Chemical Bonding and Molecular Geometry

This is Unit Six of Chemistry, Part One: bonds, Lewis structures and bond energies. Unit Five ended by naming electronegativity and deferring its trend to this unit; that trend arrives here, and it turns out to organise most of what follows. The through-line is that the familiar split between ionic and covalent bonding is not two mechanisms but two ends of one scale, measured by how unequally two atoms pull on shared electrons. By the end you will be able to say how an ionic bond differs from a covalent one; say why the two are one scale rather than two categories; say what the most common Lewis-structure error is; say what resonance describes and what it does not; and say what lattice energy measures and which factor dominates it.

Four entries. The two mechanisms as the book first presents them — one atom handing electrons over, or two atoms sharing them — plus how an ionic compound's formula follows from charge and how an ion's charge follows from group position.

Handed over.

Chem 0961 Ionic Bond Formation

An ionic bond forms when one atom transfers valence electrons to another, producing a cation and an anion held together by electrostatic attraction. Sodium transfers its one valence electron to chlorine, forming the sodium and chloride ions in sodium chloride. An ionic bond is not shared electrons: it is transfer first, then attraction between the resulting opposite charges.

Balancing to zero.

Chem 0963 Ionic Compound Formula from Charge Balance

An ionic compound's formula uses the smallest whole-number ratio of cations to anions that makes the compound's total charge zero. Calcium's two-plus charge and chloride's one-minus combine in a one-to-two ratio, giving CaCl2. Balancing charge to zero sets the formula's subscripts — it does not describe how many ions physically touch in the crystal.

Read it off the group.

Chem 0966 Predicting Ionic Charge from Group Position

A main-group metal's cation charge and a main-group nonmetal's anion charge both follow from reaching the nearest noble-gas configuration. Magnesium, in group two, loses two electrons; oxygen, in group sixteen, gains two. The rule predicts main-group ions reliably — many transition metals form more than one stable charge and resist it.

Shared instead.

Chem 0969 Covalent Bond Formation

A covalent bond forms when two atoms share a pair of electrons, each atom counting that pair toward its own valence shell. Two hydrogen atoms each contribute one electron to a shared pair, forming the bond in H2. Covalent sharing differs from ionic transfer in that both atoms keep a claim on the electrons, rather than one losing them outright.

Four entries, and the pivot of the whole part. The electronegativity trend Unit Five deferred, what a difference in it does to a bond, why that makes ionic and covalent one scale rather than two boxes, and the quantity measuring the result.

The deferred trend.

Chem 0972 Electronegativity's Periodic Trend

Electronegativity rises across a period and falls down a group, tracking the same effective-nuclear-charge mechanism as ionisation energy. Fluorine, top right, has the highest of any element; francium, bottom left, among the lowest. Electronegativity is a bonding tendency read from position rather than a measured energy the way ionisation energy is.

How unequal the pull.

Chem 0975 Bond Polarity from Electronegativity Difference

A bond's polarity grows as the electronegativity difference between its two bonded atoms grows, running from nonpolar through polar toward ionic. The carbon-hydrogen bond, a small difference, is treated as nonpolar; the hydrogen-fluorine bond, a large one, is strongly polar. Polarity is a continuum set by that difference, not a fixed line separating covalent bonds from ionic ones.

One scale.

Chem 0977 Pure Covalent, Polar Covalent, and Ionic as One Continuum

Pure covalent, polar covalent and ionic bonding are three regions of one continuous electronegativity-difference scale, not separate mechanisms. A bond between identical atoms sits at zero difference, purely covalent; sodium chloride sits at the ionic end of that same scale. Textbooks draw cutoff numbers along it for convenience, but the underlying quantity varies smoothly with no true boundary.

Charge, separated.

Chem 0980 Dipole Moment

A dipole moment is the measurable separation of partial positive and partial negative charge that a polar bond produces. The hydrogen-chlorine bond's dipole points from hydrogen's partial positive toward chlorine's partial negative end. A single bond's dipole is not the whole molecule's dipole — several bond dipoles can cancel, which Part Two takes up.

Three entries. The dot notation for an atom's valence electrons, the tendency that notation exists to track, and the step-by-step procedure for drawing a molecule — which has one error more common than all the others together.

Dots for electrons.

Chem 0985 Lewis Symbol

A Lewis symbol shows an element's symbol surrounded by dots, each dot representing one valence electron. Oxygen's shows six dots, matching its six valence electrons. A Lewis symbol counts valence electrons only — core electrons never appear as dots, however many the atom holds.

Eight around.

Chem 0988 The Octet Rule

The octet rule states that main-group atoms tend to bond so each ends with eight valence electrons, matching the nearest noble gas. Chlorine gains one shared pair in hydrogen chloride, completing its octet. The rule is a strong tendency for main-group elements rather than an absolute law, and the next topic covers its two structural exceptions.

Four steps.

Chem 0991 Drawing a Lewis Structure

Drawing a Lewis structure means arranging atoms around a central atom, connecting each with a single bond, distributing the remaining valence electrons as lone pairs to complete outer octets, then converting lone pairs to additional bonds if the central atom still lacks an octet. For carbon dioxide, filling the oxygens first leaves carbon with only four electrons; converting one lone pair from each oxygen gives two double bonds and completes every octet. Stopping before that conversion is the single most common Lewis-structure error.

Five entries. Two recognised cases where a stable structure has more or fewer than eight electrons around the central atom, and then a bookkeeping device that picks between competing structures which are all otherwise valid.

Fewer than eight.

Chem 0995 Incomplete Octet

An incomplete octet occurs when a stable Lewis structure leaves a central atom, typically beryllium or boron, with fewer than eight electrons. In boron trifluoride, boron forms three single bonds and holds no lone pair, stable with only six electrons around it. An incomplete octet is a recognised outcome for a few light central atoms, not a sign the structure is wrong.

More than eight.

Chem 0997 Expanded Octet

An expanded octet occurs when a central atom from period three or below accommodates more than eight valence electrons, using available d orbitals. In phosphorus pentachloride, phosphorus forms five single bonds and carries ten electrons. Expanded octets are limited to period three and later central atoms — a second-period atom such as carbon or nitrogen never exceeds eight.

Split every bond evenly.

Chem 1000 Formal Charge, Defined

A formal charge is the charge an atom would carry in a Lewis structure if every bonding pair were divided equally between its two atoms. Nitrogen forming four bonds with no lone pair carries a formal charge of plus one in that structure. Formal charge is a bookkeeping device for one drawn structure, not the atom's real measured charge.

Valence, minus two things.

Chem 1003 Calculating Formal Charge from a Lewis Structure

Calculating an atom's formal charge means subtracting its lone-pair electrons plus half its bonding electrons from its valence electron count. For the central sulfur in sulfur dioxide, six valence electrons minus two lone-pair electrons minus three bonding-pair electrons gives plus one — a small whole number, which is the physically sensible size for a central atom in a stable molecule. Forgetting to halve the bonding electrons produces a charge off by whole units.

Which structure wins.

Chem 1006 Using Formal Charge to Choose the Best Structure

Among valid Lewis structures for one molecule, the preferred one has formal charges nearest zero, with any negative charge on the more electronegative atom. Between two valid structures for an ion, the one placing negative formal charge on oxygen rather than carbon wins. Minimising formal charge picks among already-valid structures — it never overrides the octet rule itself.

Three entries, and the book spends all three on one idea because it is mis-pictured so reliably. A molecule whose real structure no single drawing captures, what the drawings actually are, and the specific wrong reading that measurement rules out.

No single drawing.

Chem 1009 Resonance

Resonance describes a molecule whose true electron distribution blends two or more valid Lewis structures, none of which alone is fully accurate. Ozone is drawn with two structures differing only in which oxygen-oxygen bond is double, and the real molecule matches neither. Resonance structures are not different molecules flipping back and forth — the real molecule is one fixed blend.

Each contributes.

Chem 1011 Resonance Structures Are Contributing Forms

Every resonance structure for a molecule contributes to a single averaged real structure, with bond lengths intermediate between the drawn extremes. Ozone's two oxygen-oxygen bonds measure identical in length, both between a single and a double bond. No resonance structure by itself is the molecule's true structure; each is one contributing sketch toward the average.

What measurement rules out.

Chem 1013 The Single Fixed-Bond Misreading of Resonance

Treating one resonance structure as the molecule's actual fixed bonding arrangement misreads what resonance describes. Reading ozone as truly having one fixed single bond and one fixed double bond predicts two different bond lengths — and the measured lengths are identical. The real structure is the blend itself, not a molecule rapidly flipping between the drawn forms.

Five entries closing the part. How bond length tracks bond order, how strength is measured as an energy, an estimate of a whole reaction's enthalpy built from those energies, and then the equivalent quantity for an ionic solid.

Shorter as order rises.

Chem 1018 Bond Length and Bond Order

Bond length shortens as bond order rises from single to double to triple between the same two elements. A carbon-carbon triple bond measures shorter than a double, which measures shorter than a single. Bond order predicts length reliably only between the same element pair — comparing across different pairs mixes in other factors.

What it takes to break.

Chem 1021 Covalent Bond Strength

A covalent bond's strength is measured by its bond dissociation energy — the energy required to break one mole of that bond in the gas phase. A carbon-carbon triple bond's dissociation energy exceeds a single bond's, matching its shorter length. The energy is defined for breaking a bond, which absorbs energy; forming that same bond releases the identical amount.

Broken minus formed.

Chem 1024 Estimating Reaction Enthalpy from Bond Energies

Estimating a gas-phase reaction's enthalpy from bond energies means summing the energy to break every reactant bond, then subtracting the energy released forming every product bond. For hydrogen plus chlorine giving two hydrogen chloride, the two bond-breaking energies minus twice the forming energy comes out negative, matching the reaction's known exothermic character. Average tabulated bond energies give an estimate only — real values vary slightly by molecule.

Pulling a crystal apart.

Chem 1027 Lattice Energy

Lattice energy is the energy needed to pull one mole of an ionic solid apart into free gaseous ions — a positive quantity gauging bond strength. Separating solid sodium chloride into gaseous ions requires a large positive lattice energy. Some sources define it instead as the energy released on forming the solid, a negative value; this course uses the positive, separation-energy definition.

Charge, and size.

Chem 1030 Lattice Energy Trend with Charge and Radius

Lattice energy grows larger as ionic charge increases and as ionic radius decreases, both raising the electrostatic attraction between ions. Magnesium oxide, built from two-plus and two-minus ions, has a far larger lattice energy than sodium chloride's singly charged pair. Charge affects lattice energy more steeply than radius does — doubling a charge outweighs a modest radius change.

Five things to carry out of this part. An ionic bond is a transfer followed by attraction: one atom hands valence electrons to another, and the resulting cation and anion are held together electrostatically. A covalent bond is a share, with both atoms keeping a claim on the same pair. Those two are not separate mechanisms but two ends of one scale, measured by the electronegativity difference between the bonded atoms — zero difference is purely covalent, a large difference is ionic, and the textbook cutoff numbers between them are conveniences rather than real boundaries. The most common Lewis-structure error is stopping after single bonds and lone pairs when the central atom still lacks an octet, instead of converting lone pairs into additional bonds — which is what leaves a carbon with four electrons around it in carbon dioxide. Resonance describes a molecule whose real electron distribution is a single fixed blend of several valid drawings, and what it does not describe is a molecule flipping between them: ozone's two oxygen-oxygen bonds measure identical, which a fixed single-and-double reading cannot explain. And lattice energy measures the energy needed to pull an ionic solid apart into gaseous ions, rising with ionic charge and falling with ionic radius — with charge the steeper of the two. Part Two takes up molecular shape, and what shape does to polarity.