Chapter 1

1. Chapter 1 Carbon Compounds and Chemical Bonds

2. Organic Chemistry • The chemistry of the compounds of carbon History- -Unofficially, Organic is one of the oldest sciences -Officially, it is one of the youngest

3. History • Vitalism- the belief that the intervention of a “vital force” was necessary to the synthesis of organic molecules. • Friedrich Wöhler, 1828, “Father of Organic Chemistry”

4. On your own • Review section 1.2 to make sure you understand terms such as: • Compounds, elements, atoms, isotopes, electron shells, valence shell, valence electrons, etc.

5. Structural Theory • 1800’s by Kekulé, Couper, and Butlera • Two central premises • 1) The atoms in organic compounds can form a fixed number of bonds • Valence- the measure of an atoms ability to form bonds • 2) A carbon atom can use one or more of its valences to form bonds to other carbon atoms

6. Isomers • Isomers- Different compounds with the same molecular formula. ex. • Constitutional Isomers: Isomers that differ in their connectivity, that is, in the sequence in which their atoms are bonded together

7. Two types of Bonding • Ionic Bonds- Formed by the transfer of one or more electrons from one atom to another to create ions • Covalent Bonds- a bond that forms when atoms share electrons with one another. • Octet Rule- the tendency for atoms to achieve an electron configuration where its valence shell contains eight electrons.

8. Ionic Bonds • Occur between atoms of widely different electronegativity. • Typically, that means between metals and non-metals Ex.

9. Electronegativity • Electronegativity- a measure of the ability of an atom to attract electrons. • Periodic Table Trends • Increases as you move left to right • Increases as you move bottom to top • Important Order of Electronegativity: F > O > Cl,N > Br > C,S > H,P

10. Why Do Ions Form? • Electronically identical to Nobel Gas • Stabilize each other • Note: • Electrons dictate reactivity, not charges!!

11. Properties of Ionic Compounds • Very strong • High MP • Sometimes dissolve in Polar Solvents • Usually called Salts

12. Covalent Bonds • Occur when two atoms have the same or similar electronegativity • “Share” electrons instead of complete transfer • Typically called molecules • Molecules can be represented by electron dot formulas, or dash formulas, where a dash represents a pair of shared electrons. • Examples:

13. Lewis Structures • Lewis Structures are electron dot formulas where only the valence electrons are shown • Multiple bonds are represented by multiple lines. • Ex. • Note: Ions may also contain covalent bonds! • Ex.

14. Writing Lewis Structures • Assemble the molecule or ion showing only valence electrons • Strive to give each atom an octet, except Hydrogen which receives only 2 • The number of valence electrons for an atom is equal to the group number • If the structure is an ion, we add or subtract electrons to give appropriate charge

15. Rules: • Find total number of valence electrons • Use pairs of electrons to form bonds between the atoms. (Note: remember to consider the typical valence of each atom) • Add remaining electrons as pairs to give each atom an octet.

16. Exceptions to Octet Rule • 1st period elements- only two electrons • 2nd period elements- Boron is stable with six valence electrons • 3rd period and beyond- Elements have access to d orbitals which allows them to accommodate more than 8 electrons

17. Formal Charges • Formal Charge- the charge associated with the electronic difference between the atomic state and bonded state of an atom • The sum of the formal charges on each element equals the total charge for the molecule or ion. • To calculate formal charges, simply compare the number of electrons an atom has in the bonded state to the number it has in the atomic state.

18. Counting Valence Electrons in the Bonded State • Non-shared electrons, or lone pairs, belong solely to the atom that posses them • Shared pairs of electrons are split giving half the shared electrons to each atom sharing them. • Ex.

19. Common Formal Charge States

20. Resonance Structures • Lewis structures incorrectly create an artificial location for electrons • Consider the carbonate ion, CO32-

21. Resonance Structures, cont • We can account for the experimental data by showing how the structures can be converted to the others using curved arrows. • (arrow speech!) • The overall structure, or true structure, is a mixture of the individual structures, and is called the resonance hybrid. • Note: Individual resonance structures only exist on paper!!

22. Rules for Drawing Resonance Structures • Resonance Structures only exist on paper! • You are only allowed to move electrons. • All structures must be proper Lewis structures • The energy of the actual molecule will always be lower than a single contributing structure. This is called Resonance Stabilization.

23. Rules for Drawing Resonance Structures, cont 5) Equivalent resonance structures make equal contributions to the overall structure. 6) The more stable a structure, the more it will contribute to the overall structure: a) The more covalent bonds a structure has, the more stable it is.

24. Rules for Drawing Resonance Structures, cont b) Structures in which all atoms have octets are more stable c) Opposite charge separation decreases stability d) Structures with negative charges on highly electronegative atoms are more stable than those with negative charges on less electronegative atoms

25. Section 1.9 Quantum Mechanics and Atomic Structure WHAT YOU NEED TO KNOW: • Electrons in atoms and molecules have both particle and wave characteristics • Wave properties are used to predict the shape of orbitals • Orbitals- a region of space where the probability of finding an electron is very large • The volumes we use to express the orbital shapes represents where the electron would be 90-95% of the time.

26. Atomic Orbital shape

27. Energy of orbitals 1s < 2s < 2px = 2py = 2pz < 3s • Degenerate orbitals- atomic orbitals of equal energy, like the three 2p orbitals • Using these relative energies, we can derive the electron configuration

28. Electron Configuration 1) Aufbau Principle: Orbitals are filled so that those of lowest energy are filled first. 2) Pauli Exclusion Principle: Only two electrons are allowed in each orbital and must have opposite spins 3) Hund’s Rule: When multiple orbitals of equal energy are present, each orbital receives one electron before any pairs are created.

29. Molecular Orbitals • Atomic Orbitals combine to form molecular orbitals which are used for bonding. • The number of Molecular orbitals must equal the number of Atomic orbital used. • In General Chemistry, we concentrated on the new Molecular orbitals formed from bonding

30. In organic, we use this theory to explain the combination of atomic orbitals on a single atom, called Hybridization, to form new orbitals that are then used for bonding. • Consider the simplest organic molecule, methane, CH4

31. Methane • To account for what we see, we must mix the 2s orbital with all three 2p orbitals. • This forms four new, equal orbitals called sp3 hybridized orbitals!

32. Sigma Bonds • When hybridized orbitals overlap, a sigma bond is formed • Sigma Bond- term used to describe bonds in which the greatest density of electrons lies between the two nuclei. • Sigma bonds have cylindrical symmetry along the bond axis. • As a result, there is “free” rotation about sigma bonds.

33. sp2 Hybridization • When two carbons share two pairs of electrons, the result is a carbon-carbon double bond. • Hydrocarbons whose molecules contain a carbon-carbon double bonds are called alkenes. • Pi bond- created by the overlap of p orbitals above and below the sigma bond framework.

34. Restricted Rotation • For maximum overlap, the p orbitals must be parallel • As a result, ~264 kJ/mol of energy is needed to break the p orbital overlap and allow rotation to occur • Only about 13-26 kJ/mol is needed to rotate around a sigma bond.

35. Cis/Trans Isomers • Because of the restricted rotation, a new form of isomers is created call Cis/Trans Isomers. • Ex. • These compounds are notsuperposable. • They are not constitutional isomers because the connectivity is the same

36. Stereoisomers • Stereoisomers- isomers that differ only in the arrangement of their atoms in space. • For a double bond to qualify for Cis/Trans Isomers, both carbons of the double bond must be bonded to two different groups. • Ex.

37. sp hybridization • When two carbons share three pairs of electrons, a triple bond is formed. • Hydrocarbons in which two carbons share three pairs of electrons are called alkynes. • The triple bond consists of 2 pi bonds and 1 sigma bond.

38. Bond Lengths • The shortest C-H bonds are achieved with hybridized orbitals having the most s character. sp < sp2 < sp3 % s 50% 33% 25%

39. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory is used to predict the geometry of atoms and the resulting shape of molecules. • It consists of four points: 1) We consider molecules (or ions) in which the central atom is covalently bonded to two or more groups.

40. VSEPR cont 2) We consider all of the valence electron pairs of the central atom -Shared electrons are called bonding pairs -Unshared electrons are called non-bonding or unshared pairs 3) Because electrons repel each other, they stay as far away from one another as possible Nonbonding repulsion > bonding

41. VSEPR cont 4) We arrive at the geometry of an atom by considering both bonding and nonbonding electrons but the shape of the molecule by referring to the position of nuclei. See table 1.3 on page 47.

42. Structural Formulas • Dash Formula- must show all atoms, bonds, and lone pairs. • Condensed Formula-must show all atoms but may or may not show bonds • Bond Line Formula- show bonds and all atoms except Carbon and Hydrogen

43. 3-D designations • Wedge bond- used to show bond coming out of plane towards the viewer • Dash bond- used to show bond going behind the plane away from the viewer. • These can be used with any type of structural formula.

44. Applications of Basic Principles • Page 47, section 1.17 • These sections are available at the end of most chapters. You should review these as they associate basic theories to topics in organic chemistry.

45. Other Tools • There are also other great tools at the end of the chapters such as Summaries and Review Tools, Practice Problems, Concept Maps, and Synthetic Connections.