Chemical Bonds

Chemical bonds are the strongest forces that exist between atoms. They are the forces that hold atoms together in molecules and atoms or ions together in solids. We will look at other weak bonds and forces later.

The two most important and common strong bond types in chemistry are ionic bonds and covalent bonds.

Electronegativity

So how do we determine whether two atoms will form an ionic or a covalent bond? Use a new property - electronegativity (EN). Electronegativity is a periodic measure of how electrons are shared by atoms with the highest value for F and the lowest for Cs. There are a couple of ways of determining EN's:

• Look up values on table.
  • You should memorize values for Period 2, Li (1.0) to F (4.0) in steps of 0.5 and Hydrogen (2.1)
• Use the high, low, intermediate approximation:
  • all metals are low
  • the most electronegative of the non-metals are high (N, O, F, S, Cl, Br, I)
  • all other elements are intermediate.

Bond Type: So how do we use this to predict whether a bond is covalent or ionic?

• For numbers, use a difference of 1.7 to distinguish ionic ( EN> 1.7) and covalent ( EN< 1.7).
• Easier to use the hi, lo, intermediate system where we assign approximately hi EN, lo EN or intermediate EN values:
  • If combine hi + lo, then ionic
  • Otherwise, covalent (hi + hi, lo + lo, hi + inter., lo + inter., inter. + inter.)
  • Examples:
    • Barium iodide - lo & hi, thus ionic: Ba²⁺ + 2 I ^- = BaI_{2 (s)}
    • Carbon disulfide - intermediate & hi, thus covalent: CS₂
    • Arsenic triiodide - intermediate & hi, thus covalent: AsI₃
    • Hydrogen selenide - intermediate & intermediate, thus covalent: H₂Se
    • Oxygen difluoride - hi & hi, thus covalent: OF₂
  • Note that in each case the positive ion or lower EN element is written first!

Ionic Bonds

An ionic bond is the result of the electrostatic force of attraction between ions that carry opposite electrical charges, as described by Coulomb's Law:

E = 2.31 x 10^-19J*nm (Q₁Q₂/r)

where r is the distance between ion centers in nm.

Formation of ionic bonds. We can visualize the formation of ionic bonds as the transfer of an electron from a metal atom to a non-metal atom to form an ion pair. in vacuo:

M_(g) + energy M_(g)⁺ + e^-

X_(g) + e^- X_(g)^- + energy

M_(g)+ X_(g) MX_(g)

• Example: 2 Na + Cl₂ NaCl.
  • Na + energy Na⁺ + e^-
  • Cl + e^- Cl^- + energy
  • Na⁺ + Cl^- NaCl + H

• Crystal Structure (model)

Lewis Structures for Atoms & Ions

Lewis Dot Structures are a very simple way of modeling atoms, ions, and molecules involving the representative elements (IUPAC groups 1, 2 & 13 - 18). In a Lewis Structure the nucleus and "core" electrons (all but the outermost shell) are represented by the symbol of the element, now referred to as a "kernel." (Note that kernals are not the same as Noble gas cores for atoms in Period 4 and up because of the d electrons which are included in the kernal.) Examples:

Name	Lewis Structure	Kernel electrons	Valence electrons
Sodium	Na.	1s2 2s2 2p6	3s1
Phosphorus		1s2 2s2 2p6	3s2 3p3
Bromine		1s2 2s2 2p6 3s2 3p6 3d10 (≠ [Ar] = 1s2 2s2 2p6 3s2 3p6 )	4s24p5

For ions the charge is always shown. Thus for metal ions such as calcium the Lewis Structure simply becomes the symbol for the ion. For negative ions such as we see for oxygen (2-) we enclose the ion and its electrons in brackets to indicate that the electrons are all "owned" by the oxygen - it does not share. Notice that the Lewis Structures of monoatomic ions are isoelectronic with the nearest Noble gas. Thus sodium loses an electron to leave a kernel isoelectronic with neon, whereas bromine gains an electron to become isoelectronic with krypton. Examples:

• Sodium ion: Na⁺
• Bromide ion: We need the brackets to show that the bromide ion "owns" all of the electrons rather than sharing them, and that the charge is distributed over the entire structure - it is not associated with any particular electron or locale.

Brackets are particularly important when we make ionic compounds:

• Sodium chloride
• Potassium bromide
• Aluminum chloride
  or Al³⁺ 3

• Additional Lewis Structure Examples:
• Barium iodide
• Carbon disulfide
• Arsenic triiodide
• Hydrogen selenide
• Oxygen difluoride

Structures of Ionic Crystals

Covalent Bonds

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

Let's look at the formation of HCl as an example of the creation of a covalent bond:

H₂ + Cl₂ 2 HCl

In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:

• H₂ 2 H^.
• Cl₂ 2 Cl^.These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. H:Cl (This is not a proper Lewis structure, I have only shown the bonding pair of electrons.)
• Lewis Structure:

Note Lewis Structure homework set in Discussion Manual.

Lewis Structures for Covalent Molecules

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

• Covalent Compound Lewis Structure Examples:
  • Water: first write formula = H₂O, then determine central atom, O, then add electrons to get octet and draw L.S.
  • Ammonia = NH₃
  • Ammonium ion = NH₄⁺

  Extra "practice" structures below:

  • Methane (CH₄)
  • Hydrogen sulfide
  • Arsenic triiodide
  • Hydrogen selenide
  • Oxygen difluoride

Multiple Bonds & Resonance

Recall we must show an octet (or duet for Period I) in the outer-most shell (valence electrons). When this does not occur with single electron pairs (bonds) between atoms can sometimes make it happen with multiple bonds. You might find "Clark's Method" useful for determining the bonding patterns of various molecules.

Clark's Method for determining bonding in covalent Lewis Structures

• Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge)
  • If e^- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only.
  • If e^- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = /2 (remember a triple bond is 2 multiple bonds!).
• If you can draw more than one structure, then chose the most symmetrical.

Examples:

• carbon dioxide
• carbon monoxide

• If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.

Resonance example:

• Carbonate ion

Additional exercises on Lewis Structures are available in the Lewis Structure Module.

A Quantum Picture of Bonding For a more in-depth understanding of bonding as illustrated with QuickTime movies based on quantum calculations check out the Bonding supplement.

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© R A Paselk

Last modified 19 October 2009