Bonding: General Concepts Chemical Bonds Electronegativity, Polarity Ionic Bonds Covalent Bonds: Lewis Structures, VSEPR

CHEMICAL BONDS Forces that hold groups of atoms together to form molecules. The driving force is the lowering of energy due to electrostatic attractions between the positive nuclei and the negative electrons exceeding repulsions between nuclei and between electrons.. Separated atoms have zero energy and chemically bonded atoms have negative (lower) energy. (Fig 8.1). The minimum energy or well corresponds to the bond length

Figure 8.1 a & b (a) The Interaction of Two Hydrogen Atoms (b) Energy Profile as a Function of the Distance Between the Nuclei of the Hydrogen Atoms

CHEMICAL BONDS (2) This lowering of energy is achieved when atoms achieve a noble gas electron configuration or an octet. We will see that bonds form in order that each participating atom achieves an octet. We will also see that there are exceptions.

CHEMICAL BONDS (3) Form between atoms resulting in molecules (covalent bonds, sharing of electrons). Form between ions resulting in ionic cmps (ionic bonds, electron transfer). Chemical bonding model assumes molecule consists of individual chemical bonds. Bond strength varies and is measured by bond energy (kJ/mol) = energy required to break a mole of bonds.

ELECTRONEGATIVITY Defined as the ability of an atom to attract shared electrons in a covalent bond to itself. EN > 0; Fig 8.3 EN largest in upper right hand corner of PT. This unequally sharing leads to unequal charges on the atoms. Use δ+ and δ- to indicate partial charges on the atoms.

Figure 8.3 The Pauling Electronegativity Vaules

BOND POLARITY Polar covalent bond forms when electron pair is not shared equally due to bonded atoms having different EN values. ΔEN = difference in EN –~ 0, nonpolar covalent bond. E.g. H 2, O 2 –< 2, polar covalent bond; e-pair is held more closely by atom with greater EN –> 2, bond is ionic and electron is transferred to form anion and cation (vs Sec 8.6)

Figure 8.12 a-c The Three Possible Types of Bonds

DIPOLE MOMENT When there is a separation of electron charge leading to polar bonds, the molecule may have a dipole moment. –All diatomics with polar bonds have a dipole moment. (HCl, NO, CO) –Polyatomics with polar bonds MAY have a dipole moment. (Fig 8.2). H 2 O, NH 3, SO 2 )

Table 8.2 Types of Molecules with Polar Bonds but No Resulting Dipole Moment

Figure 8.6 a-c The Structure and Charge Distribution of the Ammonia Molecule

IONIC BONDS (8.4) (Metal) Cation + (Nonmetal) Anion  Ionic Solid held together with ionic bonds. This solid has a continuous network of cations surrounded by anions and anions surrounded by cations. The formation of ionic bonds is driven by favorable energy considerations: this is illustrated by the Born-Haber cycle.

ATOMIC ION SIZE Cations shrink and anions expand as electrons are removed or added to the neutral atom. In an isoelectronic series, the number of electrons stays the same, but Z is constant. –As Z increases, the ion size decreases. –Fig 8.8 Note that

Figure 8.8 Sizes of Ions Related to Positions of the Elements on the Periodic Table

Born-Haber Cycle (Fig 8.9, 8.11) Li(s)  Li(g)Sublimation energy > 0 Li(g)  Li + (g) + e - IE, T7.6 ½ F 2 (g)  F(g)Dissociation energy > 0 F(g) + e -  F - (g) EA, T7.7 Li + (g) + F - (g)  LiF(s) Lattice energy Sum all of these rxns to get energy for Li(s) + ½ F 2 (g)  LiF(s) ΔH f o = -617 kJ/mol

Figure 8.9 The Energy Changes Involved in the Formation of Lithium Fluoride from Its Elements

Lattice Energy, U KF(s)  K + (g) + F - (g) U > 0 Electrostatic attraction between Cation and Anion. As charge increases, U increases.

COVALENT BONDS (8.7) Most common type of chemical bond. Involve electrons shared by two nuclei. The covalent bond model assumes that a molecule is an arrangement of individual bonds that form between 2 atoms because the molecule is energetically favored (i.e. energy is at a minimum) compared to the separated atoms.

DISSOCIATION BOND ENERGY Chemical bonds can be assigned average (±10%) dissociation bond energies (T8.4) and bond lengths (T8.5) D > 0 kJ/mol; measure of bond strength. AB(g)  A(g) + B(g) Note single vs double vs triple bonds D values. ΔH rxn ≈ Σ D(bonds in R) – ΣD (bonds in P) because bond breaking is endothermic and bond formation is exothermic.

Table 8.4 Average Bond Energies (kj/mol)

Table 8.5 Bond Lengths for Selected Bonds

COVALENT BONDS (2) Determine physical and chemical properties of cmps. Determine the likelihood and products of chemical reactions. Determine molecular shape (Sec 8.13).

LOCALIZED ELECTRON (LE) BONDING MODEL Valence electrons participate in the formation of chemical bonds. Electron pairs are localized between (shared or bonding pair) or on (lone pair) atoms such that each atom has an octet or duet of electrons. VSEPR model predicts molecular geometry based on LE bonding model.

LEWIS SYMBOLS and STRUCTURES Lewis symbol: picture of molecule showing arrangement of its valence electrons around atoms. Lewis structure: picture of molecule showing bonding electrons as lines and nonbonding electrons as dots or lines. Especially used for main group elements (p 357)

COVALENT BONDS (3) Form when electron pairs are shared so that each atom achieves an octet (duet). Coordinate covalent bond forms when one atom provides both bonding electrons. Multiple covalent bond forms when more than one electron pair is shared between two atoms (double bond, bond order 2 [CO 2 ] and triple bond, bond order 3 [N 2 ]).

WRITING LEWIS STRUCTURES Determine total # of valence electrons. Write skeletal structure with central atom [lowest EN]; terminal atoms [H, higher EN] Use electron pairs to form bonds. Achieve octet rule for terminal atoms Add the remaining to the central atom. Form multiple bonds if needed.

WRITING LEWIS STRUCTURES (2) Exceptions to octet rule [odd # of valence electrons (NO), free radicals, incomplete octets (B), more than 8 electrons (expanded valence shell SF 6 )]. Resonance structures showing different but equivalent distributions of electrons; note delocalization (vs localization) of electrons. Be guided by experimental observations.

FORMAL CHARGE (FC) FC = [VE in free atom] - [VE asigned in molecule] FC is a hypothetical charge for electron loss (+) or gain (-) due to bond formation. [VE] free = # valence e’s for Group A atoms [VE] assigned = all lone pair electrons on atom + 1/2 shared electrons

FORMAL CHARGE (2) Best Lewis structure has minimum FC (zero). Formal Charge method is not perfect and can lead to incorrect “best” Lewis structures. The best Lewis structure is consistent with exptal evidence (bond lengths, EN data, etc)

VSEPR MODEL VALENCE-SHELL ELECTRON-PAIR REPULSION (VSEPR) Method helps us determine molecular geometry. Molecular geometry: 3-D shape of the molecule. This method assumes that the final positions of nuclei are the ones that minimizes electron repulsions because this is the one associated with the lowest energy.

VSEPR METHOD (2) Determine Lewis structure of molecule. Count electron “pairs” around the central atom where a “pair” may be a single e, lone pair, single bond, double bond, triple bond. Determine geometry of electron pairs. Determine molecular group geometry with A = central atom; X = terminal atom; E = lone pair of electrons. T8.6, 8.7, 8.8

Table 8.6 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion

MOLECULAR GEOMETRY # e pairs e pair geometrymolecular geometry 2Linear 3Trigonal planarTrigonal planar, bent 4TetrahedralTetrah, trig pyram, bent 5Trig bipyramidalTrig bipyra, seesaw, T- shaped, linear 6OctahedralOctah, sq pyrami, sq planar

MOLECULAR GEOMETRY (2) Electron pair geometry differs from molecular geometry when there are lone electron pairs (E). Electron-electron repulsions decrease as E-A-E> E-A-X> X-A-X; X = bonded pair Resonance structures Note bond angles