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General Chemistry Exam 1

by: Aashika Kushwaha

General Chemistry Exam 1 1312

Marketplace > University of Texas at Dallas > Chemistry > 1312 > General Chemistry Exam 1
Aashika Kushwaha
GPA 3.6

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Intermolecular Forces Kinetics Gases
General Chemistry 2
Dr. Sra
Study Guide
General Chemistry
50 ?




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This 9 page Study Guide was uploaded by Aashika Kushwaha on Tuesday September 27, 2016. The Study Guide belongs to 1312 at University of Texas at Dallas taught by Dr. Sra in Fall 2016. Since its upload, it has received 6 views. For similar materials see General Chemistry 2 in Chemistry at University of Texas at Dallas.


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Date Created: 09/27/16
Exam 1 Kinetics- how fast a reaction goes from reactants to products -Related to factors such as  Pressure  Temperature  Catalysts- lowers activation energy (Ea) -do not get consumed  Concentration (molarity,M) Rate= number of collisions Seconds Reaction occurs by  Enough energy (to overcome repulsion)  Correct orientation Activation Energy (Ea)- minimum amount of energy required for reaction -must overcome activation energy to collide and form the activated complex Isomerization- molecule rearranges and forms another compound  Bond has to weaken so atoms can rearrange Rate= change of concentration Change of time (sec) -negative sign is used if concentration is decreasing, this gives the rate a positive value  As time goes on, the rate of reaction generally slows down  Reaction stops when reactants run out, because the reaction has reached equilibrium aA + bB => cC + dD rate=( -(1/a)[A]/change of time)=((-1/b)[B]/change of time)=… rate= 1/a((change of concentration)/(change of time))=1/b((change of concentration)/(change of time))=…  Replace coefficient (a) with other lower case letters from the equation. These indicate the coefficient of the reactant in the equation Average rate- gives the change in concentration over a finite time period.  Equals the slope of the curve  Linear approximation of a curve  Larger the time interval, the more that average rate will deviate from the instantaneous rate Instantaneous Rate- reaction rate at a particular instant during the reaction, given by the slope of the tangent to the curve at that instant Initial Rate- rate at t=0 Measuring Reaction Rate  Continuous monitoring of the concentration  Sampling the mixture at various time Polarimetry- degree of rotation of plane Spectrophotometry- measuring the amount light depending on the particular wavelength absorbed -Absorbs complementary color Total Pressure- mixture of partial pressures PV=nRT RATE IS DETERMINED EXPERIMENTALLY!!!! Order is determined by the experiment, not the coefficient (n,m) Overall Order- sum of all of the orders of the reactants -Order is determined by comparing the reactant concentrations in the table. You keep one of the reactants constant and then compare the change of the other with the  Compare the reactants that change and then compare the corresponding initial rates. Ratio of reactant difference and initial rate difference determines order  Doubling [A} will quadruple the rate of the reaction  Solve for k to determine value and unit of the rate constant Units for k Zero: M/s First: s^-1 Second: M^-s^- Third: M^-2s^- Nth: M^-(n-1)s^-1 Integrated Rate Law- used to know how the concentrations of the reactants and the products change with time (seconds)  Depends on concentration st Integrated Rate Law for 1 order: ln[A]t=-kt+ln[A] 0  Relates to y=mx+b  Calculate for k by  Solve for k by integrated rate law 2N 2O5 => 4NO 2+O 2 rate=k[N 2O 5]  Calculate slope -5.099-(-3.912) = -0.0017 1/s (700-0)s k= -0.0017 1/s  Slope is negative but k is positive Half-Life 1/)-length of time it takes for the concentration of the reactant to fall by ½ its initial value  Depends on order T1/=.693/k First Order: When rate: M/sec K=s^-1 Second Order Reactions 1/[A]t = kt+ 1/[A} 0 -related to mx+b=y Half life (2d Order) t1/=1/k[A] 0 -dependent on initial concentration Zero order 2nd order First order Arrhenius Equation- represents the fraction of reactant molecules that have enough energy to make it over the energy barrier.  Rate of reaction increases as temperature increases k=A(e^-Ea/RT) ln(k)=(-Ea/R)(1/T)+ln(A) ln (k2/k1)= -(a/R)(1/T2 – 1/T1) A: frequency factor (collision frequency with the correct orientation) Ea: Activation energy R: Gas constant (J/(mol)(K) T: Temperature (Kelvin) Molarity (M)=mol/liters Molality (m)= mol/kg Polarity is due to the difference in Electronegativity Electronegativity- ability of an atom to attract shared e- to itself Pure Covalent- difference in electronegativity between bonded atoms is 0 -equal sharing Nonpolar Covalent- difference in electronegativity between bondd atoms is 0.1-0. Polar Covalent- difference in electronegativity between bonded atoms is 0.5 to 1.9  Covalent- NM+NM Ionic- difference btw bonded atoms is larger thatn or equal to 2  NM+M  μ=Qxr -dipole moment  Q-charge  R-distance between the charges   μ is always positive Molecular polarity depends on shape and bond polarity London Dispersion Forces- attractive forces as a result of -weakest Ion-Dipole Forces- the result of electrical interactions between an ion and the partial charges on a polar molecule  strength determines solubility of ionic compounds Dipole-Dipole Attractions- the result of electrical interactions between neighboring molecules -exist between all polar molecules Ion-induced: When a strong cation induces a dipole in a molecule that could have been nonpolar. Ex. NaCl around F 2, can form NaF and Cl- Dipole-induced- forces responsible for solvation of gases -Double induced dipole IMFs increase as dispersion increases BP increases as IMFs increase Larger molecule: larger surface area More points for dispersive forces Polar molecule will have higher boiling point Greater molecular weight will have higher boiling point Hydrogen Bonding- Hydrogen bonds with Fluorine, Nitrogen, or Oxygen. -causes hydrogen to lose its only electron, so the H proton is exposed  strongest intermolecular force


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