Topics 4 & 14 Chemical Bonding IB Chemistry Topics 4 & 14 Chemical Bonding Jeff Venables Northwestern High School

Chemical Bonds, Lewis Symbols, and the Octet Rule Chemical bond: attractive force holding two or more atoms together. (based on “opposites attract”) Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. Ionic bond results from the transfer of electrons from a metal to a nonmetal. Metallic bond: attractive force holding pure metals together. (electron sea)

Lewis Symbols As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. The number of electrons available for bonding are indicated by unpaired dots. These symbols are called Lewis symbols. We generally place the electrons on four sides of a square around the element symbol.

Number of dots by group: 1 = 1 15 = 5 2 = 2 16 = 6 13 = 3 17 = 7 14 = 4 18 = 8

The Octet Rule All noble gases except He have an s2p6 configuration. Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). Caution: there are exceptions to the octet rule. H = 2 Be = 4 B = 6 Any element in the third period or beyond can potentially have more than 8.

Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ

Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl-. Note: Na+ has an Ne electron configuration and Cl- has an Ar configuration. (4.1.2) That is, both Na+ and Cl- have an octet of electrons surrounding the central ion. The IONIC BOND is the electrostatic attraction between oppositely charged ions. (4.1.1)

Lattice Structure (4.1.8) NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl- ions. Similarly, each Cl- ion is surrounded by six Na+ ions. There is a regular arrangement of Na+ and Cl- in 3D. Note that the ions are packed as closely as possible. Note that it is not easy to find a molecular formula to describe the ionic lattice. It is described by an empirical formula.

Energetics of Ionic Bond Formation The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g) is endothermic. Why is the formation of NaCl(s) exothermic? Most common ions: (4.1.3, 4.1.4) Group 1 = Group 15 = Group 2 = Group 16 = Group 13 = Group 17 =

Lattice energy (ionic bond strength) increases as Most common ions: (4.1.3, 4.1.4) Group 1 = +1 Group 15 = -3 Group 2 = +2 Group 16 = -2 Group 13 = +3 Group 17 = -1 Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. It represents the strength of an ionic bond. Lattice energy (ionic bond strength) increases as The charges on the ions increase The distance between the ions decreases.

Examples – Which in each pair would have the greatest lattice energy? LiF or LiCl NaCl or MgCl2 KBr or KI MgCl2 or MgO CaO or NaF

Examples – Which in each pair would have the greatest lattice energy? LiF or LiCl NaCl or MgCl2 KBr or KI MgCl2 or MgO CaO or NaF

Transition Metal Ions (4.1.5) Lattice energies generally compensate for the loss of up to three electrons. In general, electrons are removed from orbitals in order of decreasing n (i.e. electrons are removed from 4s before the 3d). Polyatomic Ions (4.1.7) Polyatomic ions are formed when there is an overall charge on a compound containing covalent bonds. E.g. SO42-, NO3-. Be sure you know your polyatomic ions.

Covalent Bonding When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. When similar atoms bond, they share pairs of electrons to each obtain an octet. (4.2.2) Each pair of shared electrons constitutes one chemical bond. The COVALENT BOND results from the electrostatic attraction between a pair of electrons and positively charged nuclei (4.2.1) Example: H + H  H2 has electrons on a line connecting the two H nuclei.

Lewis Structures Covalent bonds can be represented by the Lewis symbols of the elements: In Lewis structures, each pair of electrons in a bond is represented by a single line:

Drawing Lewis Structures (4.2.3) Add the valence electrons. Write symbols for the atoms and show which atoms are connected to which. Complete the octet for the outer atoms, then complete the octets of the central atom. Place leftover electrons on the central atom. If there are not enough electrons to give the central atom an octet, try multiple bonds.

Examples – Draw Lewis Structures for each: H2O CO2 NCl3 SO2 SO3

Lewis Dot Structures of Ions Positive ions are formed by loss of electrons, and negative ions by gain of electrons. Electrons must be subtracted from or added to the total for ions. EXAMPLES SO42- NH4+ NO2- CO32-