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Notes for Week 1

by: Bethany Lawler

Notes for Week 1 Chem 345

Bethany Lawler
GPA 3.96
Organic Chemistry 1
Dr. Crouch

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About this Document

Notes on the parts of Chpt. 1 we covered in lecture. Most of the material is focused on how and why bonds form. Not the entirety of Chpt. 1
Organic Chemistry 1
Dr. Crouch
Class Notes
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This 4 page Class Notes was uploaded by Bethany Lawler on Friday August 28, 2015. The Class Notes belongs to Chem 345 at Washington State University taught by Dr. Crouch in Summer 2015. Since its upload, it has received 49 views. For similar materials see Organic Chemistry 1 in Chemistry at Washington State University.


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Date Created: 08/28/15
Chemistry 345 Week 1 824 829 Metals and Nonmetals 1 The periodic table is made up of metals and nonmetals 2 Metals tend to give up electrons when forming a bond 3 Nonmetals tend to attract electrons when forming a bond Bonding There are two types of bonding Ionic and Covalent 1 Ionic Bonding 0 An ionic bond is formed between a metal and a nonmetal o In an ionic bond electrons are pulled from the metal and pulled to the nonmetal o The two atoms become charged 2 Covalent Bonding o Formed between two nonmetals 0 The electrons are pulled by both atoms so the electrons are shared between them 0 There are two types of covalent bonds polar and nonpolar I Polar bonds are formed when both atoms in the bond have an equal or nearly equal pull on the electron The electron is shared equally between the atoms II Nonpolar bonds are formed when one atoms in the bond pulls more than the other The atom tends to be closer to the atom with the larger pull giving it a slightly negative charge 0 Both forms of covalent bonding are dependent on the electronegativity of the atoms in the bond 0 When a bond is polar it has dipoles The side that pulls more strongly on the electrons has a partial negative charge and the side that has less pull has a slightly positive charge 0 The polarity of bonds in a molecule can also give the overall molecule a dipole Formal Charge 1 Formal charge is a way of keeping track of the overall charge on an atom 2 Formal charge can be computed in the following manner FCVE LP 12CBE FC Formal Charge VE The number of valence electrons the atom has ex Oxygen has 6 LP The number electrons in lone pairs CBE The number of electron in a covalent bond ex There are 2 in a single bond Bond Formation 1 2 The number of valence an electron has determines how many bonds a neutral atom will make Carbon 0 In order to remain neutral carbon must make four bonds 0 If carbon makes only three bonds it will have a 1 charge 0 If carbon makes five bonds it will have a 1 charge 0 If carbon has a radical is has a neutral charge Hydrogen 0 Molecular hydrogen is formed by two hydrogen which are joined by a single bond Both atoms in the molecule are neutral 0 Atomic hydrogen is neutral 0 A proton a form of hydrogen has a 1 charge 0 Hydride is a hydrogen atom with two electrons It has a charge of 1 Nitrogen 0 Neutral nitrogen forms three bonds 0 Nitrogen that forms two bonds as a 1 charge and is known as an amide ion 0 Nitrogen that forms four bonds has a 1 charge and is known as an ammonium ion Oxygen 0 Neutral oxygen forms two bonds 0 Oxygen that forms one bond has a 1 charge and is known as a hydroxide ion 0 Oxygen that forms three bonds has a 1 charge and is known as a hydronium ion 0 The suffix onium indicts that it has a positive charge Halogens o The halogens remain neutral when they only form one bond Lewis Structure 1 2 The basic symbols 0 represent a lone pair of electrons o represents a single bond 0 represents a double bond 0 E represents a triple bond Organic Conventions 0 Because carbon and hydrogen are used so frequently in organic chemistry the atomic symbols for those two symbols are often excluded from diagrams o It is assumed when two lines representing bonds meet in a corner unless otherwise specified there is a carbon atom there 0 Similarly it is assumed that unless otherwise specified all available bonds are filled with hydrogens so that the atom is neutral 0 All other conventions follow the Lewis dot model Atomic and Molecular Orbitals 1 Structure of an Atom Protons positively charged particles in the nucleus Neutrons noncharged particles in the nucleus Electrons negatively charged particles that orbit the nucleus The Atomic Number is the number of protons in the nucleus 2 Isotopes Isotopes are atoms of the same element that have different numbers of neutrons Isotopes have the same atomic number but different atomic weights Some isotopes are stable and do not emit radioactive particles Some isotopes are unstable and emit radioactive particles 3 Atomic Orbitals Electrons are distributed into shells Within each shells there are several orbitals labeled s p d and f Each of these orbitals have a different level of energy as follows Lowest energy gt s lt p lt d lt f lt highest energy Within the orbitals are orbitals for the electrons to populate There are three rules for populating atomic orbitals I Aufbau s Principle The lowest energy orbitals are filled first II Pauli Exclusion Princple There can only be two electrons in a unique orbital Hund s Rule The two electrons in a unique orbital must have opposite sp1n III Nodes are areas where there is no probability for an electron to exist All atomic orbitals with the exception of the ls orbital have nodes s orbitals have nodes between lower and higher level s orbitals p orbitals have a nodal plain between the tear drop shaped orbitals on opposite sides of the nucleus 5 Molecular Orbitals When atomic orbitals overlap molecular orbitals are formed Molecular bonds form at the lowest potential energy which is determined by the distance the electrons are from the nucleus of the atom they orbit and the atom they are bonding with and the distance the nuclei are from each other When bonding the number of molecular orbitals is the same as the number of atomic orbitals that went into making them There are bonding orbitals in which the electrons form constructive interference which creates a bond There are antibonding orbitals in which the electrons form destructive interference which keeps the atoms from forming a bond When antibonding orbitals are filled either entirely or partially they destabilize the bond that may have formed and potentially destroy it 6 Names of Bonds When two orbitals form a single bond this is called a sigma bond Sigma bonds have the ability to rotate about an internuclear axis When four orbitals form two bonds one is a sigma bond but the other is formed by a set of perpendicular orbitals These bonds are called pi bonds Pi bonds cannot rotate around an internuclear axis 7 Molecular Geometry and Hybridization When an atom such as carbon has empty orbitals that it wants to fill it will try and form bonds To form a bond the atom will form a hybrid orbital In the case of carbon one of the electrons from the s orbital will be promoted to the empty p orbital The four half empty orbitals will be mixed together so that they have the same energy The new hybrid orbitals are named after the atomic orbitals that they are made up of in the case above the four hybrid orbitals would be called sp3 orbitals The new orbitals have a different shape than the atomic orbitals that make them up The s orbital will combine with the half of the p orbital that shares the same phase via constructive interference and make it bigger The p orbital of the opposite phase will be made smaller by destructive interference The new orbitals are arranged equidistant around the nucleus which forms a tetrahedron All molecular geometry is formed by orbital hybridization


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