EMISTRY 104 elp Sheet #1 hem 103 Review hapters 6 and 7 Do topics appropriate for your lecture Prepared by Dr. Tony Jacob http://www.chem.wisc.edu/areas/clc (Resource page) Nuggets: Electronegativity (6.7), Bond Polarity (6.7), Lewis dot structures (6.2), Bond rder (6.3-6.5), Resonance (6.9, 7.2), VSEPR (7.2), Bond angles (7.2), Molecular Polarity (7.5), ybrid rbitals (7.3), Intermolecular forces (7.6), Solubility/Miscibility (13.1) APTER 6 ELETRNEGATIVITY (χ): The degree to which an atom can attract an e - in a bond. The greater the electronegativity (EN), the greater the "pull" an atom has on an e - pair. Trend: Increases toward F. increasing to the right and up. TYPES F BNDS Nonpolar ovalent: equal sharing of e - between two identical nonmetal elements (e.g., 2, and l 2 ). Electronegativity difference (ΔEN) is zero. Polar ovalent: unequal sharing of e - between two nonmetals (e.g., F and ). Electronegativity difference (ΔEN) is small. ΔEN = EN 1 EN 2 with EN 1 and EN 2 the EN for the 2 atoms in the bond BND PLARITY: as the 2 atoms get farther apart on the Periodic Table ΔEN bond polarity Polar bond: A bond in which e - are pulled to one side of the molecule with a slight positive charge residing on one side (δ ) and a slight negative charge (δ - ) residing on the other side of the bond (δ - with the more EN atom). The arrow points toward the negative side of the bond, δ -. ( ). LEWIS DT STRUTURES 1. ount valence electrons 2. Draw skeleton structure 3. Build octets 4. heck number of electrons drawn; if equals the #valence e - in Step 1 done. 5. Too many e - Step 2 & add a multiple bond; Too few e - add electron pairs to central atom. For organic molecules, connect atoms together and make sure the #bonds is correct as remaining atoms are attached: = 4 bonds; N = 3 bonds usually; = 2 bonds usually; = 1 bond BND RDER (B): Single bond B = 1; Double bond B = 2; Triple bond B = 3 Bond rder bond strength and bond length #bonds Bond rder resonance = #locations bonds are distributed Resonance and Resonance Structures: In general chemical resonance structures typically have double bonds and an equivalent location for the double bond to move to; important structures: Benzene, 6 6, B = 1.5; ; arboylate, - 2, B = 1.5 - -
APTER 7 VSEPR (7.2) (Valence Shell Electron Pair Repulsion) 1. Draw correct Lewis dot structure. 2. Add: #lone pairs of electrons around central atom #atoms bonded to central atom 3. ompare to table below. VSEPR Total = #atoms bonded # lone pairs on Electron Picture Molecular Picture Angles E Polar A B to central atom A central atom B region geometry geometry 2 2 0 Linear Linear 180 o Bel2 NP* 3 3 0 Triangular 2 1 Triangular Triangular Bent 120 o BF3 NP* 115 o or <120 o AlF 2 - P 4 4 0 Tetrahedral Tetrahedral 109.5 o 4 NP* 3 1 Tetrahedral Triangular pyramidal 2 2 Tetrahedral Bent 107.5 o or N3 P <109.5 o 104.5 o or 2 P <109.5 o P = polar molecule; NP = nonpolar molecule * = The nonpolar (NP) molecules assume the atoms around the central atom are identical; if not, the molecule will be polar Images: Public Domain from Wikipedia.org (http://en.wikipedia.org/wiki/vsepr_theory) YBRIDIZATIN (7.3, 7.4) 2 domains hybridization s orbital p orbital 2 orbitals linear (180 ) 3 domains hybridization s orbital two p-orbitals 3 orbitals triangular (120 ) 4 domains hybridization s orbital three p orbitals 4 orbitals tetrahedral (109.5 ) Sigma (σ) Bond: This bond lies on the internuclear ais between the nuclei. All bonds (single, double, triple) contain a σ bond. Sigma (σ) bonds in this course are usually created from the overlap of two hybrid orbitals,, (e.g., ) or a 1s and hybrid orbital, 1s, when is part of the bond (e.g., ). Pi (π) Bond: This bond does not lie on the internuclear ais, and is present with double and triple bonds. z y z y y y p Pi (π) bonds are made from overlap of 2 p-atomic orbitals; p p z p z π bond Single bond = σ bond; Double bond = σ bond π bond; Triple bond = σ bond 2π bond
INTERMLEULAR FRES (7.6) (IMF): Forces between molecules, not within a molecule. -bonding: riteria: bonded to N,, or F - look for groups; eamples: 3 (l) (alcohols); 2 (l); 3 (l) (carboylic acids); -bonding IMF are usually stronger than Dipole-Dipole and LDF. : : Dipole Dipole: riteria: polar molecule; eamples: F l or F Br; polar molecules: chemicals with, S, N, and P are polar; organic compounds with carbonyl group (see diagram) are polar; nonpolar molecules: hydrocarbons, y London Diersion Forces (LDF) = Induced Dipole = Instantaneous Dipole: riteria: has electrons; occurs in all molecules and individual atoms; caused by electrons shifting to one side of the molecule/atom creating an instantaneous dipole within the molecule; hydrocarbons have only LDFs #e - LDFs DETERMINE WI MLEULE AS LARGEST IMFs If two molecules have the same IMFs (both molecules have -bonding or only Dipole-Dipole or only LDFs) determine which molecule has greater IMFs as follows: a. -bonding: 1. More atoms bonded to N,, or F atoms more -bonding larger IMFs b. Dipole-dipole: 1. Polar bond has greater ΔEN greater permanent dipole-dipole larger IMFs c. LDFs: 1. one molecule is larger more electrons greater LDFs larger IMFs 2. hydrocarbons: one hydrocarbon molecule is less branched greater LDFs larger IMFs IMFs : generally -Bonding > Dipole Dipole > LDF IMFs Boiling Point ; Melting Point ; Δ vap SLUBILITY: Liquids dissolving into liquids = miscibility (page 565-567): In general: Like likes Like polar substances dissolve into polar substances (dissolve = miscible); nonpolar substances (e.g., hydrocarbons!) dissolve into nonpolar substances (dissolve = miscible); polar and nonpolar substances don t dissolve into one another (don t dissolve = immiscible) more -bonding regions more polar more soluble in water 1. What is the electron region geometry and hybridization around the bold underlined atom of the molecules below? (int: Draw the Lewis dot structure for each molecule.) a. 4 b. 3 3 c. 2 4 d. 2. Draw Lewis dot structures for each of the following molecules. rganic structures will usually have the following bonding: 4 bonds; N 3 bonds; 2 bonds a. 2 6 b. 2 4 c. 3 2 d. 3 2 2 3 e. 6 6 (benzene = 6-membered carbon ring) 3. I. Which molecule below has the strongest bond? (int: Draw the Lewis dot structure of each molecule) II. Which molecule below has the longest bond? a. 3-2 b. 3 2 c. d. 2 ( is central atom) e. Same strength/length.
4. Identify the polar regions in this molecule by labeling atoms of the molecule with δ - and δ. 5. Answer the questions below about the structure and bonding in this molecule. (int: Draw in the lone pair e -.) a. What is the angle between the -- labeled with an a? b. What is the electron region geometry around the carbon atom labeled b? a c. What is the molecular geometry around the oygen atom labeled c? d. What is the hybridization of the carbon atom labeled a? e. What orbitals overlap to make the bond between the carbon atom labeled d d and the attached hydrogen atom? f. What orbitals overlap to make the pi (π) bond between the carbon atom labeled a and the oygen atom attached to it? c g. What is the electron region geometry around the carbon atom labeled d? h. ow many lone pairs of electrons are there in the molecule? b 6. Answer the following questions for the molecule to the right. a. Draw in the lone pairs of electrons in the structure. b. What is the bond order for the bond labeled a? c. What is the hybridization on the S labeled b? d. What is the angle for the on the labeled c? e. What orbitals are used to form the bond labeled e? f. ow many σ and π bonds are in the molecule? g. What is the electron region geometry around the S atom labeled b? h. What orbitals are used to form the π bond labeled a? i. What orbitals are used to make the σ bond labeled d? b e S N d a c 7. For each set of chemicals, order them from lowest to highest boiling point. I. 4, 3 8, 2 6 II. 3 2 Br, 3 2 F, 3 2 l III. 2, 3, 3 3 8. Place the molecules in order of most soluble to least soluble in water. I. 3 2 2 II. 3 2 2 2 2 3 III. 2 2
ANSWERS 1. a. tetrahedral, b. tetrahedral, c. triangular, d. linear, 2. a. b. c..... d..... e. 3. I. c {B = 3 for structure c} II. b {B = 1 for structure b} 4. The higher electronegativity of the atoms pull the electrons towards those atoms and concurrently create a positive region at the atom and atom that is attached to the atom. 5. a. 120 b. tetrahedral c. bent d. e. σ: 1s() () f. π: p() p() g. triangular h. 6 δ δ δ δ : S: N : : 6. a. b. 2 c. d. <109.5 e. 1s() (N) f. 22σ, 2π g. tetrahedral h. π: p() p(); σ: () () 7. I. 4 < 2 6 < 3 8 {as mass LDF IMF boiling points } II. 3 2 F < 3 2 l < 3 2 Br {as mass LDF IMF boiling points } III. 3 3 < 3 < 2 {IMFs Boiling Point ; 3 3 has dipole-dipole IMFs; 3 has -bonding; 2 also has -bonding but has 2 atoms that can -bond, so it has greater -bonding while 3 only has 1 that can undergo -bonding; 3 3 has weaker IMFs and the lowest bp; 3 is net because it has -bonding but less -bonding than 2 ; 3 3 < 3 < 2 } 8. most soluble to least soluble in water: III > I > II Water is polar and substances that are polar will be soluble (miscible) in water. Both chemicals I and III are polar because of the polar groups. Since chemical III has two groups it is more polar than chemical I with only one group. The more polar a substance is the more soluble it will be in another polar substance such as water. hemical II is a hydrocarbon and is therefore nonpolar and least soluble in water.