8/24/15 Welcome to Organic Chemistry!
Chapter 1- Structure and Bonding
Ashley Piekarski, Ph.D What is organic chemistry?
•  Organic chemistry •  The study of carbon compounds
•  More than 99% of all known chemical
compounds contain carbon
2 1 8/24/15 Why is carbon so cool?
•  What group is carbon located on the periodic table? •  How many covalent bonds can it form? •  Carbon atoms can bond to one another forming long chains and rings •  Therefore, carbon can form a diversity of
compounds!
3 Why do I care, Dr. P?
•  By fully understanding the electronic structure of atoms and funcJonal groups, we can understand their reacJvity. •  We want to establish structure-­‐property relaJonships! 4 2 8/24/15 The design of the atom
•  What is located in the nucleus of an atom? •  What subatomic parJcles are located in a cloud outside the nucleus? •  An atom is neutral, so the number of electrons and protons in an atom are the _________. 5 Isotopes
•  Isotopes of an atom have the same atomic number (same number of protons) but different mass numbers (different number of neutrons) Z
A
X
Z = mass number A = atomic number X = element’s symbol 6 3 8/24/15 Where does the chemistry happen?
•  When we study chemical reacJons, we are studying the breaking and formaJon of bonds •  Which subatomic parJcle is responsible for the chemistry that occurs? 7 Orbitals
•  Wave funcJon (Ψ) is a mathemaJcal descripJon to describe the behavior of a specific electron •  The soluJon to this wave funcJon is an orbital •  If you plot Ψ2, you would obtain the three-­‐
dimensional space of the orbital surrounding the nucleus where the electron spends 90-­‐95% of its 4me 8 4 8/24/15 Orbitals
•  There are four types of orbitals: s, p, d and f •  In organic chemistry, only the s and p orbitals
are used primarily
•  What shape is the s orbital? •  What shape is the p orbital? •  Why are there 3 p orbitals?
9 Electron shells
•  Orbitals are organized into different layers, electron shells, of successively larger size and larger energy (n= 1, 2, 3 . . .) •  How many electrons can each orbital hold? 10 5 8/24/15 Nodes
•  Node is a region with zero electron density •  Each lobe has a different sign (+ or -­‐) which will have an effect on chemical bonding 11 Rules- Electron Configuration
•  Rule 1: Au8au Principle •  Lowest energy orbitals fill up first
•  Rule 2: Pauli Exclusion Principle •  Only 2 electrons occupy an orbital and they
must have opposite spins
•  Rule 3: Hund’s rule •  One electron occupies each orbital with spins
parallel until all orbitals are half-full
12 6 8/24/15 Learning check
•  Write the electron configuraJon for palladium. 13 Why do atoms bond together?
•  The resulJng molecule is lower in energy, and thus more stable, than the separate atoms •  Making bonds releases energy •  What type of reaction releases energy?
•  Breaking bonds always absorbs energy •  What type of reaction absorbs energy?
14 7 8/24/15 How do the bonds form?
•  Electronic properJes •  What type of configuration do atoms always
want to achieve?
•  What are the two types of chemical bonds you
learned about in general chemistry?
15 Ionic compounds
•  Ionic compounds achieve a noble gas configuraJon by complete electron transfer •  For example: 16 8 8/24/15 Covalent compounds
•  The energy it would take to lose or gain electrons to achieve an octet is too great for some atoms •  Therefore, some atoms share electrons (like carbon) to achieve the noble gas configuraJon •  A shared electron bond is called a covalent bond 17 Electron-dot structures
•  We can represent covalent compounds and how the electrons are shared by drawing electron-­‐dot structures •  Valence electrons are represented by the dots
•  The number of covalent bonds an atom forms
depends on how many additional valence
electrons it needs to reach a noble gas
configurations
•  How many valence electrons does carbon need?
18 9 8/24/15 Learning check
•  Draw the line-­‐bond structure for CO32-­‐. 19 Learning check
•  Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry. 20 10 8/24/15 Valence Bond Theory
•  A covalent bond forms when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom 21 Formation of H2
22 11 8/24/15 Sigma (σ) bonds
•  Sigma bonds are formed by the head-­‐on overlap of two atomic orbitals along a line drawn between the nuclei •  Bond strength is the amount of energy released to form that new bond or break that bond •  Example: H-H bond has a bond strength of
436 kJ/mol
23 Bond length
•  Bond length is the maximum distance between nuclei to achieve maximum stability 24 12 8/24/15 Hybrid orbitals
•  What is the electron configuraJon for carbon? •  How many valence electrons does carbon have and in what orbitals? •  All 4 bonds of methane are idenJcal… why? 25 sp3 orbital
•  In 1931, Linus Pauling showed mathemaJcally how s and p orbitals hybridized to form sp3 hybrid orbitals 26 13 8/24/15 sp3 orbital- How?
27 sp3 orbital- Why?
•  sp3 hybrid orbitals are unsymmetrical around the nucleus •  one of the lobes is larger than the other- why?
•  this improves overlap, which means hybrid
orbitals form stronger (more stable) bonds
28 14 8/24/15 sp3 orbital- bond angle
•  The four bonds of methane are all equal with a specific geometry, they orientate themselves around the central atom to achieve the maximum distance-­‐ bond angle •  What is the theory used to predict geometries and shapes around the central atom? 29 Ethane
•  What is the line-­‐bond structure for ethane? •  What type of bond is between the carbon-­‐
carbon bond? •  What are the hybrid orbitals orbitals in this molecule? 30 15 8/24/15 Ethane
31 Ethylene- sp2 orbital
•  What is the line-­‐bond structure for ethylene? •  2s orbital of carbon now combines with only two of the three available 2p orbitals-­‐ sp2 orbital 32 16 8/24/15 pi (π) bond
•  When two sp2-­‐hybridized carbons approach each other, they form a sigma bond (sp2-­‐sp2 head-­‐on overlap) •  The unhybridized p orbital approach for sideways overlapàpi bond 33 Ethylene
34 17 8/24/15 Ethylene
•  Bond strength: 728 kJ/mol (versus 377 kJ/mol for ethane) •  Why?
35 Acetylene
•  What is the line-­‐bond structure for acetylene? •  What is an sp orbital? 36 18 8/24/15 Acetylene
37 Acetylene
• 
• 
• 
• 
• 
sp-­‐sp sigma bond – from what overlap? pz-­‐pz pi bond py-­‐py pi bond How many electrons are shared? How do you think the strength of this bond compares to ethylene and ethane? 38 19 8/24/15 Summary
39 Learning check
•  Draw the structure for 1, 4-­‐pentadiene. •  What is the hybridization of each carbon?
•  What is the bond angle of each carbon?
40 20 8/24/15 Oh those other elements...
•  Nitrogen 41 Oh those other elements...
•  Oxygen 42 21 8/24/15 Oh those other elements...
•  Phosphorus and Sulfur •  3rd row analogs of nitrogen and oxygen
•  very common in biological molecules
•  both can be described by sp3 hybridization
43 MO Theory
•  Describes covalent bond formaJon as arising from a mathemaJcal combinaJon of atomic orbitals (wave funcJons) on different atoms to form molecular orbitals which belong to the enJre molecule rather than an individual atom 44 22 8/24/15 Rules for reading skeletal structures!
•  Rule 1: Carbon atoms aren’t usually shown. Instead, a carbon atom is assumed to be at each intersecJon of two lines (bonds) and at the end of each line. Occasionally, a carbon atom might be indicated for emphasis or clarity. 45 Rules!
•  Rule 2: Hydrogen atoms bonded to carbon aren’t shown. Since carbon always has a valence of 4, we mentally supply the correct number of hydrogen atoms for each carbon. 46 23 8/24/15 Rules!
•  Rule 3: Atoms other than carbon and hydrogen are shown. 47 Learning check
•  Draw the expanded structure for paracetamol, acJve ingredient in Tylenol. •  What is the molecular formula?
H
N
HO
O
48 24 8/24/15 Applications
•  Organic Foods: Risk vs Benefit •  PesJcides •  Advantage: help increase food production by
protecting crops from weeds, insects, and
fungi
•  Disadvantage: unknown health risks
49 Applications
•  An LD50 is a standard measurement of acute toxicity that is stated in milligrams (mg) of pesJcide per kilogram (kg) of body weight. •  An LD50 represents the individual dose required to kill 50 percent of a populaJon of test animals (e.g., rats, fish, mice, cockroaches). •  The lower the LD50 dose, the more toxic the pesJcide. 50 25 8/24/15 Applications
LD50 aspirin = 200 mg/kg (oral) LD50 atrazine = 672 mg/ kg to 3000 mg/kg (oral) The average human weighs 150 lb (68 kg) Need 124000 mg (124 g) for lethal dose What about long-­‐term effects? Can the pesJcide cause cancer or interfere with the development of a newborn? •  What do you think? Should the pesJcide be used to help grow our food? • 
• 
• 
• 
• 
51 26