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


1
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

2
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.

3
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.

4
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

5
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.

6
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.

7
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

8
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.

9
Precipitation reaction
  • Precipitate reaction- solutions that mix
    sometimes produce solids that separate from the
    solution. The solid is the precipitate.

10
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

11
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.

12
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.

13
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

14
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.

15
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.

16
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.

17
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.

18
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.

19
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)

20
Properties of Acids
  • They Burn
  • pH gt 7.00
  • Makes litmus paper turn red
  • H in chemical formula. Ex) HCl

21
Properties of Bases
  • They feel slick and/or slippery
  • pH lt 7.00
  • Makes litmus paper turn blue
  • OH- in chemical formula. Ex) NaOH

22
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-

23
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.

24
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.

25
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

26
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).

27
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.

28
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."

29
  • 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.

30
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

31
  • 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

32
  • 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.

33
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.

34
  • 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.

35
Thermodynamics
36
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.

37
Essential Definitions (contd)
  • State FunctionsDG- free energy change
  • DE- energy change
  • DH- enthalpy change
  • DS- entropy change
  • GEHS? GUESS (easy way to remember)

38
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.

39
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40
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

41
Types of Energy (contd)
  • Total Energy- the sum of a substances kinetic
    and potential energies.
  • Eq. Energy (E) potential energy (PE) kinetic
    energy (KE)

42
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

43
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.

44
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.

45
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.

46
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)

47
Gibbs Free-Energy
  • Gibbs Free-Energy Equation DG0 DH0 - T DS0
  • Equation is derived directly from the second law
    of thermodynamics.

48
Electrochemistry
49
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.

50
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51
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.

52
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

53
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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