Chapter 9 Ionic and Covalent Bonding

Chemical Bonds

• A chemical bond is a strong attractive force between atoms in a substance.

Types of Chemical Bonds

• Ionic Bonds: Result from attractive forces between oppositely charged ions (e.g., NaCl, consisting of Na^+ and Cl^-.)
• Covalent Bonds: Result from the sharing of valence electrons (e.g., Cl_2, in which there is a Cl – Cl bond).
• Metallic Bonds: Sharing of electrons between metal atoms (e.g., Na metal, consisting of Na atoms in a 3-D crystalline arrangement with a mobile “sea” of electrons).

Ionic Bonds and Ionic Compounds

• An ionic bond is formed by the electrostatic attraction between positive and negative ions.
• Large numbers of ions gather to form an ionic solid with strong attractive forces between ions.
• Ionic compounds (or salts) melt at high temperatures (e.g., $$801^
  oC$$ for NaCl).
• Molten salts conduct electrical current due to the motion of mobile ions.
• Ions form with a noble-gas electron configuration, which is particularly stable.
  • Na {[Ne]3s1} transfers its valence electron to Cl {[Ne]3s23p5} resulting in:Na^+ [Ne] and Cl^- [Ne]3s23p6

Lewis Electron-Dot Symbols

• Lewis symbols represent valence electrons as dots around the chemical symbol of an element.
• Dots are placed one to each side, until all four sides are singly occupied, then additional electrons are added as pairs.
• Example: Lewis symbols for K and As.
• After electron transfer, ions form a strong NaCl crystal.
• Example: Transfer of electrons from magnesium to fluorine atoms to form ions with noble-gas configurations.

Lattice Energy & Coulomb's Law

• Ions are assumed to be round spheres to understand lattice energy
• Energy of attraction can be calculated using Coulomb's Law:
  • E = k \frac{Q1 Q2}{r}
    • E = potential energy.
    • Q1 and Q2 are charges.
    • r is the distance between them.
    • k is a physical constant (8.99 × 10^9 J \cdot m/C^2).
• Example: Pairing between Na^+ and Cl^- ions.
  • Charge on an electron (e) = 1.602 ×10^{-19} C
  • Charge in Na^+ is +e.
  • Charge in Cl^- is –e.
  • Measured experimentally: r = 282 pm (2.82 x 10^{-10} m)
  • E = (8.99 × 10^9 J \cdot m/C^2) \frac{(1.602 ×10^{-19} C) (-1.602 ×10^{-19} C)}{2.82 × 10^{-10} m} = –8.18 × 10^{-19} J
• Lattice energy: the change in energy when an ionic solid is separated into gas-phase ions.
• Coulomb’s law explains relative lattice energies of ionic compounds.
  • MgO melts at $$2800^
    oC, NaCl melts at 801^
    oC$$.
  • The charge (Q) on the ions of MgO (+2, -2) is double the charge on NaCl (+1, -1), so the energy of attraction in MgO is four times stronger.
  • E = k \frac{Q1 Q2}{r}
• Size of Na^+ is larger than that of Mg^{2+}, and the size of Cl^− is larger than that of O^{2-}.
• Therefore, the distance (r) between Mg^{2+} and O^{2-} is smaller than the distance (r) between Na^+ and Cl^-.
• Consequently, the energy of attraction between Mg^{2+} and O^{2-} will be greater, so MgO has a much higher melting point than NaCl.

Electron Configurations of Ions

• Main-group ions gain or lose electrons to attain a noble-gas configuration, resulting in stability.
• Cations of Groups 1A to 3A:
  • Lose ns and np valence electrons.
  • Ion charges equal the group numbers.
  • d-electrons are not lost.
  • Example:
    • Na: [Ne]3s1 → Na^+: [Ne]
    • Al: [Ne]3s23p1 → Al^{3+}: [Ne]
    • Ga: [Ar] 3d104s24p1 → Ga^{3+}: [Ar]3d10
• Cations of Groups 3A to 5A:
  • Having ns2 electron configurations, the np electrons are lost, and the ion charges equal the group numbers minus two.
  • Example:
    • Tl: [Xe]4f145d106s26p1 → Tl^+: [Xe]4f145d106s2
    • Sn: [Kr]4d105s25p2 → Sn^{2+}: [Kr]4d105s2
    • Bi: [Xe]4f145d106s26p3 → Bi^{3+}: [Xe]4f145d106s2
• Anions of Groups 5A to 7A:
  • Having noble-gas configurations, the np subshell is filled, and the ion charges equal the group numbers minus eight.
  • Example:
    • N^{3-}: [He]2s22p6
    • Br^−: [Ar]4s23d104p6
• Transition metals form cations with more than 1 charge possible.
  • The atoms generally lose the ns electrons before losing the (n – 1)d electrons.
  • Most transition metals form the 2+ ion.
  • Transition elements that lose one (n – 1)d electrons as well as the two ns electrons form 3+ ions or higher.

Ionic Radius

• Defintion: size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found.
• Ionic radius can be measured using known distances between nuclei in crystals.
• Comparison of atomic and ionic radii:
  • A cation is always smaller than its neutral atom because the sodium atom loses its outer shell in forming the Na^+ ion.
  • An anion is always larger than its neutral atom because the Cl^− ion is larger than the Cl atom; the same nuclear charge holds a greater number of electrons less strongly. Valence orbitals expand due to greater electron-electron repulsion.
• Isoelectronic Species
  • Isoelectronic: different species having the same number and configuration of electrons.
  • Example: O^{2-}, F^-, Na^+, Mg^{2+}, and Al^{3+} are isoelectronic.
  • Their sizes decrease with atomic number due to increasing nuclear charge.
  • Ion Radius (pm)
    • O^{2-}: 140
    • F^-: 136
    • Na^+: 95
    • Mg^{2+}: 65
    • Al^{3+}: 50

Lattice Energy and Melting Point Trends

• To understand lattice energy, use Coulomb's Law: E = k \frac{Q1 Q2}{r}
• NaF, CsI, CaO are ordered in terms of increasing lattice energy and melting point as follows: CsI < NaF < CaO

Covalent Bonds and Molecular Substances

• Molecular substances consist of molecules strongly linked by covalent chemical bonds.
• Carbon tetrachloride (CCl_4) is a colorless liquid.
• Iodoform (CHI_3) is a low-melting yellow solid (m.p. 120^oC).
• A covalent bond is formed by sharing a pair of electrons between atoms.
• As hydrogen atoms move closer, each atom's electron is attracted to both its own nucleus and the nucleus of the second atom.
• The electron probability distribution illustrates this relationship.

Covalent Bond Formation

• The potential energy decreases—first gradually, and then more steeply—to a minimum.
• The distance between the atoms when energy is at a minimum is called the bond length.
• As the atoms continue to move closer, the energy increases dramatically due to nuclear repulsions.

Lewis Structures

• Lewis electron-dot formula (Lewis structure): a formula using dots to represent valence electrons.
• An electron pair is represented by two dots.
• An electron pair between two atoms is a bonding pair (represented by one line).
• Electron pairs that are not bonding are nonbonding, or lone pair electrons.
• The formation of a covalent bond in a hydrogen molecule (H_2) can be shown using Lewis electron-dot symbols.
• Electrons are shared between the atoms, giving each the stable electron configuration of helium.
• The formation of a covalent bond between H and Cl in a hydrogen chloride molecule (HCl).
• Electrons are shared between the atoms, giving each a stable noble gas electron configuration. Cl has 8 electrons around it, like argon.

Octet Rule

• Atoms tend toward having a full eight electrons in their valence shell when forming covalent bonds. This is called the octet rule.
• Hydrogen is an exception; it has two electrons in its valence shell.

Coordinate Covalent Bond

• A coordinate covalent bond is formed when both electrons of the bond are donated by one atom.
• Example: NH3 + H^+ \rightarrow NH4^+
• The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen
• Once shared, they are indistinguishable from the other N—H bonds.

Single, Double, and Triple Bonds

• A single bond: one pair of electrons is shared by two atoms.
• A double bond: two pairs of electrons are shared by two atoms.
• A triple bond: three pairs of electrons are shared by two atoms.

Polar Covalent Bonds and Electronegativity

• When atoms in a bond are alike (e.g., H_2), the bonding electrons are shared equally.
• A polar covalent bond (or polar bond): bonding electrons spend more time near one atom than the other.
• A polar covalent bond (e.g., HCl) is intermediate between a nonpolar covalent bond (equal sharing) and an ionic bond.
• A polar covalent bond results when the bonding pair is drawn more closely toward one atom than the other.
• Electronegativity (X): a measure of the ability of an atom in a molecule to draw bonding electrons to itself.
• Electronegativity increases from left to right and from bottom to top in the periodic table.
• F, O, N, and Cl have the highest electronegativity values.

Bond Polarity and Electronegativity Difference

• The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity.
  • Small difference (< 0.5): nonpolar bond.
  • Larger difference (0.5 or greater): polar bond.
  • Very large difference: ionic bond.
• Examples:
  • H_2, difference in electronegativity is 0.0 (nonpolar).
  • HCl, difference is 0.9 (polar).
  • NaCl, difference is 2.1 (ionic).
• Ionic bonds often form between metals and nonmetals due to the great electronegativity difference.
• Covalent bonds form primarily between nonmetals because the electronegativity differences are smaller.

Dipole Moments

• In a covalent bond, electrons are shifted toward the more electronegative atom, which can be predicted using the electronegativity scale.
• In H–Cl bond, electrons are pulled toward the Cl atom (X = 3.0) rather than the H atom (X = 2.1).
• The Cl atom acquires a partial negative charge (δ−) and the H atom acquires a partial positive charge (δ+).
• The HCl molecule is a polar molecule.
• Using electronegativities, bonds can be arranged in order by increasing polarity:
  • P—H, H—O, C—Cl.

Steps to Write Lewis Structures

1. Sum the valence electrons from all atoms, taking into account overall charge.
   • If it is an anion, add one electron for each negative charge.
   • If it is a cation, subtract one electron for each positive charge.
2. Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line representing two electrons).
3. Complete the octets around all atoms bonded to the central atom.
4. Place any remaining electrons on the central atom.
5. If there are not enough electrons to give the central atom an octet, try multiple bonds.

• Example: Write the Lewis formula for sulfur dichloride; SCl_2.
• Example: Write the Lewis formula for dibromomethane; CH2Br2.
• Example: Write the Lewis formula for the chlorite ion, ClO_2 ^−
• Example: Determine the electron-dot formula of COCl_2.
• Example: Write the Lewis formula for ozone; O_3.

Delocalized Bonding and Resonance

• The assumption that bonding electrons are strictly located in the region between 2 atoms does not always fit.
• In the Lewis structure for ozone (O_3), the double bond could be assigned to 2 possible locations, giving 2 structures.
• Experiment shows that both bonds are actually identical, and that the pair of electrons in the double bond is spread out across all 3 atoms.
• Delocalized bonding: a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than being localized between two atoms.
• A single electron-dot diagram cannot properly describe delocalized bonding.
• Resonance description (resonance structures): the electron structure of a molecule or ion having delocalized bonding is given by writing all possible electron-dot formulas, connected with a double-headed arrow.
• Example: Describe the electron structure of the carbonate ion, CO_3^{2-}, in terms of Lewis electron-dot formulas.
• Example: Describe the electron structure of sulfur dioxide, SO_2, in terms of Lewis electron-dot formulas.

Delocalized Bonding in Metals

• Metals are extreme examples of delocalized bonding.
• In sodium metal, the metal consists of positive sodium ions in a “sea” of valence electrons.
• Valence (bonding) electrons are free to move throughout the metal crystal, and are responsible for the electrical conductivity of the metal.

Exceptions to the Octet Rule

• Some molecules have electron-dot structures that do not satisfy the octet rule.
• Some have an odd number of electrons, such as NO.
• Other molecules either have more than eight or less than eight valence electrons around the central atom.
  • e.g., PF5 or BF3
• Determine:
  • xenon tetrafluoride, XeF_4.
  • PBr_5
  • boron trichloride, BCl_3.
  • ICl_5.
  • sulfur hexafluoride, SF_6.

Formal Charge

• Assigning formal charges to atoms in a Lewis formula helps in assigning the correct formula, given a few possibilities.
• Formal charge on an atom: the hypothetical charge obtained by assuming that bonding electrons are equally shared between bonded atoms and that the electrons of each lone pair belong completely to one atom.
• Formal charge on an atom = (# of valence electrons on “free” atom) ‒ ½ (# of electrons in bonds) ‒ (# of electrons in lone pairs)

Rules for Determining Most Likely Electron-Dot Formula

1. Choose the Lewis formula having the lowest magnitudes of formal charges.
2. Choose the one having the negative formal charge on the more electronegative atom (if magnitudes of formal charges are the same).
3. Choose Lewis formulas that do not have like charges (both + or both ‒) on adjacent atoms (when possible).

• Determine the Lewis electron-dot formula for phosgene, COCl_2
• Determine the Lewis electron-dot formula for BCl_3
• Determine the Lewis electron-dot formula for isocyanate ion, NCO−. Carbon is the central atom.
  • Draw all possible resonance forms, and assign formal charges to the atoms in each.
  • Choose the most likely resonance form.

Measuring Bond Length

• Bond length (or bond distance): the distance between nuclei in a bond, measured using x-ray crystallography.
• Bond lengths can be predicted ( + a few pm) from the covalent radii of the atoms.
• The sum of the covalent radii of atoms A and B predicts the approximate A—B bond length.

Predicting Bond Length

• Table of Single-Bond Covalent Radii
• Predict the arsenic-iodine bond length in the molecule AsI_3.
• Predict the bond lengths of the C-S, C-H and S-H bond in methyl mercaptan.

Bond Order and Length

• Bond order: the number of pairs of electrons in a bond (single (bond order = 1), double (2), or triple (3)).
• For a bond between the same 2 elements, a single bond is longer than a double bond, which is longer than a triple bond.
• Bond length: triple bond < double bond < single bond
• Consider the molecules N2H4, N2, and N2F_2. Determine the molecule with the shortest nitrogen-nitrogen bond and the molecule with the longest nitrogen-nitrogen bond.
• Formic acid is the irritant in ant bites. One of the carbon-oxygen bonds has a length of 136 pm, and the other is 123 pm long. What is the length of the C=O bond in formic acid?
• Which arrow points to the shortest carbon-carbon bond in the ethinyl estradiol (EE) molecule below?

Bond Enthalpy

• Bond enthalpy (BE) is the average enthalpy change (ΔH) for breaking an A—B bond in a molecule in the gas phase.
• It is a measure of bond strength: the larger the bond enthalpy, the stronger the bond.

Calculating Bond Enthalpy

• The bond enthalpy (BE) of a C-H bond can be estimated from the chemical equation below in the gas phase. BE (C-H) = ¼ x 1662 kJ = 416 kJ

Bond Enthalpies Table

• For the same nuclei, triple bonds are strongest and single bonds are weakest.
• Bond energies can be used to estimate the enthalpy change, ΔH, for a reaction.
• In general, the enthalpy of reaction is approximately equal to the sum of the bond enthalpies for bonds broken minus the sum of the bond enthalpies for bonds formed.

Estimating Enthalpy Change

• Using bond enthalpies, estimate the enthalpy change for the gas-phase conversion of methyl isonitrile to acetonitrile.
• Using bond enthalpies, estimate the enthalpy change for the gas-phase hydrogenation of ethylene to form ethane.