thermodynamics 101thermodynamics 101 first law of thermodynamics energy is conserved in a reaction...
TRANSCRIPT
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Enthalpy
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Thermodynamics 101
First Law of Thermodynamics
Energy is conserved in a reaction (it cannot be created or destroyed)---sound familiar???
Math representation: ΔEtotal = ΔEsys + ΔEsurr = 0 Δ= “change in” ΔΕ= positive (+), energy gained by system ΔΕ= negative (-), energy lost by system Total energy = sum of the energy of each part in a
chemical reaction
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Mg+ 2HCl MgCl2+ H2
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Exothermic
Temperature increase (--isolated system)
Heat is released to surroundings (--open/closed system)
q = - value
Chemical Thermal Energy
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Endothermic
Temperature decrease (--isolated system) All energy going into reaction, not into surroundings
Heat absorbed by system, surroundings have to put energy into reaction
q = + value
Thermal Chemical Energy
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Heat of Reaction
Amount of heat exchange happening between the system and its surroundings for a chemical reaction.
Temperature remains constant
Usually reactions happen at constant volume or constant pressure
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How does work factor into heat of reaction?
W = -PΔV
If volume is constant (ΔV), PΔV = 0 and no other work sooooo
If pressure (P) is constant so volume can change, work is being done soooo
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Work in terms of energy change
System DOES work------ POSITIVE work value for system, system is LOSING energy
System has work on ON it----NEGATIVE work value for system, system is GAINING energy
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Enthalpy (H) Measures 2 things in a chemical reaction:
1) Energy change
2) Amount of work done to or by chemical reaction
2 types of chemical reactions: 1) Exothermic—heat released to the surroundings, getting rid of heat,
-ΔΗ
2) Endothermic—heat absorbed from surroundings, bringing heat in, +ΔΗ
**Enthalpy of reaction—heat from a chemical reaction which is given off or absorbed, units = kJ/mol
Enthalpy of reaction Heat from a chemical reaction which is given off or absorbed At constant pressure Units = kJ/mol
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Enthalpy (H) cont.
Most chemical reactions happen at constant pressure (atmospheric pressure)—open container
Temperature and pressure are constant Only work is through pressure/volume
Sum of reaction’s internal energy + pressure/volume of system H = U + PV ΔH = ΔU + PΔV
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Properties of Enthalpy
Extensive Property Dependent on amount of substance used
State Function Only deals with current condition Focus on initial and final states
Enthalpy changes are unique Each condition has specific enthalpy value SO
enthalpy change (ΔH) also has specific value
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Example 1
CH4 + 2O2 CO2 + 2H2O ΔH = -890.3 kJ
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Example 2
2HgO 2Hg + O2 ΔH = + 181.66 kJ
HgO Hg + ½ O2 ΔH = + 90.83 kJ
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More Enthalpy
The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change
HgO Hg + ½ O2 ΔH = + 90.83 kJ
Hg + ½ O2 HgO ΔH = - 90.83 kJ
SOOO-----total ΔH = 0
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Example 1:
Based on the following:
2Ag2S + 2H2O 4Ag + 2H2S + O2 ΔH = +595.5 kJ
Find the ΔH for the reaction below:
Ag + ½ H2S + ¼ O2 ½ Ag2S + ½ H2O ΔH = ?
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Example 2:
Write a chemical equation for ice melting at 0°C through heat absorption of 334 kJ per gram.
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Stoichiometry Returns
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Example 1:
H2 + Cl2 2HCl ΔH = -184.6 kJ
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Example 2:
Calculate the ΔH for the following reaction when 12.8 grams of hydrogen gas combine with excess chlorine gas to produce hydrochloric acid.
H2 + Cl2 2HCl ΔH = -184.6 kJ
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Example 3:
Pentaborane (B5H9) burns to produce B2O3 and water vapor. The ΔH for this reaction is -8686.6 kJ/mol at 298°K. What is the ΔH with the consumption of 0.600 mol B5H9 ?
2B5H9 + 12O2 5B2O3 + 9H2O
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Calorimetry
How do we find the change in energy/heat transfer that occurs in
chemical reactions???
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Calorimetry
Experimentally “measuring” heat transfer for a chemical reaction or chemical compound
Calorimeter Instrument used to determine the heat transfer of a chemical
reaction Determines how much energy is in food Observing temperature change within water around a reaction
container
** assume a closed system, isolated container No matter, no heat/energy lost Constant volume
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Specific Heat Capacity
Amount of heat required to increase the temperature of 1g of a chemical substance by 1°C
Units--- J/g°K
Unique to each chemical substance Al(s) = 0.901J/g°K
H2O(l) = 4.18 J/g°K
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q = smΔT
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Example 1
How much heat is needed to raise the temperature of a 500g iron bar from 25° to 50°C ?
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“Coffee Cup” calorimeter
Styrofoam cup with known water mass in calorimeter Assume no heat loss on walls Initial water temp and then chemical placed
inside Final temperature recorded
Any temperature increase has to be from the heat lost by the substance SOOO All the heat lost from the chemical reaction or
substance is transferred to H2O in calorimeter
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“Coffee Cup” calorimeter (cont.)
qchemical = -qwater
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Example 2: Using the following data,
determine the metal’s specific heat.
Metal mass = 25.0g Water mass = 20.0g
Temperature of large water sample = 95°C
Initial temperature in calorimeter = 24.5°C
Final temperature in calorimeter = 47.2°C
Specific heat of water = 1.00 cal/g°C OR 4.184 J/g°K (KNOW!!!!)
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Δqrxn Heat gained/lost in experiment
with calorimeter
ΔHrxn
Heat gained/lost in terms of the balanced chemical equation
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Example 3:
A 50.0 ml sample of 0.250M HCl and 50.0 ml sample of 0.250M NaOH react in a cofee cup calorimeter. The temperature increases from 19.50°C to 21.21°C. Calculate the ΔH for this reaction.
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Homework
pp. 251-252 #25, 27, 33-35