix at body temperature but weak enough to be unzipped when the cell needs to copy or read the genetic code. Proteins fold into specific three-dimensional shapes partly because nonpolar regions cluster toward the interior (away from water) while polar and charged regions face outward — a process driven entirely by intermolecular forces.
A drug molecule must fit precisely into a biological receptor. Chemists design drugs by carefully engineering their geometry, polarity, and hydrogen-bonding capacity — all concepts from this chapter. Synthetic polymers' physical properties (hard vs. soft, rigid vs. flexible) depend on intermolecular forces between chains: polar polymers like nylon have stronger chain-to-chain attractions and higher melting points than nonpolar polymers like polyethylene. Fuels such as methane and octane are nonpolar hydrocarbons — their low polarity means they do not dissolve in water but vaporize easily enough to mix with air for combustion.
Chapter 17 (Aqueous Reactions): Polarity explains why some compounds dissolve in water. Chapter 18 (Gases): Gas molecules have such weak intermolecular forces that they essentially ignore each other. Chapter 19 (Phase Changes): Melting points, boiling points, and heats of vaporization are determined by intermolecular force strength. Chapter 23 (Acids and Bases): Water's acid-base behavior is rooted in its polarity and hydrogen bonding. Chapters 25 and 28 (Organic and Biological Chemistry): Molecular geometry and intermolecular forces are central to understanding life at the molecular level.
For each molecule: (a) draw the Lewis structure, (b) count electron groups on the central atom, (c) identify the number of lone pairs, (d) state the molecular geometry and approximate bond angle.
For each bond, identify which atom is more electronegative, label atoms with δ+ and δ−, and classify the bond as nonpolar covalent, polar covalent, or ionic.
Determine whether each molecule is polar or nonpolar. Explain your reasoning by considering both bond polarity and molecular geometry.
Identify the type(s) of intermolecular force (London dispersion, dipole-dipole, or hydrogen bonding) present between molecules of each substance.
Rank each group in order of increasing boiling point. Explain your reasoning in terms of intermolecular forces.
Carbon dioxide has very polar C=O bonds but is linear — the two bond polarities point in opposite directions and cancel perfectly. CCl₄ has four polar C–Cl bonds, but tetrahedral symmetry causes them to cancel. Both CO₂ and CCl₄ are nonpolar. You must consider both bond polarity AND geometry to determine molecular polarity.
Lone pairs have a major effect on molecular geometry because they repel bonding pairs more strongly than bonding pairs repel each other. Without counting lone pairs, you cannot correctly predict the shape of water (bent, not linear) or ammonia (trigonal pyramidal, not trigonal planar). Lone pairs must be included in the electron group count even though they do not appear in the final shape.
Intermolecular forces and covalent bonds are fundamentally different in kind, not just strength. Covalent bonds involve orbital overlap and electron sharing. Intermolecular forces arise from electrostatic interactions between partial charges and temporary dipoles. When water boils, no covalent bonds break — only intermolecular forces are overcome.
Hydrogen bonds require the hydrogen to be bonded directly to F, O, or N — not just any electronegative atom. A C–H bond does not form true hydrogen bonds even when carbon is bonded to electronegative elements. The key is that F, O, and N are electronegative enough to make the hydrogen nearly bare, which is required for hydrogen bond formation.
The 109.5° angle is the ideal for a perfect tetrahedron with four identical bonding groups and no lone pairs. When lone pairs are present — as in ammonia (~107°) or water (~104.5°) — they compress the bond angles below 109.5°. Only molecules like CH₄ and CCl₄ achieve the ideal 109.5°.
Chapter 12 completes Part III's study of bonding, structure, and molecular properties. You now understand how atoms form bonds, how molecules take on three-dimensional shapes, and how those shapes and polarities determine molecular interactions.
Chapter 13 begins Part IV by introducing chemical equations — the symbolic language chemists use to describe chemical reactions. You will learn how to read equations, balance them, and understand how they represent the conservation of atoms during chemical change. The skills you built in Chapters 9 through 12 — reading Lewis structures, understanding ionic and covalent compounds, and naming chemical formulas — will all be used actively as you begin writing and interpreting chemical equations.