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Chemical reactions that take place often THERMOCHEMISTRY

involve energy changes. Almost all chemical - study of energy or heat flow that accompanies a
reactions either release or absorb energy. chemical reaction and/or physical transformation.

HEAT
Burning pieces of wood in a bonfire shows
- form of energy transfer between two objects as
the emission of light and heat energy, a result of their difference in temperature
combustion of butane gas in LPG tank releases - energy generally flows from a hotter object to a
energy, melting of ice absorbs energy and etc. cooler one until thermal equilibrium is reached

TEMPERATURE
- measure of how hot or cold a substance is
relative to another substance
- indicator of thermal equilibrium in the sense that
there is no net flow of heat between two systems in
thermal contact that have the same temperature
- controls the type and quantity of thermal
radiation emitted from a surface

System Open System

- part of the universe being studied or to which - allows transfer of mass (matter) and energy
the attention is focused
Closed System
Surrounding - allows the flow of energy but not the mass
- include everything else in the universe (outside (matter)
the system)
Isolated System
- do not allow transfer of mass (matter) and
energy

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If two systems are made to contact with EXOTHERMIC REACTION

- heat represented as q flows from a system to its
each other at different temperatures, the two
surroundings (-q)
systems will approach a common new - heat is released from the system
temperature that is somewhere between the - initial energy of a system is greater the final energy
initial temperature of the two systems to attain -
Examples:
equilibrium.
- Combustion process which produces heat
- Freezing, Condensation, Deposition
- Cellular respiration

ENDOTHERMIC REACTION According to Law of Conservation of Energy,

- heat flows from the surroundings toward the energy can neither be created nor destroyed but
system (+q) is converted to another form. It can only be
- heat is absorbed by the system transferred between the system and the surroundings.
- initial energy of the system less than the final
energy Energy system = - Energy surroundings

Examples: Negative sign indicates the flow of energy. As the

- Melting, vaporization, sublimation system releases energy, the surroundings must
- Photosynthesis absorbed it.

The First Law of Thermodynamics can be proven The First Law of Thermodynamics states that in
by measuring the change in Internal Energy ( U), any process, the change in energy of a system is
which is a state function. equal to the heat absorbed (q) by the system and the
work (w) done on it.
U = Uf - Ui
Esys = q + w
a change in initial and final state
The net energy flow to or from any system comes in
the form of either work or heat.

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Internal energy (U) is the total energy When work is done ON THE SYSTEM, it
content of the system. gains energy and work is denoted as positive
(+w).
An exothermic reaction is characterized by When the SYSTEM DOES the work, it uses
a negative U, up or transfers some of its energy so that the
work is denoted as negative (-w)
while endothermic reaction has a positive The system may also absorb heat for which
U q is positive (+q), or it may release heat to the
surroundings for which q is negative (-q).

Directions: Answer the following. Show your

Thermodynamic
Function
Significance complete solution.
(+) value means an (-) value means a
U increase decrease 1. What is the change in internal energy of a
(+) value means (-) value means
system that absorbs 523 J of heat and does 452
q heat is absorbed heat is released J of work as a result?
(+) value means (-) value means 2. The internal energy of the system increases by
w work is performed work is performed 20 J and the quantity of work done on a system
on the system by the system is 50 J. Is heat absorbed or given off? By how
much?

 The change in internal energy is related to the

Directions: Answer the following. energy change exchanges that occur as heat (q)
1. What is the work done by a system when and work (w) shown mathematically as:
it absorbs 47 J of heat and gain 12 J of
energy? E = q + w
2. How much heat is released by a system  Because of work is force (F) multiplied by distance
when 45 J of work is done on it to (d), it follows that w is equal to the product of
pressure (P) and volume (V), derived as:
decrease its energy to 23 J?
F
W = Fd = ( ) x d3 = PV
d2

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 If the work is done in the surroundings, the  However, constant-volume conditions are often
equation: E = q + w, at constant volume impossible to do, thus chemists do experiments
becomes under constant pressure.
 This leads to another function called enthalpy, a
E = q - P V or E = qv thermodynamic quantity used to describe heat
changes taking place at constant pressure.
 where subscript v denotes that the equation can The equation is:
be applied only under constant-volume process
( V = 0). H = E + PV

 Thus, change in enthalpy is:

H = qp
H = E + (PV)  where the subscript p means the process is under
constant pressure conditions. The reaction is
 Since E = q - P V, then: when the H is while
H = (q – P V) + (PV) it is when the H is .

 If the pressure is held constant, this become  Overall, it can be stated that at constant volume,
H = q - P V + P V or the heat flow is equal to E; while under constant
H = qp pressure, the heat flow is equal to H.

A thermochemical equation provides the

 Heat of reaction is the heat associated with the following information:
transformation of the reactants to products
1. the balanced chemical equation,
 It is the amount of heat released or absorbed when 2. the moles of substances and their states of
specified amounts of substances react.
matter in the specific equation and
 It indicates also the physical state of the reactants
3. the sign and magnitude of the energy change
and products involved in the reactions which is proportional to the amount of
substance reacted or produced.

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Examples Examples

2. Burning of sulfur 3. Production of C (diamond) from carbon dioxide

S(s) + O2(g)  SO2(g) H = -296.8 kJ CO2(g)  C (diamond) + O2(g) H = 395.4 kJ

This is also an exothermic process, as expected for the This is an endothermic reaction, a process that absorbs
energy-releasing change of burning of sulfur. H is heat from the surroundings. H is positive.
negative.

Examples In writing thermochemical equations, remember

the following laws.
4. Melting of Ice
1. H is directly proportional to mass. Thus, if
H2O(s)  H2O(l) H = 6.01 kJ the coefficients in an equation are doubled,
the value H must also be doubled.
This is also an endothermic process since for the solid to
melt, its molecules must gain enough energy to
overcome the attractive forces between them.

Consider the exothermic combustion of methane Enthalpy is an extensive property. Its magnitude
(CH4) into carbon dioxide and water. is proportional to the amount of reactants and
products in the reaction.
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = - 890.4kJ

This thermochemical equation means that 890.4 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = - 890.4kJ
kJ of heat is released when one mole of gaseous
methane and two moles of gas combust to form one 2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(l) H = - 1780.8kJ
mole of gaseous carbon dioxide and two moles of
liquid water. ½CH4(g) + O2(g)  ½CO2(g) + H2O(l) H = - 445.2 kJ

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In writing thermochemical equations, remember Formation of 1 mole of liquid water from the elements
involves the evolution of 285.84 kJ of heat, when the heat of
the following laws.
reaction is measured at constant pressure. The
thermochemical equation for the formation of 1 mole of
2. H for a reaction is equal in magnitude but water can be written as:
opposite in sign to H for the reverse H2(g) + ½O2(g)  H2O(l) Horxn = -285.84 kJ
reaction.
1 mole of H2O(l) is thermochemically equivalent to 285.84 kJ
1 mole of H2(g) is thermochemically equivalent to 285.84 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = - 890.4 kJ
½ mole of O2(g) is thermochemically equivalent to 285.84 kJ
CO2(g) + H2O(i)  CH4(g) + 2O2(g) H = + 890.4 kJ

Directions: Write the thermochemical equation for the 3. Hydrogen combines with oxygen in fuel cells
following reactions to produce steam. It releases heat energy of
1. Combustion of 1 mol ethanol which yields 571.7 kJ per mole of oxygen gas.
carbon dioxide and water releasing 1 366.8 4. Combustion of 1 mol sulfur to produce sulfur
kJ of heat energy. trioxide releases 395.3 kJ of heat energy.
2. Reaction between nitrogen gas and oxygen 5. Heat energy of 566.0 kJ is absorbed when 2
gas to form nitric oxide, an exothermic mol of carbon dioxide decomposes to carbon
reaction that releases 180.5 kJ heat per monoxide and oxygen gas.
mole of nitrogen.

Sample Problem 1. The principal component of LPG is butane, C4H10(g).

If the combustion of 1 mole of butane produce 126
Given the following thermochemical equation: kJ of heat, how much heat can be generated by
complete combustion of one 11.0-kg gas tank.
C3H6O(l) + 4O2(g)  3CO2(g) + 3H2O(l)
Enthalpy of Reaction = - 1790 kJ

Calculate the amount of heat during the combustion of:

a. 0.30 mol of C3H6O
b. 100 g of C3H6O

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From the following enthalpy changes, STANDARD MOLAR ENTHALPY OF FORMATION

C (s) + O2 (g) → CO2 (g) ∆H° = -393.5 kJ ( Hof)

H2 (g) + ½ O2 (g) → H2O (l) ∆H° = -285.8 kJ - change in enthalpy when the compound is
formed from its elements under standard conditions
2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l) (1 atm, 25oC)
∆H°=-2598.8kJ - enthalpy of formation for an element in its stable
Calculate the value of ∆H° for the reaction form is zero under the same standard condition

2C(s) + H2 (g) → C2H2 (g)

STANDARD ENTHALPY OF REACTION Consider the hypothetical equation:

( Horxn) aA + bB  cC + dD

The sum of enthalpies of formation of the products is

- can be calculated by subtracting the sum (∑) of
calculated as:
the enthalpies of formation of the products and the
sum of enthalpies of formation of the reactants ∑ Hof (products) = c Hof (C) + d Hof (D)

and for the reactants

∑ Hof (reactants) = a Hof (A) + b Hof (B)

Consider the hypothetical equation: Consider the combustion of ethanol (C2H5OH) to produce
aA + bB  cC + dD carbon dioxide and water.

C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g)

The standard molar enthalpy of the hypothetical
reaction is: Calculate the standard enthalpy of the reaction.
Hof (CO2(g)) = -393.5 kJ
Horxn = ∑ Hof (products) - ∑ Hof (reactants) Hof (C2H5OH(l)) = -277.7 kJ
Hof (O2(g)) = 0
Hof (H2O(g)) = -241.8 kJ

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Determine the standard heat of reaction involved in 1. The standard enthalpy of formation of liquid
the formation of sugar from carbon dioxide and water. ethanol, C2H5OH(l) is -277.63 kJ/mol. Calculate
the heat of combustion at constant pressure and
6CO2(g) + 6H2O(l)  C6H12O6(s) +6O2(g) volume of C2H5OH(l).

Standard Enthalpies of Formation: Standard Enthalpy of formation:

Carbon dioxide gas = -393.5 kJ/mol
CO2(g) = -393.5 kJ/mol
Water liquid = -285.84 kJ/mol
H2O(l) = -285.84 kJ/mol
C6H12O6(s) = -1273.3 kJ/mol
O2(g) =0

 heat released or absorbed when solute is  heat change involved in the conversion of liquid
dissolved in a solvent to gas

 heat required to melt a substance  heat involved when gas is converted to liquid

 heat change for converting liquid to solid  heat change when acid and base react to form
salt and water

2. The standard enthalpy of the reaction:

 heat change associated when more solvent is
added to a solution CO(NH2)2(s) + H2O(l)  CO2(g) + 2NH3(g) is -133.42 kJ/mol.

Calculate the standard enthalpy of formation of urea,

CO(NH2)2(s)
 heat change associated with the dissociation of
a substance into ions H2O(l) = -285.84 kJ/mol
CO2(g) = -393.5 kJ/mol
NH3(g) = - 45.9 kJ/mol

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The flow of energy (as heat) between the system For pure substance, the heat capacity is equal to
and its surroundings is measured through calorimetry. the product of its mass(m) and specific heat (c)
The amount of heat absorbed or released by the
system is a function of its change in temperature. C = mc
Therefore,
q=C t
t = tfinal – tinitial q = mc t
The specific heat of a substance is the amount of
where C is constant as heat capacity which is the
energy needed to increase the temperature of one
amount of energy needed to increase the temperature
gram of the substance by 1oC. Its unit is J/g oC.
of a substance of material by 1oC

 Calorimetry involves the measurement of the  In a chemical reaction (for which heat of reaction
quantity of heat exchanged between a system and or heat of neutralization is measured)
its surroundings.
 The heat exchange can be mathematically
 Calorimeter is an insulator apparatus that contains expressed as
water or any liquid of known heat capacity.
qsystem = -[qcalorimeter + qwater]
 The system of interest can be a substance (for
which heat of solution and heat of dissociation is
measured).

1. An 11.5 g ethanol is heated from 25oC to 52.1oC. 2. A copper metal with a mass of 7.56 g cools from
Calculate the amount of heat absorbed by the 65.1 oC to 11.2oC. Find the heat released by the
ethanol. (specific heat of ethanol = 2.46 J / g oC metal. The specific heat capacity of copper is
0.385 J/ g oC.
q = mc (tf – ti)
q = mc (tf – ti)
q = (11.5 g)(2.46 J/ g oC)(52.1oC – 25oC)
q = (7.56 g)(0.385 J/ g oC)(11.2 oC – 65.1oC)
q = 750 J
q = -157 J

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1. A strip of aluminum weighing 0.789 g is heated

from 26.1 oC to 45.8 oC. What is the q? (Specific
heat of Al = 0.900 J/ g oC)
2. A 789.0 g of water is placed in a freezer. The
temperature of water changes from 30.5 oC to
0oC. Calculate the amount of heat released.
(Specific heat of water = 4.184 J/ g oC)
3. Heating a 15.6 g iron bar changes the
temperature by 38.6 oC. Find the q. (specific heat
of iron = 0.444 J/ g oC

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