Chem midterm 2

the enthalpy of formation

***** The standard enthalpy of formation, ΔHf°, is the enthalpy change for the reaction forming 1 mole of a pure compound from its constituent elements - the elements must be in their standard states - the ΔHf° for a pure element in its standard state = 0 kJ/mol • by definition Standard enthalpies of formation— the numbers come from tables— look at table to find which has delta of 0 gives you more systematic approach The standard enthalpy of formation, Delta Hf degrees , is defined as the enthalpy change for the reaction that produces 1 mole of the compound or 1M aqueous concentration from its element in their standard states at 25 celcius. O2(g) + C(s)èCO2(g) ΔHfo = -393.51 kJ/mol ***Because absolute enthalpies cant be measured in the laboratory, we arbitrarily define a reference state for which delta Hfdegrees=0 ******For elements we use the elements in their most stable form at the standard state of 1 atm ***For elements found as dimers (H2, O2,N2, F2,Cl2, Br2,I2)- we use the dimer form as the zero reference state ******ΔHfo of carbon as graphite is set to zero ****** The ΔHfo of H+(aq) is set to zero so that the heats of formation of other aqueous ions can be determined relative to H+.

spontanaety and it related to physical and chemical changes

A spontaneous change is one that can occur by itself without outside intervention (if someone smoking smoke will eventually come in room- now talk about reverse, if room is full os smoke, will it spontaneously move into another room— no need to hook up fans to push it out of room) once conditions have been met for its initiation. It has a direction. It need not be fast. Some physical changes are spontaneous • For example - tobacco smoke quickly diffuses into a room Some chemical changes are spontaneous For example - the reaction between baking soda and vinegar Conversion of diamond into graphite is spontaneous process! Slow rate Although spontaneous from a thermodynamic view it happens at such slow rate that they don't really ever turn into graphite but thermodynamically they will go in that direction Do not expect graphite into diamond- other way around to be spontaneous Some chemical changes that occur spontaneously appear associated with lower energy products

gas and macroscopic view description and diffusion vs effusion and 4 factors that affect a gas

A thin layer of gas is held by gravity to the earths surface. Half of its mass lies w in 5.5 km above our heads (3.4 miles) Earth as a bball- atmosphere is only 1mm thick- this layer shields us from harmful radiation and supplies us w oxygen nitrogen CO2 and H2O Macroscopic view Low density Compressible- can squeeze, pushes back Fills container Takes shape of container Solids and liquids move through a gas easily (lots of space between molecules) Mixtures of gases are always solutions (homogenous) Diffusion is faster at higher temperatures because the gas molecules have greater kinetic energy. Effusionrefers to the movement of gas particles through a small hole. Graham's Law states that the effusion rate of a gas is inversely proportional to the square root of the mass of its particles. 4 factors that affect gas Amount of substance - symbol (n) units (moles) Temperature- symbol (T) units (K) Pressure- symbol (P) units (atm,Pa) Volume- symbol (V) units- (L,ml)

absolute entropy and third law of thermo

At absolute zero (t= O k) molecules are not in motion they have no kinetic energy. they are fixed in 1 position so there is only 1 micro state available the entropy of any pure substance in its equilibrium state approaches 0 at 0K This means it is meaningful to talk about absolute entropies S as they are reference to absolute zero This is our reference point- no order which occurs at 0

measuring internal energy change diff cases

Case 1a: Constant volume (isochoric), no reaction !!! A gas is heated from 20 degrees Celsius to 50 degrees Celsius in a constant volume container such as a glass bulb No work is done since Change in V= 0 -- work is pressure times volume-- so why we know no work being done Change in U= qv=ncvChange in T (depending on units of Cv) For a monatomic ideal gas so bc no work done the delta U eq just goes to U= q (bc w =0) -- and q= mc delta T and then the delta U = KE bc no potential energy of single molecules and if molecule is just (monatonic- just has translational=3/2RT) SO - ΔU= 3/2 nR(T2-T1) = ncV ΔT Case 1b: Constant volume (isochoric), with reaction-- now going to have system and surroundings!! ΔUsystem = qV = ncV ΔT or mcV ΔT ΔUsystem = -ΔUsurroundings qV(system) = -qV(surroundings) (since here again, ΔV = 0) Case 2a: Constant pressure- no reaction Case 2a: Constant pressure- no reaction A gas is heated from 20-50 Celsius under constant pressure conditions such as in a cylinder and piston arrangement Change in U= qp+ w= n Cp Change in T - P ext Change in V If make delta V= zero then equation becomes simple- will get more hear out of constant volume reaction than out of constant pressure reaction!! If Cp is known, both terms can be determined For an ideal gas: Cp= cv + R (gas constant R) For a monatomic ideal gas: Cv= 3/2 R and Cp= 5/2 R Case 2b: Constant pressure- reaction in an open system (coffee cup calorimeter or beaker) A reaction takes place in a beaker at 1 atm room temp Change in U= qp + w= nap delta T- P ext delta V ΔV cannot be determined so ΔU cannot be determined. Since most chemistry is carried out under case 2b conditions, we need to find a way to determine the energetics of these reactions.

systems can be closed or open

Closed systems Boundaries prevent the flow of matter in and out, but energy can flow in and out Can be described by properties such as temp, pressure, volume, number of moles, which may be fixed or may vary Open systems Boundaries permit the flow of matter as well as energy in and out Can be described by properties such as temp, pressure, volume, number of moles, which may be fixed or may vary Surroundings provide external forces that can affect the properties of a system The system + surroundings form a thermodynamic universe System and surroundings must be isolated to form the thermo universe- if not isolated then surroundings are bigger than we think In a closed or open system, energy can be exchanged with the surrroundings An isolated system does not exchange matter or energy with the surroundings A system with rigid walls can not have mechanical work such as compression done on it by the sourronunidngs and cannot do the mechanical work of expansion on the surroundings A system w adiabatic walls Dows not exchange hear w surroundings

The small bags of silica gel you often see in a new shoe box are placed there to control humidity. Despite its name, silica gel is a solid. It is a chemically inert, highly porous, amorphous form of SiO2. Water vapor readily adsorbs onto the surface of silica gel, so it acts as a desiccant. Despite not knowing mechanistic details of the adsorption of water onto silica gel, from the information provided, you should be able to make an educated guess about the thermodynamic characteristics of the process. Predict the signs of ΔG, ΔH, and ΔS

Delta G less than 0 Delta H less than 0 Delta S less than 0

kinetic molecular theory

Describes what happens at the molecular level Rudolf Clausius, James Maxwell, and Ludwig Boltsmann (19 cent) developed a model based on basic laws of physics to explain the measured properties of gases They were able to relate macrospcopic properties like temp and pressure to molecular motion Assumptions A pure gas consists of a large number of identical particles separated by distances that are large compared to the size of the particles (the particles are small spheres w zero volume) The gas particles are constantly moving in random directions w some distribution of speeds The molecules exert no forces on one another bw collisions- between collisions they go in straight lines w constant velocities The collisions of molecules with the walls of the container are elastic - no energy is lost during a collision. Observed pressure is a result of molecules bouncing off walls Net result of molecules bouncing off walls when observe pressure Will calc what happens as a result of these collisions--- how you get pressure-- thats what it is the gas molecules bouncing off What happens to pressure as volume is reduced? If part traveling at same speed, more collisions w wall Larger volume- reduced pressure so inverse relationship Instead of 1 particle imagine there are countless particles The molecules have distribution of speeds PARTICLES HAVE AVG SPEED BUT SOME MOVE FASTER THAN OTHERS HIGHER temps result in higher speeds If you know the speed of a molecule at 300K you can calc the speed at 600K using a similar simplification of the form Temp should be recorded in kelvin When mix hot air w cold air eventually v of hot air will transfer energy to slower and then will become uniform middle ground v or temp

enthalpy of vaporization and enthalpy of fusion and what is true because state property

Difference in molar enthalpy between the vapor and liquid states is called the Enthalpy of vaporization Delta H vap= Hm (vapor) - Hm (liquid) For water at its boiling point, 100°C, ΔHvap = 40.7 kJ·mol-1, and at 25°C the value is ΔHvap = 44.0 kJ·mol-1. i.e. vaporization of 1.00 mol H2O(l) (18.02 g of water) at 25°C and constant pressure, requires 44.0 kJ of energy as heat.-- it is the heat (energy needed to take it from one state to the next) Water doesn't like to become vapor bc h-bonds- has higher boiling point and takes a lot of energy to break the water Freezing and boiling points correlate w vap points ** Enthalpy of condensation Opposite of vaporization Condensation is the change from a gas to liquid So this process is the reverse delta H vap-- ice sublimes directly to water vapor is a good example The difference in molar enthalpy between the liquid and solid states is called the enthalpy of fusion (melting) Delta H fus= Hm (liquid) - Hm (solid) The enthalpy of fusion of water at 0 Celsius is 6 kj mole-1: to melt 1 mol H2O (s) (18g ice) at 0 degrees we have to supply 6 kj of heat Vaporizing water takes much more energy (more than 40 kj) because to vaporize a gas, its molecules are separated completely, and KE increases dramatically. In melting, the molecules stay close together, and so the forces of attraction and repulsion are nearly as strong as in the solid. **makes sense that H vap will be higher than H fus bc breaking those hydrogen bonds completley for vapor-melting molecules still stay close ** Freezing is the change from liquid to solid Enthalpy is a state property Delta H reverse process = - delta H forward process Since the enthalpy of fusion of water at 0.0°C is 6.0 kJ·mole-1 the enthalpy of freezing for water is at 0.0°C is - 6.0 kJ·mole-1

enthalpy LOOK AT THIS

Enthalpy, H, of a system is the sum of the internal energy of the system and the product of pressure and volume H is a state function enthaply is used to find energetics under constant P!!! (why Cp is used) The enthalpy change , delta H, of a reaction is the heat evolved in a reaction at constant pressure Delta H reaction= q reaction at constant pressure Usually dealt H and delta E are similar in value, difference is the largest for reactions that produce or use large quantities of gas -- bc V would change

sign conventions and heat flow

Discharging battery is neg work Q is positive when heat is transferred from surroundings to the system (Usys increases) Q is neg when heat is transferred from system to the surroundings (Usys decreases) w is positive when work is transferred from surroundings to the system (u sys increases) W is negative when work is transferred from the system to the surroundings (u sys decreases) Sign convention replaces subscripts sys and surf Heat flow is a means by which energy is transferred from a hotter object to a colder object Heat flows from areas of higher temp to areas of lower temp- The energy can be in the form of molecular or atomic motion, in the form of chemical bonding energy, or in intermolecular interactions. When two objects at different temperatures are brought into contact with each other, heat is exchanged until they reach a common temperature.

Energy changes for reactions are determined using the law of cons of energy

Energy required to break bonds and released when make If more energy is released in new bonds in products than is required extra energy is used to heat surroundings or do work on surroundings If it takes more energy to break the bonds in the reactants than is released in the new bonds in the products, the extra energy comes from the surroundings. Ex from combustion: Correct statement about combustion System goes to lower internal energy while supplying energy to surroundings - products more stable than reactants

exothermic only way for spontaneous?? HERE IS BIG CONC

Exothermic changes have a tendency to proceed spontaneously BUT thats not the entire picture Some changes are spontaneous that require no energy (they are thermoneutral!!) Thermoneutral is where energy of reactants is equal to energy of products Some reactions require a lot of energy from the surroundings yet they still occur spontaneously the reaction of NH4SCN with Ba(OH)28H2O can easily freeze water ****HERE IS THE BIG THING- WE KNOW ENTHALPY AND WE KNOW EXOTHERMIC- GIVES OFF HEAT TO SURROUNDINGS AND SOMETIMES THOSE REACTIONS ARE SPONTANEOUS-- BUT BOTH ENDO AND EXO CAN BE SPONT SO WE NEED TO KNOW ENTROPY -- AND THEN RESULTING GIBBS- THINK EQ- TO TRULY KNOW SPONTANAEITY !!!

what is pressure? and barometer and manometer?

Force per area P= F/A= ma/A=mg/A V=Ah P=mgh/V P (funky p)=m/v P= (funkyp-density)gh Atmospheric pressure is 14.7 lbs per square inch P incomplete infinity sign h A barometer measures atmospheric pressure Mercury column falls and then measurement of pressure is made by how many mercury columns that is- if reduce pressure more mercury would fall further Manometer Measures the difference between the pressure of a gas and the atmospheric pressure Depending on diff in levels can make a linear measurement in terms of ml of mercury to establish pressure inside-

formation reactions

Formation Reactions Reactions of elements in their standard state to form 1 mole of a pure compound If you are not sure what the standard state of an element is, find the form in appendix that has a deltaHdegreesf=0 Because the definition requires 1 mol of compound be made, the coefficients of the reactants may be fractions Writing formation reactions Write the formation reaction for CO (g) The formation reaction is the reaction between the elements in the compound, which are C and O By definition can only have 1 mole of product so you can put fractions in reactants where you need- not like normal balancing bc enthalpy is a state function so you can do this *** doesnt matter how you got there- just matters is end amnt *****C + O—- CO (g) The elements must be in their standard state There are several forms of solid C but the one w the delta Hfdegrees f = 0 is graphite oxygen's standard state is the diatomic gas C(s, graphite) + O2(g) → CO(g) • The equation must be balanced, but the coefficient of the product compound must be 1 - use whatever coefficient in front of the reactants is necessary to make the atoms on both sides equal without changing the product coefficient C(s, graphite) + 1⁄2 O2(g) → CO(g) Calculating standard enthalpy change for a reaction- easier way Any reaction can be written as the sum of the formation reactions (or the reverse of formation reactions) for the reactants and products The delta ΔH° for the reaction is then the sum of the ΔHf° for the component reactions ΔH°reaction = Σ n ΔHf°(products) − Σ n ΔHf°(reactants) - Σ means sum *****If given set of diff chem reactions and asked Which of the following reactions are formation reactions at 25oC? You could determine which ones are the formation reactions are those that yield 1 mole (or 1M) of products. !!!!

types of work -- changing pressure vs not

Free expansive If the external pressure is zero ( a vaccum) then w= 0 Non expansion work is done in a vacuum bc there i son opposing force- nothing to push against (no work done) Expansion against zero pressure is called free expanision Measuring work at constant ext pressure The work done w is calc as : w= -Pex change in V Neg bc system doing work on surroundings Work expended to push other molecules back **if no pressure pushing it back- the p is constant -- work will be zero bc eq for work is dependent on pressure here so if no pressure pushing back on it then 0 * delta V= 0 for work measuring work w change in pressure simplest reversible change is isothermal- constant T exp (so now P will change)

grahams law and velocity distribution tool

Graham's law of effusion (also called Graham's law of diffusion) was formulated by Scottish physical chemist Thomas Graham in 1848. Graham found experimentally that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles. The Velocity distribution can be measured in a time of flight instrument By rotating the wheels faster, you measure faster molecules Wheels w slit cut into it and at every advancing position it will rotate by set number of gdegrees- to do exp you spin it and only molecules that go through at certain velocity will go through slits Can perform many times at diff angular velocities - then measure how many molecules arrive at detecter at given velocity and that can measure you distribution of given velocities

first law of thermodynamics and defining system and surroundings

Hard bc need to define system and surroundings- that is hard part Change in internal energy= q + w U= internal energy of the molecules (pot and kin) q= heat (thermal energy) w= work (mechanical and electrical) For gases mostly concerned w PV work The kinetic energy that plays a role in U involves all three kinds of KE Systems and surroundings System: just the reactants and products Surroundings: everything else: container, air, room, you the universe Delta U sys= qsys+wsys Delta U surr= q surr+ w surr Dellta U for whole universe must be zero so delta U sys= - delta U surr— has to be true Q sys= -q surr (this is not true if some of the heat does work ) W sys= -wsurr (not true if some of the work is converted to heat through friction) Be careful w heat and work- what happens w delta U is simple but then work and heat are not Change in U universe= change U sys + Change U surr= 0 !! This is called law of conservation of energy- what first law really is!!-- energy that goes must go somewhere- that is first law Heat and work are different for reversible and irreversible pathways change in U (forward)= - change in U (reverse) for any system that returns to its original state by any pathway

heat measurments

Heat (q) is measured by measuring the change in temp and specific heat capacity of the substance The specific teat capacity of a substance is defined as the heat required to raise temp of 1 gram of a substance by 1 degree Celsius q=mcs * delta T Where m is the mass in grams, Cs is the specific heat capacity and change in T is Tf- Ti (dont need to convert to kelvin) Cs= JK-1g-1 Molar heat capacity, Cp,Cv Expressed relative to moles not grams now JK-1mol-1 Diff is whether they are determined under constant pressure or constant volume conditions System heat capacity , Cp or Cv (JK-1)

what heat capacity means and Which types of energy are negligible compared to the total energy?

How many kj of heat is needed to increase the temp of something by 1 degree Celsius Translational and rotational because their values are so small

the origin of the heat capacities

Ideal gas, no potential energy which means we know 3/2 RT for translational part of KE but defining heat capacity w both translational and rotational so values below differ Heat capacities are helpful to find the internal energy of gases at diff temps; C p,m= C v,m +R

how are heating curves produced?

In differential scanning calorimetry (DSC) Equal masses of a sample and a reference material that will not undergo any phase changes, such as Al2O3 (which melts at a very high temp) Two separate large steel blocks act as a heat sink Electronics want to keep temp of 2 things absolutely same Separate electrical signals are generated to keep the temps exactly the same at all times Thus the output shows measures changes in heat capacity as well as total heat delta H for any phase transition IF sorting molecules into their most stable state delta H degree is equal to zero!!! So for diatomic this will be their most stable state, also H+ for H is its most stable state Paste in curve- area under curve is related to heat of fusion- fact that have gone from one level to another is bc heat capacity of water in solid and liquid state has changed

in gas law calcs- what are the three important things to remember

In gas law calculations 1. ALWAYS use K (not °C ) 2. ALWAYS use same units of pressure on both sides of the equation. 3. ALWAYS use same units of volume on both sides of the equation. Standard Temperature and Pressure (STP) 273.15 K ( =0oC) and 1 atm = 760 torr

pressure and volume changes affecting entropy and enthalpy

Incorrect. Pressure and volume changes affect entropy. As the piston is expanded, the total volume of the cylinder increases, causing an increase in the entropy of the system because the gas molecules are able to move more freely.-- think more space gives it a less defined shape which means more entropy if increase volume for enthalpy-- think of particles in a smaller volume- they are moving faster so more interactions- more energy- so if increase volume enthalpy goes down **also side note- iron rusts to iron oxide-- now increasing amnt of stuff so increasing entropy- can just be asmnont of stuff- compound vs single solid element

the standard state

Is the thermodynamically stable state at P=1 atm and a specified temp (usually but not always 25 celcius) For dissolved species, the standard state is 1M at a pressure of 1 atm For gases, the standard state is 1 atm For the the standard state, enthalpy is written as delta H degrees— tells you delta H corresponds to reaction where reactants and products are in their standard states STP from ideal gas chapter is not the thermodynamic standard state due to temperature Standard conditions The standard state is the state of a material at a defined set of conditions Pure gas at exactly 1 atm pressure Pure solid or liquid in its most stable form at exactly 1 atm pressure and temperature of interest Usually 25 degrees celcius substance in a solution with concentration 1 M ******* The standard enthalpy change, ΔH°, is the enthalpy change when all reactants and products are in their standard states

Surroundings of a system may constrain some properties at fixed values, independent of time. This constrains some of the state functions. These constraints have names

Isothermal At constant temp (ex the reaction takes place in an ice bath) Heat transferred bw ice bath and system Isochoric At constant volume- as way to prevent any expansion or compression work going on (ex- takes place in a container w rigid walls such as a glass bulb or steel box) Isobaric At constant pressure (ex the reaction takes place in a test tube or beaker open to atm) adiabatic Insulated system not allowing any heat to be transferred from system outside With no heat exchange with the surroundings (ex: the reaction takes place in an insulated container)

kinetic energy and potential energy

Kinetic energy *** all gases have same KE at given temp Like to talk about change in energy bc want to know where internal energy went Related to molecular velocites- emerge stored in motion of molecules so related to temp Energy of a moving object Air molecules moving through the room have translational kinetic energy Translational E = KE= 3/2 RT for one mol of gas ----For kinetic energy, the higher velocity of lighter molecules is offset by their lower molar mass, so they have the same kinetic energy if they are at the same temperature. KE = ½ mv Molecules have additional kinetic energy due to - other 2 besides translational Vibrations - can increase bond vibration bw molecules Rotations - rotations of molecules , rotation of molecule itself spinning other 2 types!! potential Potential energy Energy that a particle has bc it is attracted to or repelled by another particle Chem attractions (bonds and intermolecular forces) lowers the potential energy of molecules

bond enthalpy for chemical bonds -- LOOK BACK AT AVG PART

Many different molecules contain the same bonds. - For example, C-H bonds exist in CH4, C2H6, CHCl3 ... and the bond enthalpy is always about the same. We can determine "average" values for bonds. - These are the average value compared to atomization (all atoms in their gaseous, atomic state). • E.g. C(g), H(g), ... NOT their elemental form C(s), H2(g) These average values can be used to obtain approximate ΔHfo values for many molecules. ** We can determine "average" values for bonds. - These are the average value compared to atomization (all atoms in their gaseous, atomic state). • E.g. C(g), H(g), ... NOT their elemental form C(s), H2(g) These average values can be used to obtain approximate ΔHfo values for many molecules. First - write the balanced reaction for the formation of 1 mole of product from elements in their standard states. Then turn all the reactant elements to gas phase atoms. • It will require energy to do this, so all values are positive. Then determine the amount of energy released by making all the product bonds, using average bond enthalpies for all the new bonds. • Since energy is released in making bonds, these values should all be negative.

Maxwell Boltzmann distribution of molecular speeds Grahams law- Diffusion

Maxwell Boltzmann distribution of molecular speeds Lighter molecules have a greater range of speeds, and also higher average speeds Heavier molecules have speeds closer to their average speeds Very light molecules have speeds high enough to escape earths gravitational pull Molecules that travel faster have wider distribution of velocities Helium and hydrogen- most abundant element- so light can get to speeds that allow them to escape space— why don't see them a lot by themselves even if so abundant As temp increases, avg v of molecules increases What happened to number of collusions w walls as the temp increases? At higher temperatures the average speed increases.Also the spread of speeds increases. This trend is also seen in the kinetic energy of gases at higher temps. Grahams law- Diffusion The tendency of molecules to move from areas of high concentration toward areas of lower concentration until the concentration is uniform throughout, possibly through a porous boundary. Do small or large molecules diffuse faster? At same temp- Bothe molecules have same KE Larger mass means smaller V: smaller mass means larger V Does KE or V matter for diffusion? V— relative velocities is what matters **this makes sense bc KE for gas is 3/2 Kb T-- only dependent on T not mass of molecule so V is what matters -- and we know smaller mass- larger V-- the KE isnt changing based on size of molecule Smaller molecule will diffuse faster than larger- will reach equilibrium more rapidly than large molecules Treating gases s homo mixtures so uniform distribution- process of getting there is diffusion

on table Molar enthalpy of atomization vs bond enthalpy values

Molar enthalpy of atomization- how much energy it takes to put the elements into the gaseous state- estimated from gas phase reactions where remove intermolecular interaction— values on table under this is the energy needed for the diff elements — then also shows energy on that table needed to make bonds w other things (bond enthalpy) For many common bonds can look up what avg bond energy is If bonds didn't require energy to break- then they wouldn't be stable- wouldn't stay together in the first place

Kinetic energy different types

Molecules (not atoms) have 3 types of kinetic energy Potential energy assumed to be zero Translational KE: moving through 3 dimensions for all molecules Rotational KE: 2 rotating motions for linear or; *Vibrational KE: 3 rotating motions for nonlinear; *Vibrational KE: most molecules at room temp do not vibrate substantially so this is very small Rotational degree of freedom for single atom- none- want store internal energy in single atom as rotational kinetic energy The equipartition theorem (not derived here) states the average value of each quadratic contribution at a temp T is equal to 1/2 kT, so for translational KE is 3* 1/2 kT (all the same) we know that translational energy is moving through 3 dimensions- so we know the c value of the trans is 3 so why the eq is 3/2-- will be diff for other types Equipartition Means energy is shared (partitioned) equally over all modes (trans, rot, and vibrational) Says that translational KE is 3/2 kT Where k is boltzmanns constant 1.381 *10^-23 Since bolzmanns constant K is related to gas via: R= Nak We now have molar internal energy U= 3/2NakT Because RT= 2.48 kjmol-1 at 25 degrees C we find translational KE= 3/2RT For linear molecules (a&b) rotation on two axes, is 2 x 1⁄2 RT or 2.48 kJ·mol-1 plus translational 3 x 1⁄2 RT which equals 5/2 RT or 6.02 kJ·mol-1 For nonlinear molecules (a&c) the total translational (3 x 1⁄2 RT) plus rotational (3 x 1⁄2 RT) is 3RT which equals 7.44 kJ·mol-1. Conclusion: Internal energy is stored as molecular kinetic and potential energy. The kinetic contribution is estimated at 7.44 kJ·mol-1 at room temperature for nonlinear molecules and 6.02 kJ·mol-1 for linear molecules. **calculations above helping us to determine kinetic contribution to total energy

compressibility factor

PV/nRt not equal to ! We call the deviation from 1 the compressibility factor, z PV/nRT= z Attractive forces cause z<1 Repulsive forces cause z>1 The compressibility factor depends on P and T 1. The ideal gas is horizontal (z=1) 2. Deviations increase w increasing pressure (makes sense bc more molecules hitting more inelastic collis) 3.Deviations are greater at lowest temperatures (makes sense bc more mol sticking togeth) Effect of high pressure Has affect on how ideal gas is - bc increasing number density of molecules so greater interactions bw molecules and at some point will be so dense will have to apply very big volume correction bc more of space is taken up by molecules Interactions become more important bc they are inversely proportional to distance bw molecules The excluded volume is a greater percentage of the total volume Effect of low temperature The molecules are moving more slowly so attractive forces become relatively more important compared to kinetic energy. Since n/V = P/RT, the number density (n/V) becomes larger, which affects both the interactions and excluded volume in the same way as increasing pressure. (LOW TEMP AND HIGH PRESH SAME EFFECTS) The compressibility factor also depends on the identity of the gas Molecules that like to associate w one another will be the most non ideal- ex: water vapor bc extremely polar and likes to associate w each other Deviations are greater for molecules with the strongest intermolecular interactions Attractive forces Lower the z Deviations from ideal gas behavior Extremely dependent on prop of that gas Compressibility factor greater than 1- finite volume so as go to high pressure- running out of space so will have repulsive interactions influencing behavior

isothermal

Phase transitions (melting, boiling etc) are isothermal and at constant pressure They can be done reversibly by making infinitesimal changes in consditions They can be done reversibly by making infinitesimal changes in consditions DElta S vap= qrev/T= Delta H vap/ Tboiling Delta S fus= Delta H fus/ T melting Since delta H fus and delta H vap are both positive, Delta S fun and Delta S cap are also both positive

Pressure and Volume effects (2 modifications to ideal gas law)

Pressure Attractive forces act at small distances so molecules must be close together Preal<P ideal (so will be adding a quantity to make up for that) PrealV/nRT < PidealV/nRT= 1 Volume of real gas is larger than V of ideal gas The particles have a finite size, so only some of the volume is available Real gas have to account for extra molecules on top of volume- don't have to worry until volume is very crowded w molecules

volume and work

Pressure- volume work arises when a system is compressed or expanded (equations in book) For Vf > Vi, w is negative. Work is done by the system when the system expands. For Vf < Vi, w is positive. Work must be done on the system to compress it w=0 for isochoric processes(ΔV=0). heating or cooling contents of a sealed glass container would be isochoric because the delta V is zero- v is constant

state function

State function Include the properties of a system that are uniquely determined by the thermodynamic state of the system.they include volume, pressure, temp and total internal energy. They don't depend on pathway or history. Doesn't matter how you got from lib to here- location is the state function- we are here ****** For example, internal energy, enthalpy, and entropy are state quantities because they describe quantitatively an equilibrium state of a thermodynamic system, irrespective of how the system arrived in that state. Internal energy is part of enthalpy!! A thermodynamic process leads to a change in the thermodynamic state This can be a physical process such as changing the pressure or temp of a gas or dropping a hot piece of metal in a cup Can be a chemical process such as a chem reaction. Can be called a thermochemical process Changes in thermodynamic states can be reversible or irreversible Reversible processes proceed through a continuous series of thermodynamic states, each making infinitesimal changes in the thermodynamic state. A state of equilibrium is maintained through these infinitesimal changes Small amount of change of force to pull weight up pulley-- can pull little part of pulley then let go of little part and do tjat again Irreversible processes If let go of rope and let is go to ground Proceed through non equilibrium states. Aunt of work expended in both process is diff- less work to do this Abrupt large changes are made and the system is allowed to settle back into an equilibrium state-- LESS WORK FOR IRREVERSI

Enthalpy of reaction

The enthalpy change in a chemical reaction is an extensive property The more reactants you use, the larger the enthalpy change By convention, we calculate the enthalpy change for the number of moles of reactants in the reaction as written (LOOK IN BOOK) this is why we include mult the H values by number of moles it depends on how much we have bc its extensive!! When we include the delta H for a reaction- it is called a thermochemical eq— you have a balanced eq w states- then to make it a thermochemical eq you specify change in enthalpy when going from reactants to products If neg enthalpy then exothermic!!! Enthalpy change corresponds to number of moles consumed as reactants If piece of metal at temp and water at temp and then have a final temp - what is tyrue? Heat lost by metal= heat gained by water

entropy and enthalpy signs for diff phase changes liquid to gas solid to liquid gas to liquid solid to gas liquid to solid gas to solid

The entropy of a material increases with an increase in the number of possible microstates (arrangements of particles). The number of microstates increases with the mobility of the particles. The change in enthalpy of a phase change depends on whether heat is added or removed. Consider which phase changes require the addition of heat. liquid to gas: H pos S pos solid to liquid: H pos S pos gas to liquid: H neg S neg solid to gas: H pos S pos liquid to solid: H neg S neg gas to solid: H neg S neg notice has to be one sign or other none have one of each!!

microscopic vs macroscopic relating to entropy and how does phase factor in ALSO BIG

The microscopic observation is that the more energetically equivalent versions there are of the system, the larger is its statistical probability If you have 2 systems in contact, one that is hotter than other, there will be more micro states available if they could both be at same temp More possible configurations you can find system in! 2.Macroscopic observation: Heat is never transferred from a cold object to a hot body without doing work. How does phase factor in? There are more positional micro states available to gases comared to liquids There are more positional micro states available to liquids compared to solids The strength of intermolecular forces is greatest for solids, intermediate for liquids, and lowest for gases.These intermolecular forces constrain the positions of molecules.!!! To have heat spread out is state with highest entropy -- want the heat to spread like want the molecules to spread

second law of thermodynamics

The natural progression of a system and its surroundings is from order to disorder, from organized to random, from lower to higher entropy. Provided temp is constant , change in entropy is given as Delta S- q rev (reversal heat supply)/ T (J/K) T is the absolute temp at which the change takes place Change in entropy is equal to reversible heat supply divided by temp where temp must be a constant If add too much heat temp will change and this equation will fail Ex problem, Large flask of water is placed on a heater and 100 J of energy is transferred reversibly to the water art 25 Celsius what is change in entropy of water? Water is absorbing heat so entropy will be expected to increase Delta s= 100/298= +.336 J/K Adding hear will normally change the temp so need calculus— look in book for adiabatic, isochoric, isobaric **why we have the special cases and diff equations for when things are constant for S bc adding heat would usually change temp so calculus takes into account for that

limits to kinetic molec theory for real gases

There are limitations to the kinetic molecular theory for REAL gases!! - bc assuming atm pressure Molecules do interact w each other There are intermolecular forces bw molecules - requires us to modify our equations The volume of gas molecules if finite- they do occupy finite volume This is more likely to be a prob at higher densities when molecules are close together Deviations occur at very high pressures and temps -- makes sense bc will be more collisions and closer together Can rearrange to n/V= P/RT Refer to how many molecules cramming into box of given volume — higher densities- molecules closer to next-door neighvors- forces bw them become more important **** IF the collision bw molecules and the walls container were not elastic and energy was lost to the walls- would pressure of gas be higher or lower than predicted w gas law? Lower If you are gonna lose energy through collision w wall- will actually be cooling down gas itself bc some of that energy will be transferred to wall- slowing molecules- dropping pressure and so dropping pressure as well ********* 2 ways to think about this Energy loss argument- potential energy- energy can go other places so have to presume if inelastic collision we will lose energy At any given point if you have attractive forces bw molecules- wont stick together forever but will stick for a little bc attractive force-small percent when molecules associated and then reducedes net molecules available in box to collide w walls- so decrease in pressure

Troutons rule

Trouton's rule states that the entropy of vaporization is almost the same value, about 85-88 J K−1 mol−1, for various kinds of liquids at their boiling points. Delta S vap= Delta H vap/T= 88+-5 JK-1mopl-1 Empirical result True for nearly all liquids expect molecules with exceptionally strong intermolecular bonds For H20, Delta S vap= 109

What is temp? what is KE of a gas? and how does a thermometer work?

What is temperature? Temp is a measure of the random motion of the atoms (kinetic energy) KE gas (avg)= 3/2 KbT Kb is boltzmanns constant Kinetic energy is proportional to temp so when go outside and hot outside- reason why its hot is bc molecules in air gave more energy and then giving some of that energy to us so we feel hot How does a thermometer work? When liquid mercury (or alcohol) heats up, it expands (desnity decreases) and fills more of the tiny tuve in the center of the thermometer Need to calibrate so amount it expands corresponds w temp scale using

entropy definition and relation to probability

What is the probability that the molecules are in any in any specific configuration? The probability that any one molecule is in the left side is 1⁄2. The probability that any two are in the left side is 1⁄2 x 1⁄2 For 4 atoms in the left side, the probability is (1⁄2)4 = 1/16 For each additional molecule, it goes down by another 1⁄2. 1 NA For a mole of gases, the probability is ( ) where NA is 6.022 x1023. HIGHER PROB VALUES ARE RELATED TO HIGHER ENTROP BC Amy one specific configuration will have 1/16 probability - look at scenarios in book Can assign total probability of given configurations and figure out which has highest probabilities of given configurations Entropy def S- is a measure of disorder (puttingall argon in either right or left side is a more ordered configuration- where split they have the max disorder and therefore prob the The entropy of an isolated system increases in the course of spontaneous change !! of course entropy will incr bc universe tends towards more disorder example: cooling a hot metal is a accompanies by an increase in entropy as energy spreads to the surroundings Increase temp- higher entropy We expect molecules in box to go to most disordered state— so half on either side

endothermic and exothermic reactions

When delta H is neg, heat is being released by the system= exothermic reaction Chemical heat packs contain iron filings that are oxidized in an exothermic reaction- your hands get warm because the released heat of the reaction is transferred to your hands When delta H is pos, heat is being absorbed by the system= endothermic reaction Chemical cold packs contain NH4NO3 that dissolves in water in an endothermic process ─ your hands get cold because the pack is absorbing your heat The enthalpy of physical change Phase changes approximate how far apart molecules are Vaporization (boiling), requires adding energy to break and separate intermolecular attractions, an endothermic process - requiring energy to drive liquid molecules into gas phase Freezing is exothermic because energy is given off to allow new intermolecular attractions to form bw molecules Phase changes usually take place at const pressure, heat transfer is due to change in enthalpy !!! think about that freezing is exothermic bc want more heat in surroundings then in system so heat will be given off

Hess's Law

When reaction is multiplied by a factor , delta Hrxn is multiplied by that factor Because delta Hrxn is extensive (depends on how much you are reacting) If a reaction is reversed, the the sign of delta H is changes Can do basic algebraic manipulations on chemical equations- can then put things together to predict some third reaction IF a reaction can be expressed as a series of steps, then the overall reaction is the sum of the heats of reaction for each step Hess's law states: the change in enthalpy for a stepwise process is the sum of the enthalpy changes of the steps

What is energy? Where does it come from?

Work (w) is motion against an opposing force. It is the means by which energy is transferred when objects are in mechanical contact We do work when we lift and object or stretch a spring Work is done when a balloon expands and pushes against atmosphere Mechanical work- physical connection that allows work to be done Heat (q) Is a means by which energy is transferred !!from a hotter object to a colder one when the two are placed in thermal contact (thermal contact- temp related to velocity if molecules so by some mech mol of one medium have rate of v to collide w other molecules at moles level)

gas density and The physical properties of gases determine the volume of gases in chem reactions

at the same P,V, and T they have the same number of molecules. But gases weigh different amounts so they have diff mass— therefore they have diff density d=m/v m=nM, where n is the number of moles and M is the molar mass d= m/v=nM/V You can determine the density of gas if you know its pressure, temp and molar mass Density is directly proportional to molar mass !!! This has practical importance in chemical manufacturing This determines safety parameters for reactions - It determines performance for engines. - It also determines the effectiveness of an explosion in energetic materials. Bhopal gas tragedy The worst industrial accident ever was in Bhopal, India, on Dec. 3, 1984. • Due to poor maintenance, water got into a tank that was holding methyl isocyanate 2H3C-N=C=O(g) + H2O(liq)èCO2(g) + H3C-NH-CO-NH-CH3(g) This reaction produces the same number of moles of 1, 3 dimethyl urea gas as moles of reactant gas, but it is highly exothermic! • The temperature inside the tank rose to about 200oC. T1=300K, T2=500K • The pressure rose in the tank, which contained about between 20 and 40 tons of methyl isocyanate. P2=P1(T2/T1)=(5/3)P1 • The tank leaked, releasing the methyl isocyanate and products. • Since methyl isocyanate is heavier than air, it formed a dense cloud near the ground and spread through the village. Prob was generated a lot of heat which increased temp by almost factor of 2 and so increased pressure by almost factor of 2 and so was very bad

how to do problems where calorimeter has mass and then there is water in it with mass also

calculate energy individually for each and then add it would have to tell you what kettle is made out of and tell you mass-- bc if it is some sort of metal it is going to account for some of the heat so some percentage of that heat is being used to raise temp of water and some is being used to heat cal-- will ask u percent q!!

if problem asking amount of heat to do something and one of things they are asking is heat to go from one state to another what must do?

can not just use q= mc delta t from one state to another bc need the H eq in between each state

daltons law and partial pressure

daltons law:states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases. Partial pressure is what we would expect to see when gas is alone Chem often occurs in mixture of gases Partial pressure- pressure that a gas would produce if it alone was present Daltons law Ptot= Pa +Pa+Pc... Combined w avogadros hypothesis it means that P tot= n tot (RT/V)= (na +nb+ nc) (RT/V) Volume doesn't change bc gas fills those volumes Mole fraction If Xa is the mole fraction of A Xa= nA/nTot and Pa= Xa P tot Note: X is unitless and Pa has the units of pressure Can also be written as nA/ntot = PA/ Ptot A practical use for partial pressures in in the reaction chemistry where a gas is formed and it is collected by bubbling it through water Gas produces in reaction mixed in w water vapor-- the observed gas volume that is contained contains both V of the trapped gas and V of water vapor from bubbling- so will want to use partial pressure to det how much of both--- so P tot= P gas+ P water and n tot= n gas + n water *** in partial pressure probs: Problem In partial pressure problems mol fractions should must to 1 Number of moles will be proportional to pressure so can solve for pressure to and then use ratio to get to moles

how to determine spontanaity with Delta G

delta G=0 equillibrium rxn- doesnt tend either way- can keep going back and forth- double arrow (freeze then melt then freeze etC) delta G neg- spontaneous in forward direction delta G pos- spontaneous in back direction ***the only time the spontaneity is temperature dependent is when ΔΗ and ΔS have the same sign. SOO if asks how spontanaity of a reaction can be reversed and DElta S and H are same sign- it cant

if given set of values for Delta G, delta H or Delta S and then given an eq with their number of moles what is true?

dont need to use any eq to find G, H or S bc have those values so just need to take those values for 1 mol of each- mult by the correct number of moles and do prod- reactants

if problem with volume and temperature changing what are helpful things to think about for asking for q w and delta A

dont need to worry about using delta V and delta T all in one bc they effect diff things- the change in temp will be used in q eq bc change in temp effects heat content= q the v changing effects the work done on a system- p delta v and then the delta U is taking into account both

UNDERSTANDING WHEN KEEPING THINGS CONSTANT WHICH EQ TO USE FOR ENTHALPY VS ENTROPY

entropy has seperate delta S eq for the diff conditions but for enthalpy we use the eq we know!!! W=PV and then use special converting factor if volume is changing Q=mc delta t-- an eq we know but if hold diff things constant then this is where C comes in !!! have to calc diff values for Cv vs Cp ets

polarity but what also effects boiling points??

higher molar mass higher boiling point!! also larger electron radius-- lower MP-- harder to keep track of your babies

relationship bw q and H

if have reaction going on in cal-- doesnt give you q- need to solve for t and gives you H-- remember H is a value given for however many moles of certain reactants and products are in an eq-- so the H rxn is the Q rxn over the number of moles of each react bc H is specific to moles of reactant

if asks what value entropy or enthalpy is driving a reaction-- if you know G value

lets say you know G value is pos- you know that H has to be driving reaction backwards bc that has to be a bigger value than S for overlal G to \be pos!!

what maximum amount of work that can be gained from a rxn is and why this makes sense

max amnt of work gained from a rxn is gibbs- this makes sense bc in gibbs we are taking some enthalpy (heat or energy used to do work) and then subtracting out the entropy the energy that gets lost into universe- so left w max amnt of work that can be gained by a rxn

diff bw monoatomic v linear v non linear and the origin of heat capacities

monoatomic-- In physics and chemistry, monatomic is a combination of the words "mono" and "atomic", and means "single atom". It is usually applied to gases: a monatomic gas is one in which atoms are not bound to each other. --- stable as single atom linear- atoms deployed in straight line- linear electron pair geometry The origin of the heat capacities of gases The translational motion contributes to the hear capacity even at very low temperatures Rotational motion contributes significantly at above .5 K Vibrational motions contribute only at high temps (above 310 k) When the molecules dissociate the hear capacity becomes very large but settles down to the value of 2 moles of I atoms undergoing translational motion

standard enthalpy and standard entropy and determining S sys surr and universe

standard know because will have degree sign above delta H and S standard enthalpy is always 0 for element -- when element at standard state- will usually be element by itself but for C for ex its C(s)-- the enthalpy is 0 but entropy is not!! for system do moles times individual S values for each if at standard -- the surroundings is delta H of whole thing over the T and univ add them both-- but be careful bc if reaction is getting more ordered (it will have neg S value- so Ssys will be neg) but Ssurr then should be pos!!

how to do partial pressure problems

write and balance an equation and then if it is asking for partial pressure of the different products do n prod/ntot= p prod/ptot but the n tot should just be all the mols of the products