Chemical Bonding

Unit 4 builds on the periodic-table reasoning from Unit 2. The whole point is to predict what happens when atoms come together. Will they share electrons, transfer them, or pool them in a sea? Will the resulting molecule have a positive end and a negative end? Bonding is the bridge from "what atoms are" to "what stuff is made of."

By the end of this page you will be able to identify a bond as ionic, covalent, or metallic just from the two elements involved, draw the Lewis structure of any common molecule, calculate the electronegativity difference to classify a bond's polarity, and determine whether a whole molecule is polar or nonpolar based on its shape and bonds.

Atoms bond to reach a more stable configuration, usually by getting a full outer shell of electrons. That target is the noble-gas configuration: 8 valence electrons (or 2 for hydrogen and helium). This is the octet rule.

Different atoms get to a full octet different ways:

• Atoms with almost no valence electrons (metals on the left side) prefer to lose them, leaving the next-shell-down full.
• Atoms with almost full valence shells (nonmetals on the right side) prefer to gain a few to fill up.
• If both atoms want to gain, they share.

Once you know whether an element is a metal or a nonmetal, you can predict the bond type. Three combinations, three bond types.

The Regents likes to give you the formula and ask which type of bond holds it together. Look at the elements. Two metals? Metallic. Mix? Ionic. Two nonmetals? Covalent. That covers almost every question on this topic.

An ionic bond forms when one atom transfers one or more electrons to another. The atom that lost electrons becomes a positive cation. The atom that gained electrons becomes a negative anion. Opposite charges attract, and the bond is that electrostatic pull.

Ionic compounds share a recognizable set of properties. The Regents will give you a property list and ask whether the substance is ionic, covalent, or metallic. Memorize these.

All five properties trace back to the same source: ionic bonds are strong (because charges are full + and -) and the substance is built as a rigid 3D lattice.

A covalent bond forms when two nonmetal atoms share a pair of electrons. Each shared pair counts as one bond. Sometimes called molecular bonds, and the compounds formed are called molecular compounds.

Ionic and covalent bonds both involve valence electrons reaching a stable arrangement, but the mechanism is different. Use the widget to watch the two side by side. In ionic, an electron jumps. In covalent, the electrons stay between the two atoms.

Single, double, and triple bonds are exactly what they sound like.

Every covalent bond is made of one or more components called sigma and pi bonds, depending on how the orbitals overlap. The Regents will sometimes ask which is which.

The pattern. The number of sigma bonds equals the number of bonded atom pairs. The extra bonds in doubles and triples are always pi bonds. A pi bond never exists without a sigma already in place.

Sigma bonds let atoms rotate freely around them (think single bonds in ethane). Pi bonds lock the geometry in place because the p-orbitals must stay parallel to overlap (think of why $\text{C}{=}\text{C}$ bonds in alkenes are rigid).

• Often soft. Molecules are held to each other only by weak IFAs, so they slide apart easily.
• Low melting and boiling points compared to ionic. Examples: water boils at $100^\circ\text{C}$, ethanol at $78^\circ\text{C}$.
• Poor conductors of electricity. No free charges to move.
• Variable solubility. "Like dissolves like" - polar molecular compounds dissolve in water, nonpolar ones do not.

A Lewis structure is a drawing that shows every valence electron in a molecule. Each electron is a dot (or part of a line, if it is in a bond). Lewis structures are the standard way to depict covalent and ionic bonding visually.

1. Count total valence electrons across all atoms in the molecule.
2. Put the least electronegative atom (often carbon, often the only one of its kind) in the middle.
3. Connect the central atom to each other atom with a single bond (a line representing 2 shared electrons).
4. Distribute remaining electrons as lone pairs on the outer atoms first, until they have a full octet (or 2 for hydrogen).
5. If the central atom does not have an octet, convert lone pairs on outer atoms into double or triple bonds until it does.

For ionic compounds, you draw the cation alone (with its charge in brackets) and the anion alone (with its electrons and charge in brackets). No shared electrons. Example: $[\text{Na}]^+ [\,:\!\ddot{\text{Cl}}\!:\,]^-$.

Hydrogen is happy with 2 electrons (not 8). It can only ever form one bond. Helium is similarly content with 2.

Some larger atoms (like sulfur in $\text{SF}_6$) can have more than 8 valence electrons because they can use d-orbitals. The Regents rarely asks about expanded octets, so the standard "octet rule" is usually all you need.

A metallic bond forms when metal atoms pool their valence electrons into a shared "sea of mobile electrons." The metal atoms become cations (because they have given up their valence electrons), and the sea of free electrons holds them all together.

The "sea of electrons" model explains the standard list of metal properties almost perfectly:

• Excellent electrical conductors (solid or molten). The sea of electrons is already mobile - apply a voltage and current flows.
• Excellent thermal conductors. Mobile electrons also transfer kinetic energy quickly.
• Malleable and ductile. When you bend a metal, the layers of cations can slide past each other while the electron sea keeps them all held together. Ionic compounds shatter under the same pressure because layers of like charges line up and repel.
• Shiny (luster). Mobile electrons reflect light.
• Often high melting points. Many small + charges and one big sea give a strong overall bond. (Some exceptions: mercury is a liquid at room temperature.)
• Insoluble in water. The lattice is held together by a metallic bond, not by anything water can break.

Alloys (like brass, bronze, steel) are mixtures of two or more metals that still bond metallically. The "sea" works the same way even when the cations are different sizes.

Electronegativity (EN) is how strongly an atom pulls on electrons in a chemical bond. It is a property of an atom inside a bond, not the atom by itself. Higher EN means the atom hogs the shared electrons more.

• Across a period (left to right): EN increases. Nuclei have more protons and are pulling on a similar number of shells.
• Down a group (top to bottom): EN decreases. Extra inner shells shield the outer electrons from the nucleus, so the atom does not pull as hard.

Fluorine is the most electronegative because it is at the top-right corner of the periodic table. Cesium (bottom-left) is the least.

In a covalent bond, both atoms pull on the shared electrons, but if one atom has higher EN, it pulls harder. The bigger the EN difference, the more lopsided the sharing.

A bond between two atoms of the same element (like H-H, O=O, N≡N) is always nonpolar. $\Delta$EN is zero, so the electrons are shared perfectly evenly.

A bond can be polar even when the whole molecule is not. Molecular polarity depends on two things together: the polarity of the bonds, and the shape of the molecule.

1. Are the bonds polar? Use Table S to check $\Delta$EN. If $\Delta$EN $\lt$ 0.5 in every bond and there are no lone pairs, the molecule is nonpolar. Stop.
2. Is the molecule symmetric? If the polar bonds all point at each other and cancel out, the molecule is nonpolar overall. If they do not cancel, the molecule is polar.

• $\text{CO}_2$ (linear, nonpolar). Two polar C=O bonds, but they point in exactly opposite directions and cancel out. Net dipole = 0.
• $\text{H}_2\text{O}$ (bent, polar). Two polar O-H bonds, but the bent shape means they do NOT cancel. The oxygen end is $\delta^-$, the hydrogen side is $\delta^+$.
• $\text{NH}_3$ (pyramidal, polar). Three N-H bonds plus a lone pair sticking up on top. Lopsided. Nitrogen end is $\delta^-$, hydrogen side is $\delta^+$.
• $\text{CH}_4$ (tetrahedral, nonpolar). Four C-H bonds, all very slightly polar but perfectly symmetric. Dipoles cancel.
• $\text{H}_2$, $\text{O}_2$, $\text{N}_2$ (diatomic elements, nonpolar). Two atoms of the same element. $\Delta$EN = 0. No polarity to begin with.

Inside a molecule, atoms are held together by chemical bonds (covalent or ionic). Those are strong. But molecules also pull on each other, much more weakly. Those between-molecule pulls are called intermolecular forces of attraction, or IFAs. Sometimes called "weak bonds" or "non-bonds" because they are not real chemical bonds, just attractions.

Three types of IFAs, weakest to strongest:

LDFs exist between all molecules, polar or nonpolar. They come from random, momentary fluctuations in electron position. At any instant, an atom's electron cloud might be slightly lopsided, creating a temporary partial charge. That instant charge induces a matching partial charge on a neighbor, and the two attract briefly.

Dipole-dipole forces exist between polar molecules. The $\delta^+$ end of one molecule attracts the $\delta^-$ end of a neighbor. Stronger than LDFs because the partial charges are permanent, not fleeting.

A hydrogen bond is an extra-strong dipole-dipole interaction that happens only when a molecule contains a hydrogen atom bonded directly to F, O, or N. The exposed proton on that hydrogen attracts to a lone pair on a neighboring molecule's F, O, or N.

Hydrogen bonds are still much weaker than a covalent bond inside a molecule (about 5-10% as strong). But they are several times stronger than ordinary dipole-dipole forces. That is why molecules that hydrogen-bond have dramatically higher boiling points than you would otherwise predict from their size.

Real chemical bonds (top two) are 5-50× stronger than even the strongest IFAs (bottom three). But IFAs are everywhere, and they decide the macroscopic properties you actually observe.

Water is the most-tested substance on the Regents because it is the most useful and also the most chemically unusual. Almost every weird thing about water traces back to one fact: water molecules hydrogen-bond to each other, aggressively.

• Very high boiling point ($100^\circ\text{C}$). $\text{H}_2\text{S}$, water's heavier cousin, boils at $-60^\circ\text{C}$. Water "should" boil around $-90^\circ\text{C}$ based on molecular weight alone, but hydrogen bonds hold the molecules together so tightly that you need much more energy to vaporize them.
• High specific heat (4.18 J/g·°C). Hydrogen bonds absorb a lot of energy before they let go and let molecules speed up. This is why water resists temperature changes and is used in cooling systems (radiators, computer water-cooling loops, hot-water heating). Detailed coverage in Unit 5.
• Very high heat of vaporization ($H_v = 2260$ J/g). Breaking all the H-bonds to vaporize water takes huge energy. This is why sweat is such an effective cooling mechanism: each gram that evaporates carries away 2260 J of heat.
• Ice floats. Ice's crystal lattice forces molecules into an open, hexagonal pattern (so that every O can sit at the right H-bond angles to four neighbors). The lattice has empty space, making ice less dense than liquid water. Lakes freeze top-down, not bottom-up. Fish survive winter.
• High surface tension. Surface molecules can H-bond inward and sideways but not upward. The net inward pull creates a "skin" strong enough for insects to walk on water.
• Capillary action. Water can pull itself up narrow tubes (xylem in plants, paper towels) because H-bonding lets it grip both itself (cohesion) and other polar surfaces (adhesion).
• Universal polar solvent. Water dissolves a huge range of polar and ionic compounds because its dipoles can pry apart ionic lattices and surround polar solutes. More on this in Unit 5.

Picture a glass of liquid water as a constantly-rearranging mesh. Each H atom on a water molecule reaches out for a lone pair on a neighboring O. Each O lone pair is being approached by an H from a different molecule. The bonds are dotted lines (not full covalent bonds) but there are billions of them per drop. That mesh is why water flows, sticks to things, dissolves things, and resists temperature changes.

Unit 5, Physical Behavior of Matter, takes the IFA story further: how H-bonds set water's specific heat, why "like dissolves like" works the way it does, and how water cooling exploits all of this.