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Exam 1 Study Guide

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by: Jessica Brown

Exam 1 Study Guide chem 10061-001

Jessica Brown

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Important topics for exam 1
general chemistry 2
David bowers
Study Guide
Chemistry, Science, General Chemistry
50 ?




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This 9 page Study Guide was uploaded by Jessica Brown on Wednesday February 17, 2016. The Study Guide belongs to chem 10061-001 at Kent State University taught by David bowers in Summer 2015. Since its upload, it has received 125 views. For similar materials see general chemistry 2 in Chemistry at Kent State University.


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Date Created: 02/17/16
General Chemistry 2 Exam 1 Study Guide Chapter 15 Organic Chemistry: the study of carbon compounds  Mostly deals with hydrocarbons Rule of Thumb for Organic Compounds 1. Carbon will always have 4 bonds 2. Nitrogen typically has 3 bonds a. 4 bonds makes it cationic i. Ex.) Ammonium 3. Oxygen typically contains 2 bonds 4. Halogens only contain 1 bond Uniqueness of Organic Molecules 1. Organic Molecules have structural complexity a. Recall: Carbon 2. Organic Molecules have chemical diversity Carbon  It is not energetically favorable for Carbon to form ions  Has the ability to catenate  Short enough bonds to be able to pi bond o This allows larger variety of compounds that can be formed  Form stable bonds with heteroatoms o Form linear and ring structures o Can contain single, double, triple bonds  Functional Groups: Specific combination of atoms, typically C-C multiple bonds or C-X bond, that reacts in a characteristic way no matter what molecule it occurs in o Most reactions of organic compounds tend to happen where the functional group occurs  Hydrocarbon: an organic compound consisting of only C and H o Ex.) methane, ethane, benzene etc.  Alkene: Hydrocarbons that contain only single bonds o Referred to as saturated o Formula is CnH2n+2 o Tetrahedral o Sp3 hybridizes Naming Organic Compounds Prefix + Root + Suffix Root: determined by the amount of carbons are in the longest continuous carbon chain Suffix: indicates the type of organic compound. Is found after the root. Prefix: Identifies any group that is attached to the main chain and the position of that group Steps to Naming 1. Name the longest chain of carbons (root) 2. Name the type of compound based on what type of bonds you have a. Single bonds: -ane b. Double bonds: -ene c. Triple bonds: -yne 3. Name Branches (groups attached to the main chain a. End of branches are named with –yl b. Branch names come after the chain name. i. When there are two or more branches the names must go in alphabetical order c. Specify where the branch is located by numbering your carbon chain to determine where the branch is i. Make sure the carbons are numbered so that the branch has the lowest possible number Geometrical Isomers: cis-trans versions of a molecule  Cis and trans will have different properties from each other o When deciding if something is trans or cis you must look at the bond directly before and after the double bond Alkynes: hydrocarbons that contain at least one C-C triple bond  Known as unsaturated carbons  Formula: CnH2n-2  Rotation is restricted at the triple bond  Linear shape  Sp hybridized  Electron rich and therefore act as a functional group Aromatic Hydrocarbons: cyclic molecules with delocalized pi electrons  Single and double must alternate for them to be considered an aromatic hc  Resonance forms exist—remember because of alternating single/double bonds  Do NOT need to know how to name these Functional Group  Many functional groups exist  Functional groups determine o Physical properties o Chemical properties o Reacitivity  Determine the polarity of a compound o Polarity determines what intermolecular forces are present  Intermolecular forces depend on polarity and vice versa  Regions of high and low electron density is determined by functional groups Common abbreviations in organic Chemistry  You’ll use these when looking at functional groups I. R—an organic group of atoms bound by carbon a. Exception: when carbon is bound to oxygen II. R/H—used for substituents a. Can be either R or H III. X—halide a. F,Cl,Br,I *Know the functional groups and their suffixes!!!!!* Alcohol  Carbon bound to –OH  Named by replacing –e at the end of the hydrocarbon name with an –o  High melting point o Due to the fact that they H bond  Less acidic and basic compared to H2O Haloalkanes  Carbon bound to a halogen Amines  Contains a nitrogen atom  Viewed as a derivative of NH3  Weak bases o Due to lone pair on N  Three Types o Primary: NRH 2 o Secondary: NR H2  Primary and secondary can hydrogen bond. Tertiary can not. o Tertiary: N3 Carboxyl (not a functional group)  They are parts of functional groups o Aldehyde, ketones, carboxylic acids, ester and amides  Contain a C=O double bond o Partial positive charge is located on the carbon. o Partial negative charge is on oxygen  Always has a dipole Aldehydes and Ketones  Aldehydes have a hydrogen at the terminal end  Ketones have two terminal carbons  Easy to confuse so be careful! Carboxylic Acid  Contains the functional group –COOH  Weak acid Amides  Contains a nitrogen  In most drugs  Also a peptide in biology General Chemistry 1 concepts to apply to Organic Chemistry 1. Pi and sigma bonds 2. Hybridization 3. Molecular polarity a. Very important Chapter 12 Matter occurs as a solid liquid or gas depending on the conditions it’s in  Phase: a physically distinct homologous part of a system o Solids and liquids are considered to be a “condensed” state  Intermolecular forces: form of potential energy that holds particles together o Intermolecular Force: weak forces that exist between two separate molecules. High charge—stronger the force will be o Intramolecular Force: bonding forces (ionic, covalent, metallic)  Understand that these are different  Kinetic Energy is responsible for phase changes o If kinetic outweighs potential  substance melts or evaporates o If potential outweighs kinetic  substance condenses or freezes Potential and Kinetic Energy in Phases Potential Energy: energy of attraction (intermolecular forces) Kinetic Energy: random motion of individual particles  Gas molecules o High Kinetic Energy o Low Potential Energy  Liquid Molecules o Similar of kinetic and potential  Solid o Low kinetic energy o High potential energy Phase changes and enthalpy (delta H)  Temperature increases  kinetic energy increases o Endothermic (delta H= +)  Heat is absorbed  Causes intermolecular forces between molecules to be interrupted  Particles moving faster makes it easier to overcome potential energy  Temperature decreases  potential energy increases o Exothermic (delta H= -)  Particles slow down making the intermolecular forces stronger Key equations 1. Within a phase change a. q=(moles)(molar heat capacity)(∆temp) i. q=mc∆t 2. During phase change a. q=(moles)(∆H phase change) b. Used to quantify heat change i. q= heat removed Liquid-Gas Equilibrium Equilibrium: condition existing when a process and its reverse process proceed at equal rates  Liquids have a vapor pressure above them—this causes it to have a vapor pressure  Vapor pressure : pressure exerted by the vapor at equilibrium o Remember flame demonstration Open System  Non equilibrium process o H2O in a glass eventually disappears Closed System:  Equilibrium process  Increase temperature means an increase in the amount of molecules turning to vapor o Higher temps are helping overcome intermolecular forces o Increased vapor pressure=lower intermolecular forces Quantifying the effect of temperature  Nonlinear relationship between pressure and temperature can be converted to a linear relationship −∆Hvap 1  L nP= R (T +C lnP2 =∆Hvap 1 − 1  P1 R (T 1 T 2) *Know how to do problems using these equations! Practice problems can be found on learn and in the book!* Intermolecular Forces 12.3 Intramolecular Forces: attractive forces within the same molecule  Attraction between cation and anion in ionic bonding  Attraction between nuclei and electron pair in covalent  Attraction between metal cations and delocalized electron in metallic bonding Intermolecular Force: attractive forces between molecules atoms and ions  Can arise due to partial forces  Can also arise between ions and molecules  Weak compared to bonding forces Van Der Whal distance: Closest two molecules can get to each other Intermolecular forces Ion Dipole Forces  Strongest intermolecular force  Interaction between an ion and a polar molecule  Electrostatic motion o Ionic compound dissolved in H2O Dipole Dipole  Attraction between two dipoles of separate molecules  Forces are more orderly in a solid compared to a liquid Ion Dipole  Ion and a molecule with a dipole Hydrogen Bond  Force involving H bound to F, N, or O directly and another very electronegative element o X—H-----Y  X and Y are most electronegative  X and Y can also be bound to other atoms  H bonded compounds have higher boiling points then others Ion Induced Dipole and Dipole induced Dipole  A nearby electric field can “induce” or create a distortion in an electron cloud o Pulls electron density toward positive pole and pushes it from negative pole  Only temporary dipole  Nonpolar molecules o It will induce them to have a temporary dipole moment  Polar Molecules o Enhances (or strengthens) an already existing dipole moment Polarizability: The ease at which an electron cloud is distorted (ability to form pole) Polarizability Trends  Smaller atom = harder to polarize o Electrons have less space to move away in a smaller atom  Larger atoms = easier to polarize  Polarizability increases down a group o Follows same trend as atomic size  Polarizability effects ALL intermolecular forces London Forces  Temporary and instantaneous reaction o Constantly interacting and then breaking  Most universal intermolecular force o Present in all molecules and ions  Cylindrical chain molecules have more London forces then compact molecules o More points of contact  Very weak  Large Molecule= large amount of London forces Deciding stronger dispersion forces 1. Polarizability 2. Surface Area Uniqueness of H2O Liquids: combine the ability to flow with strong intermolecular forces Gases: Random arrangement is same at any place in a container 2 Surface Tension: The energy required to increase the surface area (J/m ) Types of Liquid Molecules 1. Interior Molecules a. Are in the middle of water (not toward the top) i. Therefore, they have attractions on all sides 2. Exterior a. Located at the surface of water i. Only attracted from molecules below and on the sides 1. Causes it to have a downward net force ii. In order to increase its attractions and become more stable like the interior molecules, it must break it’s intermolecular forces (this requires energy)—surface tension iii. Surfactants: decrease the strength of H2O by collecting at the surface and disrupting the H bonds Water Properties Capillarity: Riding of a liquid against the pull of gravity through a narrow space Viscosity: The resistance of a fluid to flow Surface Tension: described above Chapter 12.6 Types of Solids Crystalline Solids: well defined shapes because of their particles (atoms, molecules, ions) occur in an orderly shape Amorphous Solids: non crystalline structures; lacking defined shape due to particles not having an orderly arrangement Unit Cell: Smallest portion of a crystal which gives you the total crystal when repeated in all 3 dimensions  Coordination number: number nearest neighbors of a particle o Higher the CN, greater the particles in a given volume  Body Centered Cube o 1/8 atoms at 8 corners o 1 atom at center o CN: 8  Face Centered o 1/8 atoms at 8 corners o ½ atoms at 6 faces o CN:12  Packing Efficiency: measure of the total volume occupied by spheres o Simple Cubic: 52% o Body Centered: 68% o Face Centered: 74%  Best we can do, most common in nature With unit cells, try to practice visualization. Do not stress over names.


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