Bonding: General Concepts Chapter 8

Bonds Forces that hold groups of atoms together and make them function as a unit.

Bond Energy It is the energy required to break a bond. It gives us information about the strength of a bonding interaction.

Ionic Bonds Formed from electrostatic attractions of closely packed, oppositely charged ions. Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

Ionic Bonds This is a statement of Coulomb’s Law where: Q1 and Q2 = numerical ion charges r = distance between ion centers (in nm) When E is positive (+), repulsion is indicated. When E is negative (-), attraction is indicated.

Bond Length The distance where the system energy is a minimum.

Interaction of two hydrogen atoms.

:N ≡ N: Bond Order :F-F: : O=O An indication of bond strength and bond length Single bond: 1 pair of e- shared Ex: F2 Longest, weakest •• •• :F-F: Double bond: 2 pairs of e- shared Ex: O2 O=O : Triple bond: 3 pairs of e- shared Ex: N2 Shortest, strongest :N ≡ N:

Covalent Bonding - covalent bonds are formed by sharing electrons between nuclei. H. + .H ----> H-H - coordinate covalent bonds are bonds where both shared electrons originate on the same atom H3N: + H+ ----> H3N-H+

Types of Covalent Bonds Polar covalent bond -- covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. A dipole moment exists. Nonpolar covalent bond -- covalent bond in which the electrons are shared equally between both atoms. No dipole moment exists.

 = (H  X)actual  (H  X)expected Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X)actual  (H  X)expected

Pauling Electronegativity Values

Percent Ionic Character where xA is the larger electronegativity and xB is the smaller value. Watch significant figures!!! Ionic Bond % IC > 50 % Polar Covalent % IC 5 - 50 % Nonpolar Covalent % IC < 5 %

Three Possible Types of Bonds Nonpolar Covalent (Electrons equally shared.) Polar Covalent (Electrons shared unequally.) Ionic (Electrons are transferred.)

Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

The Effect of an electric field on hydrogen fluoride molecules.

Dipole Moment for the water molecule.

Dipole moment for the ammonia molecule.

Nitrogen Trichloride Does nitrogen trichloride exhibit a dipole moment? Yes. It has three nonpolar bonds but, also, has a lone pair of electrons which makes it assymetrical and therefore, polar.

Nonpolar molecule--zero dipole moment.

Table 8.2 on page 356 in Zumdahl.

Achieving Noble Gas Electron Configurations (NGEC) Two nonmetals react: They share electrons to achieve NGEC. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC. See Table 8.3 on page 361 in Zumdahl.

Sizes of ion related to position on the periodic table.

O2> F > Na+ > Mg2+ > Al3+ Isoelectronic Ions Ions containing the same number of electrons (O2, F, Na+, Mg2+, Al3+) O2> F > Na+ > Mg2+ > Al3+ largest smallest

Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M+(g) + X(g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

Formation of an Ionic Solid 1. Sublimation of the solid metal M(s)  M(g) [endothermic] 2. Ionization of the metal atoms M(g)  M+(g) + e [endothermic] 3. Dissociation of the nonmetal 1/2X2(g)  X(g) [endothermic]

Formation of an Ionic Solid (continued) 4. Formation of X ions in the gas phase: X(g) + e  X(g) [exothermic] 5. Formation of the solid MX M+(g) + X(g)  MX(s) [quite exothermic]

Q1, Q2 = charges on the ions r = shortest distance between centers of the cations and anions

Comparison of the energy changes in the formation of sodium fluoride and magnesium oxide. .

H = D(bonds broken)  D(bonds formed) Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic). H = D(bonds broken)  D(bonds formed) energy required energy released Draw the Lewis Structure for each reactant and product before doing any calculations!

Single, Double, & Triple Bonds Single bonds -- one shared pair of electrons. Double bonds -- two shared pairs of electrons. Triple bonds -- three shared pairs of electrons. See bond energy Tables 8.4 & 8.5 on pages 372 in Zumdahl.

Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Bond enthalpy: Amount of energy required to break a particular bond between two elements in gaseous state. Given in kJ/mol. Remember, breaking a bond always requires energy! Bond enthalpy indicates the “strength” of a bond. Bond enthalpies can be used to figure out Hrxn . Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? 1 C-H & 1 Cl-Cl bond are broken (per mole) 1 C-Cl & 1 H-Cl bond are formed (per mole) Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Note: this is the “opposite” of Hess’ Law where Hrxn = DHproducts – Dhreactants Bond Enthalpy link

Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? Bond Ave DH/mol Bond Ave DH/mol C-H 413 Cl-Cl 242 H-Cl 431 C-Cl 328 C-C 348 C=C 614 Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Hrxn ≈ [(4(413) + 1(242)] – [1(328) + 3(413)+1(431)] Hrxn ≈ -104 kJ/mol Hrxn = -99.8 kJ/mol (actual) Note: 2 C-C ≠ 1 C=C 2(348) = 696 kJ ≠ 614 kJ

Ex: CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g) DHrxn=? *CH3(g) + H(g) + 2 Cl(g) Absorb E, break 1 C-H and 1 Cl-Cl bond Release E, form 1 C-Cl and 1 H-Cl bond H CH4(g) + Cl2(g) Breaking a bond decreases stability, therefore releasing energy. Creating a bond increases stability, therefore absorbing energy. The formation of ADP from ATP has a net release of energy, but the breaking of the peptide bond itself does not release energy. CH3Cl (g) + HCl (g) DHrxn Hrxn =  (Hbonds broken) +  (- Hbonds formed) Hrxn =  (Hbonds broken) -  (Hbonds formed)

Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

Fundamental Properties of Models A model does not equal reality. Models are oversimplifications, and are therefore often wrong. Models become more complicated as they age. We must understand the underlying assumptions in a model so that we don’t misuse it.

Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Localized Electron Model 1. Description of valence electron arrangement (Lewis structure). 2. Prediction of geometry (VSEPR model). 3. Description of atomic orbital types used to share electrons or hold lone pairs.

Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

Lewis Structures Covalent Compounds Ionic Compounds In ionic compounds, electrons are transferred and ions are formed. In covalent compounds, electrons are shared to form a molecule.

Electron Deficient Molecules Beryllium chloride -- BeCl2 -- is electron deficient with four electrons. It forms a linear molecule. Boron trifluoride -- BF3 -- is electron deficient with six electrons. It forms a trigonal planar molecule. See page 380 for the reaction between boron trifluoride and ammonia.

Electron Rich Molecules Expanded octet: occurs in molecules when the central atom is in or beyond the third period, because the empty 3d subshell is used in hybridization (Ch. 9) PCl5 SF6

Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Rules for Writing Lewis Structures Sum the valence electrons from all the atoms. Use a pair of electrons to from a bond between each pair of bound atoms. Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row elements.

Lewis Structures NO+ 5 e- + 6 e- - 1 e- = 10 e- [:NO:]+ Each atom has an octet and is satisfied.

Formal Charge Movie on Formal Charge For each atom, the numerical difference between # of valence e- in the isolated atom and # of e- assigned to that atom in the Lewis structure. To calculate formal charge: Assign unshared e- (usually in pairs) to the atom on which they are found. Assign one e- from each bonding pair to each atom in the bond. (Split the electrons in a bond.) Then, subtract the e- assigned from the original number of valence e-. #VALENCE e- in free atom – #NON-BONDING e- – ½(#BONDING e-) FC

Used to select most stable (and therefore most likely structure) when more than one structure are reasonable according to “the rules”. The most stable: Has FC on all atoms closest to zero Has all negative FC on most EN atoms. FC does not represent real charges; it is simply a useful tool for selecting the most stable Lewis structure.

Examples: Draw at least 2 Lewis structures for each, then calculate the FC of each atom in each structure. SCN1- N2O BF3 Can differentiate between equivalent atoms and bond distances using “tagging” with isotopes. Refer to Pauling’s book p. 171.

Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures called a resonance hybrid. See the resonance structures for the nitrate ion on page 384 in Zumdahl.

8.7: Resonance Structures Equivalent Lewis structures that describe a molecule with more than one likely arrangement of e- Notation: use double-headed arrow between all resonance structures. Ex: O3 Note: one structure is not “better” than the others. In fact, all resonance structures are wrong, because none truly represent the e- structure of the molecule. The “real” e- structure is an “average” of all resonance structures.

Odd-Electron Molecules NO2 contains 17 electrons. cannot satisfy the octet rule. a more sophisticated model is needed- the molecular orbital model.

Stereochemistry The study of the three- dimensional arrangement of atoms or groups of atoms within molecules and the properties which follow such arrangement.

VSEPR Model Valence Shell Electron Pair Repulsion -- The structure around a given atom is determined principally by minimizing electron pair repulsions.

Predicting a VSEPR Structure 1. Draw Lewis structure. 2. Put pairs as far apart as possible. 3. Determine positions of atoms from the way electron pairs are shared.(Parent Geometry) 4. Determine the name of molecular structure from positions of the atoms.(Actual Geometry)

Molecular Geometry Parent Geometry is electron pair arrangement about the central atom. linear trigonal planar tetrahedral trigonal bipyramidal octahedral Actual Geometry is the arrangement of atoms about the central atom. linear bent trigonal pyramid seesaw T-shaped square pyramid square planar

Lone pair of electrons on the ammonia molecule.

Lone pairs on the water molecule. 08_143 Lone pair Bonding pair Bonding O pair O H H Lone pair (a) (b) O H H (c) Lone pairs on the water molecule.

Octahedral structure for phosphorus hexachloride.

Octahedral structure for xenon.

Parent and actual geometry for xenon tetrafluoride.

Three possible arrangements of the electron pairs in triiodide ion.

VSEPR Model Summary Determine the Lewis structure(s) for the molecule. For molecules with resonance structures, use any of the structures to predict the molecular structure. Sum the electron pairs around the central atom to determine the parent geometry. The arrangement of the pairs is determined by minimizing electron-pair repulsions.(Actual Geometry)

VSEPR Model Summary (Continued) Lone pairs require more space than bonding pairs since they are tightly attracted to only one nucleus. Lone pairs produce slight distortions of bond angles less than 120o.