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TAMU CHEM 227 - Exam 1 Study Guide

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Chem 227 SantanderExam # 1 Study Guide Lectures: 1-6Chapter 1What is organic chemistry?-The study of carbon compounds-Carbon is a group 4A element, it can share 4 valence electrons and form 4 covalent bonds1.1 Atomic structure1.2 Atomic structure: orbitals- Atomic orbitals: The mathematical function that describes the wave-like behavior of eitherone electron or a pair of electrons in an atomo In other words: the areas in an atom where electrons can be found around the nucleusShapes of atomic orbitals:- In organic chemistry, we focus mainly on s and p orbitals- Electron shells: orbitals in an atom are organized into different “shells,” centered around the nucleus and having successively larger size and energy.o The first shell contains only a single s orbital, denoted 1s, and holds only 2 electronso Second shell contains one 2s orbital (2 electrons) and three 2p orbitals (6 electrons) and thus holds a total of 8 electrons- P orbitalso Three different p orbitals are oriented in space along mutually perpendicular directions, denoted px , py , and pz.o The two lobes of each p orbital are separated by a region of zero electron densitycalled a node.1.3 Atomic Structure: Electron configurations- Atoms will always favor the lowest energy arrangement, or ground-state electron configuration.- We can predict arrangement of electrons with these ruleso 1. Aufbau principle: The lowest energy orbitals fill up first, according to order 1s2s2p3s3p4s3d, o 2. Pauli Exclusion principle: Electrons act in some ways as if they were spinning around an axis… This spin can have two orientations, an up or down arrow. Only two electrons can occupy an orbital, and they must be of opposite spin.o 3. Hund’s rule: if two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full.1.4 Development of Chemical Bonding Theory- Valence shell: the outermost shell- Covalent bond: shared electron bond- Molecule: neutral collection of atoms held together by covalent bonds- Lewis Structures (electron dot): show valence electrons as dots- Kekule Structures (line-bond): have a line drawn between 2 atoms indicating a 2 electron covalent bond- Valence electrons not used are called non-bonding electrons1.5 The Nature of Chemical Bonds: Valence bond theory- Valence bond theory: a covalent bond forms when two atoms approach eachother closelyand a singly occupied orbital on one atom overlaps a singly occupied orbital on the otheratom. The electrons are now paired in the overlapping orbitals and are attracted to the nuclei of both atoms, thus bonding the atoms together.- Sigma bonds (σ)o The head-on overlap of two atomic orbitals along a line drawn between the nuclei- Bond strengtho H-H bond in an H2 molecule has a bond strength of 436 kJ/mol, in other words, we would have to put 436 kJ/mol of energy into the H-H bond to break the H2 molecule apart into H atoms- Bond lengtho How close are the two nuclei in the H2 molecule? If they are too close, they will repel eachother, but if they’re too far apart they won’t be able to share the bonding electrons.o The optimum distance between nuceli that leads to maximum stability is called the bond lengtho The distance is 74 pm in the H2 molecule1.6 sp3 Hybrid Orbitals and the structure of Methane- Because carbon uses two kinds of orbitals for bonding, 2s and 2p, we might expect Methane (CH4) to form two kinds of C-H bonds- In fact, all four C-H bonds in methane are identical and spatially oriented toward the corners of a regular tetrahedron- An s and 3 p orbitals can hybridize to form four equivalent atomic orbitals with tetrahedralorientationo These are called sp3 hybrids *Note that superscript 3 in name sp3 tells how many of each type of atomic orbital combine to form the hybrid, not how many electrons occupy it- The concept of hybridization explains how carbon forms four equivalent tetrahedral bonds– unsymmetrical about the nucleus. One of the two lobes is larger than the other and can therefore overlap more effectively with an orbital from another atom to form a bond- Bond angle: The angle formed in the bond o Each H-C-H is 109.5º or the tetrahedral bond angle1.7 sp3 hybrid orbitals and the structure of ethane- Same kind of orbital hybridization that accounts for methane structure also accounts for bonding together of carbon atoms into chains and rings- Structure of ethane with sp3 orbitals:1.8 sp2 Hybrid orbitals and structure of Ethylene- Occur when there is a C-C double bond involved - sp2 Hybrid orbitals are ina plane with 120º angles- remaining p orbital remains unchanged is perpendicular to the plane- when two carbons of sp2 hybridization approach eachother, they form a strong sigma bond by sp2-sp2 head-on overlap. At the same time, the unhybridized p orbitals interact by sideways overlap to form what is called a pi (π) bond. Results in formation of carbon-carbon double bond.- Carbon-Carbon double bond in ethylene is both shorter and stronger than the single bond in ethane because it has four electrons bonding the nuclei together rather than two.1.9 sp Hybrid Orbitals and the structure of Acetylene- Occur when there is a C-C triple bond involved- two sp hybrid orbitals are obtained - two p orbitals are not used in this hybridization- sp orbitals are linear, 180º- two sp orbitals form sp-sp sigma bond- pz orbitals form a pz-pz π bond by sideways overlap and py orbitals overlap similarly- net effect: sharing of six electrons and formation of C-C triple bond1.10 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur- H-N-H bond angle in ammonia (NH3) is 107.3º- One sp3 orbital is occupied by two non-bonding electrons- It is like this in N, O, P, SEx: MethylamineEx: WaterMore examples:1.11 Describing Chemical Bonds: Molecular Orbital Theory- Molecular Orbital Theory (MO): describes covalent bond formation as arising from a mathematical combination of atomic orbitals (wave functions) on different atoms to form molecular orbitals, so called because they beong to the entire molecule rather thanto an individual atom.o Where are the electrons around the entire molecule??- Bonding MO: Additive combination and is lower in energy- Antibonding MO: subtractive combination and is higher in energyEx: H2 - Pi bonding MO: form combining p orbital lobes with the same sign- Pi antibonding MO: form combining lobes with opposite signs- Only bonding MO is occupiedEx: Ethylene


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