Chem 241a 1nd EditionExam # 1 Study Guide Lectures: 1-7Test Format: the format of the chemistry tests are all short answer. They consist of 5-7 (maybe more) pages of questions that you can answer. Example tests will be at the end of this study guide. There are review sessions Mondays in Koffler 218 at 3 and this Wednesday from 6-8 in the Steward Observatory N210. Also, Dr. Dollinger has office hours Tuesdays and Thursdays from1:00-2:30 pm or by appointment in Koffler 335. Good luck!Lecture 1-3 (August 26, 28, September 2)What is organic chemistry?: Organic chemistry is the study of carbon based compounds, typically involving nitrogen and oxygen atomsRecall the electron configuration for Carbon: 1s22s22p2 – carbon has four valence electrons, but wants an octet. Therefore it must bond.I. Bondinga. Ionic bond: the loss/gain of electrons (e-)i. Ex: NaCl: sodium completely loses its e- and gives it to chlorine, so with the charges sodium chloride is: Na+Cl- ii. This occurs with the alkaline metals (first two groups) and the non-metals (except noble gases).b. Covalent bond: equal sharing of e-i. Ex: CH4 -- With its tetrahedral shape (more on that later), there is no polarity in the molecule, thus all e- are shared equallyii. This occurs with bonding of two non-metalsc. Polar covalent bonds: unequal sharing of e-i. Ex: HCl: refer to the period table: when the differences of electronegativity are high, the dipole moment causes one atom (Cl) to pull the charge closer to itself, giving it a partially negative charge. The other molecule (H) then gets a partially positive charge because the e- is being pulled away from it slightlyii. This separation of charge is the dipole.iii. Any molecule that contains one polar bond is a polar molecule, thus operating through polar covalent bondingiv. Dipoles are additive/cancel each other out – this is where geometry comes inv. Geometry can be seen through Lewis StructuresII. Lewis Structures: A diagram that depicts the bonding between atoms and the lone pairs that exista. Ex: Formic acid – CH2O2i. First we must create a Bonding Scheme1:ii. Second we count the valence e-:1. H - 1 valence e- × 2 H’s = 2 e-2. C – 4 valence e- × 1 C = 4 e-3. O – 6 valence e- × 2 O’s = 12 e-4. Total e- = 2+4+12= 18 e-iii. Third take the Total number of e- and subtract the bonds2 to get the leftovers: 18-10=8 e-iv. Fourth, assign leftovers to fill octet to most electronegative atom firstv. Fifth, if all lone pairs are assigned but there are still unfilled octets, move lone pairs to form double and triple bonds with atoms, like between O and C in the picture on the previous pagevi. Last, calculate formal charge of base atoms (subtract the actual amount of electrons from the valence electrons)1. C: 4-4=02. O: 6-6=0vii. There are other ways that this molecule could exist, and many molecules exists in several forms that are all stable. This is called resonanceb. Resonance Theory3 – the same molecular formula has the same connectivity but places the π e- differently (pi electrons to follow). In theory, these forms stabilize a moleculei. Ex: CH2O2 (formic acid) 1 Please refer to Dr. Dollinger’s “Lewis Structure Handout” for more details on typical bonding schemes: http://bit.ly/1uDmkx4 (Through D2L)2 Two per bond3 Please refer to Dr. Dollinger’s “Resonance Hand Out” for more detail and how to draw resonance: http://bit.ly/1rIJHpF (Through D2L)ii. In order to truly grasp Lewis Structures, we must understand the VSEPR theoryIII. Valence Shell Electron Pair Repulsion (VSEPR) Theory: helps us to predict e- geometrya. “Center of electrons”/”Electron group” is defined as = lone pairs, single bonds, double bonds, triple bondsb. Three geometries toknow:i. Atom with 2 centers of e- (coe):180° - linearii. Atom with 3 coe: 120° - trigonalplanariii. Atom with 4 coe: 109° - tetrahedralc. Sometimes VSEPR does not reflect the real geometry, so we turn to hybridizationIV. Orbital Hybridizationa. The atomic orbitals present in carbon are s orbitals in the shape of a sphere, and 3 p orbitals in the shape of a dumbbellb. Bonding occurs when these orbitals overlapc. Two types of bonding:i. Sigma (σ) bonds – head on overlapping of orbitals (which is stronger because it bonds together stronger)ii. Pi (π) bonds – side to side overlapping of p orbitalsd. NOTE: single bonds contain one σ bond, double bonds contain one σ andone π bond, and triple bonds contain one σ bond and two π bondse. Atomic orbitals cannot describe molecular bonds, so Linus Pauling introduced hybridization to account for the geometry of molecules and their overlap4f. Atoms with 4 e- groups contain 4 sp3 hybrid σ bonds – the s part taking up 25%g. Atoms with 3 e- groups contain 3 sp2 hybrid σ bonds with 1 p π bond – 33% s component; thus, the atoms have to be in the same plane (trigonal planar) in order to have π bonds( ethylene shown above)h. Atoms with 2 e- groups contain 2 sp hybrid σ bonds with 2 p π bonds – 50% s component(acetylene shown below)4 Please refer to “Orbital Hybridization” for more detail on the math of hybridization: http://bit.ly/1qn5GCsLecture 4 (September 4)I. Bronsted-Lowry: came up with a definition of protein acids and basesa. This is a narrow definition that states:b. All reactions (rxns) are proton (H+) transfers. For example:c. When talking about acids and bases, the following sentences all mean the same thing (and their opposites):i. The equilibrium lies to the right (left)ii. Hx (the acid in the reactants) is a stronger (weaker) acid than Hy (the acid in the products)iii. Y (the base in the reactants) is a weaker (stronger) base than X (the base in the products)d. All acid/base rxns proceed towards the weaker acid/base pairi. This relates to thermodynamics and stabilitye. This is expressed by the equilibrium constant, or more specifically, Ka:i. Keq = [A-][H3O+]/[H2O][HA]; thereforeii. Ka = Keq[H2O] = [A-][H3O+]/[HA]iii. However what we are concerned with is the pKa, which equals -logKaiv. When the pKa is large, acid dissociates more  bigger pKa = weaker acidf. The acidic element – atom directly bound to the H+ of an acidi. Ex: H3C-O∷-H  The oxygen is the acidic elementii. If the acidic element contains a positive charge, is becomes more acidic1. H3O+ is more acidic than H2O, etc.g. General trend: when the same acidic element becomes more positive (either from – to 0 or 0 to +), it becomes more acidic as wellh. We can use