Thermochemistry Energy and Chemical Change
Energy Energy can change for and flow, but it is always conserved.
The Nature of Energy Energy the ability to do work or produce heat Potential energy Kinetic energy
Chemical systems contain both kinetic and potential energy Kinetic energy Potential energy
Chemical potential energy Energy that is stored in a substance because of its composition
Heat Heat = energy Symbol = q Always flows from warmer object to cooler object
Measuring Heat calorie the amount of energy needed to raise the temperature of one gram of pure water by one degree Celsius Calorie vs. calorie
Joule the SI unit of energy 1 calorie = 4.184 J A chemical reaction releases 213 J of energy. How many calories has this reaction released?
A big mac contains 550 Calories, express this energy in Joules.
Specific Heat Amount of heat needed to raise the temperature of 1 g of a substance by 1 o C Symbol = c Unit = J/g o C High specific heat = absorbs a lot of energy w/small changes in temperature Water = 4.184 J/g o C Concrete = 0.84 J/g o C
Calculating changes in heat q = (m)(c)( T)
If the temperature of 34.4 g of ethanol increases from 25.0 o C to 78.8 o C, how much heat has been absorbed by the ethanol? (specific heat of ethanol = 2.44 J/g o C)
The temperature of a 10.0 g piece of iron changed from 50.4 oc to 25.0 oc, how much energy was release by the iron? (specific heat of iron is 0.449 J/g o C)
Calorimeter insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process. Heat
Chemical Energy and the Universe Thermochemistry the study of heat changes that accompany chemical reactions and phase changes System specific part of the universe that contains the reaction or process you want to study Surroundings everything else in the universe other than the system
Law of conservation of energy 1 st law of thermodynamics energy cannot be created or destroyed, only transferred
Energy being absorbed/released is indicated from the system s point of view - q = energy released by system(exothermic) + q = energy absorbed by system(endothermic) q system = -q surroundings (m 1 )(c 1 )( Δ T 1 ) = - (m 2 )(c 2 )(Δ T 2 )
A 125 g sample of iron at 93.5 o C is dropped into an unknown mass of water at 25.0 o C. The final temperature of the mixture is 32.0 o C. The C of iron is 0.451 J/g o C, the C of water is 4.18 J/g o C. What is the mass of the water?
A 118 g piece of tin at 85 o C is dropped into 100 g of water at 35 o C. The final temperature of the mixture is 38 o C. C of water is 4.18 J/g o C. What amount of heat is absorbed by the water? What amount is released by the tin? What is the specific heat of tin?
Enthalpy and enthalpy change Enthalpy (H) heat content of a system at a constant pressure Cannot measure enthalpy directly but can measure change in enthalpy (heat absorbed or released in a chemical reaction) Enthalpy of reaction ( H rxn ) change in enthalpy for a reaction
ΔH rxn = H products H reactants Exothermic reaction Endothermic reaction
Thermochemical Equations Thermochemical equations express the amount of heat released or absorbed by chemical reactions
Writing Thermochemical Equations Thermochemical equation balanced chemical equation that includes the states of all reactants and products and the energy change CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) ΔH comb = -891 kj
Enthalpy (heat) of combustion (ΔH comb ) the enthalpy change for the complete burning of one mole of the substance Enthalpy (heat) of vaporization (ΔH vap ) heat required to vaporize one mole of a liquid Enthalpy (heat) of fusion (ΔH fus ) heat required to melt one mole of a solid
How much heat is released when 54.0 g of glucose is burned? ΔH comb = -2808 kj/mol
H 2 O(l) H 2 O(g) ΔH vap = 40.7 kj H 2 O(s) H 2 O(l) ΔH fus = 6.01 kj
Phase Change Diagram C steam = 1.70 J/g o C Liquid steam: 40.7 kj/mol C Water = 4.18 J/g o C solid liquid: 6.01 kj/mol C ice = 2.10 J/g o C
How much heat must be absorbed to melt 150.0 g of water?
How much heat is released when 50.0 g of steam cools to 40 o C?
Challenge!!! A 39.0g sample of ice at -125 o C changes into steam at 125 o C. How much energy is absorbed during this process?
Calculating Enthalpy Change Hess s Law if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction
What is the energy change for the following reaction: 2S(s) + 3O 2 (g) 2SO 3 (g) a. S(s) + O 2 (g) SO 2 (g) ΔH = -594 kj b. 2SO 3 (g) 2SO 2 (g) + O 2 (g) ΔH = 198 kj
What is the energy change for the following? 2H 2 O 2 (l) 2H 2 O(l) + O 2 (g) a. 2H 2 (g) + O 2 (g) 2H 2 O(l) ΔH = -572kJ b. H 2 (g) + O 2 (g) H 2 O 2 (l) ΔH = -188kJ
Standard heat of formation The change in enthalpy that accompanies the formation of one mole of the compound in its standard state ΔH f ΔH f of an element = 0
ΔH rxn = sum of ΔH f products sum of ΔH f reactants
What is ΔH rxn for the following equation: CH 4 (g) +2O 2 (g) CO 2 (g) + 2H 2 O(l)
What is the ΔH rxn for the following: 4NH 3 (g) + 7O 2 (g) 4NO 2 (g) + 6H 2 O(g)
Reaction Spontaneity Changes in enthalpy and entropy determine whether a process is spontaneous
Spontaneous processes Any physical or chemical change that once begun, occurs with no outside intervention Iron rusting Paper burning Often some energy from the surroundings must be supplied to get the process started
Entropy Entropy(S) a measure of the number of possible was that the energy of a system can be distributed Determined by the freedom of the systems particles to move and the number of ways they can be arranged Disorder or randomness of a system
Second law of thermodynamics spontaneous processes always proceed in such a way that the entropy of the universe increases
Predicting changes in entropy +ΔS = entropy increases - ΔS = entropy decreases
Changes resulting in +ΔS Changes in state that allow more molecule movement (s) (l) (l) (g) Number of particles increases in a reaction CaCO 3 CaO + CO 2 Solid or liquid dissolves in solvent Increase in temperature
Changes resulting in -ΔS Phase changes that decrease molecule movement (g) (l) (l) (s) Number of particles decreases in a reaction Dissolving of gas in a solvent Decrease in temperature
Predict the sign of ΔS for the following: ClF(g) + F 2 (g) ClF 3 (g) NH 3 (g) NH 3 (aq) CH 3 OH(l) CH 3 OH(aq) C 10 H 8 (l) C 10 H 8 (s)
Gibbs Free Energy ΔG = ΔH TΔS - ΔG = spontaneous reaction + ΔG = nonspontaneous reaction
For a process, ΔH = 145 kj and ΔS = 322 J/K. is the process spontaneous at 382K?
ΔH ΔS ΔG Reaction Spontaneity Negative Positive Negative Always spontaneous Negative Negative Negative or positive Positive Positive Negative or positive Spontaneous at low temperatures Spontaneous at high temperatures Positive Negative Positive Never spontaneous