General Chemistry

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  1. Ionic Bonding
    Due to the electrostatic charge between two atoms.

    The atoms form a crystal lattice to minimize repulsive forces and maximize attractive forces. They have a high electronegative difference, high bp/mp, and make great conductors in polar solvents.
  2. Difference between coordinate covalent and covalent bond.
    • Coordinate Covalent: both of the shared electrons come from one atom
    • Covalent Bond: an elctron pair is shared between two atoms
  3. Bond Order
    Number of electrons shared between atoms.

    • Single bond - 1
    • Double bond - 2
    • Triple bond - 3
  4. Electronic Geometry vs Molecular Geometry
    Electronic geometry takes into account all pairs of electrons while molecular takes iin only bonded atoms.

    • Extra Info:
    • Linear - 2
    • Trigonal Planar - 3 (120)
    • Tetrahedral -4 (109.5)
    • Trigonal Bipyramidal -5 
    • Octahedral - 6
  5. Change in electronegativity for:
    Nonpolar Covalent
    Polar Covalent
    • Nonpolar Covalent: 0.5
    • Polar Covalent: 0.5-1.7
    • Ionic: 1.7+
  6. Gram Equivalent Weight
    Grame Equivalent Weight: amount of compound measured in grams, that produces on equivalent of the particle of interest.

    Gram equivalent weight = (molar mass) / n

    n = number of protons, hydroxide ions, electrons, or ions produced or consumed by solute
  7. Equivalents
    Equivalent: How many moles of things we are interested in (protons, hydroxide ions, electrons, ions) will one mole of a given compound produce?

    Equivalents: (mass of compound)/(gram equivalent weight)

    • Side Note
    • Gram equivalent weight = (molar mass) / n

    n = number of protons, hydroxide ions, electrons, or ions produced or consumed by solute
  8. Molarity based on Normality
    Normality: measure of concentration based on equivalents/L

    Molarity = (Normality)/n

    n = number of protons, hydroxide ions, electrons, or ions produced or consumed by solute
  9. Combustion Reactions
    A special reaction that involves a fuel (hydrocarbon) and an oxidant (O2). It forms CO2 and H2O as a product.
  10. Single Displacement Reaction vs Double Displacement Reaction
    Single Displacement Reaction: when an atom or ion in a compound is replaced by an atom, ion, or another element. Commonly classified as oxidation - reduction reactions.

    Double Displacement Reaction: (metathesis reactions) elements from two different compounds swap places to form two new compounds
  11. Neutralization Reactions
    Specific types of double-displacement reactions where an acid reacts with a base to produce a salt.

    HCl + NaOH -> NaCL + H2O
  12. Naming monatomic anions
    • The ending is dropped and replaced with -ide
    • H- : hydride
    • F- : fluoride
    • O2- : oxide
  13. Naming polyatomic anions with a hydrogen attached
    You would add a hydrogen before the name or the bi- at the beginning.

    • HSO4- : hydrogen sulfate or bisulfate
    • HCO3- : hydrogen carbonate or bicarbonate
  14. MnO4-
  15. SCN-
  16. BO33-
  17. Chromate and Dichromate
    • CrO42-
    • Cr2O72-
  18. Electrolyte
    Electrolytes: solutes that enable solutions to carry electricity.

    Ions by themselves are not conductive. When they are put in water, the ion lattice is broken and it allows for the free movement of electrons which allows it to conduct electricity.
  19. Phosphate
  20. Gibbs Free Energy (ΔG)
    Spontaneous or Non-spontaneous; determine whether a reaction will occur by itself without assistance.

    Just because a reaction is spontaneous, does not mean it has a fast reaction rate.

    • ΔG = (+) = endergonic
    • ΔG = (-) = exergonic
    • If it is between reactants (products) and the transition state, then it is referring to th activation energy.
  21. Rate Determining Step of a Reaction
    Each reaction has intermediates (molecules that are part of the reaction but do not appear in the overall reaction) form.

    From these layers of mini reactions, there is one that is the slowest which determines the rate of the entire reaction.
  22. Collisional Theory of Chemical Kinetics
    Reaction rate is proportional to the number of collisions per second.
  23. Energy Barrier
    (AKA activation energy, Ea) The minimum amount of energy of collision necessary for reaction to take place.
  24. Rate of Reaction Equation
    rate = Z*f

    • Z= total number of collisions per second
    • f= fraction of effective collisions
  25. Arrhenious Equation
    k = Ae((-Ea)/(RT))

    • k= rate constant of reaction
    • A = frequency factor
    • Ea = activation energy of reaction
    • R = ideal gas constant
    • T = temperature in kelvin

    For this, know the equation to be able to find the relationship of things! Solving these equations is rare on the MCAT
  26. Frequency Factor
    (AKA attempt frequency) measure of how often molecules in a certain reaction collide
  27. At what energy should a molecule collision be to weaken old bonds to form new bonds
    The molecule must collide with energy equal or greater than the activation energy.
  28. Transition State
    (AKA activated complex) Point in the reaction with the greatest amount of energy than the reactants and the products, denoted as ⇟ without the arrow.
  29. What is temperature a measure of?
    Average kinetic energy. So higher the temperature, higher the energy in the molecule, easier it is for the molecule to overcome the activation energy.
  30. Why are polar solvents preferred medians to increase reaction rates?
    Their molecular dipole polarizes bonds of reactants, causing them to lengthen and weaken such that the reaction proceeds faster.
  31. Homogeneous vs Heterogeneous Catalyst
    • Homogeneous: catalyst in same phase as reactant
    • Heterogeneous: catalyst in a different phase as the reactant

    Catalysts have no effect on Keq of reaction.
  32. Rate of a reaction over time.

    aA + bB → cC + dD
    rate = -Δ[A]/aΔt = -Δ[B]/bΔt = Δ[C]/cΔt = Δ[D]/dΔt

    Since A and B are getting consumed, the rate is (-)
  33. Rate of a reaction proportional to concentration

    aA + bB → cC + dD
    • rate = k[A]x[B]y
    • k = reaction rate coefficient (M/s)
    • x & y = orders of reaction so zero, first, second, or third order but they are determined experimentally
  34. For the following expression:

    rate = k[A]x[B]y

    What is x and y? What are some noob mistakes in determining the equation values?
    x & y: orders of reaction so zero, first, second, or third order but they are determined experimentally

    • 1) It is NOT determined stoichiometrically unless it is a one step reaction or the rate determing step is given and balanced. Even then, you must use Keq to determine concentrations.
    • 2) Keq is NOT the same equation. Keq involves BOTH products and reactants. Rate involves ONLY reactants.
    • 3) k is NOT constant
    • 4) rate has to be measured at the beginning to remove reverse errors
  35. How do you find x, y, and k?

    rate = k[A]x[B]y
    • Image Upload 1
    • 1) Looking at this image, first locate 2 equations in which concentration for one does not change. 
    • 0.126 = k[0.2]x[0.2]y
    • 0.252 = k[0.4]x[0.2]y

    • 2) Then find out what number you have to get to get the product, and what number you have to get to get the second concentration.
    • 0.126g = 0.252
    • 0.1h = 0.2

    • 3) Then plug it into the following equation.
    • g = hx

    4) Solve for x

    5) Follow through for everything else

    Note: rate order is a x+y+z
  36. Zero order reaction

    Plotting Concentration vs Time
    The rate of formation of C is independent of A and B concentrations. The only way to change rate of C is by adding a catalyst, and changing k.

    • rate = k[A]0[B]0
    • Plotting Concentration vs Time
    • Slope: -k
    • Line: linear
  37. First order reaction

    Plotting Concentration vs Time
    The rate of formation of C is directly proportional to ONLY ONE reactant, such that doubling the concentration of that one, doubles the rate.

    • rate = k[A]1[B]0
    • rate = k[A]0[B]1

    • Plotting [A] vs Time
    • Slope: -slope but NOT constant
    • Line: nonlinear, with a negative decay

    • Plotting ln[A] vs Time
    • Slope: - slope = k
    • Line: linear
  38. Concentration of Radioactive Substance Over Time
    [A]t = [A]0e-kt

    • k = rate constant
    • t = time allotted
    • [A]0 = initial concentration
    • [A]t = after t, final concentration
  39. Second order reaction

    Plotting Concentration vs Time
    The rate of formation of C is directly proportional to TWO reactants. Normally thought to be due to a physical collision.

    • rate = k[A]1[B]1
    • rate = k[A]0[B]2
    • rate = k[A]2[B]0

    • Plotting [A] vs Time
    • Slope: -slope but NOT constant
    • Line: nonlinear, with a deep negative decay

    • Plotting 1/[A] vs Time
    • Slope: + slope = k
    • Line: linear
  40. Entropy
    Whether a reaction is spontaneous or not. Depicted by ΔS. It is a measure of distribution of energy throughout a system or between a system and its environment.

    Free energy diagrams show this.
  41. Enthalpy
    Change in heat of a reaction. Depicted by ΔH.
  42. Dynamic Equilibrium
    The rate of the forward reaction and the reverse reaction is exactly the same causing the concentration levels to go unchanged. BUT reactions are still occurring. If they weren't that would be static equilibrium.
  43. Relating equilibrium constants to rates of a reaction
    2A ⇌ B + C

    • ratef = kf[A]2
    • rater = kr[B][C]

    • when...
    • kf[A]2 = kr[B][C]
    • then the system is at equilibrium such that...

    kf/kr = [B][C]/[A]2
  44. How can you get the overall equilibrium constant when you are only given the steps of the reaction and those equilibrium constants?
    Multiply individual reaction constants
  45. Gibbs free energy trend
    • (-) favored, spontaneous reaction
    • (+) unfavored, non-spontaneous reaction
  46. How doe the system react if the pressure increased (volume decreased)?
    The system favors the product/reactant with the least amount of moles to decrease the pressure
  47. How doe the system react if the pressure decreased (volume increased)?
    The system favors the product/reactant with the most amount of moles to increase the pressure
  48. How does heat affect equilibrium constants?
    It does not affect Q, but it does affect K.

    If the reaction is endothermic, then adding heat favors more products.

    If the reaction is exothermic, then adding heat favors more reactants.
  49. System vs Surroundings
    • System: the object, or the stuff beiing observed(like the reactants and products of a chemical reaction)
    • Surroundings: everything outside the system
  50. First Law of Thermodynamics (equation)
    Energy can neither be created or destroyed, thus:

    ΔU = Q - W

    The heat added to the system, minus the work done by the system produces the change in energy of the system.
  51. Isothermal Process (effect on the ΔU)
    Isothermal: a constant temperature is kept

    Temperature is directly proportional to internal energy, so if ΔT = 0, then ΔU = 0

    • This means:
    • ΔU = 0 = Q - W

    • Which means:
    • Q = W in isothermal processes
  52. Graph of Isothermal Expansion and Adiabatic Expansion
    Image Upload 2
  53. Adiabatic Process (effect on the ΔU)
    Adiabatic: when there is no heat exchange between system and environment

    • This means that:
    • Q = 0

    • Which means
    • ΔU = 0 - W = -W

    So the change in internal energy is equal to the work done on the system (W is negative here, so not done by)
  54. Isobaric Process
    That is when the pressure of the system is constant. In a P vs V diagram, it is a horizontal line with only change in volume.
  55. Isovolumetric (Isochoric)
    Isovolumetric: the volume is kept constant.

    • This means in a P vs V diagram, there is no work being done. This means:
    • ΔU = Q

    The internal energy is equal to heat added to the system.
  56. What is one method of providing energy to a nonspontaneous reaction (chemically)?
    Coupling: supplying energy by pairing up nonspontaneous reaction with a spontaneous reaction
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General Chemistry
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