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Exam 1 Study Guide

by: Dallin Rigby

Exam 1 Study Guide Chemistry 1220

Dallin Rigby

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Study Guide for the test on Chapters 11 - 13 in Chemistry, The Central Science
Principles of Chemistry
Dr. Weaver
Study Guide
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This 7 page Study Guide was uploaded by Dallin Rigby on Wednesday January 27, 2016. The Study Guide belongs to Chemistry 1220 at Southern Utah University taught by Dr. Weaver in Winter 2016. Since its upload, it has received 49 views. For similar materials see Principles of Chemistry in Chemistry at Southern Utah University.


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Date Created: 01/27/16
Exam 1 Study Guide Chapter 11 1 Be able to identify what types of intermolecular forces exist in a substance: Van der Waals forces: London/dispersion "Induced Dipole-Dipole"  Nonpolar molecules and a diple is created for just a moment. Random.  The greater the molar mass, the greater the "polarization" of the molecule, the greater the London Dispersion force.  Temporary  Instantaneous  Weakest Dipole/dipole  Consists of two polar molecules.  Hydrogen bonds also have dipole-dipole. Hydrogen bonding Hydrogen with Fluorine, Oxygen, or Nitrogen.  Permanent bonds.  Strongest intermolecular forces.  Creates dipole.  HP, H 2, NH 3 Ion-Dipole o Consists of an ion with a polar molecule.  Na and H O2 Ion induced Dipoles and Dipole induced Dipoles Ionic bonding Metal bonding 1 Be able to predict boiling points based upon the ionic forces in the molecule. Trends Relative Strengths of Intermolecular Forces Hydrogen Bonding > Dipole-Dipole > London Dispersion The influence of each of these attractive forces will depend on the functional groups present. Boiling points Increase with Molecular Weight Van der Waals forces are proportional to surface area. As chain length increases, surface area increases. This increases the individual molecules' ability to attract one another. Symmetry Branching decreases boiling point. Example: What could you stack better: spaghetti or macaroni noodles? 1 Understand surface tension and viscosity with regards to intermolecular forces. Viscosity Viscosity is when liquids are resistant to flowing. Greater viscosity = slower flowing. Lesser viscosity = faster flowing.  For all substances, viscosity decreases as temperature increases.  When comparing related molecules, viscosity increases with molecular weight. Think: To measure viscosity, time how long it takes for a liquid to flow through a vertical tube. Substanc Formula Viscosity e (kg/m-s) Hexane CH 3H C2 CH2CH 2H 2 3 3.26 * 10 -4 Heptane CH CH CH CH CH CH CH 4.09 * 10 -4 3 2 2 2 2 2 3 -4 Octane CH 3H C2 CH2CH 2H C2 CH2 2 3 5.42 * 10 Nonane CH 3H C2 CH2CH 2H C2 CH2C 2 2 7.11 * 10 -4 H3 Decane CH 3H C2 CH2CH 2H C2 CH2C 2 2 1.42 * 10 -3 H2CH 3 Surface Tension The energy required to increase the surface area of a liquid by a (unit) amount.  Molecules within water (not on surface) are equally attracted in all directions.  This net force pulls molecules on the surface toward the interior. This reduces the surface area and makes the molecules pack closer together. -2 2  Surface tension of water at 20 °C is 7.29 * 10 J/m  This means 7.29 * 10 Joules must be used to increase water surface area by 1 m .2 1 Understand energy changes that accompany phase changes. (Energy versus temperature). Things to Remember:  Energy must be supplied when;  The change involves the disruption of intermolecular forces.  Disruption of intermolecular forces;  State becomes less ordered.  As intermolecular forces increase;  More energy must be supplied to disrupt them. Fusion The melting process for solids.  The change in enthalpy associated with this is called the heat of fusion (ΔH )fus  Ice ΔH fus 6.01 kJ/mol Vaporization  The heat needed to vaporize a liquid is called the heat of vaporization (ΔH ) vap  Water Δh vap= 40.67 kJ/mol  Requires input of heat (energy). 1 Be familiar with heating curves and be able to calculate energy changes in multiple phase changes. Heating Curves A graph showing temperature versus amount of heat added is a heating curve. 1 What is a critical temperature and pressure? Substance Critical Temperature Critical Pressure (K) (atm) Nitrogen, N 2 126.1 33.5 Argon, Ar 150.9 48.0 Oxygen, O 2 154.4 49.7 Methane, CH 190.0 45.4 4 Carbon dioxide, 304.3 73.0 CO 2 Phosphine, PH 3 324.4 64.5 Propane, 370.0 42.0 CH 3H C2 3 Hydrogen sulfide, 373.5 88.9 H 2 Ammonia, NH 3 405.6 111.5 Water, H 2 647.6 217.7 Critical Temperature The highest temperature at which a distinct liquid phase can form.  Highest temperature at which a liquid can exist.  Greater intermolecular forces give a substance a higher critical temperature. Critical Pressure The pressure required to bring about liquefaction at critical temperature. Chapter 12 1 Be able to calculate the length of unit cells, given the radius of an atom. Know whether the unit cell is simple, body centered, or face centered. Atom Number of Unit Fraction of Location Cells Sharing Atom Within Atom Unit Cell 8 1/8 or 12.5% Corner 4 1/4 or 25% Edge Face 2 1/2 or 50% Anywhere 1 1 or 100% else Volume Face Body Simple Centered Centered Cubic 3 2√2r 4 r 2r √ 3 Mass Face Centered Body Centered Simple Cubic 4∗Molar Mass 2∗Molar Mass 1∗MolarMass ' ' ' Avogadro s Number Avogadr o s Number Avogadr o s Number Packing Efficiency Simple Body Face 0.52 0.68 0.74 1 Understand the molecular Orbital and Electron Sea model of metallic solids. 1 Understand what a semiconductor is (including n and p- types). Semiconductors N-Type P-Type 1 Understand polymers, nanomaterials, and fullerenes. Polymers Molecules of high molecular weight formed by polymerization of monomers.  Coined in 1827 by Jons Jakob Berzelius Nanomaterials Materials that have dimensions of 1 - 100 nanometers. Fullerenes Molecules composed of carbon, shaped as hollow spheres, ellipsoids, or tubes.  Discovered in 1985 by Buckminster Fuller.  Spherical fullerenes are called "buckyballs"  The C60variant is referred to as "buckminsterfullerene"  Cylindrical fullerenes are referred to as "buckytubes"  Fullerene structures are similar to graphite. Chapter 13 1 Understand: a The principles of solvation (entropy and enthalpy). Solvation Helps stabilize ions in solution, prevents cations and anions from recombining. ← Solute+SolventCrystallizeDissolveSolution → Entropy Gradual decline into disorder. Enthalpy Total heat content of a system. a The principles of saturated, supersaturated, and unsaturated. Saturated A solution that is in equilibrium with undissolved solute. Supersaturated Under specific conditions, solutions can contain a greater amount of solute than needed to form a saturated solution.  Supersaturated solutions are unstable because the solute concentration is higher than equilibrium concentration. Unsaturated Less solute is dissolved than the amount needed to form a saturated solution. a How temperature affects solubility of gasses and solids in liquids. a How pressure affects the solubility of gases. (Henry's Law) 1 Be able to calculate molarity, molality, percent by mass, molar fractions and ppm. Expressing Solution Concentration Molarity Percent Volume molessolute Volumesolute ∗100 Literssolution Volume solution Molality Parts Per Million molessolute gsolute 6 Kgsolvcent gsolution∗10 Percent Mass Parts Per Billion gsolute∗100 gsolute ∗10 9 gsolution gsolution Mole Fraction (Χ) Molessolute moleA ∨ Totalmoles∈solution moleB+moleA 1 Convert between different units of concentration. Colligative Properties 1 Be able to calculate freezing points, boiling points, vapor pressures, and osmotic pressures.  For molecular compounds (nonvolatile).  For ionic compounds. 1 Be able to calculate molecular mass as the result of a colligative effect.


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