ch 5 thermochemistry Flashcards
thermodynamics
study of energy and its transformations. Thermochemistry: relationships between chemical reactions and energy changes
energy definition
capacity to do work or to transfer heat
kinetic energy3
energy of motion. Ek=1/2mv2. m=mass, v=velocity. Kinetic energy increases as speed increases. Atoms/molecules have mass and are in motion=have kinetic energy
potential energy2
energy by virtue of position relative to other objects. Arises when there is a force (push or pull) acting on an object. Ep=mgh. m=mass, g=gravitational constant=9.8m/s2. h=height
electrostatic potential energy4
arises from interaction between charged particles. Eel is proportional to the electrical charges on the two interacting objects and inversely proportional to the distance between themEel=(kQ1Q2)/d. k=constant of proportionality, =8.99*109J–m/C2. (C =coulomb, unit of electrical charge. J=joule). When Q1 and Q2 have the same sign, they repel one another, and Eel is positive. When they have opposite signs, they attract and Eel is negative.Lower energy of system=more stable.
SI unit for energy3
Joule, J. 1J=1kg–m2/s2. Energy associated with chemical reactions is kilojoules.
non SI unit for energy4
calorie (cal)=amount of energy to raise temperature of 1g of water from 14.5ºC to 15.5ºC. 1 cal=4.184J. 1Cal=1000cal=1 kcal
closed system
can exchange energy but not matter with its surroundings, does not lose or gain mass
work2
energy used to cause an object with mass to move. w=F*d
heat
energy used to cause the temperature of an object to increase. Transferred from hot object to cold objects
first law of thermodynamics
energy is conserved. Any energy lost by the system must be gained by the surroundings, and vice versa.
internal energy defs
sum of all kinetic and potential energies of components of system. Represented as E. When E is positive=system gained energy from surroundings. When E is negative=system lost energy to surroundings. Initial state=reactants, final state=products.
internal energy formulas2
E=Eproducts-Ereactants
E=q+w.
endothermic3
system absorbs heat. (melting of ice). Feels cold because heat transferred from our hands (surroundings) to system. E>0
exothermic3
system loses heat. (combustion). Feels hot as temp of system drops and enters surroundings.
E<0
state function
property of a system that is determined by the specifying the system’s condition/state. Value of state function depends only on present state of system, not on path system took to reach that state
pressure-volume work3
work involved in the expansion or compression of gases. When pressure is constant, w=-PV. P=pressure, V=change in volume (Vfinal-Vinitial). When volume expands, V is positive and w is negative, meaning energy leaves system as work and vice versa for when a gas is compressed.
enthalpy5
thermodynamic function that accounts for heat flow in processes occurring at constant pressure when no forms of work are performed other that P-V work. Denoted by symbol H.
H=E+PV.
H=E+PV=(qP+w)-w=qp. qp emphasizes changes at constant pressure.
Change in enthalpy =heat gained/lost at constant pressure
H>0(meaning when qP>0), heat is gained=endothermic
enthalpy of reaction/heat of reaction
Hrxn=Hproducts-Hreactants
thermochemical equations
balanced equations that show the associated enthalpy change where the coefficients represent the number of moles of reactants/products producing the associated enthalpy change.
enthalpy diagram
another way of representing enthalpy change through a diagram. If reaction is exothermic, reactants are written higher than products with an arrow drawing down to indicate loss of heat
facts about enthalpy change5
The enthalpy change of a reaction is equal in magnitude but opposite in sign to H for the reverse reaction.
When H is large and negative, it is spontaneous/thermodynamically favored
When H is large and positive, it is spontaneous in the reverse direction
Gases have greater internal energy than liquids, so if liquids were changed to gas on the reactant side, H would decrease but if liquid changed to gas on product side, H would increase
calorimetry
measurement of heat flow. Calorimeter: device used to measure heat flow
heat capacity2
represented as C, the temperature change experienced by an object when it absorbs a certain amount of heat. Amount of heat required to raise its temperature by 1 K (or 1ºC). Greater heat capacity, the greater the heat required to produce a given increase in temp
molar heat capacity
Cmolar, heat capacity of one mole of a substance.
specific heat3
heat capacity of 1g of a substance. Represented as s.
s=q/(mT).
q=sm*T
coffee cup calorimeters4
since they are not sealed, the reaction occurs under constant pressure of the atmosphere.
Heat gained/lost by the reaction (system) must be gained/lost by the solution(surroundings). qsoln= -qrxn. qsoln=ssolnmsolnT=
-qrxn.
A temp increase(T>0) in the solution means the reaction is exothermic (qrxn<0)
bomb calorimeters5
combustion reactions.
Substance to be studied is placed in a small cup within a sealed vessel called a bomb which can withstand high pressures.
Bomb is sealed and pressurized with oxygen and placed in insulated container (calorimeter) when a measured quantity of water.
Combustion reaction occurs, causing water temp to rise. qrxn=-Ccal*T. (Ccal is total heat capacity of calorimeter)
Because reactions in bomb calorimeters are carried out out under constant volume conditions, heat transferred corresponds to change in internal energy, E, rather than change in enthalpy, H
hess law
if a reaction is carried out in a series of steps, Hfor the overall reaction will equal the sum of the enthalpy changes for the individual steps.
H is the same whether the reaction takes place in one step or in a series of steps
enthalpy of formation/heat of formation
Hf, enthalpy change associated with formation of a compound from its elements. Subscript f indicates the substances has been formed from its component elements.
standard state
state of a substance at its pure form at atmospheric pressure, 1 atm, and 25ºC/298 K. Standard enthalpy change of a reaction is the enthalpy change when all reactants and products are in their standard states. Denoted as Hºwhere superscript º indicates standard-state conditions
standard enthalpy of formation
Hºf, is change in enthalpy for the reaction that forms one mole of the compound from its elements, with all substances in their standard states. They are reported in kJ/mol. Each reactant is an element in its standard state and product is one mole of the compound.
The standard enthalpy of formation of the most stable form of any element is zero because there is no formation reaction needed when the element is already in its standard state.
standard enthalpy of reaction formula
Hºrxn= { nHºf(products)-
{ mHºf(reactants). = { sum of, n and m=stoichiometric coefficients of the chemical equation. First term represents formation reactions of products written in forward direction (elements reaction to form products), second term represents reverse of formation reactions of reactants (reactants decompose to form elements)
fuel value
energy released when one g of a material is combusted. Reported as positive numbers
Greater percentage of C and H in a fuel=higher its fuel value
fossil fuels
coal, petroleum, natural gas which have formed over millions of years from decomposition of plants and animals and are being depleted faster than being formed, nonrenewable
natural gas
gaseous hydrocarbons
petroleum
liquid composed of hundreds of compounds, mostly hydrocarbons
coal
solid, contains hydrocarbons of high molecular weight
Coal gasification: coal is pulverized and treated with superheated steam a produces mixture of gaseous hydrocarbons (syngas)
nuclear energy
energy released in either splitting or fusion of nuclei of atoms, nonrenewable. Produce radioactive waste products
renewable energy
solar, wind, geothermal, hydroelectric, biomass energy
