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					Solutions, Acids and Bases, Thermodynamics,
  Electrochemistry, Precipitation, Chemical
     Equilibrium, and Chemical Kinetics




   By: Karl Lewis, Mark Liv, Kevin Mahon, Doug Reed
                    AP Chemistry-3A
                      2OO3-2004
     Solutions- What are they?
• Substances 3 states of matter- SOLID,
  LIQUID, GAS
• Solution- Basically a mixture of solvents in
  solutes
• EXAMPLE- Salt water, Brass, etc.
• Solutions can move from the 3 states of
  matter but solubility is best undergone in
  the liquid stage of matter.
               Definitions
• Pure Substance- Substance with constant
  composition
• Ideal Solution- Solution‟s vapor pressure
  directly proportional to mole fraction of
  solvent present.
• Solubility- amt. Of substance that dissolves
  in a given volume of solvent at a given
  temperature.
                         Equations
• Enthalpy- E+PV At constant pressure, the change is equal to the
  energy flow of the heat. E- internal energy P- pressure V-Volume.
• Entropy- Randomness of disorder.
• Molarity- M Unit of concentration. # of moles of solute dissolved in 1
  L of solution. #Moles/1L of solution
• Molality- m #moles dissolved in 1Kg. #Moles/1Kg of solution.
• Mole Fraction- # compound moles/T.Mole.
• Mass/Weight Fraction- Mass A/ Mass B+ Mass C+….
• Boiling Point Evaluation- DT= Kbm
• Freezing Point Evaluation- DT= -Kfm
            Factors in solubility
• Structure- Arrangement of crystalline structure matters.
  Re-arrangements of structure happens during solubility.
• Volatility- Readiness to become a gas.
• Pressure- Pressure effects gas solubility in rate of entry
  and exit.
• Temperature- Effect for aqueous. Usually solubility rises
  with temp. rise.
• Process- 1- NRG breaks attraction of solute bonds. 2-
  Solvent molecules break. 3- Molecules combine.
       Distillation Separation
• Distillation depends on volatility. In a
  device, a solution is heated and the liquid
  with the most volatility turns into a gas at
  the lowest temperature. It passes through a
  cool tube and condenses back into a liquid
  into a beaker or so, thus separating the 2
  substances.
                Filtration
• Filtration works with a solid and a liquid.
  Simply, you pour the water through a mesh
  and the liquid goes through and the solid
  stays behind. Solid must be of a good size
  in order to get caught in the mesh
           Chromatography
• This deal with 2 states of matter. A mobile
  phase and stationary.
• Stationary- Solid Mobile- Gas/liquid
• The mixture moves through phases at
  different rates because of their affinities.
  Paper chromatography simply has a sample
  of liquid on paper and reacts to a mobile
  phase.
     Precipitation reaction

• Precipitate reaction-
  solutions that mix sometimes
  produce solids that separate
  from the solution. The solid
  is the precipitate.
         Acid- Base Reaction
• Acid- Proton donor
• Base- Proton acceptor
• Deals with net ionic and spectator ions to
  predict what type of reaction will happen.
• 1- list species present 2- write balanced
  net ionic equation 3- find mole of reactant
  4- find LR 5- convert
       Oxidation- Reduction
• This is also known as redox reaction.
• Oxidation state- imaginary charges an atom
  has if the shared electrons were divided
  equally between identical atoms that are
  bonded.
• Oxidation- increase charge, loss of elc.
• Reduction- decrease charge, gain elc.
    Colligative boiling/freezing
• Colligative- means collective and is the
  change in physical properties of a solution
  after formation.
• Boiling property- nonvolatile solutes
  elevate boiling points.
• Freezing property- when mixtures of
  solutions have a lower freezing point
  because of vapor pressure changes.
          Osmotic pressure
• Osmosis- flow of the solvent into the
  solution through the semipermeable
  membrane that only lets the solvent pass
  through.
• Osmotic pressure- a hydrostatic pressure on
  the solution other then the pure solvent.
• Equilibrium- equal pressure/flow
           Reverse Osmosis
• Reverse Osmosis- The semipermeable
  membrane acting to remove solute particles
  as a molecular filter.
• Isotonic solutions- solutions with similar
  osmotic pressures.
• Dialysis- when the membrane allows
  transfer of solute and small solvent
  particles.
         Electrolyte Solutions
• Ion Pairing- When to particles come
  together to form a single particle.
• Electrolytes dissociate into two- ions when
  dissolved in water. Have effects on
  pressure and points.
• Tyndall effect- scattering of light particles
  to help distinguish between a suspension
  and true solution.
                 Calloids
• Calloid- Suspension of tiny particles in
  some medium.
• These are classified by dispersed phase
  states and mediums. Electrostatic repulsion
  is a factor that helps particles remain
  suspended instead of precipitation out.
  Coagulation is destruction of a calloid.
                Henry‟s Law
• This is a relationship between gas pressure and the
  concentration of dissolved gas. P=kC
• P- partial pressure
• k- is constant characteristic.
• “The amount of a gas dissolved in a solution is
  directly proportional to the pressure of the gas
  about the solution”
• This is obeyed most accurately by dilute solutions
  of gases that don‟t dissociate/react with the
  solvent.
               Raoult‟s Law
• Psoln=Xsolvent P0solvent Psoln = observed vapor
  pressure P0solvent = vapor pressure of pure
  solvent. This is a linear equation of the
  form Y=MX+B.
• Negative deviation when observed vapor
  pressure is lower than the value predicted
  by his law.
           Vant Hoff‟s Law
• Relationship between the moves of a solute
  dissolved and the moves of particles in a
  solution. I= (mole of particle in
  solution)/(moles of solute dissolved)
           Properties of Acids
•   They Burn
•   pH > 7.00
•   Makes litmus paper turn red
•   H+ in chemical formula. Ex) HCl
          Properties of Bases
•   They feel slick and/or slippery
•   pH < 7.00
•   Makes litmus paper turn blue
•   OH- in chemical formula. Ex) NaOH
     Nature of Acids and Bases
• Ahhrenius Concept
  – Acids produce H+ in aquaeous solutions
  – Bases produce OH- in aquaeous solutions

• Brønsted-Lowry Model
  – Acids are proton (H+) donors
  – Bases are proton acceptors

• pH = -log [H+]
• pOH = -log [OH-]
    Nature of Acids and Bases
• Conjugate base - everything that remains of
  the acid molecule after a proton is lost.
• Conjugate acid - formed when the proton is
  transferred to the base.
• Conjugate acid-base pair - two substances
  related to each other by the donating and
  accepting of a single proton.
             Acid Strength
• Involves the percentage of the initial
  number of acid molecules that are ionized.
• Strong acids (I.e. HCl) have nearly 100%
  ionization.
• Weak acids (I.e. HF) have only 1-5%
  ionization.
• Ka - the acid dissociation constant. Will be
  seen again in „equilibrium‟ section.
                 Acid Strength
• The strength of an acid is defined by the
  equilibrium position of its dissociation
  (ionization) reaction:
  – HA(aq) + H2O(l) <=> H3O+(aq) + A-(aq)

• In a strong acid, almost all the original HA
  is dissociated
• In a weak acid, most of the acid originally
  placed in the solution is still present as HA
  at equilibrium
                    Acid Strength
• Common strong acids:
   –   Sulfuric Acid: H2SO4
   –   Hydrochloric Acid: HCl
   –   Nitric Acid: HNO3
   –   Perchloric Acid: HClO4
• Most acids are oxyacids, in which the acidic proton is
  attached to an oxygen atom. The above acids are all
  examples of oxyacids, except for Hydrochloric Acid
  (HCl).
        Water: Acid and Base
• Amphoteric- if a substance can behave as an
  acid or base; I.e. water (H2O). This
  definition came from our textbook.
• An interesting side-note not covered by our
  textbook (taken from the internet):
  – …water is said to be amphiprotic. Water is often
    incorrectly termed amphoteric. An amphiprotic species
    like water can either donate or accept a proton.
    Amphoteric species can both donate and accept
    hydroxide ions, as water cannot.
          Basics of Precipitation
– Precipitation Reactions
    • A precipitation reaction is a reaction in which soluble ions in separate
      solutions are mixed together to form an insoluble compound that settles out of
      solution as a solid. That insoluble compound is called a precipitate.
– Predicting Precipitation Reactions
    • Solubility rules can be used to figure out whether ions that are already in
      solution will come together to form an insoluble compound, that is,
      precipitate.
    • You must use solubility rules to predict precipitation reactions.
    • For Example, Because the solubility rule for "hydroxides" says that sodium
      hydroxide is soluble, sodium ions and hydroxide ions will not come together
      out of solution to form a solid material.
    • On the other hand, the rule for "chlorides" says that lead(II) chloride is
      insoluble. Therefore lead(II) ions and chloride ions already in solution will
      come together to form a solid material that we say "precipitates out of
      solution."
– Writing Equations for Precipitation Reactions
    • Precipitation reactions can be represented using several types of chemical
      equations: complete-formula equations (also known as "molecular" equations),
      complete ionic equations, and net ionic equations. Each provides a different
      perspective on the chemicals involved in the reaction.
– Precipitation Titration
    • In a precipitation titration, the stoichiometric reaction is a reaction which
      produces in solution a slightly soluble salt that precipitates out.
    • For Example, In a precipitation titration of 46.00 mL of a chloride solution of
      unknown concentration, 31.00 mL of 0.6973 molar AgNO3 were required to
      reach the equivalence point. The molar concentration of the unknown solution is
      calculated as follows: 31.00 mL x 0.6973 molar = 21.62 mmol Ag+ = 21.62
      mmol Cl-
    • 21.62 mmol Cl-/46.00 mL Cl- = 0.4700 molar Cl-
– p Notation
    • It is inconvenient to the point of being impractical to plot, or even to compare,the
      changes in ionic concentrations which take place over the course of a
      precipitation titration because the values of the concentrations cover so many
      orders of magnitude in range.). The logarithmic p notation is commonly used not
      only in titration but for the general expression of solution concentrations. In other
      sections this notation, in the form of pH, is extensively used to express the acidity
      of solutions.
           What is Chemical Equilibrium?
–   Le Chatelier’s Principle
      • Le Chatlier's principle allows us to predict the direction a reaction will take when we perturb the equilibrium by
        changing the pressure, volume, temperature, or component concentrations.
      • Simply stated, the principle says that if an external stress is applied to a system at equilibrium, the system will
        adjust itself to minimize that stress.
      • A good non-chemical analogy is two people on a see-saw. If their masses are equal then the see-saw balances. If
        we stress the system by adding weight to one side, the only way we can return to balance is by having the
        heavier person move closer to the fulcrum.
–   The Equilibrium Constants
      • Value that expresses how far the reaction proceeds before reaching equilibrium. A small number means that the
        equilibrium is towards the reactants side while a large number means that the equilibrium is towards the
        products side.
      • The equilibrium constant, Keq is defined as:
      •       [C]c [D]
      • Keq = ---------
      •       [A]a [B]b
      • Products are always in the numerator.
      • Reactants are always in the denominator.
      • Express gas concentrations as partial pressure, P, and dissolved species in molar concentration, [].
      • The partial pressures or concentrations are raised to the power of the stoichiometric coefficient for the balanced
        reaction.
      • Leave out pure solids or liquids and any solvent
– The Reaction Quotient
    • Reaction Quotient is a ratio of molar concentrations of the reactants to those of the products, each
      concentration being raised to the power equal to the coefficient in the equation.
    • Q can be used to determine which direction a reaction will shift to reach equilibrium. If K > Q, a
      reaction will proceed forward, converting reactants into products. If K < Q, the reaction will proceed
      in the reverse direction, converting products into reactants. If Q = K then the system is already at
      equilibrium.
– Mole Fractions
    • The number of moles of a particular substance expressed as a fraction of the total number of moles.
– Spontaneous Reactions
    • A reaction that will proceed without any outside energy.
– Equivalents and Normality
     • Equivalent
           – An equivalent is the amount of substance that gains or loses one mole of electrons in a redox
             reaction, or the amount of substances that releases or accepts one mole of hydrogen ions in a
             neutralization reaction.
     • Normality
         – Normality can only be calculated when we deal with reactions, because normality is a function
            of equivalents.
     • Equivalent weight = molar mass/(H+ per mole)
     • Equivalent = mass of compound / Equivalent weight
     • And Normality = (equivalents of X)/Liter
– Dissociation, self-ionization of water, Kw
   • Pure water is not really pure. The purest water contains some hydronium ions and hydroxide
      ions. These two are formed by the self-ionization of two water molecules. This happens
      rarely. The process is an equilibrium where the reactants, intact water molecules, dominate
      the mixture. At equilibrium the molarities for the hydronium ion and hydroxide ion are equal.
      [H3O+] = [OH-]
   • The equation is
   • H2O + H2O <---> H3O+ + OH-
   • The equilibrium expression is the normal products over reactants.
   • K = [H3O+] [OH-] / [H2O] [H2O]
   • The molarity for the water is a constant at any specific temperature. This means the equation
      can be rewritten as
   • K[H2O] [H2O] = [H3O+] [OH-]
   • The quantity on the right hand side of the equation " K[H2O] [H2O] = Kw " is formally
      defined as Kw. The numerical vale for Kw is different at different temperatures.
   • At 25ºC Kw = 1.0 x 10-14
   • Kw = K[H2O] [H2O]
   • Kw = [H3O+] [OH-] = 1.0 x 10-14
– Acid/Base Dissociation Constants
   • an equilibrium constant (kd) for the dissociation of a complex of two or more biomolecules
      into its components; for example, dissociation of a substrate from an enzyme.
   • the dissociation constant of an acid (ka); or base (kb), describing its dissociation into its
      conjugate base and a proton; or conjugate acid and a hydroxide ion.
                              Chemical Kinetics is...
Reaction Rates
• Reaction Rate
     – A reaction rate is the speed at which reactants are converted into products in a chemical reaction. The reaction rate
        is given as the instantaneous rate of change for any reactant or product, and is usually written as a derivative (e. g.
        d[A]/dt) with units of concentration per unit time.
•   Rate Law
      – A rate law or rate equation relates reaction rate with the concentrations of reactants, catalysts, and inhibitors. For
         example, the rate law for the one-step reaction A + B C is d[C]/dt = k[A][B].
•   Catalyst
      – A substance that increases the rate of a chemical reaction, without being consumed or produced by the reaction.
         Catalysts speed both the forward and reverse reactions, without changing the position of equilibrium. Enzymes are
         catalysts for many biochemical reactions.
•   Enzyme
      – Protein or protein-based molecules that speed up chemical reactions occurring in living things. Enzymes act as
         catalysts for a single reaction, converting a specific set of reactants (called substrates) into specific products.
         Without enzymes life as we know it would be impossible.
•   Arrhenius Equation.
      – In 1889, Svante Arrhenius explained the variation of rate constants with temperature for several elementary
         reactions using the relationship
•   Order
      – The order of a reaction is the sum of concentration exponents in the rate law for the reaction. For example, a
         reaction with rate law d[C]/dt = k[A]2[B] would be a third order reaction. Non-integer orders are possible.
           – k = A exp(-Ea/RT)
           – where the rate constant k is the total frequency of collisions between reaction molecules A times the fraction
              of collisions exp(-Ea/RT) that have an energy that exceeds a threshold activation energy Ea at a temperature
              of T (in kelvin). R is the universal gas constant.
      • Zero Order Reaction
           – A reaction with a reaction rate that does not change when reactant concentrations change
      • First Order Reaction
           – The sum of concentration exponents in the rate law for a first order reaction is one. Many radioactive decays
              are first order reactions.
      • Second Order Reaction
           – A reaction with a rate law that is proportional to either the concentration of a reactant squared, or the product
              of concentrations of two reactants.
      • Half Life
           – The half life of a reaction is the time required for the amount of reactant to drop to one half its initial value.
–   Theories
      • Collision Theory
           – A theory that explains reaction rates in terms of collisions between reactant molecules.
      • Activated Complex
           – An intermediate structure formed in the conversion of reactants to products. The activated complex is the
              structure at the maximum energy point along the reaction path; the activation energy is the difference
              between the energies of the activated complex and the reactants.
      • Integrated Rate Law
           – Rate laws like d[A]/dt = -k[A] give instantaneous concentration changes. To find the change in concentration
              over time, the instantaneous changes must by added (integrated) over the desired time interval. The rate law
              d[A]/dt = -k[A] can be integrated from time zero to time to obtain the integrated rate law ln([A]/[A]) = -kt,
              where [A]o is the initial concentration of A.
Thermodynamics
           Essential Definitions
• System- the part of the universe that is under
  study.
   – Open System- a system that can transfer both energy
     and matter to and from the surroundings. An open
     bottle of perfume is an example of an open system.
   – Closed System- a system where energy can be
     transferred to the surroundings but matter cannot. A
     well-stoppered bottle of perfume is a closed system.
   – Isolated System- a system where there is no transfer of
     energy or matter to or from the surroundings. A
     thermos is a close example of an isolated system.
  Essential Definitions (cont‟d)
• State FunctionsDG- free energy change
  – DE- energy change
  – DH- enthalpy change
  – DS- entropy change
       GEHS GUESS (easy way to remember)
          More Definitions
• Standard State- when the pressure is one
  atmosphere, the temperature is 25º C and
  one mole of compound is present. When the
  thermodynamic quantities are at standard
  state they are represented with a zero
  power. Ex. DH0
• Calorimeter- the device used for measuring
  the heat energy produced by chemical
  reactions and physical changes.
                Types of Energy
• Kinetic energy- energy that matter possesses
  because of its motion.
   – Eq: KE=1/2mv2
      • M=mass in kilograms, v=velocity in meters per second, and
        KE= kinetic energy in joules.
• Potential energy- stored energy. The two types of
  potential energy are gravitational energy and
  electrostatic attraction.
   – Eq: PEgrav= Kgrav(m1m2/r)
         PEelect= Kelect(q1q2/r)
      • m=mass in kilograms, q=charges, r=distance, k=a
        proportionality constant that is different for each type
     Types of Energy (cont‟d)
• Total Energy- the sum of a substance‟s
  kinetic and potential energies.
  – Eq. Energy (E)= potential energy (PE) + kinetic
    energy (KE)
      Measurement of Energy
• Specific heat- the amount of heat needed to
  raise one gram of a substance one degree.
  – Eq. q (heat energy)= C (specific heat)* g (mass
    in grams)* DT (change in temperature)
     • Specific heat of water= 4.184Jg ºC
• Dulong and Petit Law:
  – (Specific heat)(Molar mass)= 25Jmol ºC
  First Law of Thermodynamics
• First Law of Thermodynamics- energy is always
  conserved.
   – Eq. DE= q (heat) + w (work)
• Work= Force x Distance moved
  – Force can be defined as pressure exerted over a given
    area, so. . .
• Work= Pressure x Area x Distance moved
  – Multiplying the area by the distance results in volume
    units, so. . .
• Work= Pressure x Volume change
  – Work is the product of the pressure and the change in
    volume that occurs during a chemical reaction.
           First Law (cont‟d)
• Therefore, DE can be defined as:
  – DE= DH - P DV
    • For many reactions DH is very large and the
      value of P DV is relatively small, so that DE
      and DH are approximately equal.
                   Hess‟s Law
• Hess‟s Law- whatever mathematical operations
  are performed on a chemical reaction the same
  mathematical operations are applied also to the
  heart of reaction.
   – If the coefficients of a chemical reaction are all
     multiplied by a constant, the Dh0react is multiplied by
     that same constant.
   – If two or more reactions are added together to obtain an
     overall reaction the heats of these reactions are also
     added to give the heat of the overall reaction.
Second Law of Thermodynamics
• Second Law of Thermodynamics- any
  physical or chemical change must result in
  an increase in the entropy of the universe.
  – Entropy- the degree of randomness in a sample
    of matter
     • All motion ceases at 0 K or absolute zero and there
       is perfect order, thus 0 entropy.
     • As the temperature of 1 mole is increased from
       absolute zero, the entropy increases.
     • Standard Entropy: S0= qrev (heat added)/T
       (temperature in Kelvin)
          Gibbs Free-Energy
• Gibbs Free-Energy Equation: DG0 = DH0 -
  T DS0
  – Equation is derived directly from the second
    law of thermodynamics.
Electrochemistry
         Essential Definitions
• Electrolysis- a non-spontaneous chemical
  reaction is forced to occur when two
  electrodes are immersed in an electrically
  conductive sample, and the electrical
  voltage applied to the two electrodes is
  increased until electrons flow.
   Essential Definitions (cont‟d)
• Electrolytic cell- a device in which electrolysis can
  be produced, usually consisting of an electrolyte,
  its container, and electrodes.
• Electrolyte- a chemical compound that
  separates into ions in a solution or
  when molten and is able to conduct
  electricity
• Cathode- the negative electrode of an electrolytic
  cell.
• Anode- the positive electrode in an electrolytic
  cell.
Stoichiometric Electrochemistry
• Faraday found that 96, 485 coulombs
  is equal to 1 mole of electrons.
  – 1 coulomb = 1 ampere x 1 second
• mol X = I (current) x t (time) / (n)96,
  485
  – Current measured in amperes and time
    measured in seconds n is the number of
    moles of X
                Example
• A current of 2.34 A is delivered to an
  electrolytic cell for 85 min. How many
  moles of Au from AuCl3 will be obtained?

				
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