5.1: Measuring Energy Changes
Temperature: the average kinetic energy of molecules (K).
Heat: the amount of energy exchanged due to a temperature difference between two substances
(J).
Enthalpy (H): the amount of energy or heat content of a substance. The energy is stored in the
chemical bonds. It includes kinetic and potential energy and is not measured directly but changes
are measured (J).
Standard enthalpy change of reaction (ΔH°): the difference between the enthalpy of the
products and the enthalpy of the reactants
- 298K and 1.00 x 105 Pa
Change in enthalpy (ΔH): enthalpy of products less than reactants. It is the heat energy change
per mole.
Exothermic Reactions: the temperature of the surroundings increases and ΔH is negative,
energy is given out.
Endothermic Reactions: The temperature of the surroundings decreases and ΔH is positive,
energy is absorbed.
- The higher the enthalpy the less stable the substance.
-
The energy is measured by temperature. We can only calculate the energy change. Not
the energy stored in the resultant products or reactants.
Exothermic Reactions
Examples include:
1. Combustion reactions
2. Neutralization reactions
In an exothermic reaction products:
- Have less stored energy
- Are more stable
- Have stronger bonds (harder to break)
Heat Energy Change (Q)
m = mass in grams
c = specific heat capacity of the substance JK-1g-1
ΔT = change in temperature K
Q = Heat change
ΔH = Heat energy change per mole
Q = mcΔT
ΔH = Q/n
Calorimetry
- A technique used to measure temperature change
Heat absorbed, Q = (m x c x ΔT)liquid
cwater =4.18 JK-1g-1
mwater = volume in cm3, 1cm3 = 1g
Breaking bonds is endothermic.
- Bond dissociation energy
- Absorb heat from their surroundings
- Positive ΔH
- Enthalpy of the products is greater than that of the reactants
- Products are less energetically stable than the reactants
Forming bonds is exothermic.
- Heat transferred from the system to the surroundings
- Negative ΔH
- Enthalpy of products is lower than the reactants
-
Products are more energetically stable than the reactants
If the overall reaction is exothermic, the bonds have lost energy, are more stable, and have
higher bond enthalpies
If the overall reaction is endothermic, the bonds have gained energy, are less stable, and
have lower bond enthalpies.
At absolute zero, all motion of the particles theoretically stops
Entropy: A measure of the distribution of available energy among the particles. The
more ways energy can be distributed, the higher the entropy.
System: where the chemical reaction is taking place
- All the reactants, products, and any solvents
Surroundings: the apparatus that contains the reaction, thermometers or other measuring
devices, the laboratory, everything external to the reacting substances
Thermochemistry: the study of heat changes that occur during chemical reactions
At constant pressure, the change in enthalpy ΔH is defined as the heat transferred by a
closed system during a chemical reaction
ΔH: “change in enthalpy” or “heat of reaction”, measured in kJ
Specific heat capacity, c: The amount of heat needed to raise the temperature of 1g of the
substance by 1°C or 1K. Units are kJ kg-1K-1.
- Lower specific heat capacity = higher temperature rise achieved for the same amount of
heat transferred to the sample
- Intensive property
- Q=mcΔT, m in kg
5.2: Hess’s Law
In a chemical reaction the total change in chemical potential energy (enthalpy change) must be
equal to the energy lost or gained by the reaction system.
Hess’s Law: Regardless of the route by which a chemical reaction proceeds, the enthalpy change
will always be the same providing initial and final states of the system are the same.
Standard enthalpy change of formation (ΔH°f): The energy change upon the formation of 1
mol of a substance from its constituent elements in their standard state. The enthalpy change that
results when one mole of a compound is formed from its elements at 298K and 1.00 x 105Pa. All
reactants and products are in their standard states.
Eg.
C(s) + 2H2(g) → CH4(g)
6C(s) + 3H2(g) + ½O2(g) → C6H5OH(s)
Standard enthalpy change of combustion (ΔH°c): the heat evolved upon the complete
combustion of 1 mol of substance. The enthalpy change that results when one mole of a
compound is formed from its elements at 298K and 1.00 x 105Pa. All reactants and products are
in their standard states.
Eg.
C(graphite) + O2(g) → CO2(g)
H2(g) + ½O2 (g) → H2O(l)
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
5.3: Bond Enthalpies
Average bond enthalpy: the amount of energy required to break one mole of bonds in the
gaseous state averaged across a range of compounds containing that bond.
- Bond breaking is an endothermic process and has a positive enthalpy value
- Average values in section 11 of the data booklet in table 1
- The bond enthalpy for molecules in the gaseous state does not take account of the
intermolecular forces that exist (this is not significant as the average bond enthalpy is an
accepted value to use for enthalpy reaction calculations).
Bond Length
- As you move down a group, ionic radius increases and so does bond length
Bond Strength
- The bond enthalpy reflects the strength of the covalent bond
- Triple bonds stronger than single bonds
Bond Polarity
- Difference in electronegativity of the bonded atoms
- Partial charges = ionic character
- Partial charges attract one another, increasing the strength of the bond
The reason or the difference is that when calculating the enthalpy
change using bond enthalpy values, it is assumed that the reaction
takes place in the gaseous state, with no intermolecular forces involved.
However, the experimentally derived enthalpy of combustion involves
octane and water in their standard states, namely liquid. Additionally,
as mentioned above, all bond dissociation enthalpy values are averaged
across a wide range of related compounds so they represent only an
approximation of the true value.
Ozone is both created and destroyed in the stratospheric layer of Earth's atmosphere. Ultraviolet
(UV) rays from the sun are absorbed by oxygen, O2, splitting the molecule into single oxygen
atoms. These oxygen atoms can then combine with oxygen molecules to form ozone:
UV
O2(g) → O.(g) + O.(g)
O2(g) + O.(g) → O3(g)
Ozone is very effective at absorbing harmful long- and short-wavelength UV radiation. This
absorption breaks down the ozone molecule to reform molecular oxygen and a single oxygen
atom. Without the presence of ozone in the stratosphere, lie on Earth would change forever, as
harmful UV radiation would damage cells in both plants and animals.
The bond dissociation enthalpy of an oxygen molecule is 498 kJ mol-1. In comparison, the energy
required to break an oxygen oxygen bond within an ozone molecule is 364 kJ mol-1.
The consequence of this is that an ozone molecule is decomposed by UV rays more readily than
an oxygen molecule. The ozone photolysis reaction described above is an endothermic reaction
with the required energy coming from the UV radiation. Examination of this energy profile
reveals that the oxygen molecule and oxygen atom have a greater combined energy than the
reactant ozone molecule. The products of this reaction are said to be less stable, as they exist at a
higher energy.
LAB
Assumptions/Sources of Error:
- That no heat is lost to or gained from the surroundings
- Significant errors are associated with reactions
- Involving the evolution of gas
- Involving combustion where hot gas is used to heat liquid in a calorimeter
- Heat rises & ΔT under these conditions are reported to be less than the literature
stipulates
- Thermometers often have precision uncertainty ⁤± 0.1°C or greater
Measuring Enthalpy Changes Graphically
- Temperature is constantly changing to return to room temperature so lines are
extrapolated to calculate the value of ΔT.
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In a chemical reaction total energy is conserved
Chemical potential energy is stored in the chemical bonds of reactants and products
Heat can be transferred through the processes of conduction, convection and radiation
Heat has the ability to do work
System: all the reactants, products and any solvents
Surroundings: the apparatus that contains the reaction, thermometers or other measuring
devices, the laboratory, and everything external to the reacting substances
literature-actual/literature
Calorimetry
Change of formation
Change of combustion
Bond Enthalpies
Hess’s Law