Chemical Bonding
Start here if bonding feels like too many separate topics. Building from stoichiometry, this unit connects the whole story: why atoms bond, how to tell ionic from covalent, how to draw Lewis structures, and how shape and polarity explain real properties — setting up thermochemistry.
What you'll learn
10.1 Start Here: Why Atoms Bond and What Type of Bond to Expect
Atoms bond because bonding lowers their total energy. A bond forms when the bonded atoms are more stable together than they are apart.
The instinct is to memorize bond types as isolated facts. Start instead with element type: metal + nonmetal, nonmetal + nonmetal, or metal + metal. Then use electronegativity as support, not as your first guess.
In introductory chemistry, you should recognize four main bonding pictures: ionic, metallic, polar covalent, and nonpolar covalent.
Use this order when you classify bonding in introductory chemistry:
- If both elements are metals, use the metallic bonding picture.
- If one element is a metal and the other is a nonmetal, use the ionic bonding picture.
- If both elements are nonmetals, use the covalent bonding picture.
Then use the difference in electronegativity values (ΔEN) to decide whether a covalent bond is nonpolar or polar.
Electronegativity measures how strongly an atom attracts shared electrons in a bond. Fluorine (EN = 4.0) is the most electronegative element, while cesium and francium are among the least electronegative metals.
For metallic bonding, the key idea is different. Metal atoms pool their valence electrons into a shared, mobile electron cloud rather than using ΔEN cutoffs.
The three columns below show how electron behavior differs across the main bonding types — use this as your classification reference.
Electrons transferred
from metal to nonmetal
Forms crystal lattices
Conducts electricity when dissolved
Electrons shared unequally
δ+ and δ– partial charges
Forms discrete molecules
Usually does not conduct
Electrons shared equally
No partial charges
Forms discrete molecules
Does not conduct
Metallic bonding
- Valence electrons are delocalized, not trapped between one pair of atoms
- Metals are pictured as positive ions in a mobile “sea of electrons”
- This explains conductivity, heat transfer, and why metals can bend without shattering
- Examples include Cu, Fe, Zn, and alloys such as brass
About ΔEN cutoffs
- The cutoffs are guidelines, not hard rules
- Some textbooks use slightly different boundaries such as 0.5 and 1.8
- The key idea is the trend: larger ΔEN means more ionic character
- Na-Cl has ΔEN = 2.23, so it is strongly ionic
- O-H has ΔEN = 1.24, so it is polar covalent
- Do not use ΔEN to override the obvious metal/nonmetal pattern.
10.2 Lewis Symbols: Showing Valence Electrons Clearly
A Lewis symbol, or electron-dot symbol, shows an element's chemical symbol surrounded by dots. Each dot represents one valence electron.
Start here if Lewis structures feel hard before they even start. If you cannot see valence electrons quickly, the later structure work will feel much harder than it needs to.
Valence electrons are the electrons in the outermost energy level, and they are the ones that participate in bonding. For main-group elements, the number of valence electrons matches the group number.
Place the first four dots one at a time on the four sides of the symbol before you start pairing them. This mirrors Hund's rule for the valence shell.
1 valence e⁻
4 valence e⁻
5 valence e⁻
6 valence e⁻
7 valence e⁻
8 valence e⁻
Lewis symbols for ions
- Cations lose electrons, so their dot count goes down
- Na⁺ has 0 dots because it lost its 1 valence electron
- Anions gain electrons, so their dot count goes up
- Cl⁻ has 8 dots and is shown in brackets with the charge outside
- The ion charge tells you exactly how many electrons to add or remove
10.3 Drawing Lewis Structures Step by Step with NASBU
NASBU helps you make one Lewis-structure decision at a time. First count how many electrons the atoms want. Then count how many valence electrons they actually have. From those two counts, you can find how many electrons must be shared in bonds and how many are left as lone pairs.
Do not miss the point of this method: it keeps you from guessing where the bonds go. Use it when a structure feels messy, especially with polyatomic ions and multiple bonds.
Quick example: for H₂O, A = 8 valence electrons and S = 4, so the structure has 2 bonds and 4 unshared electrons.
Count how many electrons are needed for every atom to have a full shell. Most atoms want 8 electrons, while hydrogen only needs 2. This gives the target total for a complete Lewis structure.
Add the valence electrons the atoms actually bring. For ions, add electrons for a negative charge and subtract electrons for a positive charge. Think of A as your real electron budget.
Find how many electrons must be shared in bonds: S = N − A. These are the electrons that need to sit between atoms so everyone can get as close as possible to a full shell.
Convert shared electrons into bond lines: B = S2. This tells you the total number of covalent bonds to draw. After that, sketch the skeleton by choosing a sensible center atom: hydrogen is never central, carbon is usually central when present, and halogens usually stay terminal. Then use double or triple bonds if needed so the total number of bond lines matches B.
Find the electrons left for lone pairs: U = A − S. Place these electrons on outer atoms first, then on the center if any remain. After that, check common exceptions such as hydrogen, boron, resonance, and expanded octets.
Exceptions to the octet rule
- Hydrogen is satisfied with only 2 electrons
- Boron is often stable with 6 electrons
- Period 3 and heavier central atoms can have expanded octets
- P, S, and Cl are common examples with 10 or 12 electrons around the center
- Never force an expanded octet on a Period 2 element
10.4 Lewis Structures: Worked Examples to Study the Pattern
The diagrams below show the final Lewis structures for six common molecules, drawn correctly with lone pairs, bonding pairs, and multiple bonds where needed. Study each one and notice how NASBU predicts both the total number of bond lines and the electrons left over as lone pairs.
10.5 Resonance Structures: When One Lewis Drawing Is Not Enough
Some molecules cannot be accurately described by a single Lewis structure. When the electrons in a molecule can be arranged in two or more equally valid structures — differing only in where double bonds and lone pairs are placed, not in which atoms are connected — those structures are called resonance structures.
Notice the common mistake: redrawing atoms in different positions and calling it resonance. That is not resonance. The atoms stay fixed. Only the electrons move.
The real molecule is not one structure or the other. It is a resonance hybrid — a single blended structure where the electron density is spread out (delocalized) across all the relevant bonds. The double-headed arrow (⟺) between resonance structures is not a chemical equilibrium; it means the two drawings represent the same real molecule.
The nitrate ion (NO₃⁻) — three equivalent resonance structures:
In the real NO₃⁻ ion, all three N–O bonds are identical — bond order ≈ 1.33. No single structure captures this; the hybrid does.
How to recognize resonance
- If a lone pair can move into a bond and create an equivalent structure, resonance is present
- The atoms stay in the same places; only electrons move
- Common examples include O₃, SO₂, NO₂⁻, CO₃²⁻, NO₃⁻, and benzene
10.6 VSEPR Theory: Predicting Molecular Shape from Electron Groups
VSEPR stands for Valence Shell Electron Pair Repulsion. The core idea is simple: electron pairs (both bonding pairs and lone pairs) around a central atom repel each other and spread out as far apart as possible. This repulsion determines the geometry around the central atom.
If this feels shaky, start by counting electron groups, not by memorizing shape names. One double bond still counts as one group. One triple bond still counts as one group.
There are two different geometry labels for every molecule:
- Electron geometry — the shape described by all electron pairs, including lone pairs. This is what VSEPR theory predicts first.
- Molecular geometry — the shape described by only the atoms. Lone pairs are invisible in the molecular geometry name.
The Seven Fundamental Electron Geometries:
Lone pairs compress bond angles
- Lone pairs repel more strongly than bonding pairs, so they push bonded atoms closer together.
- 4 bonding + 0 lone pairs: tetrahedral, 109.5°
- 3 bonding + 1 lone pair: trigonal pyramidal, about 107°
- 2 bonding + 2 lone pairs: bent, about 104.5°
- That is why H₂O has a smaller bond angle than NH₃, and NH₃ has a smaller bond angle than CH₄.
10.7 Molecular Polarity: Do the Bond Dipoles Cancel?
A bond can be polar (unequal sharing) without the whole molecule being polar. For a molecule to have a net dipole moment (to be a polar molecule), two conditions must both be true:
Stop here before moving to shape: a molecule with polar bonds is not automatically a polar molecule. Check the 3D shape and ask whether the dipoles cancel.
- The molecule must contain at least one polar bond (ΔEN ≥ 0.4).
- The bond dipoles must not cancel each other out in the molecule's actual 3D shape.
Ask the question this way: do the bond dipoles cancel in this shape? If they cancel, the molecule is nonpolar. If they do not cancel, the molecule is polar.
Bond dipoles are vectors. If a molecule's geometry makes the dipoles point in equal and opposite directions, the net dipole is zero — the molecule is nonpolar even though it has polar bonds.
Dipole arrows: red arrows show bond dipoles; blue arrows show the overall molecular dipole.
Because the O-H bond dipoles point into a bent shape, they add together instead of cancelling, so water has a net dipole.
In NH₃, the N-H bond dipoles do not cancel. The lone pair changes the 3D shape, so the molecule has a net dipole.
These six molecules show the full decision chain — polar bonds, shape, cancellation, and the final verdict. Study the 'Dipoles cancel?' column; that is the step most students skip.
| Molecule | Bonds | Shape | Dipoles cancel? | Polar? |
|---|---|---|---|---|
| CO₂ | 2 polar C=O | Linear | Yes — opposite, equal | No (nonpolar) |
| H₂O | 2 polar O–H | Bent | No — the bent dipoles do not cancel | Yes (polar) |
| CCl₄ | 4 polar C–Cl | Tetrahedral | Yes — perfect symmetry | No (nonpolar) |
| CHCl₃ | 3 polar C–Cl; C–H is nearly nonpolar | Tetrahedral | No — H ≠ Cl | Yes (polar) |
| NH₃ | 3 polar N–H | Trigonal pyramidal | No — lone pair tilts net dipole up | Yes (polar) |
| BF₃ | 3 polar B–F | Trigonal planar | Yes — symmetric | No (nonpolar) |
10.8 Intermolecular Forces: Why Molecular Shape Affects Properties
Intermolecular forces are attractions between molecules — they are not the covalent or ionic bonds within a molecule. IMFs are much weaker than intramolecular bonds, but they govern physical properties: boiling point, melting point, viscosity, and solubility. The stronger the IMF, the more energy is needed to separate molecules, and the higher the boiling point.
This is where the bonding unit starts paying off. Lewis structure, shape, and polarity are not separate topics anymore. They help you explain why one substance boils higher, dissolves better, or interacts differently than another.
Quick ΔEN review for IMFs
- First decide the bonding picture from element types, then use ΔEN to judge bond polarity in covalent bonds.
- Polar bonds can create a polar molecule if the bond dipoles do not cancel.
- Polar molecules can have dipole-dipole forces, and molecules with H bonded to F, O, or N can have hydrogen bonding.
- Nonpolar molecules do not have dipole-dipole forces, but they still have London dispersion forces.
When you identify IMFs, use this order:
- Every molecule has London dispersion forces.
- Polar molecules also have dipole-dipole forces.
- Molecules with H bonded to F, O, or N can also have hydrogen bonding.
Every molecule has London dispersion forces
- Polar molecules have LDFs too, not just dipole-based forces
- When comparing molecules, add up all IMFs present
- A large nonpolar molecule such as I₂ can have stronger LDFs than a small polar molecule
- That is why some nonpolar molecules can still have surprisingly high boiling points
10.8.1 London Dispersion Forces
London dispersion forces (LDFs) exist in all molecules — polar and nonpolar alike. They arise because electrons are always in motion. At any instant, electrons may be unevenly distributed, creating a fleeting instantaneous dipole. That temporary δ+/δ− on one molecule distorts the electron cloud of a nearby molecule, inducing an induced dipole in it. The two temporary dipoles attract each other momentarily.
LDF strength increases with:
- More electrons (larger molar mass) — bigger electron clouds are more easily distorted (more polarizable).
- Larger surface area — longer, flatter molecules have more surface contact and stronger LDFs than compact spherical ones of the same mass.
Why I₂ is a solid but F₂ is a gas
- Both are nonpolar diatomic molecules, so both rely only on LDFs
- I₂ has 106 electrons, while F₂ has only 18
- The larger electron cloud in I₂ is more polarizable, so its LDFs are much stronger
- That same idea explains why larger alkanes are liquids while very small ones are gases
10.8.2 Dipole-Dipole Forces
Polar molecules have a permanent, built-in dipole (δ+/δ−) from their asymmetric bond polarity. When polar molecules get close, they orient so that the δ+ end of one molecule aligns with the δ− end of a neighbor. This consistent attractive alignment is a dipole-dipole force. It is stronger than LDF for similarly sized molecules because the attraction is permanent — not just a fleeting fluctuation.
Only polar molecules experience dipole-dipole forces (in addition to LDFs).
In dipole-dipole attractions, the partially positive end of one polar molecule lines up with the partially negative end of a neighboring molecule.
Dipole-dipole vs. LDF
- For molecules of similar size, the polar one usually boils at a higher temperature
- That is because it has both LDFs and dipole-dipole forces
- Propane is nonpolar and boils at −42 °C
- Dimethyl ether is polar, has similar mass, and boils at −24 °C
10.8.3 Hydrogen Bonding
A hydrogen bond is a special, extra-strong dipole-dipole interaction that forms when hydrogen is covalently bonded to one of three highly electronegative atoms: fluorine (F), oxygen (O), or nitrogen (N). These atoms are so electronegative and so small that the bond is extremely polar, leaving H with a large partial positive charge (δ+). That exposed H is then powerfully attracted to a lone pair on an F, O, or N atom of a neighboring molecule.
The dashed line marks a hydrogen bond: a strongly δ+ hydrogen on one water molecule is attracted to a lone pair on the δ− oxygen of a neighboring molecule.
Why water is so unusual
- Water’s hydrogen bonds are about 20 kJ/mol each, much stronger than typical LDFs
- Each H₂O molecule can form up to four hydrogen bonds
- This bonding network gives water a very high boiling point for such a small molecule
- H₂S is heavier but cannot hydrogen-bond, so it boils at only −60 °C
IMF Summary and Comparison
| Force | Present in | Relative strength | Key requirement | Example molecules |
|---|---|---|---|---|
| London Dispersion | All molecules | Weakest (but largest for big molecules) | Any electrons (all molecules qualify) | He, N₂, CH₄, I₂, C₈H₁₈ |
| Dipole-Dipole | Polar molecules only | Moderate | Permanent dipole (polar molecule) | HCl, SO₂, acetone, CHCl₃ |
| Hydrogen Bonding | Molecules with H–F, H–O, or H–N | Strongest IMF | H bonded to F, O, or N; lone pair on F, O, or N nearby | H₂O, HF, NH₃, ethanol, DNA base pairs |
IMFs and boiling point
- Stronger IMFs lead to higher boiling points
- Ne has only weak LDFs, so its boiling point is very low at −246 °C
- HCl has LDFs plus dipole-dipole forces, so it boils higher at −85 °C
- H₂O has especially strong hydrogen bonding, so it boils much higher at 100 °C
- The overall order is Ne < HCl < H₂O
Next step after Unit 10
Bonding explains structure and properties. The next shift is energy: move into thermochemistry to study how chemical and physical changes absorb or release energy.