Chapter 8 Bonding: General Concepts Forces that hold groups of atoms together and make them function as a unit. How Often Does The Topic Appear On AP Exam? 12% MC Questions Every Year On FR

8.1 Types of Chemical Bonds Bond Energy: The energy required to break a bond. It gives us information about the strength of a bonding interaction. When a bond forms, the system achieves the lowest possible energy

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 (Metal + nonmetal) . System achieves the lowest possible energy by behaving this way. All bonds occur because of electrostatic attractions. These electrostatic forces are governed by Coulomb’s law so the entire study of bonding comes down to understanding how the law applies to different chemical situations.

Ionic Bonds Coulomb’s Law – used to calculate the energy of interaction between a pair of ions. Q1 and Q2 = numerical ion charges r = distance between ion centers (in nm) A negative sign indicates an attractive force between opposite charged ions and the ion pair has a lower energy than the separated ions. Ex. Na+ Cl- Coulomb’s law can also be used to calc. a repulsive force in which the sign will be positive. Two points from analyzing Coulomb’s law: 1. Bigger charges mean stronger bonds; smaller charges mean weaker bonds 2. Charges close together mean stronger bonds; charges far apart mean weaker bonds.

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

Interaction of Two Hydrogen Atoms The system will position themselves so that they achieve the lowest possible energy. The distance at this point is called bond length. Zero point of energy is defined with the atoms at infinite separation. Energy terms involved are the net potential energy that results from the attractions and repulsions among the charged particles and the kinetic energy due to motions of the electrons.

Covalent Bonding Type of bonding in H2 molecule Electrons are shared by nuclei Polar Covalent Bond: Intermediate cases in which atoms are not so different that e- are not transferred like ionic bonds but different enough where e- are not shared equally.

Bond 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 lowercase  (delta) symbol is used to indicate a fractional charge. The arrow indicates the direction of the polarity.

= (H  X)actual  (H  X)expected 8.2 Electronegativity The ability of an atom in a molecule to attract shared electrons (covalent bond) to itself. = (H  X)actual  (H  X)expected This is Linus Pauling’s equation for determining the relative electronegativities of H and X atoms Paulings process listed in slide has been used to determine relative electronegativities for most elements

Electronegativity & Bond Type Electronegativity Differences Bond Type Character of Bond Zero Covalent  Intermediate Polar Covalent Large (3.3) Ionic

Relative Bond Polarity Order the following according to polarity: H-H, O-H, Cl-H, S-H, F-H Use electronegativity values in fig 8.3 p. 334 to calculate polarity values Covalent Bond Polar Covalent Bond 0.4 0.9 1.4 1.9 11

8.3 Bond Polarity & Dipole Moments A molecule of HF has a center of positive charge and a center of negative charge and is referred to as having a dipole moment. The arrow always points to the negative end.

Electrostatic Potential Diagram 13

Bond Polarity and Dipole Moment For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3, CH4, H2S Sample Ex. 8.2 page 337

8.4 Ions: Electron Configurations & Sizes Electron Configuration of Compounds Two nonmetals react: They share electrons to achieve Noble Gas Electron Configurations (NGEC). A nonmetal and a representative group metal react to form a binary ionic compound. The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal attracts electrons to achieve NGEC.

Predicting Formulas of Ionic Compounds Rules: Solid Ionic Compounds are most stable Noble Gas Electron Configurations are generally the most stable. Chemical compounds are electrically neutral Using Electron Configurations, you can predict the bond formation: Ca[Ar]4s2 O[He]2s22p4 You can see that Ca will give up its valence electrons to O. Ca+2 + O-2

O2> F > Na+ > Mg2+ > Al3+ Ion Size Isoelectronic Ions: Ions containing the the same number of electrons O2> F > Na+ > Mg2+ > Al3+ largest smallest Note that they all have the same electron configuration but the size differs, general for all isolectronic ions the size get smaller as you increase the atomic number because of the increase in number of protons (nuclear attraction).

Relative Ion Size Choose the largest ion in each of the following groups. Li+, Na+, K+, Rb+, Cs+ Ba+2, Cs+, I-, Te-2 Sample Ex. 8.4 p. 342

8.5 Energy Effects in Binary Ionic Lattice Energy is the change in energy when separated gaseous ions are packed together to form an ionic solid. Lattice energy is negative (exothermic) from the point of view of the system. Negative values indicate energy is given off which is what happens when you form bonds.

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] Follow the steps on p343 as it goes through the entire process of forming an ionic solid. Point out the individual steps are either endothermic ( have to put energy into the system) or exothermic (system gives off energy). The net result is exothermic

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]

Lattice Energy of CsF

Lattice energy can be represented by a modified Coulomb’s Law k = a proportionality constant that depends on structure of the solid & electron config. Q1, Q2 = charges on the ions r = shortest distance between centers of the cations and anions Lattice energy will have a negative sign because your joining cations and anions.

Relationship in Electronegativity Difference & Percent Ionic Character

Covalent Bond: A Model Chemical bonds are forces that cause a group of atoms to behave as a unit. Bonds result from the tendency of a system to seek its lowest possible energy. Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

Fundamental Properties of Models Fundamental Properties of Models p350 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.

8.8 Covalent Bond Energies & Reactions Process Energy KJ/mol CH4(g)  CH3(g) + H(g) 435 CH2(g) 453 CH(g) 425 C(g) 339 Total 1652 Average 413 This table shows that the actual bond energy is sensitive to its environment. Chemists use the average values. Bond energy values can be used to calculate approximate energies for a reaction.

Covalent Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic). ∆H = enthalpy, a measure of the total energy of a system. Unit is Joules D = Sum of bond energies(D) per mole, always has a positive sign. Enthalpy is a measure of the total energy of a thermodynamic system. It includes the internal energy, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure. Enthalpy is a thermodynamic potential. It is a state function and an extensive quantity. The unit of measurement for enthalpy in the International System of Units (SI) is the joule, but other historical, conventional units are still in use, such as the British thermal unit and the calorie. Go over the process of adding up bond energy using sample exercise 8.5

H from Bond Energies Using bond energy values (Table 8.4) to estimate H for the following reaction in the gas phase. H2(g) + Cl2(g)  2HCl Bonds Broken Bonds Formed 1 H-H (432kJ/mol) 2 H-Cl (427 kJ/mol 1 Cl-Cl (239 kJ/mol

Calc. ∆H from Bond Energies CH4 + 2Br2 + 2F2 → CF2Br2 + 2HBr + 2HF Bonds Breaking Bonds Forming 4 C-H 413kJ/mol 1652kJ 2 C-F 485kJ/mol 970 kJ Br-Br 193kJ/mol 386kJ C-Br 276kJ/mol 552 kJ F-F 154kJ/mol 308kJ H-Br 363kJ/mol 726 kJ H-F 565kJ/mol 1130 kJ Sum of Bonds Broken 2346kJ Sum of Bonds Formed 3378 kJ 30

Calc. ∆H from Bond Energies The total ∆H is -1032 kJ of energy for the reaction. It is an exothermic reaction and this is the amount of heat released in the formation of the the 3 products. Do #53b on p. 384 31

8.9 Localized Electron Model Assumes a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electron pairs are: Localized on atom = Lone Pair Shared between atoms = Bonding Pair

Localized Electron Model – Three Parts 1. Description of valence electron arrangement (Lewis Structures 8.10). 2. Prediction of geometry (VSEPR Model 8.13). 3. Description of atomic orbital types used to share electrons or hold lone pairs. (Ch 9)

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.

Steps for Writing Lewis Structures Sum the valence electrons from all atoms. Add e- for neg. polyatomic ions Subtract e- for pos. polyatomic ions Draw skeletal structure using pair e- to form bonds, least electronegative atom is usually central position Add pairs e- to surrounding atoms Add remaining e- to central atom p.42 Cracking AP Chemistry 2013

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. Go over several problems in book on writing Lewis structures: p357 sample exercises 8.6, p 385 exercise 67 & 68 Also look at section 8.11 exceptions to octet rule

8.12 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.

Formal Charge Used to help determine the correct Lewis structure when more than one structure is possible. The difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule. We need to know: 1. # VE on free neutral atom 2. # VE “belonging” to the atom in the molecule #VEbelonging = (# Lone Pair) + (1/2 # shared E)

Formal Charge Not as good Better

8.13 VSEPR Model 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 to minimize repulsion. 3. Determine positions of atoms from the way electron pairs are shared. 4. Determine the name of molecular structure from positions of the atoms. Work through the examples in section 8.13 and #89-98 at end of chapter. Use molecular model kits.