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

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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

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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.

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Definitions

• Pure Substance- Substance with constant
  composition
• Ideal Solution- Solutions 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.

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Equations

• Enthalpy- EPV 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

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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.

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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.

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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

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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.

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Precipitation reaction

• Precipitate reaction- solutions that mix
  sometimes produce solids that separate from the
  solution. The solid is the precipitate.

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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

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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.

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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.

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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

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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.

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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.

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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.

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Henrys Law

• This is a relationship between gas pressure and
  the concentration of dissolved gas. PkC
• 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 dont dissociate/react
  with the solvent.

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Raoults Law

• PsolnXsolvent P0solvent Psoln observed
  vapor pressure P0solvent vapor pressure of
  pure solvent. This is a linear equation of the
  form YMXB.
• Negative deviation when observed vapor pressure
  is lower than the value predicted by his law.

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Vant Hoffs 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)

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Properties of Acids

• They Burn
• pH gt 7.00
• Makes litmus paper turn red
• H in chemical formula. Ex) HCl

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Properties of Bases

• They feel slick and/or slippery
• pH lt 7.00
• Makes litmus paper turn blue
• OH- in chemical formula. Ex) NaOH

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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-

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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.

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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.

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Acid Strength

• The strength of an acid is defined by the
  equilibrium position of its dissociation
  (ionization) reaction
• HA(aq) H2O(l) ltgt 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

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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).

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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.

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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."

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• 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.

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What is Chemical Equilibrium?

• Le Chateliers 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
• Cc D
• Keq ---------
• Aa Bb
• 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

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• 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 gt
  Q, a reaction will proceed forward, converting
  reactants into products. If K lt 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

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• 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 lt---gt 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
• KH2O H2O H3O OH-
• The quantity on the right hand side of the
  equation " KH2O 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 KH2O 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.

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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. dA/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 dC/dt kAB.
• 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
  dC/dt kA2B would be a third order
  reaction. Non-integer orders are possible.

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• 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 dA/dt -kA 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 dA/dt
  -kA can be integrated from time zero to time to
  obtain the integrated rate law ln(A/A) -kt,
  where Ao is the initial concentration of A.

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Thermodynamics
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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.

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Essential Definitions (contd)

• State FunctionsDG- free energy change
• DE- energy change
• DH- enthalpy change
• DS- entropy change
• GEHS? GUESS (easy way to remember)

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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.

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Types of Energy

• Kinetic energy- energy that matter possesses
  because of its motion.
• Eq KE1/2mv2
• Mmass in kilograms, vvelocity 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)
• mmass in kilograms, qcharges, rdistance, ka
  proportionality constant that is different for
  each type

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Types of Energy (contd)

• Total Energy- the sum of a substances kinetic
  and potential energies.
• Eq. Energy (E) potential energy (PE) kinetic
  energy (KE)

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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

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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.

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First Law (contd)

• 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.

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Hesss Law

• Hesss 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.

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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)

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Gibbs Free-Energy

• Gibbs Free-Energy Equation DG0 DH0 - T DS0
• Equation is derived directly from the second law
  of thermodynamics.

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Electrochemistry
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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.

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Essential Definitions (contd)

• 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.

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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

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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?