Premed · Premed · General Chemistry 1
Lecture 10: Chemical Bonding I: Ionic and Covalent
General Chemistry I
Learning Objectives
By the end of this lecture, students will be able to:
- Explain why chemical bonds form in terms of energy stabilization
- Describe the formation of ionic bonds and predict which elements form ionic compounds
- Calculate lattice energy and use the Born-Haber cycle
- Describe the formation of covalent bonds and distinguish between polar and nonpolar covalent bonds
- Use electronegativity differences to classify bonds as ionic, polar covalent, or nonpolar covalent
- Define and calculate bond energy, bond length, and bond order
Lecture Content
I. Why Do Chemical Bonds Form?
Atoms form chemical bonds because the bonded state is lower in energy -- and therefore more stable -- than the separated atoms. The octet rule captures this drive: main-group atoms tend to gain, lose, or share electrons until they achieve a noble gas electron configuration with 8 valence electrons (or 2 for hydrogen, the "duet rule"). The energy lowering that stabilizes bonds arises either from electrostatic attraction between oppositely charged ions (ionic bonding) or from the sharing of electron density between nuclei (covalent bonding).
II. Ionic Bonding
An ionic bond is the electrostatic attraction between oppositely charged ions arranged in a crystal lattice. Ionic bonds typically form between elements with a large electronegativity difference, most commonly a metal and a nonmetal. The metal loses electrons to become a cation, achieving a noble gas configuration, while the nonmetal gains electrons to become an anion. Sodium, for example, has the configuration [Ne]3s^1 and loses one electron to form Na+ with the [Ne] configuration. Chlorine has the configuration [Ne]3s^2 3p^5 and gains one electron to form Cl- with the [Ar] configuration. Together, Na+ and Cl- attract each other in a three-dimensional crystal lattice to form NaCl.
A. Properties of Ionic Compounds
Ionic compounds have high melting and boiling points because of the strong electrostatic forces holding the lattice together. They are hard but brittle -- displacing one layer of ions relative to another brings like charges into contact, causing the crystal to fracture. In the solid state, ions are locked in place and cannot conduct electricity, but when melted or dissolved they become free to move and conduct readily. Ionic compounds tend to be soluble in polar solvents, especially water.
III. Lattice Energy
Lattice energy (U) is the energy released when gaseous ions come together to form one mole of an ionic solid. A large negative value indicates a very stable compound. Equivalently, it can be defined as the energy required to completely separate one mole of ionic solid into gaseous ions (an endothermic process). The magnitude of lattice energy depends on two factors: the charges of the ions (U is proportional to q+ * q-) and their sizes (U is proportional to 1 / (r+ + r-)). This follows from Coulomb's law. Among the sodium halides, for instance, lattice energy decreases in the order NaF > NaCl > NaBr > NaI as the anion grows larger. MgO has a much larger lattice energy than NaCl because Mg^2+ and O^2- carry higher charges than Na+ and Cl-.
IV. The Born-Haber Cycle
The Born-Haber cycle applies Hess's law to calculate lattice energy indirectly by breaking the formation of an ionic solid into a series of measurable steps. For NaCl(s) forming from Na(s) and 1/2 Cl2(g), the steps are: (1) sublimation of Na(s) to Na(g), with enthalpy delta_H_sub; (2) dissociation of 1/2 Cl2(g) to Cl(g), requiring 1/2 the bond dissociation energy; (3) ionization of Na(g) to Na+(g) + e-, requiring the first ionization energy; (4) electron capture by Cl(g) + e- to form Cl-(g), releasing energy equal to the electron affinity; and (5) formation of the lattice as Na+(g) + Cl-(g) combine to form NaCl(s), releasing the lattice energy (-U). Summing these steps gives the overall enthalpy of formation, and rearranging allows the lattice energy to be solved for.
<image>A Born-Haber cycle energy diagram for the formation of NaCl. The diagram is arranged as a series of steps with horizontal energy levels and vertical arrows. Starting level: Na(s) + 1/2 Cl2(g). Step 1 (upward arrow): sublimation of Na to Na(g), labeled delta_H_sub = +108 kJ. Step 2 (upward arrow): bond dissociation of 1/2 Cl2 to Cl(g), labeled 1/2 D = +122 kJ. Step 3 (upward arrow): ionization Na(g) -> Na+(g) + e-, labeled IE = +496 kJ. Step 4 (downward arrow): electron affinity Cl(g) + e- -> Cl-(g), labeled EA = -349 kJ. Step 5 (large downward arrow): formation of ionic solid Na+(g) + Cl-(g) -> NaCl(s), labeled U = -787 kJ. The overall arrow from start to NaCl(s) is labeled delta_H_f = -411 kJ. Each energy level is clearly marked with the chemical species present.</image>
V. Covalent Bonding
A covalent bond forms when two atoms share one or more pairs of electrons, and it typically occurs between nonmetal atoms with similar electronegativities. Electron sharing allows both atoms to achieve an octet (or a duet for hydrogen). A single bond consists of one shared pair (2 electrons) and is represented by a single line (as in H-H). A double bond involves two shared pairs (4 electrons) and is drawn as two lines (as in O=O). A triple bond comprises three shared pairs (6 electrons) and appears as three lines (as in N≡N).
VI. Bond Properties
A. Bond Energy (Bond Dissociation Energy)
Bond energy is the energy required to break one mole of a particular bond in the gas phase. A higher bond energy signifies a stronger bond. For bonds between the same pair of atoms, triple bonds are stronger than double bonds, which are stronger than single bonds: C-C is 347 kJ/mol, C=C is 614 kJ/mol, and C≡C is 839 kJ/mol.
B. Bond Length
Bond length is the distance between the nuclei of two bonded atoms. Stronger bonds pull nuclei closer together, so triple bonds are shorter than double bonds, which are shorter than single bonds: C-C is 154 pm, C=C is 134 pm, and C≡C is 120 pm.
C. Bond Order
Bond order is the number of bonding pairs shared between two atoms -- 1 for a single bond, 2 for a double bond, and 3 for a triple bond. A higher bond order corresponds to a shorter bond length and a higher bond energy.
VII. Electronegativity and Bond Polarity
Electronegativity (EN) quantifies an atom's ability to attract shared electrons in a bond. On the Pauling scale, representative values are F (4.0), O (3.5), N (3.0), C (2.5), H (2.1), with metals having low values. The difference in electronegativity between two bonded atoms (delta_EN = |EN_A - EN_B|) determines the character of the bond.
When delta_EN is zero or very small (less than about 0.4), the bond is nonpolar covalent -- electrons are shared equally. Examples include H-H, Cl-Cl, and the nearly nonpolar C-H bond. When delta_EN falls between roughly 0.4 and 2.0, the bond is polar covalent -- electrons are shared unequally, creating partial charges (delta+ on the less electronegative atom and delta- on the more electronegative atom). H-Cl, O-H, and N-H bonds are polar covalent. When delta_EN exceeds about 2.0, the bond is essentially ionic, involving near-complete electron transfer, as in NaCl and KF. These boundaries are approximate; bond character exists on a continuum.
VIII. Dipole Moments
A dipole moment (mu) arises whenever there is a separation of charge, calculated as mu = q * d (charge times distance), with units of Debye (D). For a polar bond, the dipole moment arrow conventionally points from the partially positive end to the partially negative end. A molecule may contain polar bonds yet be nonpolar overall if the individual bond dipole moments cancel due to molecular symmetry. Carbon dioxide, for example, is linear, and its two C=O dipoles point in exactly opposite directions and cancel. Molecular polarity depends on both the polarity of individual bonds and the three-dimensional geometry of the molecule, a topic developed fully in Lecture 12.
<image>A comparison diagram of three bond types along an electronegativity difference spectrum. A horizontal arrow at the top shows delta_EN increasing from 0 to >2.0. Three panels below: Panel A (Nonpolar covalent, Cl2): an electron density map showing symmetrical distribution between two Cl atoms, delta_EN = 0. Panel B (Polar covalent, HCl): an electron density map showing electron density shifted toward Cl, with delta+ on H and delta- on Cl marked, a dipole moment arrow pointing from H to Cl, delta_EN = 0.9. Panel C (Ionic, NaCl): separated Na+ and Cl- ions with full charges, electron density entirely on Cl-, delta_EN = 2.1. Each panel includes the electron density surface color-coded from blue (electron-poor) to red (electron-rich).</image>
IX. Exceptions to the Octet Rule (Preview)
Not all molecules obey the octet rule. Electron-deficient compounds have fewer than 8 electrons around the central atom, as in BF3 (6 electrons around boron) and BeCl2 (4 electrons around beryllium). Expanded octets, found only for elements in Period 3 and beyond that have accessible d orbitals, feature more than 8 electrons around the central atom, as in SF6, PCl5, and XeF2. Odd-electron species, also called free radicals, are molecules with an odd number of electrons, leaving at least one atom with an unpaired electron; NO and NO2 are common examples. These exceptions are explored in greater detail in the Lewis structures lecture.
