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Chapter 1 Notes

by: Rebeka Jones

Chapter 1 Notes CHMY 321-001

Rebeka Jones
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These are the notes from Chapter 1 in the textbook. Includes table to help understand electronegativity and bonding types. Also a table to help with identifying hybridization and geometry.
Organic Chemistry I
Holmgren, Steven
Class Notes
Organic Chemistry




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This 10 page Class Notes was uploaded by Rebeka Jones on Monday September 19, 2016. The Class Notes belongs to CHMY 321-001 at Montana State University taught by Holmgren, Steven in Fall 2016. Since its upload, it has received 4 views.


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Date Created: 09/19/16
Chapter 1 Constitutional isomers have different physical properties and different names but same molecular formula. Tetravalent – forms 4 bonds (C) Trivalent – forms 3 bonds (N) Divalent – forms 2 bonds (O) Monovalent – forms 1 bond (H, F, Cl, Br, or I) Covalent bond – Two atoms sharing electrons -ΔH = lowering of energy The nucleus of an atom is comprised of protons and neutrons. Each proton has a charge +1 and neutrons are neutrally charged. For a neutral atom the number of protons is balanced by an equal number of electrons which have a charge of -1 and exist in shells the first shell (closest to the nucleus) holds 2 electrons and the second shell can hold up to 8 electrons. The electrons in the outer most shell of an atom are called the valance electron. The number of valance electrons in an atom is identified by its group number in the periodic table. The Lewis dot structure of an individual atom indicates the number of valence electrons. Step 1: Determine the number of valance electrons Step 2: Place on valance electron by itself on each side of the atom. Step 3: If the atom has more than 4 valance electrons the remaining electrons must be paired with the already existing electrons. Lewis Dot Structures of atoms can be combined to create small molecules. When doing so follow the octet rule. Step 1: Draw all individual atoms. Step 2: Connect atoms that form more than one bond. Step 3: Connect single bonding elements. Step 4: pair any paired electrons to form an octet. A formal charge is associated with any atom that does not exhibit the appropriate number of valence electrons. When this exists in Lewis dot structures the charge must be drawn. Step 1: Determine the appropriate number of valance electrons for an atom. Step 2: Determine whether an atom exhibits the appropriate number of electrons. If there is an atom that has more electrons than it should assign a formal charge. Bonds are classified into 3 groups Covalent Polar Covalent Ionic These groups arise from the electronegativity of the atoms sharing a bond. 2 Electronegativity generally increases as you go across the periodic table left and up. If the difference in the electronegativity of the two atoms is less than 0.5 the electrons are said to be equally shared resulting in a covalent bond. If the difference in electronegativity is between 0.5 and 1.7 the electrons are not shared equally resulting in a polar covalent bond. The withdrawal of electrons toward one atom is called induction and is indicated by an arrow showing where the electrons are coming from and where they are going. Induction causes areas of positive changes and negative charges. This is symbolized by delta. If the difference in electronegativity is greater than 1.7 the electrons are not shared at all resulting in an ionic bond. Ionic bonds are the result of the force of attraction between two oppositely charged ions. These values are simply guide lines many bond fall between groups. Step 1: Identify all polar covalent bond. Step 2: Determine the direction of each dipole. Step 3: Indicate the location of partial charges. H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 3 K Br 0.8 2.8 An orbital is a region of space that can be occupied by one electron -Electron clouds only come in a small number of shapes and sizes -An electron cloud is a single entity and is not comprised of billions of particles. It can be thicker in some places and thinner in others. -an electron cloud does not have defined edges. Electron density is associated with the probability of finding an electron in a particular region in space. The shape of an orbital refers to the region that contains 90-95% of the electron density. Beyond this region the other 5-10% of the electron density tappers off but never ends. If we want to consider 100% of the density, we must consider the entire universe. In summary, we must think of an orbital as a region of space that can be occupied by electron density. An occupied orbital must be treated as a cloud of electron density. This region is called an atomic orbital. Examples of orbitals are s, p, d, and f orbitals. Orbitals with the same energy level are called degenerate orbitals. The order in which the orbitals are filled by electrons is determined by just three simple principals. 1. The Aufbau Principle – the lowest energy orbital is filled first. 2. The Pauli Principle – each orbital can accommodate a maximum of two electrons that have opposite spins. 3. Hund’s Rule – When dealing with degenerate orbitals such as p orbitals, one electron is placed in each degenerate orbital first, before electrons paired up 4 Constructive interface produces a wave with larger amplitude. In contrast destructive interference results in waves canceling each other out which produces a node. According to the valence bond theory, a bond is simply the sharing of electron density between two atoms as a result of the constructive interface of their orbitals. A sigma bond is characterized by a circular symmetry with respect to the bond axis -all single binds are sigma binds Molecular orbital theory also describes a bond in terms of constructive interference between two overlapping atomic orbitals. But it also says that the orbitals are mathematically combined to produce new orbitals called molecular orbitals. Atomic orbitals are regions of space associated with an atom whereas molecular orbitals are associated with the while molecules. The atomic orbitals are combined mathematically to produce two molecular orbitals. Bonding MO is the result of destructive interference. The high energy orbitals are the antibonding MO. It is the result of destructive interference. Methane Pauling mathematically average, or hybridized, the 2s orbitals, giving four degenerate hybridized atomic orbitals. This gives us four orbitals that were produced by averaging one s orbital and 3 p orbitals and therefore we refer to these atomic orbital as Sp -hybridized orbitals. The four sp -hybridized 5 orbitals are equal in energy and therefore position themselves as far apart as they can to create tetrahedral geometry. Hybridized orbitals have a larger front lob and smaller back lobes. The larger front lobes enable hybridized atomic orbitals to be more efficient than p orbitals in their ability to form bonds. Ethylene In ethylene’s case each carbon only needs to make three bonds. Therefore, each carbon only needs 3 hybridized orbitals. So we will mathematically average the s orbital with only two of the 3 p orbitals. The remaining p orbital will remain unaffected. This results in a carbon atom with one p p 2 orbital and 3 sp -hybridized orbitals. When p orbitals overlap they create a pi bond. The bond forms above the plane of the molecule and below the plain. Sigma bonds experience free rotation at room temperature whereas pi bonds do not. A triple bind is formed by sp-hybridized carbon atoms one s orbital is mathematically averaged with only one p orbital. This lea ves 2 p orbitals unaffected. Therefore, there is 2 sp orbitals and two p orbitals. The sp orbitals can form sigma bonds and the two p orbitals form pi bonds. A triple bond between two carbon atoms is therefore the result of three separate bonding interaction one sigma bond and two pi. 6 3 A carbon atom with four single bond will be sp hybridized. A carbon atom with three sigma bonds and one pi bond will be sp hybridized. A carbon atom with two sigma bonds and two pi bond will be sp hybridized. In order to predict the geometry of a bond we must count the number of sigma bonds and lone pairs. This is the steric number. It represents the number of electron pairs that are repulsing each other. This is VESPER theory. There are 3 different types of geometry arising from sp hybridization tetrahedral, trigonal pyramidal, bent. In all cases the electrons were arranged in a tetrahedron but the lone pair were ignored when describing geometry. Sp -hybridization orbital achieves maximal separation in trigonal planer. Can also be bent. Sp hybridized orbitals have maximum separation when they are linear. If steric number is 4 If the steric If steric number is number is 3 2 3 2 Sp Sp Sp Tetrahedral arrangement of Trigonal Planar Linear electron pairs arrangement of arrangement of electron pairs electron pairs No lone One Two One No lone Linear pairs lone pair lone lone pairs pairs pair Tetrahedral Trigonal Bent Trigonal planar planar 7 It is not needed to describe the hydrogen atoms because they are monovalent Things are said to exhibit a dipole movement because the center of negative charge and the center of positive charge are separated from one another by a certain distance. The dipole movement is used as an indicat or of polarity, where it is defined as the amount of partial charge on either end of the dipole multiplied by the distance of separation. Most compounds will have a dipole movement of 10 -18 so it’s more convenient to report dipole moment with a new unit called a debye 1 debye = 10 esu x cm measuring the dipole moment allows us to calculate the % ionic character of that bond. When dealing with a compound that has more than one polar bond, it is necessary to take the vector sum of the individual dipole moments. This is called the molecular dipole moment, and it takes into account both magnitude and the direction of each individu al moment. The presence of lone pairs has a significant effect on the dipole moment. There is a dipole moment connected to each lone pair. Step 1: predict the molecular geometry – use steric number Step 2: Identify the direction of hall dipole moments. Step 3: Draw the net dipole moment. 8 Intermolecular forces- attractive forces between individual molecules. All intermolecular forces are electrostatic meaning that they occur as a result of the attraction between opposite charges for neutral molecules there are -dipole-dipole -hydrogen bonding -fleeting dipole dipole interactions In solid phase molecules align to attract each other whereas the liquid phase molecules are free and move more often to attract then to repel. More attraction = higher boiling point. Hydrogen bonding is a type of dipole-dipole interactions. Due to individual hydrogen has a partial positive charge. To understand fleeting dipole-dipole interactions consider the electrons to be in constant motion so the center of negative charge coincides with the venter of positive charge, resulting in a zero dipole moment. But the venters might not co-inside. This resulting transient dipole movement can create separate transient dipole moment in other molecules – this is called London dispersion forces. Larger molecules experience this more. Branching molecules have small surface areas. When comparing boiling points of compounds, we look for -are there any dipole-dipole interaction in either compound - will either compound form hydrogen interactions in either compound how much branching is in each compound. 9 Polar compounds are solvable in polar solvents and vice versa. 10


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