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- T= Kbm

• Freezing Point Evaluation- T= -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 FunctionsG- free energy change– E- energy change

– H- enthalpy change

– S- entropy changeGEHS 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. H0

• 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)* T (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. E= 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, E can be defined as:– E= H - P V

• For many reactions H is very large and the value of P V is relatively small, so that and H 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 h0react 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: G0 = H0 - T S0

– 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?