Back to CHEM 1411 Concept Reviews

Lewis Symbol: A symbol that represents the valence electrons of an element as dots which are placed on the four sides of that element’s chemical symbol.

Example: CHEM 1411 Ch 9 img 1.png

Octet Rule: Atoms will tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

Ionic Bonding: Bonding which involves the complete transfer of electrons from one element to another.

Lattice Energy: The energy change associated with the formation of 1 mole of an ionic solid from its gaseous ions.

Example: $LaTeX: NaCl\left(s\right)\longrightarrow Na^+\left(g\right)+Cl^-\left(g\right)$ $LaTeX: \Delta H_{lattice}=-788\:kJ/mol$

Covalent Bonding: Bonding which involves atoms sharing one or more pairs of electrons in order to allow the atoms involved to achieve a more stable state. A single bond is formed when one pair of electrons is shared between 2 atoms, a double bond is formed when two pairs of electrons are shared between 2 atoms, etc. Generally, atoms will share as many electrons as needed to obtain an octet or 8 valence electrons.

Lewis Structures: Drawings that show how atoms are bonded together and electron placement in a molecule using lines (dashes) to represent shared pairs (i.e. bonds) and dots to represent unshared electrons. For a systematic method of drawing Lewis Structures, please see p. 356 in the textbook Principles of Chemistry: A Molecular Approach, 3^rd ed. or the handout “How to Draw Lewis Structures for Covalent Compounds” available at the SLC.

Electronegativity and Bond Polarity

Electronegativity: The ability of an atom in a molecule to attract electrons to itself.

Trend: The higher the effective nuclear charge, the more strongly an atom in a molecule will draw electrons to itself due to the attractive forces of the positively charged nucleus and negatively charged electrons. Since effective nuclear charge (Z_eff) increases going left to right across a period, electronegativity increases going left to right across a period. The attractive force gets weaker as the distance between the nucleus and the electrons gets bigger. Therefore, since atomic bonding radius gets larger going down a group, electronegativity gets smaller going down a group.

The size of the electronegativity difference between two atoms can help you more accurately classify the bond as non-polar covalent, polar covalent, and ionic.

Example: One may expect SnCl₄ to be an ionic compound based on its formula and the fact that Sn is a metal and Cl is a non-metal. However, the difference in electronegativity is only 1.2, making it a polar-covalent bond. This explains the observations that SnCl₄ is a liquid at room temperature and has no ionic characteristics.

Polar Covalent Bonds: Bonds that involve the unequal sharing of electrons between the two atoms. This occurs when the atoms have different electronegativities. The atom with the higher electronegativity will have a slight negative charge whereas the atom with the lower electronegativity will have a slight positive charge. A few examples would be the H-F bond, the C-H bond, and the S-O bond.

Non-Polar Covalent Bonds: Bonds that involve equal sharing of electrons between the two atoms. This occurs when the atoms have the same (or nearly the same) electronegativity. Some examples would be the bonds of the diatomic elements (Br₂, I₂, N₂, Cl₂, H₂, O₂, and F₂)

Dipole Moment (μ): a measure of the degree of separation of positive and negative charge between bonded atoms.

$LaTeX: \mu=qr$

where “q” is the charge and “r” is the distance between the atoms’ nuclei. Dipole moments are usually measured in debyes (D). $LaTeX: 1\:D=3.34\times10^{-30}\:C\cdot m$ Complete separation of 1 electron is about 6.2 D.

Formal Charge in Lewis Structures

Formal Charge: The charge that a given atom would have if all the atoms in a molecule had the same electronegativity (equal sharing of electrons in all bonds). In molecules that have multiple resonance structures (different Lewis structures for the same molecule), the greatest resonance contributor (most likely Lewis structure) will be the one that has all formal charges as close to zero as possible. Formal charges can be calculated using the following equation:

$LaTeX: Formal\:Charge=valence\:electrons-nonbonding\:electrons-\frac{1}{2}bonded\:electrons$

where the Valence Electrons are the number of valence electrons that the given atom would normally have if it was neutral and by itself, and the bonded and non-bonded electrons of the atom are determined from the Lewis Structure.

Example:

Atom	N	C	S
Valence Electrons	5	4	6
1/2 Bonded Electrons	-3	-4	-1
Nonbonding Electrons	-2	-0	-6
Formal Charge on the Atom	0	0	-1

In cases where more than one Lewis structure can be drawn for a given compound, these different Lewis structures are referred to as resonance structures, where the actual molecule exists as a hybrid of the different resonance forms. However, not all resonance structures are equal. The most important (dominant) resonance structures are the ones where all of the atoms of the molecule have formal charges as close to zero as possible. If a negative charge must be present on the Lewis structure, the best resonance structure will put the negative charge on the most electronegative element.

Exceptions to the Octet Rule

Odd Number of Electrons: When a compound has an odd number of valence electrons, one of the atoms in the compound will have an unpaired electron or “free radical”. This usually results in one atom having only 7 valence electrons surrounding it instead of the “normal” 8 predicted by the Octet Rule.

Less than an Octet of Valence Electrons: These major exceptions are summarized in the table below.

Atom H He Be B

Alternative to Octet 2 valence electrons 2 valence electrons 4 valence electrons 6 valence electrons

More than an Octet of Valence Electrons: The largest group of exceptions. Basically, any atom that is in the third period or below is capable of having an “expanded octet” (i.e. being hypervalent). An expanded octet will usually occur in order to minimize formal charges in a molecule or create compounds with minimized formal charges. This usually occurs when a larger central atom bonds to small electronegative outer atoms.

Bond Enthalpies

Bond Enthalpy: The enthalpy change (∆H) for breaking a particular bond of one mole of a gaseous substance.

Bond enthalpies can be used to calculate enthalpies of reaction (∆H_rxn) using the following equation:

$LaTeX: \Delta H_{rxn}=\sum\left(enthalpies\:of\:bonds\:broken\right)-\sum\left(enthalpies\:of\:bonds\:formed\right)$

To simplify this process, one can assume that all reactant bonds are broken and all product bonds are formed.