CHEM 112 study guide 1
CHEM 112 study guide 1 CHEM 112000
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This 10 page Study Guide was uploaded by Olivia Lee on Monday February 1, 2016. The Study Guide belongs to CHEM 112000 at Purdue University taught by Abu-Omar in Winter 2016. Since its upload, it has received 31 views. For similar materials see Chemistry in Chemistry at Purdue University.
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Date Created: 02/01/16
CHM 112 1nd Edition Exam # 1 Study Guide Lectures: Lectures Jan. 13 – Jan. 27, Ch.9-10 of textbook (Bauer 3rd Ed.) Lecture 1 -The universe is composed of all types of gases. The atmosphere is a thin layer of gasses surrounding the Earth. The gas molecules have little effect on each other. -Hot gases have a lower density than cold gases -To calculate any change in gases, use: PV=nRT P= Pressure, T= Temperature, V =volume, n= number of atoms, R= constant Lecture 2 -as temperature increases, volume increases -Avogadro’s hypothesis: volume occupied by a gas at a given temperature and pressure is directly proportional to the number of gas particles. • As the moles of gas increase, the volume increases (at constant T and P). Boyles Law states V= 1/P (constant pressure) -Volume is inversely related (proportionally) to a constant pressure. - Boyle's law is an experimental gas law which describes how the pressure of a gas tends to decrease as the volume of a gas increases. Charles’ Law state V/T = constant -Volume and temperature are at a constant pressure - Charles' law is an experimental gas law which describes how gases tend to expand when heated. Lecture 3 Ideal Gas Law PV=nRT We can solve for density in the Ideal Gas Law using substitution d= m= Px MN v RT Two balloons filled with CO2 & O2 held at equal temperature and pressure. They have… -same number of molecules -same number of moles -CO2 has the greatest mass -CO2 has the greatest density Dalton’s Law of Partial Pressures -gases in a mixture behave independently and exert the same pressure they would exert if they were in the container alone. -when a gas is collected about a liquid, such as water, the liquid’s vapor adds to the gas and total pressure. P(total) = P(gas) + P (H20 vapor) Example problem: Suppose 2.25 L of H2 gas is collected over water at 18.0 degrees Celsius and 722.8 torr. How many moles of H2 are produced in this reaction? 1. Convert to K: 18°C +273=291 Kelvin 2. Convert to atm & subtract water pressure value 722.8torr−15.5torr 760 =.931atm 3. Plug into PV=nRT (.931atm)(225L)=n(0.08206)(291K) Solve for n. Kinetic-Molecular Theory of Gasses -a model that explains experimental observations about gases under normal temperature and pressure conditions that we encounter in our environment. -the pressure of a gas arises from the sum of the collisions of the particles with the walls of the container. -the average kinetic energy of gas particles depends only on the absolute temperature. Compare the gasses Ar, CO2, and H2 all at the same temperature. Which has average greatest kinetic energy? Same because ‘T’ is the same. Which has the greatest average velocity? Hydrogen because it’s the lightest Lecture 4 Intermolecular forces: The attractive forces between molecules o London Dispersion Forces Relative boiling points tell us about intermolecular force strength The stronger the forces, the higher the temperature needed to overcome them o Dipole Dipole forces One of the elements polarizes the other attractive forces between the positive end of one polar molecule and the negative end of another polar molecule o Hydrogen Bonding Occurs with Nitrogen, Oxygen, and Fluoride. Hydrogen bonding is the strongest thing that holds DNA helixes together Hydrogen bonding in H20 causes ice to be less dense than the liquid state Properties of Liquids Vapor pressure Density Viscosity Surface tension Capillary action Properties of Solids The particles in the solid state are held together in specific positions—there is no translational motion Superconductors: Unique in that they offer no resistance to the conduction of electrical current Repel magnetic fields Low temperatures are requires Lecture 5 A Solution is: o A homogeneous mixture with uniform composition throughout o Composed of: Solute: substance being dissolved Solvent: substance doing the dissolving o Solubility How readily or completely a solute dissolves in a solvent. The solubilities for ionic compounds do not follow a predictable pattern. Must be recalled form experimental data Soluble Na+, K+, NH 4 + NO 3 - Soluble Ag+ Insoluble-except for AgNO3 S Insoluble except Na+, K+, and NH+ 2 - Electrolytes o A solution that conducts electricity o Two kinds: Strong electrolyte: solute that dissociates completely into ions in aqueous solution Weak Electrolyte: a solute that dissociates partially into ions in aqueous solutions o Nonelectrolyte: a solute that does not dissociate into ions in aqueous solution o Like dissolves like When bonding in a solvent is similar to the bonding in a solute, then the solute will dissolve in the solvent Solute Solvent Solvent solute Polar Nonpolar Ionic Soluble Insoluble Polar Soluble Insoluble Nonpolar Insoluble Soluble Solution Process o When an ionic compound dissolves in water: Ionic bonds in the solute break Hydrogen bonds between water molecules break Ion-dipole forces form between ions and water molecules Entropy o A measure of the tendency for matter to become disordered or random in its distribution Matter spontaneously changes from a state of order to a state of disorder
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