CHEM 117 1st EditionExam #2 Study Guide Topics 4-6Chapter 9: Chemical Bonding I: The Lewis ModelTypes of Chemical BondsDescribe the three types of chemical bonds.Ionic Bond:- An interaction between a metal and nonmetal element. Electrons are transferred. The metal atom becomes a cation and the nonmetal because an anion. Covalent Bond:- An interaction between a nonmetal and nonmetal elements. Electrons are shared. The shared electrons interact with the nuclei of both of the bonding atoms.Metallic Bond: - An interaction between a metal and metal element. Electrons are pooled. **Not discussed much in lecture.Section 9.3: Representing Valence Electrons with Dots1. What do we mean by valence electrons? :-Valence electrons are in the outermost principal energy level. -Represented by dots for the Lewis Model.-The number of valence electrons is equal to the group number of the element. (Ex. Carbon has 4 valence electrons).2. What is the octet rule? :-When you try to not have more than 8 electrons around any atom; the configuration is usually eight electrons in the outermost shell. Section 9.4: Ionic Bonding1. Lattice Energy:-The energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions. -Lattice energy is calculated with the Born-Haber cycle.Example (page 387 in text):Na(s) + 1/2Cl2 NaCl(s) ; ΔHf= -411 kJ/mol1. The first step is the formation of a gaseous sodium from solid sodium.Na(s) Na(g) ; sublimation energy of Na = +108kJ 2. Formation of a chlorine atom from a chlorine molecule.½ Cl2(g) Cl(g) ; = +122 kJ (remember to multiply bond energy of Cl2 by ½)3. Ionization of gaseous sodium. The enthalpy change for this step is the ionization energy of sodium. Na(g) Na+(g) + e- ; = +496 kJ4. Addition of an electron to gaseous chlorine. The enthalpy change for this step is the electron affinity of chlorine. Cl(g) + e- Cl-(g) ; = -349 kJ5. Formation of the crystalline solid from the gaseous ions. The enthalpy change for thisstep is the lattice energy, the unknown quantity. -411 kJ = 108 + 122 + 496 – 349 + H5** -Lattice energies become less exothermic (less negative) with increasing ionic radius.- Lattice energies become more exothermic (more negative) with increasing magnitude of ionic charge. - Recall Coulombs law: F= q1q2/r Section 9.5: Covalent Bonding-A shared pair of electrons is called a bonding pair.-A pair that is associated with one atom is called a lone pair.-Lone pair electrons are also called nonbonding electrons. -When two electron pairs are shared between two atoms, it is called a double bond.How does a polar covalent bond differ from a nonpolar covalent bond? : Polar is not shared equally, where nonpolar are shared equally.Section 9.6: Electronegativity-The ability of an atom to attract electrons to itself in a chemical bond.-Electronegativity increases across a period in the periodic table.-Electronegativity decreases down a column in the periodic table.-Fluorine is the most electronegative element.-Francium is the least electronegative element.Section 9.7Writing Lewis StructuresSection 9.8: Resonance and Formal Charge-Resonance is used when two or more valid Lewis Structures can be drawn for the same compound.-An example is with O3, a double bond can alternate sides. -The formal charge is the charge it would have if all bonding electrons were shared equally between the bonded atoms.*Formal Charge= number of valence electrons – (number of nonbonding electrons + ½ number of bonding electrons)FOUR RULES FOR FORMAL CHARGE:1. The sum of all formal charges in a neutral molecule must be equal zero.2. The sum of formal charges in an ion must equal the charge of the ion.3. Small formal charges on individual atoms are better than large ones.4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom. (Example would be for the compound “cyanate,” OCN-, the structure with the lowest formal charge is the best, Oxygen is the most electronegative so the specific structure would be best fitting)Section 9.9: Exceptions to the Octet RuleThree categories:1. Odd number of electrons (Example- NO2, 17 valence electrons.2. Molecules or ions with fewer than 8 electrons around atom. (Example- BF3)3. Molecules or ions with more than 8 electrons around an atom.Topic 5, Chapter 10Section 10.2: VSEPR Theory: The Five Basic ShapesGoals:- Shapes, electron geometry, molecular geometry- Effect of lone pairs- Predicting molecular geometries- Bond anglesVSEPR Theory: based on the simple idea that electrons groups (lone pairs, single bonds, multiple bonds, and single electrons) repel one another through coulombic forces. Main Concept of VSEPR Theory: Repulsions, between electron groups on interior atoms of a molecule determine the geometry of the molecule. Different Geometries:1. Two Electron Groups: Linear Geometry, 180 degree bond angle. 2. Three Electron Groups: Trigonal Planar Geometry, 120 degree bond angles. 3. Four Electron Groups: Tetrahedral Geometry, 109.5 degree bond angles. 4. Five Electron Groups: Trigonal Bipyramidal Geometry, the angles between theequatorial positions are 120 degrees, and the angle between the axial positions is90 degrees. 5. Six Electron Groups: Octahedral Geometry, 90 degree angles. Section 10.5: Molecular Shape and PolaritySteps to determine molecular polarity:1. Find molecular geometry2. Consider dipole moment3. Find vector sumsSection 10.6: Valence Band Theory: Orbital Overlap as a Chemical BondValence Bond Theory: electrons reside in quantum-mechanical orbitals localized on individual atoms.-When two atoms approach each other, the electrons and nucleus of one atom interact with the electrons and nucleus of the other atom.-If the energy of the system is lowered because of the interactions, then a chemical bondforms. If the energy is raised by the interactions, then a chemical bond does NOT form.-A chemical bond results from the overlap of 2 half filled orbital with spin-pairing of the 2 valence electrons. **Orbitals of similar energy have better overlap and therefore form stronger bonds (example: 2s-2s overlap stronger than 2s-5s)Section 10.7: Valence Bond Theory: Hybridization of Atomic Orbitals*The concept of hybridization in Valence Bond Theory is to help recognize that the orbitals in a molecule are not necessarily the same as the orbitals in an atom.Hybridization: mathematical procedure in which the standard