Thermochemistry
Specific heat of water
1 cal /g C or 4.184 J / g C
Cold-pack equation
NH4NO3 (s) + H2O(l) → NH4NO3 (aq). H= 27 KJ
Joule
SI unit of energy and heat
Heat (energy) transfer
When 2 objects of different temp. are brought into contact to reach equilibrium.
Condensation
When a vapor molecule loses energy, its velocity is reduced and colliding more with other molecules to form a liquid
Heat-pack equation
4Fe (s) + 3O2 (g) → 2Fe2O3 (aq). H= -1625 KJ
absorb
When the cooler objects _____ energy, its temp. rises
Specific heat of metals
x < 1 J / goC
c = 125g water x 4.184 x 1.7 degrees / (11.8 g unknown * 84 degrees) = 0.897 aluminum
125 grams of water are placed in a foam cup. The initial temperature is 22.3 degrees Celsius. An 11.8 gram sample of metal is placed in the water at 108 degrees Celsius. The final temperature is 24.0 degrees Celsius. What is the specific heat of the metal, and what metal might the sample be?
Set the energy gained by water equal to the energy lost by aluminum and the sum of the changes in temperature equal to 80 degrees. Solve these two simultaneous equations for change in temperature
A 5g piece of aluminum foil at 100 degrees C is dropped into a 25g container of water at 20 degrees C. What is the final temp. of the aluminum? 23 degrees Celsius
Hess's Law
A way to calculate the heat of a reaction that too slow or too fast to collect data. Add together several reactions that will result in the desired reaction. Add the ΔH for these reactions in the same way. ΔH (total) = ΔH (products) - ΔH (reactants)
Energy
Ability to do work or produce heat. SI unit is Joules (J).
Calorie
Amount of energy required to raise the temp. of 1g of pure H2O by 1 degree C.
Thermochemical equations
An equation that includes the heat change. Ex: CaO(s) + H2O(l) Ca(OH)2(s) + 65.2 kJ
Increases
As temp. increases, the particles' motion ________
Part I increasing the temperature of ice from -15.0 °C to 0.0°C. Heat 1 = mass x specific heat x temperature change. Heat 1 = 2.09 J/(g•°C) x 10.0 g x (0.0°C -(-15.0°C)) = 313.5 J. Part II changing solid water to liquid water. Temperature does not change. Heat 2 = 10.0 g / 18.015 g/mol x 6.01 kj/mol = 3330 J. Part III increasing the temperature of water from 0.0°C to 50.0°C. Heat 3 = c x m x T. Heat 3 = 4.18 J/(g•°C) x 10.0 g x (50.0°C - 0.0°C) = 2090 J. Total heat = heat 1 + heat 2 + heat 3 = 313.5 J + 3330 J + 2090 J = 5730 J =5.73 kJ (convert)
Calculate the amount of heat absorbed by 10.0 g of ice at -15.0°C as it is converted to liquid water at 50.0°C. The specific heat of H2O(s) = 2.09 J/(g °C). The specific heat of H2O(l) = 4.18 J/(g °C). The molar enthalpy (heat) of fusion is 6.01 kJ/mol.
The enthalpy of the reaction is the sum of the enthalpies of the products less the sum of the enthalpies of the reactants. Enthalpy of reaction = [3 mol FeO × (-272 kJ/mol) + (1 mol CO2 × (-393.5 kJ/mol))] - [1 mol Fe3O4 × (-1118 kJ/mol) + (1 mol CO × (-110.5 kJ/mol))] = +19 kJ
Calculate ΔH° for Fe3O4(s) + CO(g) →3FeO(s) + CO2(g). The standard enthalpies of formation for the two reactants and two products are, in order, -1118, -110.5, -272, -393.5.
Endothermic
Decrease in kinetic energy=decrease in temperature= heat transfers from the surrounding to the system resulting in a cooler surrounding. Absorbs heat. Positive value for q. ΔH = q >0. ΔH products > ΔH reactants (potential energy). Energy is reactant
zero
Elements in their standard state always have standard enthalpies of ________.
Q (heat absorbed or released)= C (specific heat) times M (mass in g) times T (change in temp.)
Equation for calculating heat
HCl is a reactant in the final equation but a product in the original equation. In addition, the amount of HCl in the final equation is twice that of the original. Therefore, the first original equation must be reversed and multiplied by 2. The second original equation does not need to change. The enthalphy of the final equation is -2 (-185 kJ) + (-483.7 kJ) = -114 kJ
H2 (g) + Cl2 (g) →2HCl(g) ΔH°v = -185 kJ. 2H2 (g) + O2 (g) →2H2O(g) ΔH° = -483.7 kJ. Calculate ΔH° for 4HCl(g) + O2 (g) →2Cl2 (g) + 2H2O(g)
Molar Heat of Fusion (ΔHfus)
Heat absorbed by 1 mole of a substance during melting. Constant temp. Ex: H2O(s) -> H2O(l). ΔH = +6.01 kJ/mol
Calorimety math
Heat gained by the water = q. Heat lost by the system = -q. mCΔT = q. q gained by water = q lost by system; q water = - q system. mCΔT = -mCΔT. Ex: (mass H2O)(spec. heat H2O)(ΔT H2O) = - (mass sys)(spec. heat sys)(TΔsys)
Molar Heat of Solidification (ΔHsolid)
Heat lost when 1 mole of a liquid solidifies. Constant temp. Hfus= -Hsolid. EX: H2O(l)->H2O(s). ΔH = -6.01 kJ/mol
Molar Heat of Vaporization (ΔHvap)
Heat needed to vaporize 1 mole of a liquid. Ex: H2O(l) ->H2O(g). ΔHvap= +40.7 kJ/mol
Molar Heat of Condensation (ΔHcond)
Heat released when 1 mole of vapor condenses. Ex: H2O(g) -> H2O(l). ΔHcond = -40.7 kJ/mol. ΔHvap= -ΔHcond
Heat flows from the surrounding to the system
How heat flows in endothermic reaction?
Heat flows from the system to the surrounding
How heat flows in exothermic reaction?
Electrons release light energy when they fall from a high energy level (excited state) to a lower energy (ground state).
How is light energy produced?
Convert grams to moles and multiply by 624.7 to give an answer of 31.2 kJ.
How much heat is absorbed in the complete reaction of 3.00g of SiO2 with excess carbon in the reaction below? ΔH° for the reaction is +624.7 kJ/mol. SiO2(g) + 3C(s) → SiC(s) + 2CO(g)
negative
If object feels hot, it is undergoing change with H that is ______
exothermic
If the object feels hot, it is undergoing ______ reaction
Law of conservation of energy
In chemical reaction or physical changes, energy can be convert from one form to another, but it cannot be create or destroy. In a chemical reaction potential energy is transferred to kinetic energy. First law of thermodynamics
The water changes temperature as heat is either released or absorbed by the reacting substances.
In the calorimeter, what is the function of the water?
Molar mass of ammonium nitrate is 80.043 g/mol. 50 grams is 0.62 moles. 27 kJ/mol x 0.62 mol = 17 kJ.
In the cold pack process, 27 kJ are absorbed from the environment per mole of ammonium nitrate consumed. If 50 g of ammonium nitrate are consumed, what is the total heat absorbed? 27 kJ + NH4NO3(s) → NH4+(aq) + NO3-(aq)
Exothermic
Increase in kinetic energy= increase in temp. = heat released to the surrounding resulting in a hotter surrounding. Negative value for q. ΔH = q < 0. ΔH products < ΔH reactants (potential energy). Energy is a product
Phase change
Occurs when energy is added or removed from a system and the substance can go from one physical phase to another
Sublimation
Process by which a solid changes directly to a gas without first becoming a liquid. Dry ice (CO2)
Deposition
Process by which a substance changes from a gas or vapor to a solid without first becoming a liquid
phase diagram
Shows in which phase a substance exists under different conditions of temperature and pressure.
Chemical potential energy
Stored in the bonds of the reactants and the products. The excess energy is released as heat
Universe
System (part that contains the reaction that desire to study) and surrounding (everything else)
boiling point
Temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure
heat
The burning of fuel always produce _____
Standard enthalpy (heat) of formation (ΔHf)
The change in enthalpy that accompanies the formation of one mole of a compound in its standard state from its elements in their standard states. Ex: S (s)+ 3/2 O2 (g)→ SO3 (g). ΔHf=-396 KJ
Melting
The energy absorbed to melt a solid is not used to raise the temperature of that solid. The energy instead disrupts the bonds holding the solid's molecules together and cause the molecules to move into the liquid phase. Depends on the strength of the forces that hold the solid together
activated complex
The highest energy point of a reaction where the full rearrangement of reactants occurs. Also known as transition state. The unstable arrangement of reactants.
vapor pressure
The pressure exerted by a vapor over a liquid
Amount of heat was constant, initial temp. was constant, and mass was constant, the relationship between the specific heat and the temperature is an inverse proportion. As specific heat decreases, final temp. increases. Therefore, gold, with the smallest specific heat, will reach the highest temperature.
The same amount of heat is added to a 10-g sample of each of the following metals. If each metal is initially at 20.0°C, which metal will reach the highest temp? beryllium 1.82 J/(g°C) calcium 0.653 J/(g°C) copper 0.385 J/(g°C) gold 0.129 J/(g°C)
Kinetic and potential energy
Two major types of energy
lose
When the warmer objects _____ energy, its temp. decrease
evaporation
When vaporization takes place only at the surface of the liquid. Method by which the human body maintains and controls its temperature
Phase change graph
Where the graph inclines, potential energy is at its greatest and temperature is increasing. Where the graph plateaus (flat region) kinetic energy is at its greatest, but the temperature remains constant.
spontaneous process
a physical or chemical change that occurs without outside intervention and require energy from the surrounding to be supply to begin the process
Heat of Reaction (ΔHr)
amount of energy or heat absorbed or released in a reaction
Specific heat
amount of heat required to raise temp. of 1g of pure H2O by 1 degree C.
Enthalpy (heat) of reaction (ΔHrxn)
change in enthalpy in reaction; the difference between enthalpy of substances at end and at the start of reaction; ΔHrxn=ΔH (final) - ΔH (initial)
Potential energy
energy based on the composition or position of an object. Depends on the type of atoms in the substance, the number and type of chemical bonds joining the atoms, and the way the atoms rearranged.
Heat (q)
energy flowing from warmer to cooler object (changing temperature). Depends on the mass of the sample.
Kinetic energy
energy of motion. Related to constant random motion of particles and temp.
Qp
energy released or evolved from reactions carried out at constant pressure
Chemical potential energy
energy stored in the chemical bonds of a substance by arrangement of atoms and molecules. It released or absorbed as heat during chemical reaction.
Enthalpy (heat) of combustion (ΔHcomb)
enthalpy change for the complete burning of 1 mole of substance. Ex: C6H12O6 (s) + 6O2 (g) -> 6CO2 (g) + 6 H2O (L)
Enthalpy (H)
heat content of system at constant pressure; ΔH=q= mCT
critical point
indicates critical pressure and temperature above which water cannot exist as a liquid
Calirometer
insulated device used for measuring the amount of heat absorbed or released in chemical or physical process at constant pressure. Mass of water placed in insulated chamber to absorb energy released from reacting system or provide energy absorbed by the system. The data collected is change in temp. of mass water
Temperature
measure of the average kinetic energy; it does not depend on the amount of matter in the sample
activation energy
minimum amount of energy required to start a chemical reaction
Vaporization
process by which a liquid changes into a gas or vapor.
Standard ethalpy (ΔH0)
reactants and products at standard condition (1atm pressure and 25 C)
triple point
represents the temperature and pressure at which three phases of a substance can coexist
Thermochemistry
study of heat changes that accompany chemical reactions and phase changes.
Freezing Point
temperature at which a liquid is converted into a crystalline solid The same temperature as the melting point of a given substance.
Melting point of a crystalline solid
temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid
