Lewis Theory of Bonding

LEWIS THEORY OF BONDING (Section 4.1) Key Ideas: Lewis (1916) produced the 1st clear understanding of chemical bonding: atoms are stable if they have a noble gas-like electron structure; i.e. stable octet of electrons (isoelectronic) electrons are most stable when paired atoms form chemical bonds to achieve a stable octet of electrons a stable octet may be achieved by an exchange of electrons between1 metal & non-metal atoms (ionic bonding)   a stable octet of electrons may be achieved by the sharing of electrons between non-metal atoms (covalent bonding) the sharing of electrons results in the formation of covalent bonds LEWIS STRUCTURES Lewis Structures represent a chemical compound in which shared pairs of electrons are represented by a dashed line (-) and unpaired electrons are represented by two dots (..) Review - Ionic Compound Structures e.g. NaCl [Na]+ [Cl]- metal + non-metal Review - Covalent Compounds e.g. H2O H-O-H Lewis Structures and Quantum Mechanics Putting Lewis structures and quantum mechanics together, a connection between the two theories becomes apparent In a Lewis structure the electrons around the core element symbol represent the s & p energy sublevels Valence e- = the s + p of the highest n# Let’s take a look at the electron configuration for an atom of Magnesium (Mg): [Diagram of electron orbitals for Mg, showing 2 valence electrons in the 3s orbital] For N: [Diagram of electron orbitals for N, showing 5 valence electrons in the 2s and 2p orbitals] The Unity of the Atomic and Bonding Theories Element Magnesium, Mg Nitrogen, N Sulfur, S Valence 2 3 2 Lewis symbol -Mg- :N· :S· Energy-level diagram [Diagram of Mg's electron orbitals, showing 2 valence electrons in the 3s orbital] [Diagram of N's electron orbitals, showing 3 valence electrons in the 2p orbital] [Diagram of S's electron orbitals, showing 2 valence electrons in the 3p orbital] Electron configuration 1s² 2s² 2p⁶ 3s² or [Ne] 3s² 1s² 2s² 2p³ or [He] 2s² 2p³ 1s² 2s² 2p⁶ 3s² 3p⁴ or [Ne] 3s² 3p⁴ Extending the Lewis Theory of Bonding Lewis' theory explained many simple molecules but some molecules (including polyatomic molecules) could not be explained using this theory Work by Sidgwick showed Lewis structures can work if we don't require that each atom contribute 1 electron to a shared pair in a covalent bond i.e. one atom could contribute both electron pairs that are shared (co-ordinate covalent bond)   Sidgwick also realized that an octet of electrons around an atom may be desirable but is not necessary in all molecules & polyatomic ions   exceptions to octet rule (see powerpoint for exceptions) *Using the method that is on the following page can be applied to creating Lewis Structures for any molecule including polyatomic molecules. Drawing Lewis Structures Draw the Lewis Structure of the sulfate ion SO₄²⁻ Add the valence electrons for each atom in the molecule. Include any additional electrons from overall negative charges, and subtract electrons for each overall positive charge: S = 6 ve⁻ O = 6 ve⁻ × 4 (24 ve⁻) 30 e⁻ + 2 e⁻ (charge of -2) = 32 e⁻ total Arrange the elements in a central peripheral pattern. The central atom is usually the LEAST electronegative element. Distribute electron pairs around the central atom (Sulfur) and each of the surrounding atoms (oxygen): Place e- in bonds first, then to side atoms. Any remaining e- are placed on the central atom as a lone pair of electrons. NOTE - if there are any electrons left over AFTER the distribution, they are placed on the central atom as a lone pair of electrons. Double bonds are between two atoms are possible if ALL of the following conditions are met: i) The atoms are Carbon, Nitrogen, Oxygen, or sometimes Sulfur ii) There are insufficient electron pairs to fill the shells of the two atoms iii) The double bond does not cause the total number of valence electrons around either of the two atoms to exceed 8 Use the electron dot diagram to draw the Lewis Structure: