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P Chem test 1
| Question | Answer |
|---|---|
| gases assume the ____ and ____ of their containers | volume and shape |
| when pressure is applied to a gas, its volume | decreases; gases are highly compressible |
| gases form what type of mixtures | homogeneous |
| A gas's density is what compared to liquids and solids | lower |
| characteristic properties of gases arise because | individual molecules are relatively far apart (don't feel inter molecular attractions as much) |
| pressure equals | force/area = mass*acceleration/(length*width) = kgm/s^2m^2 therefore 1 N/m^2 = 1 Pa = kg/ms^2 |
| atmospheric pressure equation | P = h(density)g |
| standard atmospheric pressure | 1 atm; the pressure that supports a column of mercury exactly 760 mm high at 0 degrees Celsius at sea level |
| how can a measure of volume be a measure of pressure? | P=h(density)g = 760 X10^-3 m*(1.36X10^4 kg/m^3)*9.81 m/s^2 = 1.01X10^5 Pa because density mercury = 13.6 cm^-3 |
| perfect gas equation of state | PV=nRT |
| perfect gas | the absence of molecular interactions |
| ideal gas | mixtures in which all the molecular interactions are the same but not necessarily zero |
| Boyle's law | PV = constant; P proportional to 1/V |
| isotherm | constant temperature |
| isobar | constant pressure |
| Charles's Law | V proportional to T |
| absolute zero | 0 K; -273.15 degrees Celsius; theoretically lowest attainable temperature |
| why does some ideal gas data have to be extrapolated | gas-liquid phase change |
| Avogadro's Law | V proportional to n |
| density and the ideal gas equation | PV = nRT; d = M/V ; d= PM/RT |
| One of the most variable constituents of air is water vapor, and the humidity it causes. The presence of water vapor results in a lower density of air at a given temperature and pressure. Why? | P,V,n,T; P & T constant; n increase; % of H2O Molar Mass in air decrease (lighter than other constitutes of air); d = PM/RT; RTP constants therefore M decreases |
| Dalton's law of partial pressure | the total pressure of a mixture of gases = the sum of the pressure that each would exert if it were present alone; Pt = Pa + Pb +... |
| The pressure exerted by a particular component of a mixture of gases is called the | partial pressure of that gas |
| Pj = xjP is what | partial pressure of a gas equation where xj = nj/ntotal |
| How is the pressure exerted by N2 gas affected when some O2 is introduced into a container if the temperature and volume remain constant? | pressure of N2 won't change but the total pressure will |
| Kinetic Molecular Theory assumptions 1 & 2 | 1. gases consist of large # of molecules in continuous, random motion; 2. combination of volumes of all the molecules of the gas is negligible relative to the Vtotal in which the gas is contained (molecules can be considered as 'points' |
| Kinetic Molecular Theory assumptions 3-5 | 3. intermolecular forces are negligible; 4. collisions are perfectly elastic (PE = 0 b/c of #3); 5. average kinetic energy is proportional to the absolute temperature |
| pressure of a gas is caused by | collisions of the molecules with the walls of the container; container in different areas has varying levels of pressure but if you take the average it feels constant; magnitude determined by how often & how forcefully they hit the wall |
| explanation of temeprature | absolute T of a gas is a measure of average KE of its molecules; 2 different gases are at the same T their molecules have the same average KE; if the absolute T doubles, the average KE doubles |
| Consider three samples of gas: HCl (at 298K), H2 (at 298K), and O2 (at 350K). Compare the average KE of the molecules in the three samples. | HCl and H2 will have the same average KE, O2's will be higher |
| Kinetic model pressure volume equation | pV = (nMc^2)/3 |
| root mean square speed | c= <v^2>^0.5 |
| Maxwell Distribution of Speeds | we can't determine an exact speed of a molecule we need a Gaussian Distribution |
| Normal Gaussian Distribution equation | e^(-ax^2) |
| Why is the Maxwell distribution of speed a shifted Gaussian curve | v^2 term |
| do all molecules travel at the same speed | nope that's why we have a distribution |
| Maxwell distribution of speeds | f(v) = 4pi (M/2piRT)^3/2 v^2 *e^(-Mv^2/2RT) |
| Interpretations of Maxwell Distribution | M & T are the only 2 things that can change; v is just your placement n x-axis |
| At high temperatures, a __________ fraction of the molecules can be expected to have high speeds than at low temperatures so the scewed Gaussian shifts right | higher |
| will heavy molecules be likely to be found at very high speeds | not really; light molecules are more likely to move faster than heavy molecules; the fraction of heavy molecules will move slwer than the fraction moving @ a faster speed |
| how do you use Maxwell's distribution of speed? | calculate the fraction in a range of speeds that is too wide to be treated as infinitesimal; fraction range v1 to v2 use an integral |
| mean speed | c = v1 + v2+.../N ; sqr(8RT/piM) |
| most probable speed | sqr(2RT/M) ; speed for which f(v) is a maximum (peak of a curve) |
| Diffusion | the spread of a gas throughout open space; the process by which the molecules of different substances mingle with each other |
| solids and diffusion | the atoms of two solids diffuse into each other when the two solids are in contact, but the process is very slow |
| liquids and diffusion | the diffusion of a solid through a liquid solvent is much faster but mixing normally needs to be encouraged by stirring or shaking the solid in the liquid (the process is then no longer pure diffusion) |
| gases and diffusion | gaseous diffusion is much faster; it accounts for the largely uniform composition of the atmosphere |
| effusion | the escape of a gas through a small hole, as in a puncture in an inflated balloon or tire |
| Graham's Law of Effusion | at a given T and p, the rate of effusion of a gas is inversely proportional to the square root of its molar mass |
| rate of effusion is proportional | 1/M^0.5 |
| Graham's Law | r1/r2 = sqr(M2/M1) at constant temperature and ressure in containers with identical pinholes |
| Graham's Law explained | the only way for a molecule to escape from its container is for it to 'hit' the pinhole. The faster the molecules are moving, the greater is the likelihood that a molecule will hit the pinhole and effuse |
| which will effuse first He or N2 | He has smaller M, therefore effuses faster than N2 |
| Diffusion and Mean Free Path | open a vial of perfume at one end of a room. a few minutes elapse before the scent is detected at the other end of the room. this is because of mean free path |
| which to dyou smell first ammonium and HCl | HCl is heavier than NH4, but HCl is moving faster than the NH4 particles |
| RMS Speed | sqr(3RT/M) |
| Molecular Collisionsnm | kintetic thery can be used to determine the collisional frequency between gaseous particles. Understanding the frequency of olecular collisions is important in describing chemical phenomena like rates of chemical reactions |
| mean free path | average distance traveled by a molecule between collisions. Varries with pressure; V is constant, P is changing, n can change, T is usually held constant; ex. shopping mall @ Christmas? on a Tuesday? |
| z-collision frequency | the average rate of collisions made by one molecule in a given time interval divided by the length of the interval; ex 10 collisions per second, z = 10 s^-1 |
| time of flght | 1/z; average time that a molecule spends in flight between collisions |
| assumpion made for collision frequency and time of flight | hard spheres with radius d and two molecules will hit another molecule |
| collision cross-section | sigma=pid^2 =pi(r1 + r2)^2 |
| mean free path | lambda = (RT)/(2^0.5 NasigmaP |
| collision frequency for one type of particle | (1/2^0.5)(P1Na/RT)^2sigma(8RT/piM1)^0.5 |
| effect on the mean free path of the gas molecules in a sample of gas: a. increasing pressure; b. increasing T | a. mean free decreases; b.no change b/c T/P = constant and V & n are constant and T and P are changing proportionally |
| Real Gases | rel gases show deviations from the erfect gas law because molecules ineract with one another and actually have volume; thereefore they have some PE |
| Real Gases: at intermediate separations, where the PE is | negative, the attractive interactions dominate |
| Real Gases: at large separations the PE is | 0 and there is no interaction between the molecules |
| For a perfect gas PV/RT is wat at all pressure if n =1 | 1 b/c other terms will change to balance it out |
| Real Gases: at high pressures, deviation from ideal behavior is | large and is different for each gas |
| at low pressure and high temperature | real gases follow ideal gas laws |
| Would you expect He gas to deviate from ideal behavior at | low temp and hight pressure |
| Understanding Deviations from ideal behavior | V of the gas particles has an effect on the V of the gas; interactions between gas particles have an effect on the P of the gas (there is PE); real gases are not t masses; they have V |
| Effect of Volume | n&Tconstant; real molecules have V; at high p, the combined V iof the gas molecules is a larger fraction of the containerV; the non-0 V of molecules themselves makes the V available for the molecules to move in less than V |
| effect of interaction between gas particles | p depends on z w/ walls & force of each collision; attractive forces from nearby molecules; interactions lessen impact w/ molecules hit the wall therefore pressure appears > than it is; attractive forces and volume effects are more dominant under high p |
| effects oof temperature | T decreases,KE decrease; intermolecular attractions reamin constant; cooling a gas deprives the molecules of energy; more KE means molecules have more energy to break away from intermolecular forces & hit the wall as opposed to being pulled back |
| van der Waals coefficients | a and b; characteristic of each gas but independent of the temperature |
| why is pressure term added in van der Waals equation | attractive forces reduce the pressure |
| why is volume term subtracted in van der Waals equation | particles actually have volume |
| van der Waals equation can only predict above | critical temperature |
| Redlich-Kwong euation state's a and b are ________________ van der Waal's equation | different from |
| compression factor | Z= PVm/RT =Vm/Vm perfect; the ratio of a gas of its measured Vm to Vm of a perfect gas at the same pressure and temperature; deviation from Z is a measure of departure from perfect behavior |
| at high pressures Z | Z>1 therefore Vm > Vm perfect gas |
| at intermediate pressures Z | Z< 1 Vm <Vm perfect attractive forces reducing molar volume relative to that of perfect gas; lots of attractive forces compress the gases (force the molecules to be closer together) |
| PV = ZnRT is true for | a real gas |
| critical temperature | the T at and above which vapor of the substance cannot be liquefied, nomatter how much pressure is applied; at and above Tc the sample has a single phase that occupies the entire V of the container; such a phase is a gas; lq hase doesn't form above the Tc |
| unphysical region | van der Waal's flaws; the region where both P & V are increasing which can't happen due to Boyle's Law |
| Supercritical Fluid | T>Tc and P>Pc ; density is more of that of a liquid; fluid nt strictly a liquid b/c it never possesses a surface that sepparates it from a vapor phase; too dense to be a gas; used as solvents; commercially used for decaffeinated coffee |
| Tc of CO2 (31 degrees Celcius ) is only a few degrees above room T. Therefore, it can be used to treat organic compounds with little danger of thermal degraton. CO2 is an excellent solvent for | nonpolar compounds but not as good for polar compounds |
| Pr | P/Pc |
| Vr | Vm/Vc |
| Tr | T/Tc |
| reduced state | basically proportionizing them; pretty good approximation all gases show the same P-Vm-T behavior |
| Law of Corresonding States | an approximation; it works best for gases compoed of sppherical molecules, fails sometimes badly whenthe molecules are non-pherical or polar |
| Law of Corresponding States graph | P-V-T data for gases within a few %, excet for compounds with large dipole moments |
| thermodynamics | the study of the transformations of energy and, in particular, the transformation of heat into work and vice versa; originally formulated by physicists and engineers interested in the efficiency of steam engines |
| thermodynamics of steam engines | water is heated and boils and the expanding steam does work to ropel the locomotive. example of a practical benefit of termodynamics |
| why we care about thermodynamics | reactions reach equilibrium; composition at equilibrium, and how reactions in electrochemical (and biological) cells can be used to generate electricity |
| Classical thermodynamics | `provides useful relations between observable properties of bulk matter |
| statistical thermodynamics | theory of the connection between atomic and bulk thermodynamic properties |
| work | done to achieve motion against an opposing force; energy used to cause an object with mass to move |
| energy | capacity to do work |
| heat | energy used to raise temperature of an object |
| thermodynamic system | could be a reaction vessel, an engine, an electrochemical cell, a biological cell, etc; ex. popping popcorn |
| thermodynamic process | a prcess where there are changes in state of a thermodynamic system |
| surroundings | region outside the system and are where we make our measurements; so huge it either has constant volume or constant pressure regardless of any changes that take place to the system; remain effectively the same size |
| open system | exchange both energy and matter with its surroundings |
| closed system | exchange energy but not matter with its surroundings |
| isolated system | can exchange neither matter nor energy with its surroundings |
| universe | system and surroundings |
| flask that is not stoppered and to which various substances can be added is an example of an | open system |
| stoppered flask, energy can be exchanged with the contents of the flask b/c the walls may be able to conduct heat is an example of a | closed system |
| insulated thermos containing hot coffee is an imperfect example of an | isolated system |
| diathermic | walls that permit heating as a mode of transfer of energy; e.g. metal container |
| adiabatic | walls that don't permit heating even though there is a difference in temperature; e.g. double walls of a vacuum flask (approximate) |
| exothermic | a process in a system that releases energy as heat; e.g. combustion of coal; qp = - |
| endothermic | a process n a system that absorbs energy as heat. e.g. endothermic dissollution of ammonium nitrate in water in instant cold-packs that are included in sme first-aid kits; qp = + |
| expansion work | many chemical reactions produce gas. Work is done when a system expands against an opposing pressure. Let us find out the work done when a system expands through a volume (deltaV) against a consntant Pex. |
| work done equation at constant pessure | w=-Pex(deltaV) |
| free expansion | expansion against 0 opposing force (Pex=0 therefore w=0) occurs in a vacuum when gas expands |
| work = max when | a reversableprocess |
| reversible change | a process that can be reversed by an infinitesimal change in a variable (pressure); Pex = Pgas equilibrium after each change we have to wait for them to reach equilibrium before moving on (mechanical equilibrium) |
| max work is obtained when the expernal pressure is | only infinitesimally less than the pressure of the gas in the system |
| for isothermal, reversible expansion of an ideal gas work = | w= -nRT ln(Vf/Vi) |
| irreversible work | boom and done |
| an irriversible process is an | idealized process in real life |
| Internal energy | sum of KE + sum of PE of all particles; canot measure total energy but we can measure then change in |
| internal energy equation | deltaU = q + w |
| the system is blind to the mode employed in | internal energy |
| First Law of Thermodynamics | if a system is isolated from its surroundings, then no change in internal energy takes place; q=0 w=0 |
| +q | system gains heat |
| -q | system loses heat |
| +w | work done on system |
| -w | work done by system |
| +delta U | net gain of energy by the system |
| -delta U | net loss of energy by the system |
| intensive property | does not depend on the amount oof sample being examined;; ex. T, mp, density |
| extensive properties | depend on quantity of sample; ex. mass, vollume, internal energy |
| state function | doesn't care how yyou got there, just cares about being there; ex. internal energy; depends on initial and final |
| heat and work _______ state functions | aren't |
| atmosphere conditions | constant pressure |
| enthalpy | H = U + PV (if only PV work); state function; a nonsense thermo term that has no physical meaning; temperature dependent |
| enthaply at constant pressure and PV work only | deltaH = qp |
| molar enthalpy | Hm = H/n =Um + RT |
| enthalpy is more useful that internal energy b/c | qp is easily measured or calculated b/c many physical and chemical changes of inerest occur at constant pressure |
| for most reactions the difference between delta H and delta U is | small because PdeltaV is small especially in solid or liquid phase |
| under constant V delta U = | qv |
| specific heat | the amount of heat that must be added to one gram of a substance to raise its temperature by 1K or 1 degree Celcius; property of a material; intensive; difference between K and degrees Celcius is the same; s=q/(mdeltaT) |
| heat capacity | C=ms=q/deltaT the amount of heat required to raise its temperature by 1K or 1 degree Celcius |
| calculate the work done when 50g of iron reacts with HCl to produce FeCl2 (aq) and H in a closed vessel of fixed volume | 0 b/c volume is constant w=-PdeltaV |
| calculate the work done when 50 g of iron reacts with HCl acid to produce FeCl2 (aq) and H in an open beaker at 25 degree Celcius | w=-PVf = -nRT = (50g/55.85g)*8.314 J/mol*K *298.15K = -2.2 kJ |
| calorimetry | the study of heat transfer during physical and chemical processes |
| calorimeter | a device for measuring energy transferred as heat; heat capacity must be known |
| to ensure adiabaticity, the calorimeter is | immersed in a water bath w/ T continuously readjusted to that of the calorimeter at each stage of the combustion; makes them in equilibrium; feedback loop |
| heat capacity of the calorimeter is determined by | combustion a sample that releases a known quantity of heat and measuring the temperature change; C may be determined by burning a known mass of substance (usually benzoic acid) that has a known heat output |
| constant pressure calorimetry | atmosphere is not changing; qsoln=-qrxn ; acid-base neutralization; heat of solution |
| ex. of constant pressure calorimetry | T of solution is rising b/c acid-base rxn is exothermic & you're left w/ aq soln since all you have in the cup is the solution the heat given off is absorbed by the solution the reaction is happening in |
| enthalpies of reaction | deltaH=Hf-Hi (can't actually measure); enthalpy change that accompanies a reaction is called the deltaHrxn under constant pressure |
| what effect does reversing a reaction have on deltaH for a reaction | change the sign of deltaH |
| what effect does multiplying coefficients by 2 have on deltaH for a reaction | deltaH doubles |
| Hess's Law | if a reaction is carried out in a series of steps, deltaH for the overall reaction will equal the sum of the enthalpy changes for the individual steps; state function |
| standard enthalpy of formation | deltaH(nod)f; enthalpy change that results when one mole of a compound is formed from its elements at a pressure of 1 atm (1 bar); usually reported at 298 K |
| standard enthalpy of formation is most stable for any element when it equals | zero |
| enthalpy of reaction | deltaH(nod)rxn = sum of n*deltaH(nod products - sum of m*deltaH(nod)reactants (n & m are coefficients) |
| bond dissociation energy | the enthalpy change required to break a particular bond in one mole of gaseous molecules |
| all bond enthalpies are_______charged. Why? | positively; takes energy to break a bond so you have to put energy into it |
| chart values for bond dissociation energy are bad. Why? | they don't take into account the chemical environment and their geometric shape (ex. equatorial vs. axial) |
| computer-aided molecular modelling | estimates standard enthalpies of formation but it also has to make assumptions |
| most reliable method for determining enthalpies of formation | calorimetry |
| Can you develop a solution of anions or cations alone? | no; the charge has to be neutralized |
| which ion is defined to have a zero standard enthalpy of formation at all temperatures | H+ (aq) |
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