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Energy and Chemical Change

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Presentation on theme: "Energy and Chemical Change"— Presentation transcript:

1 Energy and Chemical Change

2 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 functions volume, pressure, temperature, density, refractive index .

3 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

4 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

5 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.

6 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

7 Calorimetry The amount of heat evolved or absorbed in
a chemical reaction is meaured by an apparatus known as a calorimeter.

8 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

9 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.

10 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.

11 Example @ Exothermic reaction is associated with a decrease in enthalpy while endothermic reaction increases the enthalpy

12 Thermochemical Equations
A balanced equation with the states specificed and the heat flow listed. Note: coefficients = # of moles state, l, s, or g specify temperature

13 Laws of Thermochemistry
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.

14 Hess's Law The third statement above is known as Hess's Law

15 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) +

16 Example Note that the for Al and Fe is 0 respectively
since they are elements. Calculate

17 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) The above reactions are exothermic and spontaneous.

18 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

19 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.

20 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.

21 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.

22 Example a. Change in state

23 Example b. Dissolving a gas in a solvent Solution

24 as the number of mol gas decreases
Example c. Chemical reaction as the number of mol gas increases. as the number of mol gas decreases 2SO3(g) = # mol gas products - # mol gas reactants Therefore, ( increases)

25 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.

26 The Gibbs-Helmholtz Equation
G = H TS 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 TS 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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