1 Exam
Gases o States of matter- solid, liquid, gas o Gas material with no definite shape or volume move more readily N2 is the most common gas (air) Air: N2, O2, H2O, Ar, He, CO2, O3, CH4 o What kind of substances are gases? Mostly small molecules Some free elements H2, O2, Ar, Noble gases, N2, F2, Cl2, Br2, Hg Smells are molecules in their gas state o Microscopic characteristics Particles are randomly moving Particles will collide- elastic collisions (don't lose energy) Particles exert pressure by bouncing off walls Simple relationship between moles, pressure, volume, and temperature (PV=nRT) o Microscopic Properties Gases are compressible Lower density than solids of liquids Completely mix- no "immiscible" gas o Units of Pressure Pressure = F/A (Force/Area) SI units of Area = m2 SI units of Force = N N/m2 = 1Pa Atmospheric Pressure at sea level = 101 kPa = 1atm 760 torr = 1atm = 760 mm Hg Relationship of physical properties of gas V P (inverse proportion) o V 1/P o V = k(1/P) where k=constant o VP = k o Boyle's Law V1P1 = V2P2 V T o VT o (V1/T1) = k = (V2/T2) No change in pressure o Charles' Law (V1/T1) = (V2/T2) V n (n = # of moles) o Vn o Avogardro's Law
(V1/n1) = (V2/n2) No change in P or T The number of moles does not change by substance! PV=nRT o R = ideal gas constant = 0.082057 Latm/molK o See Figure 1 o Density of a Gas Much lower than liquids or solids g/L or g/cm3 D = PM/RT M = molar mass o Dalton's Law of Partial Pressure (1801) Total pressure of the mixtue of gases is equal to the sum of their individual pressures Implies that P depends on the total number of moles, not the chemical nature Ptot = Pa + Pb + Pc +..... Ptot = ntotRT/V If you have two gases, ntot = na + nb Ptot = (na + nb)RT/V o Kinetic Molecular Theory of Gases Gas molecules are separated by large spaces Random motion with collisions Average kinetic energy of gas particles is proportionate to the gas's temperature U = (3RT/M) Where U = average of the square of speeds and M = molar mass Average speed depends on T and molar mass RMS (root mean square) o Gas Diffusion Gradual mixing of 2 gases randomly Graham (1832)- at a constant pressure and temperature, the rate of diffusion of gases is inversely proportional to the square of M r1/r2 = (M1/M2) r = rate M = molar mass Lighter gas diffuses more quickly o Gas Effusion Process by which gas escapes out of a tiny opening in a container Graham's law works here too o Ideal Gas Assumptions no attractions between molecules ("stickiness") assume that the molecules have no volume This is where the real gas law comes in Intermolecular Forces o Forces that hold molecules together
o Not bonds! Only forces/attractions Dipole-dipole Attraction between - and + on different molecules (two polar molecules) Dipole-induced dipole Between 2 different substances: 1 polar and 1 non-polar Ex) the dipole of water repels electrons to make a dipole on C2H6 London Dispersion Forces (LDF) Occur between all molecules Aka: induced dipole-induced dipole and Van der Waals Forces Weak Relatively insignificant Hydrogen Bonding (not a bond!) The IMF between and H atom on an N, F, or O and its attraction to the lone pairs of electrons on another N, F, or O atom o N, F, and O are the smallest and most electronegative If there is hydrogen bonding, there are dipole-dipole attractions C-H: never involve hydrogen bonding because they are non-polar and there is no attraction o Molecular Level Solids- little movement and lots of strong IMFs Liquids- molecules close together, but can move Gases- molecules are far apart, no IMF Bigger molecule = more IMF's o Polarizability Ability of a molecule to react to a dipole o Polarity Similar polarities mix Similar polarity = similar IMFs o Transition from liquids to gas Liquids- lots of IMF Gas No IMF Liquid to gas = evaporation (boiling) High IMF = slower evaporation = higher boiling point As you go down the periodic table the IMFs increase Vapor Pressure increases as the temperature of the liquid goes up Boiling point is when vapor pressure equals external pressure o Viscosity Resistance to flow in a liquid As IMFs go up, the viscosity also goes up o (IMFs and large molecules affect viscosity) Capillary effect pulls a liquid up into a small tube (paper and marker) Non-polar molecules will move along paper Polar molecules will stay attached to the paper
Graduated Cylinder- read the the bottom of the meniscus- there is an attraction (IMF) between the water and the glass (what makes the water creep up the sides) Adhesive forces- attraction of water to a solid Cohesive forces- IMF between liquid In water and glass cohesive<adhesive In mercury and glass cohesive>adhesive Surface Tension o Molecules near the surface are closer together Takes extra energy to pierce a liquid surface This is because of hydrogen bonding between water molecules o Phase Diagram
Pressure-Temperature phase diagram for CO2 Triple point- all three phases exist in equilibrium Critical Point- where the liquid phase of the matter ceases to exist o Supercritical Fluid Can't tell if it is a liquid or a gas Caffeine extraction uses supercritical CO2 Supercritical water Flammable Used for extractions (to remove contaminants) Hard to reach (225atm) Boiling Point- the point where vapor pressure equals atmospheric pressure o Stronger IMFs = higher boiling point o Stronger IMFs = lower melting point Ionic bonds are the strongest If the slope of the line between liquid and solid is positive than the solid is more dense than the liquid If the slope of the line between liquid and solid is negative (like in water), than the solid is less dense than the liquid o Reason why ice floats in water Solids o Structural materials Building
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Transportation Electronic materials Optical materials (fiber optics) Thermoelectrics Solids that get cold when conducting electricity HUGE area of research Kinds of Solids Molecular solid (ice and CO2) Solid made of molecules Held together by IMFs Relatively weak Low melting points Sugar, dry ice, water Non-conducting Ionic solids (ions) Made of ions NaCl (Na+ and Cl-) Very high melting points Very brittle Non-conducting Metallic solids (metals) Held together by metallic bonds Positive cations in a "sea of electrons" Delocalized e- are responsible for metallic properties o Electrical conductivity, heat conductivity and reflectivity Non-directional bonding o Ductile (wire), malleable o Bonds rearrange but don't break Covalent networks Held together by covalent bonds into 2D infinite structures (graphite) or 3D infinite structures (diamonds) Very high melting point Very hard Usually an insulator or semi-conductor Two Categories of Solids Amorphous Lacking form or shape Lacking a well-defined long range molecular level structure Glass- random chains of corner sharing Crystalline Posses long-range well defined atomic level structure Diamond, NaCl, graphite, ceramics The smallest repeating unit = unit cell Unit cell can be used to recreate the entire structure
Exam 2
Solids- crystalline solids o Unit cell- repeating unit present in crystalline structures o Unit cell matches formula as a multiple Simple cubic cell A=B=C All angles are 90 Tetragonal unit cell A=BC All angles are 90 Orthorhombre unit cell ABC All angles are 90 o Special locations 8 corner atoms Atom split by 3 planes x x = 1/8 8 corner atoms at 1/8 each = 1 atom/unit cell 12 edge atoms Atom split by 2 planes x= 12 corner atoms at each = 3 atoms/unit cell 6 face atoms Atom split by 2 planes 6 atoms at each = 3 atoms/unit cell Body atom Atom located completely within the unit cell 1 body atom 1 body atom = 1 atom/unit cell o Only have one type of atom present (element) Simple cubic structure 1 atom/unit cell Very rare Only 2 elements that are simple cubic- Po and Hg Very inefficient o A lot of space o 52% full Body centered cubic 8 corner atoms at 1/8 each = 1 atom/unit cell 1 body atom 2 atoms/unit cell 68% full Fe, alkali metals Face centered cubic Corner atoms + face atoms
8 corner atoms at 1/8 each = 1 atom/unit cell 6 face atoms at each = 3 atoms/unit cell 4 atoms/unit cell 74% full o Highest possible for a pure element Al, Ag, Au, Cu, Ni, Pt Hexagonal Closely packed For body centered, face centered, and hexagonal close pacjed Have planes of atoms which can slide past each other Ex.) steel = Carbon + Iron o Carbon sits in holes between iron o Can't shift New way to look at bonding o Old way doesn't work well for metals o O2 (l) is paramagnetic (has unpaired electrons) o New concept called molecular orbitals Modern view o Basic Rules of Molecular Orbital Theory Total number of Molecular Orbitals (MOs) = number of atomic orbitals used Bonding orbitals will always be lowering in energy than parents Antibonding orbitals will always be higher in energy than parents Electrons will fill orbitals from lowest to highest energy Orbitals combine best with orbitals of the same energy See Figure 2 o Bond Order (e- in bonding orbitals e- in antibonding orbitals) See Figure 3 o How to combine p-orbitals e- are waves that want constructive interference in order to build on each other o Reactions occur in the HOMO- highest order molecular orbital o LUMO- lowest unoccupied molecular orbital Solid Materials o Metals, covalent networks o Each atom brings orbitals to make molecular orbital o Metals Promote electrons to higher energy, you need very, very little energy (conductors) Electrons once excited can easily move through the entire metal, due to MOs Therefore metals conduct electricity See Figure 4 o Insulator Takes a lot of energy to promote electrons to higher energy
Insulators do no conduct electricity See Figure 5 o Semi-conductors Requires energy, but is possible Can switch between conducting and not conducting by applied voltage See Figure 6 o Conductivity as a Function of Temperature For metals As T increases conductivity decreases Higher T = more Energy (move nuclei) o Impedes flow of electrons For Semi-conductors As T increases conductivity increases Additional energy naturally excites more electrons into conducting band For Special Metals- Super Conductors Below a certain T, conduct electrons without any resistnce Solutions o What makes up a solution? Solvent- liquid (component in greater amount) Solute- what is dissolved (component in lesser amount) Solution will have the same phase as the solvent o Consider solubility with liquids as a solvent Solubility- how much solute dissolves Effect of pressure doesn't really effect solubility on liquid and solid solutions However, gases do see an effect o As pressure increases solubility increases (in a liquid) Henry's Law (gas) o Sg = KHPg o Solubility = constant (pressure) o Ex.) scuba diver goes down pressure goes up O2 and N2 become more soluble in the blood When diver comes up, pressure decreases, solubility decreases and therefore the gases come out of solution Break capillaries The solution is to use a mix of O2 and He He is non-soluble at high pressure Effect of temperature on solubility As temperature increases Sg decreases Boiling H2O o First get bubbles because the gas is coming out of solution Solubility of liquids Depend on IMFs Liquids that mix are miscible
Liquids that don't mix are immiscible As the temperature increases the solubility of a liquid increases o Overcome differences in IMFs Solubility of solids When solids dissolve o If heat is released H<0 Exothermic reaction Liquid IMFs are greater than solid IMFs (lower energy) o If energy is absorbed H>0 Endothermic reaction Solid IMFs are greater than liquid IMFs o As temperature increases, solubility increases Colligative Properties o Properties of solutions depend only on the concentration of solute particles, not on the chemical nature Vapor Pressure Boiling Point Elevation Freezing Point Depression Osmotic Pressure o Vapor Pressure Fewer molecules can escape because some surface sites are taken by solute See Figure 7 Psolvent = (Xsolvent)(Psolvent) P = Vapor Pressure X = mole fraction P = pressure of pure solvent Xsolvent = nsolvent/[(nsolvent + nsolute)(i)] o i is the dissociation constant o See Figure 8 o Boiling Point Elevation The boiling point of the solution is always higher than the boiling point of the solvent T = Kbpmi (where m = molality) As V increases, T increases and Molarity decreases Concentration is dependent of T See Figure 9 o Freezing Point Depression Solid of pure substance Maximize IMF As solution Not quite as "pretty" of an arrangement Don't "fit" right Lower IMFs than pure solvent Freezing point goes down
Expect lower IMFs to result in a lower freezing point Example: salt on the roads lowers the freezing point T = Kfp(m)(i) where m = molality See Figure 10 o Osmotic Pressure See Figure 11 Osmosis-water moves to try to equilibrate concentrations Difference in pressure = osmotic pressure = MRTi M = molarity R = constant T = temperature i = dissociation constant Chemical Kinetics o Chemical kinetics- the study of rates of reactions Going into the arrow of the reaction o Rate of chemical reaction- rate of change of concentration of reactants See Figure Depends on which time period Reaction rate changes over time Average rate of reaction o Over time, rate of reaction slows, so average goes down Instantaneous Rate- rate of reaction at any moment o Derivation of concentration vs. time gives instantaneous rate What effects rate of reaction? o Temperature o Solubility o Pressure o Concentration of product and reactants o Quantity o IMFs o Collision Theory of Reactions Reactants must collide for reaction to occur Reactants must collide with "sufficient energy" Reactants must collide in the proper orientation to react Examples o As concentration increases more collisions occur and therefore the rate of the reaction increases o As temperature increases, speed increases, kinetic energy increases and therefore the rate of the reaction increases o Catalyst- increases the rate of reaction but isn't used up o In general... aA + bB cC + dD Rate of reaction = k[A]x[B]y x and y (exponents) must be determined experimentally
ex: 2NO2 2NO + O2 o rate of the reaction = k [NO2]2 Determining Rate Laws o Integrated rate laws First Order Rate Law ln[R]t/[R]o = -kt linear Second Order Rate Law 1/[R]t = kt + 1/[R]o Slope = k Zero Order [R]t = -kt + [R]o The change in concentration has no effect on the rate of reaction for zero order reactions In order to determine the correct order, graph all of them and look for the best R2 value o Pseudo Rate Laws- when one reactant is vast in excess Half Life (Radioactive Decay) o 1st order: t1/2 = ln2/k o 2nd order: t1/2 = [R]o/2k o 3rd order: t1/2 = 1/k[R]o Reaction Mechanisms o Step by step process of breaking and making bonds o Sum of steps of mechanism must equal reaction NO + O3 NO2 + O2 Single step mechanism o All bonds occur at once Elementary steps- individual steps of mechanisms Intermediate steps- species that are formed and then subsequently used in a mechanism Unimolecular steps- involves 1 molecule Bimolecular steps- involves 2 molecules Termolecular steps- involves 3 molecules Catalyst- a species that is used then subsequently reformed o How do we distinguish between mechanisms? Individual steps have their own rate laws in these rate laws, exponents are determined by coefficients rate of any reaction is defined by the slowest step of the mechanism if first step is slow than mechanism #1 is the rate law However, you do not want any intermediates in the rate law (if slow step is not the first step) All steps before the slow step are assumed top be in equilibrium (reaction occurring at equal rates in both directions) Rate of forward reaction = rate of reverse reaction Intermediates after the slow step always have a very low concentration (will be used up as soon as it's formed)
Intermediates before the slow step can build up in concentration Exam 3 Kinetics (con't) o See Figure 11 Equilibrium o Rates of forward and reverse reactions are equal o No macroscopic change Example: H2O + SO3 H2SO4 Rate = k[H2O][ SO3] In equilibrium... o k1[H2O][SO3] = k-1[H2SO4] So far we've only looked at kc (indicates concentration) o Kp = equilibrium using pressure Used only for gases See figure 12 All k's are unitless Relationship between kc and kp o kp = kc(RT)n Reaction Quotient- not in equilibrium o Q = [C]c[D]d/[A]a[B]b o Compare Q to keq (kc or kp) If Q<k then there needs to more products so the reaction shifts to the right If Q > k then there needs to be more reactants so the reaction shifts to the left If Q = k then the reaction is at equilibrium Le Chatelier's Principle o A system at equilibrium will try to counter any stress For gases, if volume decreases then pressure increases It is important to know whether the reaction is endothermic or exothermic Steps to Solve Equilibrium Problems o Write out equilibrium o Use Q to determine the direction of shift, if not in equilibrium (if one of the products or reactants is missing, reaction will shift to from it) o Set up ICE table (use balanced equation in order to find relative change) o Plug "E" values into our k expression o Solve algebraically for "x" (change) o Use x to find equilibrium values o See Figure 12 Acid/Base o Proton transfer chemistry o Brnstead Definition Acid: donates H+ in aqueous solutions Base: accepts H+ in aqueous solutions o Acid Strong acids
HCl, H2SO4, HBr, HI, HNO3 HClO4 Which means they completely dissociate in water Conjugate acid- species formed when the base gains H+ Conjugate base- species formed when the acid loses H+ o Base Strong bases OH- (soluble OH- salts), H-, RO-, R-, NH2 They will completely dissociate in water All other bases are "weak" and will be mostly organic compounds with an N atom or O containing anions o H2O is amphiprotic meaning that it can donate or accept H+ (it can be an acid or a base) o kw is the equilibrium of water with itself It is often called the autoionization constant with water At 25 C kw = 1.0 x 10 -14 This equilibrium will be true for all aqueous solutions + o [H3O ] > [OH-], then the solution is acidic o [H3O+] < [OH-], then the solution is basic o Equations: pH = -log[H3O+] therefore [H3O+] = 10-pH pOH = -log[OH-] therefore [OH-] = 10-pOH kw = [H3O+][OH-] 14 = pH + pOH o pH is a logarithmic scale [H3O+] = 0.1 then pH = 1 [H3O+] = 0.01 then pH = 2 [H3O+] = 0.001 then pH = 3 o See figure 14 Weak Acids o HA + H2O A-(aq) + H3O+(aq) o ka = [A-][H3O+]/[HA] ka represents an acid reaction with H2O o log(ka) = pka As ka increases the stronger the acid is As ka increases, pka decreases As pka decreases, the stronger the acid is Weak Bases o B(aq) + H2O BH+(aq) + OH-(aq) o kb = [BH+][OH-]/[B] kb = equilibrium constant for base reacting with water o log(kb) = pkb As kb increases, the base strength increases As pkb decreases the base strength increases Few things about kb and ka o If ka increases, the acid is stronger If kb decreases, for CB, weaker base
o If ka decreases, the acid is weaker If kb increases, CB is stronger Conjugate bases of strong acids have basically no base strength and vice versa o Largest ka value = strongest acid o Strongest CB = smallest ka value Predicting acidity/basicity of salts o NaNO3 Na+ No effect on pH All alkali metals have no effect NO3 HNO3 is a strong acid and therefore NO3- has no base strength Therefore NaNO3 is neutral o KNO2 K+ No effect NO2 HNO2 is a weak acid Therefore KNO2 is a weak base Predicting the Direction of an Acid/Base reaction o Strongest acid reacts with strongest base Polyprotic Acids: acids with more than one acidic hydrogen o Release 1 H+ at a time o pH usually defined only by the 1st equilibrium (if pure) Buffers o Buffer is a solution that resists changes in pH (when acid/base is added) o A buffer must be able to react with H3O+ and OHo Must contain an acid and base o Acid and base must not react with each other Conjugate acid/base pair See Figure 15 o How do you make a buffer? Add a weak acid and conjugate base to the same solution Ass a weak acid and less strong base Add a weak base and less strong acid o How do you find the pH of a buffer? Do equilibrium problem (will ALWAYS work) Henderson-Hasselbach Equation (H-H equation) pH = pka + log([A-]/[HA]) o gives pH of a Buffer ONLY pH remains constant even when volume changes Assumption: assume that there is a significant concentration of each the weak acid and weak base pOH = pkb + log([HB+]/[B]) not used often due to not caring about pOH Titrations- acid/base
Quantitative reaction between an acid and base Neutralization reaction Slowly adding 1 reagent to the other Go to http://www.shodor.org/UNChem/basic/ab/#titra for more information on acid/base chemistry and extra problems and how to do them Salt Solubility o What happens when you put salt in water? It dissolves Ex. NaCl Na+(aq) + Cl-(aq) ksp = solubility product ksp = [Na+][Cl-] small ksp values = not very soluble (favor reactants) high ksp values = very soluble Insoluble- doesn't dissolve in water, much o See Figure 16 Information for Final What you need to know about electrochemistry o How to balance a redox equation o How to determine voltage (Eo) of an electrochemical cell o How to draw an electrochemical cell o Batteries Redo Chemistry- requires the transfer of electrons from one atom to another o Identifiable by change in oxidation state o Oxidation numbers change One will gain an electron One will lose an electron o Look at hand-out "Balancing Redox Reactions" for steps In principle, we can separate oxidation and reduction processes o Electrons move through a wire (electricity) o Need to separate two different half reactions Voltaic cell Named after Count Alessandra Volta (1745-1827) See Figure 17 Batteries o AA battery (alkaline- basic) o See figure 18 for picture of battery Cathode- reduction Anode- oxidation The battery dies because you run out of MnO2 Rechargeable batteries Lead storage battery (car battery) o 12-V battery o Alternate anode/cathode Pb(s)- Anode Oxidation Pb(s) + SO42- PbSO4 + 2eo o o o
PbO2- Cathode Reduction PbO2(s) + 4H+ + 2e- + SO42- PbSO4(s) + 2H2O(l) 6 units (cathode/anode, cathode/anode, cathode/anode...) and 2V each = 12-V total In order to make reaction go backward, you need to apply electrical voltage greater than the voltage of the cell (Eo) Voltage Voltage shows difference in potential energy o Always measure potential difference o Remember: electrons want to go down (from high to low energy levels) o Can't find absolute potentials so we define a "zero point" H2(g) + 2H2O 2H3O+ + 2eE = 0.00V o Voltage Eo = Eox + Ered Eox = -Ered Standard Reduction Potential Table (p.967) As voltage increases, reduction is more likely As voltage decreases, oxidation is more likely o If E <0 reaction is NOT spontaneous If Eo>0 reaction IS spontaneous Electrochemistry o Thermodynamics H- enthalpy (heat) When H<0 the reaction is favorable (giving off E) Endothermic- gets cold Exothermic- gets hot st 1 Law of Thermodynamics E=0 Energy is neither created nor destroyed nd 2 Law of Thermodynamics In a spontaneous process, the entropy of the universe increases Spontaneous- process that proceeds on its own (does not necessarily mean fast) Entropy- molecular order or randomness (denoted as S) o Favor more random ...