Learning Objectives for Final Exam – Spring 2015Burdge & Overby Atoms First -2nd Edition

Chapter 11 Students should be able to: Describe the properties of gases

o Expand to fill container + take container’s shapeo Compressibleo Much smaller density/variable densityo Homogenous mixtures (solutions)

Use the assumptions of kinetic molecular theory to explain the behavior of gases

o Four Basic Assumptions A lot of empty space between the gas particles

compared to the size of the particles Attraction or repulsion between particles is

negligible The particles of the gas are constantly moving in

straight paths with perfectly elastic collisions (KE lost by one is gained by other)

The average KE of the gas particles is directly proportional to K (temp)

As the temp increases, the average speed increases, but not all particles move at the same speed (AVERAGE!!!)

o Gases in same container have same temp, so have the same avg KE

o If particles have different masses In order to have same kinetic energy the average velocities vary

Light particles have faster average velocityo Equations:

Where NA*mass= molar mass in kg/mol Where R= 8.314 J/mol*K Urms has the units m/s (velocity)

Root mean square speed is the speed of a

molecule w/ the average KE in a gas sample

Understand the relationship between effusion and molecular mass

o Effusion The escape of gas molecules from a container to

a vacuum

o Graham’s Law The rate of diffusion/effusion is inversely

proportional to the square root of its molar mass

Define the standard temperature and pressure (STP)

o Standard atmospheric pressure is 1 atm= 760 mmHg= 760 torr

o Standard temperature =273 K (0 C) List factors that may cause a gas to deviate from ideal

behavioro Ideal gas law assumes

No attraction between molecules Gas molecule volume insignificant

o 2 Conditions where gases deviate At low temp and high pressure

High pressure: molecules are close together so volume is significant

o Real gases take up space so molar volume is larger than predicted

Low temp: molecules are slower so intermolecular forces are significant

o Molar volume is smaller than predicted

Use the relationships among P, V, T, and n in calculationso Boyle’s Law

Pressure is inversely proportional to volume P1V1=P2V2

o Charles’s Law Volume is directly proportional to temperature V1/T1=V2/T2

o Avogadro’s Law Volume is directly proportional to number of gas

molecules V1/n1=V2/n2

Equal volumes=Equal # of moleculeso Combined Gas Law

Use the ideal gas law to determine density or molar masso Ideal Gas Law

R= 0.08206 L*atm/ mol*K

o The volume of 1 mole of gas at STP= molar volume 22.4 L

o To calculate density

where M is the molar mass At standard conditions the density is the (molar

mass/ 22.4 L) Perform gas mixture calculations using Dalton’s Law of

Partial Pressures and mole fractionso The total pressure of a mixture is the sum of partial

pressures exerted by each gaso PV=nRT so Pgas=nRT/V where R=0.08206 Latm/molKo Ptotal= PA+PB+PC…. o Mole Fraction

3 Things to Remember

Mole fraction of component always less than 1 Sum of mole fractions =1 Mole fractions are dimensionless

Perform stoichiometric calculations for reactions with

gaseous reactants and/or products.o Amount of gas is given as a Volume

Convert volume to moles Use coefficients in balanced equation as

mole ratioso STP1 mole=22.4 L

Chapter 12 Students should be able to: Describe the physical properties of liquids (surface tension,

viscosity, vapor pressure, boiling point) and relate these to the types of attractive forces experienced by the sample of matter

o Intermolecular forces Ion-dipole(ionic mixtures)>Hydrogen bonding

(molecules containing H,F,O,N) > dipole-dipole (polar molecules) > dispersion (every molecule/atom)

Magnitude of dispersion force increases with molar mass

o Surface tension Amount of energy required to STRETCH or

INCREASE the surface of a liquid by a unit area The STRONGER the intermolecular forces

the HIGHER the surface tension Factors Affecting surface tension

Raising the tempReduces surface tension (increases KE making it easier to stretch surface

Reducing the temp would have opposite effect

Cohesion Attraction between like molecules

Adhesion Attraction between unlike molecules

(meniscus is seen when adhesion >

cohesion and convex shape seen when cohesion > adhesion)

o Viscosity Measure of a fluid’s resistance to flow

The higher the viscosity the more slowly a liquid flows

The GREATER the intermolecular forces the HIGHER the viscosities

o Vapor Pressure The equilibrium pressure of vapor above its liquid

(rate of vaporization) Vaporization is endothermic

The GREATER the intermolecular forces the LOWER the vapor pressure

THE WEAKER the intermolecular forces the GREATER the vapor pressure

Vapor pressure INCREASES with TEMP Dynamic Equilibrium

When two opposite processes reach the same rate

o Boiling Point The temperature at which the vapor pressure of

the liquid=1 atm (or the external atmospheric pressure)

The lower the external pressure, the lower the boiling point

Predict the physical properties of various solids (melting point, vapor pressure, and amorphous vs. crystalline structures)

o Melting Point Temperature at which particles can break apart

from fixed positions o Vapor Pressure

Solids have low vp at room temperature E.g. iodine, carbon dioxide

o Amorphous Solids Lack 3D arrangement, melt over a range of

temperatureso Crystalline Solids

Possess rigid and long-range order Unit cells

o Basic repeating structure

o

Describe the basic types of crystalline solids (ionic, covalent, molecular, metallic)

o Ionic Held together through ionic bonds Hard, brittle, high MP, poor conductors

o Covalent Held together through covalent bonds Very hard and brittle, very high MP, poor

conductorso Molecular

Held together through weak intermolecular interactions such as Van der Waals, H-bonding, dipole-dipole and dispersion

Low MP, poor conductorso Metallic

Held together by metallic bonds Variable hardness and MP

Identify phase changes and their associated enthalpies

Interpret heating/cooling curves

o

Calculate heat(q) changes with temperature and/or phase changes

o Solid Warming q=mass X Cs X delta T Cs is the specific heat

o Solid to Liquid in Equilibrium q= n* delta Hfus

o Liquid Warming q=mass X Cs X delta T Cs is the specific heat

o Liquid to Gas in Equilibrium q= n* delta Hvap

o Gas Warming q=mass X Cs X delta T Cs is the specific heat

qtotal= q1+q2+q3+q4+q5 (KJ/mol)

Interpret a phase diagram

o Triple point= all three phases in dynamic

equilibrium Critical temperature= temp above which its gas

cannot be liquified Critical pressure= min pressure that must be

applied to liquify a substance at critical temp

Chapter 13 Students should be able to: Describe types of solutions, including all phases

o Unsaturated Contains less solute than the solvent has capacity

to dissolve

o Saturated Contains the maximum amount of solute that will

dissolve in a solvento Supersaturated

Contains more dissolved solute than is present in a saturated solutionunstable

Describe changes in entropy and enthalpy when solutions are formed

o Solvation depends on Solute-solute interactions

Must overcome (endothermic) Solvent-solvent interactions

Must overcome (endothermic) Solute-solvent interactions

New solute-solvent attractions provide energy (exothermic)

o Entropy (S) A measure of how dispersed or spread out the

energy is in a system Natural tendency for entropy to increase

Describe how environmental factors (T and P) affect the solubility of a gas in a liquid

o Solubility of one substance in another depends on Types of intermolecular forces Nature’s tendency towards mixing (increase in

entropy)o Solubility usually increases with temperatureo Henry’s law

The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution

More pressure increased solubilityo Molarity and Molality

Molarity is moles of solute/L of solution Molality is moles of solute/kg of solvent

Molality does not vary with temperatureo Parts solute in Parts Solution

0.9% by mass= 0.9 g of solute in 100 g of sol’n 36 ppm by volume= 36 mL of solute in 1 million

mL of sol’n Use Henry’s law to calculate solubility of a gas in a solution

o Explain and calculate different physical properties between

solutions and pure solvent (colligative properties) o Colligative properties depend on # of solute particles

in solution vapor pressure lowering

vapor pressure of solution is lower than vapor pressure of pure solvent

boiling point elevation elevation

solutions boil at a higher temperature than the pure solvent

freezing point depression

solutions freeze at a lower temp than pure solvent

osmotic pressure

osmosis is movement of water high to low concentration

osmotic pressure is the pressure required to stop osmosis

Chapter 14Students should be able to: Distinguish between a spontaneous and a nonspontaneous

processo Exothermic (energy released) ΔH negativeo Endothermic (energy absorbed) ΔH positiveo Spontaneous reactions occur under specific set of

conditions Most release energy (exothermic) high PE to low PE Some are endothermic (like melting ice)

o Nonspontaneous reactions require energy input and do not occur under a specific set of conditions

o Spontaneity Compare chemical potential energy before and

after reactiono Less PE after reaction means it’s favored

Favored by an increase in entropy Define entropy (S) and Gibbs free energy (G)

o Entropy (S) Measure of how spread out or how dispersed a

system’s energy iso Gibbs free energy (G)

Since measuring the surroundings is impractical, we use Gibbs free energy change to describe spontaneity in a system

Recognize the conditions necessary for standard entropy (S°) Identify trends in standard entropy values

o Changes that increase S Products are in a more random state (mobility)

Solid to liquid to gas Larger numbers of product molecules than reactant

molecules-More Moles Increase in temperature Increase in molecular complexity Solids dissociating into ions upon dissolving

(sometimes) Molecular solutes S increases Ionic solutesS can increase or decrease

Calculate the standard entropy change (ΔS°) for a given reaction

o

o

Predict the sign of ΔS for a given process and use the sign to indicate whether the system has undergone an increase or decrease in entropy

o Calculate ΔSsurr given ΔHsys and temperature Determine whether a process is spontaneous given ΔSsurr and

ΔSsys State in words or symbols the second law and third law of

thermodynamics o Second Law of Thermodynamics

For a process to be spontaneous, ΔSuniverse must be positive

o Third Law of Thermodynamics states that the entropy of a perfect crystalline substance is ZERO at absolute zero (S increases as Temp increases)

Define Gibbs free energy (G) Calculate ΔG from temperature, ΔH, and ΔS

o

Use the sign of ΔG to indicate if a process is spontaneous

o

Predict the sign of ΔG given the signs of ΔH and ΔS at high and low temperatures

o

Calculate the standard free energy change (ΔG°) of a given reaction

o Calculate the temperature at which a process becomes

spontaneous o Use ΔG = ΔH – TΔS, substituting delta G for 0

Calculate ΔS then substitute into equation and solve for T (K) Convert to C if needed

Chapter 15Students should be able to: Define equilibrium Distinguish between reversible and irreversible processes Write the equilibrium constant (K) expression for a given

reactiono K=[Products]/[Reactants]

Explain the relationship between the equilibrium constant (K) and the reaction quotient (Q)

Calculate K given equilibrium concentrations of reactants and products

Predict the relative amounts of reactants and products at equilibrium given the equilibrium constant (K)

Calculate the equilibrium concentration of reactants or products given initial concentrations.

Differentiate between heterogeneous and homogeneous equilibria

o Heterogeneous equilibriao Homogeneous equilibria

Convert between Kc and KP for a reaction involving gases Calculate the reaction quotient (Q) and predict the direction

of a reaction given initial concentrations of reactants and products and the equilibrium constant (K).

Calculate ΔG and ΔG° of a reaction at a specified temperature given Q or K.

Use an ICE table and, if necessary, the quadratic formula to determine equilibrium, initial or final concentrations of reactants and products

Employ and explain the “x is small” approximation Predict the shift of a reaction using Le Châtelier's principle

given a change in one of the following: removal or addition of reactant or product, change in volume or pressure, and temperature change.

Chapter 16Students should be able to: Write base/conjugate acid and acid/conjugate base pairs Define through words and examples Brønsted acids/bases

and Lewis acids/bases o Bronsted acido Bronsted baseo Lewis acido Lewis base

Define amphoteric and give examples of substances that are amphoteric

Determine the relative strength of acids based on their composition and structure

o Strength of Acids depends on Write the equilibrium expression for water (Kw) and use it to

determine whether a solution is acidic, basic, or neutral Classify a solution as being acidic, basic, or neutral using

the pH scale o Acidic solutionso Basic solutionso Neutral

Calculate pH, pOH, concentration of hydroxide ion, or concentration of protons

Identify an acid or base as being strong or weak o Strong Acido Strong Baseo Weak Acido Weak Base

Calculate the pH of a weak acid or base using an ICE table and, the acid or base-dissociation constant (Ka or Kb) and, if necessary, the quadratic equation

Employ the “x is small” approximation Explain the relationship between Ka, Kb, and Kw Calculate the percent ionization of a weak acid or base Use the pH of a weak acid or base solution to calculate

the Ka or Kb, respectively. Calculate the Ka or Kb of conjugate pairs using Kw  Identify polyprotic acids and calculate their pH at a given

concentration Classify a salt as being basic, acidic, or neutral based upon

the acid and base used to form the salto Basic salto Acidic salto Neutral salt

Calculate the pH of a salt solution

Chapter 17Students should be able to:

Recognize that a solution containing a weak acid and a salt containing its conjugate base is a buffer solution

Write the chemical reactions for a buffer reacting with strong acid or strong base

Calculate the pH of buffers given initial amounts or concentrations of a weak acid and a salt containing its conjugate base or a weak base and a salt containing its conjugate acid

Select an appropriate acid and salt to prepare a buffer of a specific pH

Predict whether the products of an acid-base neutralization (titration) will be acidic, basic or neutral

Calculate the pH (initially, at the mid-way point and at the equivalence point) for a titration between an acid and a base - where either both are strong, or one is strong and the other weak

Describe or identify the different profiles of titration curves depending upon the acid and base strengths

Select an appropriate indicator (given a chart of indicators and pH range) for a given acid-base titration

Write Ksp expressions given the name or formula of a sparingly soluble salt

Calculate molar solubility from Ksp values and vice versa Predict how factors such as a common ion or pH will affect

solubility of a specific salt

Chapter 18Students should be able to: Balance oxidation-reduction reactions taking place in

neutral, acidic or basic solutions Predict the direction of the flow of electrons, identify the

anode and cathode, and calculate the cell potential for a galvanic cell

Use line notation to write the spontaneous redox reaction of a galvanic cell

Identify the various components of a galvanic cell from the line notation

Given two half cells, predict the spontaneous reaction, and calculate the standard cell potential (E°), free energy (ΔG) and the equilibrium constant (K). [Equations will be provided.]

Use the Nernst equation to calculate cell potentials under non-standard conditions

Explain the driving force of concentration cells

Chapter 19Students should be able to: Describe the factors (concentration, temperature, catalyst)

that affect the rate of reactions, and describe how these factors affect reaction rates

Express the rate of a reaction in terms of rate of loss of reactants or rate of formation of products given the balanced chemical reaction

Write rate laws for first-order, zero-order and second-order reactions

Determine the rate law and calculate the rate constant for a reaction using the method of initial rates

Use graphical treatment of kinetic data to determine reaction order, rate law, and the rate constant

Use integrated rate laws to calculate the concentration of reactant at a given time

Use half-life and reaction order to calculate the concentration of reactant at a given time

Use the Arrhenius equation or an Arrhenius plot (ln k vs. 1/T) to determine activation energy of a reaction

Interpret reaction profile (energy vs. reaction progress) graphs to find the activation energy of forward, reverse or catalyzed reactions

Write rate laws for elementary reactions Given several reaction mechanisms, be able to determine

which are consistent with the stoichiometry and rate law for the reaction.

Describe how a catalyst (or an inhibitor) affects reaction rates

Identify reaction intermediates in a given reaction mechanism

Please Note: Nomenclature and the writing of inorganic formulas will be included throughout the exam.