Chemistry Video 11 Flashcards
Energy
Capacity to do work. Displacement of an object. Chemical reactions require energy to occur. Measured in Joules (J)
Kinetic energy
Energy of motion
Potential energy
Energy of location
First law of thermodynamics
Conservation of energy. Energy is not created nor destroyed, it only changes forms.
delta U = q + w
q = heat w = work U = internal energy; sum of the energies present in the system
Internal energy increases if heat is absorbed or if work is done on the system by the surroundings
Internal energy decreases if heat is releases or work is done by the system
+ q = heat flows in
- q = heat flow out
+ w = work done on the system
- w = work done by the system
Expansion work
A type of work. Pressure-volume work.
Occurs when a system pushes against the pressure of its surroundings. i.e. internal combustion engines
State function
Depends on current state of system and not how the system got there.
i.e. internal energy.
Not state function
Depends on how the system got there.
i.e. heat and work
delta Enthalpy
delta H; The sum of a system’s internal energy and the product of its pressure and volume
delta H = delta U + (P)*(delta V)
Internal energy, pressures and volume are all state functions. Therefore, enthalpy is also a state function.
(P)*(delta V)
Represents the expansion work done by the system.
(P)*(delta V) = -w
Change in enthalpy is also equal to…
delta H = q
change in enthalpy is the same as heat exchanged when a system remains at constant pressure
delta H = (delta H of reactants) - (delta H of products)
Bonds breaking vs forming
Bond breaking needs energy. Bond forming releases energy.
Negative delta H
Exothermic, energy released
Positive delta H
Endothermic, energy absorbed
Stoichiometric value of enthalpy
Specific to the molar quantities listed in an equation.
If the coefficients in an equation double, then the change in enthalpy also needs to double
If the direction of the reaction is reversed, then the sign on the change in enthalpy must reverse
Dynamic equilibrium
Changing from 2 phases at the same rate. i.e. going from water to liquid and from liquid to water at the same rate
Vapour pressure
The pressure exerted by dynamic equilibrium between liquid and gas. Increases with increasing temperature
Boiling point
When the vapour pressure is equal to the surrounding pressure of a liquid
Vaporization
Endothermic process. Liquid to gas. Heat energy needed.
Enthalpy of vaporization
The energy associated with the vaporization process.
Enthalpy of vaporization of water at room temperature is 40.65 kJ/mol.
Condensation
Exothermic. Gas to liquid. Heat energy released.
40.65 kJ of energy released per mole of water.
Heat added to solid
As heat is added to solid, its temperature increases.
When the melting point is reached, more heat will NOT go towards raising the temperature of the solid. The heat will disrupt the lattice structure of the solid. The temperature remains constant through the process of melting.
Once completed melted, additional heat will continue to raise the temperature of the substance
Melting point/freezing point.
Temperature at which there is equilibrium between solid and liquid phases. They are both the same temperature. Freezing and melting occurs at equal rates.
Enthalpy of fusion
energy needed to melt a solid.
Weak interactions in solid result in low melting point and smaller enthalpy of fusion.
Enthalpy of fusion is 6.02 kJ/mol for water.
Freezing
Exothermic. Liquid to solid. Heat energy released
Melting
Endothermic. Solid to liquid. Heat energy needed
Sublimation
Solid to gas. Endothermic. Heat energy needed
Enthalpy of sublimation for carbon dioxide is +26 kJ/mol
Deposition
Gas to solid. Exothermic. Heat energy released.
Enthalpy of deposition for carbon dioxide is -26 kJ/mol
Calorie
cal, amount of energy needed to raise the temperature of one gram of water by 1 degrees Celsius
Joule
SI unit for energy, the amount of energy utilized when a 1 N force moves an object a distance of 1 meter
1 Joule = 1 kg m^2/s^2
1 cal = 4.184 J
Heat capacity
Extensive property that depends on amount of substance
Energy needed to raise the temperature of an object by 1 degrees Celsius
OR
Energy released when the temperature of an object is reduced by 1 degrees Celsius
C = q / delta T
q = heat C = heat capacity in J/degrees Celsius units
Specific heat capacity AKA specific heat
Intensive property that does not depend on amount of substance.
Energy required to raise the temperature of 1 gram of a substance by 1 degrees Celsius
q = mcdelta T
c units is J/g degrees C
Calorimeter
Calibrated device that measures heat transferred (and associated values like temperature)
i.e. coffee cup calorimeter with styrofoam
Heating and cooling curves
Shows ice, water and steam as having a non-zero slope because temperature is changing as energy is added.
Ice & water and water & steam are flat horizontal lines because temperature is not changing as energy is added.
Hess’ Law
Change in enthalpy for a reaction is a state function. Can write out a reaction as the sum of several other equations with known enthalpy change. The sum of the enthalpy change of the several equations is equal to the unknown enthalpy change of the full reaction
Enthalpy of formation
Energy associated with the formation of one mole of a compound from its respective elements in their most stable states and under standard conditions.
Symbol is delta H subscript f superscript degrees
Standard state elements have 0 enthalpy of formation
Change in enthalpy equation with enthalpies of formation
Change in enthalpy = (sum of enthalpies of formation of products) - (sum of enthalpies of formation of reactants)
Adiabatic process
Occurs without any transfer of heat or mass between system and surroundings.
Isobaric process
Pressure is constant
Isochoric process
Volume is constant
Isothermal process
Temperature is constant
