General Chemistry Exam 2 Complete Study Guide PDF

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Comprehensive study guide for Exam 2 based off Dr. White's lectures and advisement for the exam...

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GENERAL CHEMISTRY EXAM 2 COMPREHENSIVE STUDY GUIDE UNIT 2 PART 1: BONDING 8.1 TYPES OF CHEMICAL BONDING A. BONDS 1. Forces that hold atoms together 2. 3 types of bonds: a. Ionic bonds b. Metallic bonds c. Covalent bonds i. Covalent bonds: electrostatic attraction between positive nuclei and shared electrons ii. Nonpolar covalent bonds (NPCB): electrons shared equally between atoms iii. Polar covalent bonds (PCB): electrons not shared equally, they spend more time near one atom iv. Covalent bond between atoms- repulsions minimized, attractions maximized v. Most stable state, lowest potential/bond energy vi. High bond energy: atoms to close, repulsive forces cause high potential energy vii. Bond energy: the E required to break a covalent bond viii. Bond length: the distance between the 2 nuclei in a covalent bond ix. Bond order (BO): the number of bonds between 2 atoms x. The smaller the bond order, the longer the bond xi. BO=1 (single bond) BO=2 (double bond) BO=3 (triple bond) 8.2 ELECTRONEGATIVITY (EN) - The attraction by an atom for electrons in a covalent bond - Analogy: how selfish is that atom? How well does it share its electrons in a bond? A. Linus Pauling’s EN Scale a. In general, EN increases from L to R and from bottom to top b. F is the most EN atom (the most selfish, so to speak) B. The change in EN indicates the bond polarity a. 0.0: electrons shared equally (nonpolar covalent bond) ex. H2 b. 2.0: electrons not shared equally (PCB), ex. HCl c. 3.0: electrons not shared at all (ionic bond) ex. NaCl d. As change in EN increases, the bond becomes more polar. The % ionic character of the bond increases (becomes more ionic) i. Ex. O-H, change in EN= 3.44-2.20 = 1.24 8.9 THE LOCALIZED ELECTRON BONDING MODEL A. Localized Electron (LE) Model: a. Used to describe covalent bonds in molecules

b. Describes molecules that are formed by atoms sharing electrons c. Types of electron pairs: i. Lone pairs: electron pairs localized on one atom ii. Bonding pairs: electron pairs localized between 2 atoms (not shared) (the ones that form the bond are shared) 8.10-8.12 LEWIS STRUCTURES A. Octet rule: tendency of an atom to have 8 electrons in its outer (valence) shell. Exception: H (a small atom) will have 2 (duet) B. Rules for drawing Lewis Structures 1. Determine the # of valence electrons 2. Determine the arrangement of atoms a. Least EN atom goes in the center (never H) 3. Place bonds between atoms a. Each bond is 2 electrons 4. Place remaining electrons around outside atoms to make octets 5. Place any remaining electrons around central atom to make octets 6. If necessary, make double or triple bonds to satisfy octets 7. Minimize formal charge C. Exceptions to the octet rule 1. Reduced octet a. Molecules with B or Be as central atom have less than an octet around the central atom 2. Expanded octet a. Elements in periods 3-6 can have more than an octet around them i. Ex. P, S, Xe as central atom D. Resonance: this concept is used because the Lewis structure model is incomplete when describing bonding in some molecules 1. Consider nitrite NO2a. Double bonds are shorter than single bonds but in NO2-, experiments show that both N-O bonds are the same length because the ion is a hybrid of 2 Lewis structures b. Has characteristics of both structures c. It’s not switching back and forth between the 2 structures d. The electrons in a double bond are delocalized (spread out over both bonds) e. Bond order in a resonance structure = sum of BOs for the bonds divided by the # of resonance structures E. Formal Charge 1. Used to determine the dominant (most important) resonance structure 2. FC= ve- on free atom - number of bonds - # of lone e-s 3. Sum of FCs on all atoms in a species = 0 for a molecule, = the charge of the ion for a polyatomic ion 4. Rules for selecting the dominant structure (in order of importance:) a. Look for FC = 0 on all atoms- this is best

b. Sometimes it’s possible to minimize FCs (make all of them zero) c. Most electronegative atom should have the nonzero formal charge if it exists F. Organic molecules 1. Organic- contains C atoms 2. Functional groups- specific groups of atoms or bonds within a molecule CLASS

FUNCTIONAL GROUP

Alkane

C-C single bond

Alkene

C-C double bond

Alkyne

C-C triple bond

Alcohol

C bonded to -OH

Amine

C bonded to N (likely to have lone pair)

Carboxylic acid

C with double bond to O and single bond to -OH

8.13 VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY - Predicts the shape of molecules and polyatomic ions by assuming the electron pairs are as far apart as possible to minimize electron-electron repulsions A. Electron pair arrangement and molecular shape (geometry) a. To determine electron pair arrangement, count total number of electron groups (bonded atoms and lone pairs) around central atom b. To determine shape, differentiate between number of bonded atoms and number of lone pairs on central atom c. From shape bond angle is determined. Bond angle: angle formed between 3 bonded atoms Electron Pair No. and Arrangement

0 lone pairs Molecular Geometry (shape)

1 lone pair Molecular Geometry (shape)

2 lone pairs Molecular Geometry (shape)

2 Linear

180° Linear

3 Trigonal Planar

120° Trigonal Planar

< 120° Bent*

4 Tetrahedral

109.5° Tetrahedral...