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Energy and Chemical Change
Definitions
Energy is the ability to do work.
Energy is conserved (Law of Conservation of Energy).
Energy is made up of two parts: Heat and Work.
State function: independent of the path, or how you get from point A to B.
Examples of state functionsvolume, pressure, temperature, density, refractiveindex .
Definions
Path functions: properties that depend on
the path (e.g. work).
Work is a force acting over a distance.
Work = Force (N) x Distance (m)
Heat is energy transferred between objects
because of temperature difference
Heat Energy
• Every energy measurement has three parts.
• a unit ( Joules or calories).
• a number ( how many).
• a sign to tell direction.
• negative - exothermic
• positive- endothermic
The Universe
• is divided into two halves.• the system and the surroundings.• The system is the part you are concerned with.• The surroundings are the rest outside the
system.• Exothermic reactions release energy to the
surroundings.• Endothermic reactions absorb energy from
the surroundings.
Quantity of Heat
Heat- energy transferred between objects
because of temperature difference is
expressed as: q = m x c x
Heat lost = Heat gained
-q metal = + q water
T
Calorimetry
The amount of heat evolved or absorbed in
a chemical reaction is meaured by an
apparatus known as a calorimeter.
Calorimetry
q = specific heat x mass x change in
temperature
q = c x m x
where q = heat flow
C = specific heat
m = mass of the substance in grams
= change in temperature =Tfinal - Tinitial
T
T
Specific Heat
• The specific heat is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.
• Molar Heat capacity is the amount of heat required to raise the temperature of one mole of a substance by one degree Celsius.
The Enthalpy Change
• The heat content of a material is called enthalpy.
• It is given the symbol H.
• The heat flow is exactly equal to the
difference between the enthalpy of the
products and that of the reactants.tsreacproductsreaction HHH tan
Example
lOHgO2
1gH 222
kJ 286 pressure, atm 1 C,25o H@
kJ 286O of mole of HH mole 1 of HlOH mole 1 of H 221
22 H
0H
Exothermic reaction is associated with a decrease in enthalpy
while endothermic reaction increases the enthalpy
Thermochemical Equations
• A balanced equation with the states specificed and the heat flow listed.
lOHgOgH 2221
2 kJ 286H
kJ 7.09 ;gOlHgsHgO 221 H
Note: coefficients = # of moles state, l, s, or g
specify temperature
Laws of Thermochemistry
1. is directly proportional to the amount of substance that reacts or is produced in a reaction.2. for a reaction is equal in magnitute but
opposite in sign to for the reverse rxn..3. If a reaction can be regarded as the sum
of two or more other reactions, for the overall reaction must be the sum of the enthalpy changes for the other reactions.
H
H
H
Hess's Law
The third statement above is known as
Hess's Law
Add________________________
kJ7.349 ;sSnClgClsSn 122 H
kJ4.195 ;lSnClgClsSnCl 2422 H
kJ 545 ;SnClCl2sSn 42 Hg
Heats of Formation
• The molar heat of formation of a compound, , is equal to the enthalpy change, , when one mole of the compound is formed from the elements in their stable forms at and 1 atm.
• for any reaction is equal to the sum of the heats of formation of the products compounds minus the sum of the heats of formation of the reactant compounds.
• Note: Any elementary substance in its stable form has zero heat of formation.
S(s) +
fH
H
C 25 o
H
kJ396)g(SOO2
332 o
fH
Example
Note that the for Al and Fe is 0 respectively since they are elements.
)reactants()products( of
of
orxn HHH
Calculate sFe9sOAl4sOFe3s8Al :reaction for the 3243 orxnH
sOFe3sOAl 4 4332 ff HHH kJ 3326kJ 0.11173kJ 8.16694 H
ofH
Spontaneous Reactions/Processes
• A physical or chemical change that occurs without outside influence or intervention. That is, it occurs without obvious reason.
CH4(g) + 2O2(g)
4Fe(s) + 3O2(g)
kJ891)(OH2)g(CO 22 Hl
kJ1625)(OFe2 32 Hs
The above reactions are exothermic and spontaneous.
Spontaneous Reactions/Processes
H2O(s) H2O(l);
NOT exothermic but spontaneous.
So, is NOT the sole determinant of spontaneity
Entropy plays an important role in
determining whether a reaction or process is
spontaneous or not
kJH 01.6
H
Free Energy Change, G
• The capacity of a spontaneous reaction, at constant temperature and pressure, to produce useful work is known as the free energy, G.
Free Energy Change, G
1. If G is , the reaction at constant temperature and pressure is capable of producing useful work and hence is spontaneous.
2. If G is +, work must be supplied to carry out the reaction at const T and P and it is non spontaneous. The reverse reaction is spontaneous.
3. If G = 0, the reaction system is at equilibrium.
Entropy, S
• a measure of the disorder or randomness in a system.
• the tendency towards randomness or disorder Law of Disorder
• spontaneous processes tend to always proceed to increase the entropy of the universe.
Example
a. Change in state
)()()( sSlSgS
0)(OH)(OH 22 systemSgl
0)(OHHC)(OHHC 5252 systemSls
Example
b. Dissolving a gas in a solvent Solution
0 systemS
0)(CO)(CO 22 systemSaqg
Example
c. Chemical reaction
0 systemS as the number of mol gas increases.
as the number of mol gas decreases 0 systemS
2SO3(g) )(O)(SO2 22 gg )(gn = # mol gas products - # mol gas reactants
123)( gn
Therefore, 0S increases)(
Example
d. Solution Formation Solution formation is always accompanied by
increase in entropy
e.Temperature Effect Increase in temperature is always accompanied by
increase in entropy since molecular motion always increase with increase in temperature.
0)(C)(Na)(NaCl Saqlaqs
0)(C)(Na)(NaCl Saqlaqs
The Gibbs-Helmholtz Equation
G = H TS• 1. A negative value of H. Exothermic
reaction H 0, will tend to be spontaneous in as much as they contribute to a negative value of G.
• 2. A positive value of S. If the entropy change is positive, S 0, the term TS will make a negative contribution to G. Reactions tend to be spontaneous if the products are less ordered than the reactants.
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