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CHE 131: General Chemistry, Notes for Exam 2

by: Robert_Smith

CHE 131: General Chemistry, Notes for Exam 2 CHE 131

Marketplace > Stony Brook University > Chemistry > CHE 131 > CHE 131 General Chemistry Notes for Exam 2
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These notes cover the material for our second exam.
General Chemistry IB
Roy Lacey
Study Guide
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This 5 page Study Guide was uploaded by Robert_Smith on Sunday March 20, 2016. The Study Guide belongs to CHE 131 at Stony Brook University taught by Roy Lacey in Winter 2016. Since its upload, it has received 25 views. For similar materials see General Chemistry IB in Chemistry at Stony Brook University.

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Date Created: 03/20/16
03/20/2016  CHE 131: General Chemistry­ Exam 2 Study Guide   ____________________________________________________________________________ __  Limiting Reactants   ­ Limiting Reactan ­ the reactant that limits how long the reaction takes  ­ It will always be used up by the end of the reaction  ­ You can check for the limiting reactant by testing which reactant produces the  least amount of a certain product, using their molar ratios from a balanced  equation  ­ Percent Yield the amount of product created in real life divided by how much  theoretically should have formed  ____________________________________________________________________________ __  Solutions  ­ solutions ­ a homogeneous mixture of two or more substances  ­ Can be a gas, solid, or liquid  ­ A solvent exists in greater quantity and dissolves a smaller am​olute  s ­ Ionic substances dissolve into their ions; molecular substances do not  disassociate  ­ Concentration ­ ​mount solute / amount solution  ­ Molarit ­ a measure of concentration,   ­ (moles solute) / (liters solution)  ­ Unit is M  ­ Dilution involves lowering the molarity of a solution by adding stock solution to  a solute  ­ M​diluV​dil= M​concentra​econcentrated   ____________________________________________________________________________ __  Electrolytes, Acid­Based Reactions, & Precipitation   ­ Solubility soluble substancewill dissolve in a solution and not ​recipitat​(a  solid formed by the combination of two insoluble ions)  ­ Soluble cations: alkali metals, ammonia  ­ Soluble anions: nitrate, acetate, halides, sulfates  ­ Insolubilit insoluble substancewill form a precipitate in solutions  ­ Soluble compounds: hydroxides, sulfides, carbonates, phosphates  03/20/2016  ­ Electrolytes a solute that makes ions in a solution  ­ These ions ​harge carriers)can run an electric current through the solution  ­ Non­electrolytes ­ does not ionize or run a current (e.g. sugar)  ­ Strong electrolytes ­ breaks down almost completely into ions, conducts  electricity well  ­ Weak electrolytes ­ only partially dissociate into ions, don’t conduct electricity well  ­ The strength of an electrolyte isn’t related to its solubility  ­ Ionic Equation ­an equation that shows the ions formed by the reactants and the  products  ­ A Net Ionic Equation​does not includ​pectator ions (ions that don’t change  during the reaction and are the same on either side of the equation)  ____________________________________________________________  Acid­Base Reactions, Stoichiometry, & Titrations  ­ Amphiprotic ­ substances that accept or donate protons (e.g. water)  ­ Neutralization reaction​ an acid and base combine to create a salt and water  ­ Equivalence point ­​point in a redox titration when the number of electrons lost by the  oxidized substance equals the number gained by the reduced species  ­ Titration  when one adds an unknown substance to a known solution until the  equivalence point has been reached (usually marked by a change in the solution’s color  ­ One usually does this to try to find the concentration of the unknown solution  using thetitran(the concentration of the known solution)   ____________________________________________________________  Redox Equations  ­ Redox Reaction (or Oxidation­Reduction Reaction) ­ ​involve the transfer of electrons  ­ The electron transfer is marked by the change inoxidation numbers ( the  charge of the ion formed by part of the substance)  ­ Oxidation numbers are 0 for neutral atoms, usually +1 for H except when it forms  hydrides with metals (then it is ­1), it is usually ­2 for O  ­ Oxidation ­increases a substance’s oxidation number (it becomes more positive as it  loses electrons)  ­ Reduction ­ decreases a substance’s oxidation number (it becomes more negative as it  gains electrons)   03/20/2016  ­ Oxidizing agent ­ ​does oxidation, is reduced  ­ Reducing agent ­  ​does reduction, is oxidized  ­ Balancing Redox Equations ­  ­ Split equation into half­reactions of oxidation and reduction  ­ Balance all elements, using water to balance O and H+ (for acidic solutions) or  OH­ (for basic solutions) to balance H  ­ Balance charge with electrons  ____________________________________________________________  Gasses, Pressures, & Gas Laws  ­ Ideal Gas Law ­ P​V = nRT  ­ R is the gas constant, which is given during exams  ­ T must always be measured in Kelvins, not celsius. This is true for all of the  following gas laws  ­ P is pressure in atm, V is volume in liters, and n is the number of moles  ­ Boyle’s Law ­ V​ = k/P  ­ k is a constant   ­ Charles’ Law ­ ​V = bT  ­ b is a constant  ­ Avogadro’s Law ­ V ​  = a n  ­ a is a constant, and n is the number of moles  ­ Pressure ­ ​P = F / A  ­ Pressure equals force per unit area  ­ Units: 1 atm = 760 mmHg = 760 torr = 101.325 kPa  ­ A m​anometer​  measures pressures of various gases  ­ Standard temperature and pressure (STP) ­  ​at STP, T = 273 K and P = 1 atm  ­ At STP, l mole of gas takes up 22.4 L  ­ Density ­​ (PM) / (RT)  ­ M is the molar mass of the gas  ­ R is the gas constant, which is given during exams  ____________________________________________________________  Dalton’s Law, Kinetic Molecular Theory, & Real Gases  03/20/2016  ­ Dalton’s Law ­is a way to calcul​artial pressur​(pressure exerted by one gas in  a mixture)  ­ P​ =(niRT)/V  ­ Ptot = P1 +P​2 + .Pn =n​tot((RT)/V)  ­ Mole Fraction  the proportion of the mixture that the gas makes up  na a ­ χ =a na+b +..n n ntotal   ­ Pa =χaP otal  ­ Kinetic Molecular theory a model that explains gas behavior  ­ Pressure is the impulse per collusion times the frequency of collisions with the  walls of the container  ­ PV = Nm u2 u = 3RT    M 2 ­ M is the molar masu. s the mean speed of all the molecules squared.  ­ Average Kinetic Energy (Temperature) ­   ­ KE​avg= (3/2)RT  ­ Diffusion  the mixing of gases, happens more quickly for light molecules  ­ Effusion ­the rate at which gas molecules pass through a small hole in a chamber  ­ Real vs. Ideal Gases​­eal gases don’t acct ideally at high pressures because of  intermolecular forces and the finite volume of gas molecules  ­ Adjusted Ideal Gas Law­   2 ­ (Pobs+ a(n/v) (V­nb) = nRT  ____________________________________________________________  Forms of Energy & Enthalpy  ­ Energy ­ is stored in chemical bonds, is used and released during chemical equations  ­ Potential energ­ energy based on position  ­ Chemical Energy ​stored in bonds  ­ Kinetic Energy ​k = (1/2) m v , m is mass and v is velocity  ­ Work = force times distance  ­ Units ­ 1 J = (1 kg m )/s1 cal=4.184 J 1 Cal = 1  cal ­ System ­ ncludes everything involved in reactions, bu​urroundings  03/20/2016  ­ Open system ­​  allows mass and temperature changes  ­ Closed system ­ a ​ llows temperature changes  ­ Isolated system ­ ​doesn’t allow any changes  ­ State function ­ measures only the difference between the starting and ending state, the  path between the two states does not matter  ­ Equations ­   ­ ΔE = q + w  ΔE is the energy change, q is heat, w is work done on system  ­ W = ­P ΔV     ΔP is the volume change, w is work, P is pressure  ­ ΔH = ΔE ­ V    ΔH is the enthalpy change  ____________________________________________________________  Calorimetry  ­ Exothermic ­ ​heat is released to surroundings  ­ Endothermic ­ ​ heat is absorbed from surroundings  ­ Enthalpy ­ H = E + PV ΔH = ΔE + PΔV  ­ Calorimeter ­ a device used to measure heat exchanges during reactions  ­ Specific heat capacity ​­ the heat required to raise the temperature of one gram of a  substance by 1 degree Celsius at constant pressure 


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