CHEM 301 1st Edition Lecture 1 Outline of Current Lecture I. Defining Organic Chemistry and Atomic StructuresII. Electron Configurations:a. Hund’s ruleIII. Bond Formationa. Electronegativityb. Bond PolarityIV. Calculating Formal Chargea. ResonanceV. Acids and Basesa. Arrhenius Acids and Basesb. Bronsted-Lowry Acids and Basesc. Lewis Acids and BasesCurrent LectureI. Defining Organic Chemistry and Atomic Structuresa. Organic Chemistry: Chemistry of carbon compoundsb. Carbon: Its ability to form strong carbon-carbon bondsc. Atomic Structures:i. Protons: positive (determines the identity of the element)ii. Neutrons: neutraliii. Electrons: negativeII. Electron Configurationsa. Electron Configurations: place electrons in lowest energy orbital firstb. Hund’s Rule: Equal energy orbitals are half filled, then filledc. Degenerate orbital: orbbitals with identical energiesd. Valence electrons: are those electrons that are in the outermost shell. The number of bonds an atom usually forms is called it valencei. ***Always: carbon has four bonds, nitrogen three, oxygen two and hydrogen one. ii. ***Exceptions: structures with formal chargese. Octet Rule: when forming compounds, atoms gain, lose, or share electrons so that the number of their valance electrons is the same as that of the nearest These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.noble gas. For the elements carbon, nitrogen, oxygen, and the halogens this number is 8. III. Bond formationa. Ionic bonding: electrons are transferredb. Covalent bonding: electron pair is sharedc. Lewis structures: each valence electrons is symbolized by a dotd. Nonbonding electrons or lone pairs: valance shell electrons are NOT shared between 2 atomse. The sharing of one pair: between two atoms is called a single bondf. Sharing two pairs:: double bond. Three pairs::: triple bondg. Carbon: normally forms “four bonds” in neutral organic compoundsh. Nitrogen: “three bonds” oxygen “two bonds” hydrogen and halogens “one”i. Electronegativity: A measure of the ability of an atom to attract the electrons in a covalent bond toward itself. Fluorine is the most electronegative element. Greater electronegativy means greater polarityj. Polarity: The atom with the higher electronegativity is the negative end of the dipole. IV. Calculating Formal Chargea. Each atom in a valid Lewis Structure:i. Count the number of valence electronsii. Subtract all its nonbonding electronsiii. Subtract half of its bonding electronsb. Calculating Formal Charge=[group number]-[nonbonding electrons]-1/2[shared electrons]c. Resonance: when two or more valence bond structures are possible, differing only in the placement of electrons.i. Only electrons can be moved (usually lone pairs or pi electrons)ii. Nuclei positions and bond angles remain the sameiii. The number of unpaired electrons remains the sameiv. Resonance causes a delocalization of electrical chargev. ***The real structure is a resonance hybrid. All the bond lengths are the same.d. Major resonance form: i. Has as many octets as possibleii. Has as many bonds as possibleiii. Has the negative charge on the most electronegative atomiv. Has as little charge separation as possibleV. Acids and Basesa. Arrhenius Acids and Basesi. Acids dissociate in water to give H3O ionsii. Bases dissociate in water to give OH ions.iii. pH=-log[H3O]iv. Strong acids and bases are 100% dissociated.b. Bronsted Lowry Acids and Basesi. Acids can donate a proton/Hydrogenii. Bases can accept a proton/Hydrogeniii. Conjugate acid-base pairsiv. Spontaneous acid-base reactions proceed from stronger to weakerc. Structural Effects on Acidityi. Electronegativity: As the bond to H becomes more polarized, H becomes more positive and the bond is easier to break.ii. Size-As size increases, the H is more loosely held and the bond is easier tobreak.1. A larger size also stabilizes the anioniii. Resonance stabilization of conjugate base: 1. Delocalization of the negative charge on the conjugate base will stabilize the anion so the substance is a stronger acid.2. More resonance structures usually mean greater stabilization. More resonance more acidicd. Lewis Acids and Basesi. Acids accept electron pairs=electrophileii. Bases donate electron