Lect_W9_151-1 Weso (short)

Lect_W9_151-1 Weso (short) - General Chemistry CHEM 151...

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Unformatted text preview: General Chemistry CHEM 151 Week 9 UA GenChem REVIEWING CHEMICAL BONDING In a chemical reaction between two atoms, their valence electrons are reorganized. Atoms lose or share electrons because that leads to a more stable state (full electron shells). During the process, a net attractive force– A CHEMICAL BOND- occurs between atoms Li+ F- Ionic bond Electrostatic force UA GenChem + CHEMICAL BONDING + F-F Covalent bond The atoms in a molecular compound are connected by covalent bonds The number of covalent bonds that each atom forms is determined by the number of electrons that the atom must share to achieve a noble gas configuration (full shell) The tendency in molecular compounds is to have a structure in which each atom has eight valence electrons (OCTET RULE), there are exceptions (OCTET UA GenChem Electron Electron Distribution in Molecules Molecules • Electron distribution is Electron depicted with Lewis electron dot structures dot • Electrons are Electrons distributed as shared or BOND PAIRS and BOND unshared or LONE PAIRS. PAIRS. G. N. Lewis G. 1875 - 1946 1875 UA GenChem Bond and Lone Pairs • Electrons are distributed as Electrons • shared or BOND PAIRS and unshared or shared BOND . H .. Cl : .. LONE PAIRS. LONE This is called a This LEWIS ELECTRON DOT STRUCTURE DOT STRUCTURE UA GenChem Building a Lewis Structure Ammonia, NH3 1. Decide on the central atom; never H. 1. never Central atom tends to be the atom with the lowest Electronegativity. the Therefore, N is central (H cannot) 2. Count valence electrons 2. valence H = 1 and N = 5 Total = (3 x 1) + 5 (neutral) charge = 0 (neutral) = 8 electrons or 4 pairs electrons pairs UA GenChem Building a Lewis Structure 3. Form a single bond between the central atom and surrounding atoms. H NH H Each bond line represents two electrons 4. Remaining electrons form LONE 4. PAIRS to complete octet as needed. PAIRS (Start with terminal or outside atoms, but Start not if H; place any leftover electrons on the central atom) the H NH H UA GenChem .. •• 3 BOND PAIRS and 1 LONE PAIR. and LONE Lewis Structure: ClF21+ Step 1. Central atom Step 2. Count valence electrons Step Cl = 2xF = (Positive) charge = ? (Positive) TOTAL = Step 3. Form single bonds F Cl F 8 pairs of electrons pairs are now left. are UA GenChem Lewis Structure: ClF2+ Remaining pairs become lone pairs, first on Remaining outside atoms and then on central atom outside .. :F .. .. Cl F : .. .. .. + Ions need Brackets And the Correct Charge. Each atom now has 8 electrons. UA GenChem Hint: since H is always on the3 outside 2 making one bond try CH NH making Step 1. Central atom (s) Step 2. Count valence electrons Step C= N= N= 5xH= charge = TOTAL = TOTAL Step 3. Form single bonds Lewis Structure: CH5N H HCNH HH 1 pair of electrons pair are now left. are UA GenChem Lewis Structure: CH3NH2 Remaining pairs become lone pairs, first on Remaining outside atoms and then on central atom outside H HCNH HH Each atom now has 8 electrons. UA GenChem .. Carbon Dioxide, CO2 1. Central atom = 1. 1. Valence electrons = __ or 1. Valence __ pairs __ Charge = Charge 3. Form single bonds. 4. Place lone pairs on outer 4. atoms; place any leftover electrons on the central atom. atom. O C O .. :O .. C .. O: .. UA GenChem Carbon Dioxide , CO2 5. So that C has an 5. octet, we shall form DOUBLE BONDS DOUBLE between C and O. .. :O .. C .. O: .. Each atom has 8 electrons .. :O .. .. O .. C C .. O: .. .. O .. UA GenChem Nitrogen, N2 1. Central atom = 1. 1. Valence electrons = __ 1. Valence or __ pairs or Charge = Charge Total electrons= Total 3. Form single bonds. 4. Place lone pairs on outer 4. atoms; place any leftover electrons on the central atom. atom. NN : N N: .. UA GenChem .. Nitrogen, N2 3. Form single bonds. 1. Place lone pairs on outer 1. Place atoms; place any leftover electrons on the central atom. central 2. Arrange electrons to make an octet around each atom. each : N N: .. .. : N N: .. .. :N N: UA GenChem Need TRIPLE BOND Double and even triple bonds Double are commonly observed for C, N, P, O, and S and H HH H HC H CCCCCCH H :N .. O .. C N: C2F4 .. O .. UA GenChem Even more . H2CO C H▬O▬H O H▬C▬H Possible skeletons C▬H▬O▬H Formal Charge= group #– unshared e– # bonds= Rules for formal charge – 1) Keep as close to zero as possible 2) If greater than zero spread out as far apart as possible 3) Keep more positive element with pos. formal charge keep more negative elements with neg. formal charge UA GenChem Which structure? 1. 2. 3. 4. A B C Can’t choose UA GenChem Hydrazine . N2H4 UA GenChem Is there a pattern for C, N, O and H? H HC H HH H H CCCCCCH Typically carbon makes ___ bonds with ___ unshared pairs Typically hydrogen makes ___ bonds with ___ unshared pairs : O: HC .. O .. H UA GenChem Typically oxygen makes ___ bonds with ___ unshared pairs Is there a pattern for C, N, O and H? H N N N C N: H Typically Nitrogen makes ___ bonds with ___ unshared pairs Element H O N C Bonds (pairs) 1 2 3 4 Lone pairs 0 2 1 0 UA GenChem .. .. .. Types of Bond Names The first bond between atoms is a sigma (σ) bond .. :O .. C .. O: .. .. :O .. C .. O: .. .. O .. C .. O .. The second/ third bond between two atoms is a pi (π ) bond UA GenChem Identify the total number of σ and π bonds :C H N: H .σ = ____ .π = ____ : F: H H .σ = ____ .π = ____ .. CCCC H .σ = ____ .π = ____ CH3NH2 UA GenChem Ozone, O3 1. Central atom = 1. 1. Valence electrons = __ or 1. Valence __ pairs __ Charge = Charge Total electrons= 3. Form single bonds. 4. Place lone pairs on outer atoms; place any leftover electrons on the central atom. atom. OOO .. .. .. : O O O: .. .. UA GenChem Ozone, O3 Ozone, .. .. .. : O O O: .. .. .. .. .. : O O O: .. .. . .. .. .. : O O O: .. .. OR .. .. .. : O O O: .. .. .. .. .. O : .. O O : UA GenChem .. .. .. : O O O: .. OR Experimental data Single bonds ( ― ) are longer than double bonds ( = ) which are longer than triple bonds (≡) .. .. .. : O O O: .. OR Experimental evidence .. .. .. O : .. O O : Identical Bond Lengths Longer than single bonds but shorter than double bonds! UA GenChem Resonance . .. .. .. .. .. .. O : O O O : ↔ : .. O O : .. O3 has two equivalent structures These equivalent structures are called These RESONANCE STRUCTURES. The true RESONANCE The electronic structure is a HYBRID of the two. HYBRID Neither of the individual structures is correct, but these are the best we can do with dots and sticks. UA GenChem Resonance In resonance structures the placement of the atoms is the SAME, but the distribution of the electrons changes. 3 resonance structures NO3 1- .. :O .. N .. O: .. .. O .. N .. O: .. .. : .. O N O: .. : O: : O: .. : O: .. In the actual molecule, the electrons are “delocalized” over the entire molecule. One could say there are 4/3 bonds between each pair of atoms. How many resonance structure does the carbonate (CO32-) ion have? UA GenChem The Answer .. .. 2: O : : O: C :O : .. .. O .. C .. O: .. .. :O .. C .. O .. : O: : O: C .. .. :O : .. Average C to O bond order? :O : .. :O : UA GenChem 4 bonds/ 3 locations = 1.33 or 4/3 Organic chemists’ shorthand for carbon and Hydrogen • C2H4 • CH3CH2CH2OH UA GenChem Resonance When electrons become delocalized, they occupy a larger volume, which reduces electron-electron repulsion and thus stabilizes the molecule Experimental Electron Density Map Benzene C6H6 H H H H C C C C H C C H H H H C C H C C H H C C H H H C C C C H C C H H Does the Electron Density map show single and double bonds? UA GenChem Resonance Super Shorthand Super, super Shorthand Organic chemist’s shorthand UA GenChem Violations of the Octet Rule Usually occur with Boron, elements of higher periods, or compounds of noble gases. noble SF4 Too many electrons .. :F .. B .. F: .. .. Too few electrons :F: .. BF3 F F Free Radical S F F O-N=O Odd # e- NO2 UA GenChem . Do Lewis for SF4 and SF6 SF4 Too many electrons .. F F S F F UA GenChem .. :F .. B F : .. : .. F: .. BF3 . UA GenChem Charge Distribution Once molecular structure is established, chemist are very interested in analyzing electron distribution in a molecule because it affects its stability and its reactivity. How do they do it? 1. Determine bond polarity and assign partial charges 2. Calculate charge distribution in a bond in two extreme cases (hypothetical): When electrons are equally shared (Formal Charge) When the most electronegative atoms takes all the electrons in a bond (Oxidation number). UA GenChem Formal Charge Consider the Ozone molecule (O3): O-O=O Formal Charge Compare the number of valence electrons the atom would have by itself to the number of electrons it actually has (Assuming the bonding electrons are equally divided) 6-7=-1 Valence electrons from group number on periodic table 6-6=0 O-O=O 6-5=+1 Total formal charge -1 + 1 + 0 = 0 UA GenChem Formal Charge Consider the ion NCO-: 5-7=-2 4-4=0 6-5=+1 5-6=-1 4-4=0 6-6=0 5-5=0 4-4=0 6-7=-1 [ N-C=O ]-1 [ N=C=O ]-1 [ N=C-O ]-1 Assign formal charges to each atom in each resonance structure. Determine the most stable structure based on these criteria: Smaller formal charges (whether positive or negative) are preferable to larger ones; Like charges on adjacent atoms are not desirable; A more negative formal charge should reside on a more electronegative atom. UA GenChem Formal Charge Consider the ion NCO-: 5-7=-2 4-4=0 6-5=+1 5-6=-1 4-4=0 6-6=0 . 5-5=0 4-4=0 6-7=-1 [ N-C=O ]-1 [ N=C=O ]-1 [ N=C-O ]-1 Which structure is the most stable? Why? UA GenChem Oxidation Number [ N-C=O ]-1 [ N=C=O ]-1 [ N=C-O ]-1 Oxidation Number Compare the number of valence electrons the atom would have by itself to the number of electrons it would have if the most electronegative atom in a bond takes both electrons. 5-2-6=-3 6-6-2=-2 The oxidation numbers are equal to the expected charges if the compound was ionic UA GenChem [ N=C-O ]-1 4-0-0=+4 Putting it All Together Build the Lewis structures; Identify number of bonding and lone pairs; Identify possible resonance structures; Identify Assign Bond Polarity; Assign formal charges and oxidation numbers; SCNI3 UA GenChem ...
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This note was uploaded on 03/30/2009 for the course CHEM 151 taught by Professor Staff during the Spring '08 term at University of Arizona- Tucson.

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