5.1 measuring energy changes - ms. nordstrom's...
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TOPIC 5 ENERGETICS AND THERMOCHEMISTRY
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5.1 Measuring Energy Changes
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Thermodynamics the study of energy and how it is interconverted
Law of Conservation of Energy also called the first law of thermodynamics, states that energy can be converted from one form to another BUT IT CAN NOT BE CREATED OR DESTROYED!
This law holds true for chemistry in a chemical reaction, total energy is conserved.
Chemical potential energy is stored in the chemical bonds of reactants and products, while the temperature of a reaction is a function of the kinetic energy of the particles involved.
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Units for energy most common is JOULES, J.
A joule of energy is essentially the work done when you exert 1 N of force over a distance of 1 m.
One joule of energy is used every time your heart beats
Other units include: calories (think about energy content of food), although one Calorie in food is actually a kilocalorie or 1000 calories
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WHAT IS HEAT (q, in J)?
Heat, q, is a form of energy that is always transferred FROM a warmer object to a cooler object until the temperature of the objects matches.
Heat can be transferred by conduction, convection or radiation
When heat is transferred to an object, it does WORK by increasing the kinetic energy of the particles in the material, resulting in an increase in the temperature of the object
Absolute zero, 0 K (273.15 C) is the point at which particles theoretically STOP MOVING
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To analyze energy changes in a chemical reaction, we divide the "universe" into two parts
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A chemical system can be open (matter and energy can be transferred), closed (only energy can be transferred), or isolated (nothing is transferred)
Note: the arrows could also point the other way for the first two systems.
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Bond Dissociation Energy During a chemical reaction, chemical bonds in the reactants are BROKEN (requires energy, endothermic) and the atoms are rearranged. The amount of energy needed to break a specific bond is called bond dissociation energy, and it can be measured for every different bond.
New bonds are FORMED (releases energy, exothermic) between these newly rearranged atoms to form products.
Thermochemistry is the study of heat changes that take place during a chemical reaction!
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Enthalpy (H, kJ or kJ/mol) the heat content of a system. We typically are more interested in how the enthalpy of a system CHANGES during a chemical reaction, ΔH.
At constant pressure, we assume q = ΔH.
Otherwise, q = nΔH
Enthalpy is a state function, meaning any change in the value is independent of the pathway between the initial and final value
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Endothermic Reaction a chemical reaction where more energy is absorbed from the surroundings to break chemical bonds in reactants than is returned to the surroundings when chemical bonds of the products form.
Exothermic Reaction a chemical reaction where more energy is returned to the surroundings when product bonds are formed than was required to break the bonds of the reactants
+ ΔH ΔH
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Every substance has a property called specific heat capacity (c) that measures how easily the temperature of that substance can be raised.
Specific heat capacity measures how much heat is needed to raise 1 g of a pure substance by 1 oC or 1 K.
Specific heat capacity is an INTENSIVE property, meaning it does not depend on the size of the sample. 12 g of water and 475 kg of water both have a specific heat capacity of 4.18 kJ kg1 K1
This formula is frequently rearranged to allow us to calculate the heat absorbed or released by a pure substance:
q = mcΔT
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Larger c means more energy is required to heat the subtance.
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Heat capacity is slightly different than specific heat capacity it only measures how much heat is needed to increase the temperature of an OBJECT (so mass and size are now important) by 1 oC or 1 K.
The heat capacity of a swimming pool is much larger than the heat capacity of a tea kettle.
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Example: How much heat is released when 10.0 g of copper with a specific heat capacity of 0.385 J g1 oC1 is cooled from 85.0 oC to 25.0 oC? Is this endothermic or exothermic?
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Example 2: 180.0 J of heat is transferred to a 100.0 g sample of iron, causing its temperature to increase from 22.0 oC to 26.0 oC. Calculate the:
a) heat capacity of the iron sample
b) specific heat capacity of pure iron
c) amount of heat needed to cause the same temperature change in a 325.0 g sample of pure iron
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Homework
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The change in heat of the surroundings is the opposite as the change in heat of the system: qsystem = qsurroundings or ΔHsystem = ΔHsurroundings
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Calorimetry a process for measuring the amount of heat released or absorbed during a chemical reaction or phase change.
One method two styrofoam cups stacked, inner one has water inside it. Initial temperature of water is recorded, and then some sort of reaction or phase change is performed in the water. The reaction releases heat to or absorbs heat from the water, and the temperature change of the water is noted.
qsystem = qsurroundings
qrxn = qwater
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Example: You drop a 385 K piece of hot iron (45.0 g) into 275 g of 298 K water. If the specific heat capacity of iron is 0.450 J g1 K1, what temperature will the iron and the water reach when the heat transfer is complete?
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What are some assumptions we make when we say that in a coffee cup calorimeter, qrxn = qwater?
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Example: You have an aqueous solution of copper (II) sulfate to which you add solid zinc after a short period of time, according to the following equation:
CuSO4(aq) + Zn(s) ZnSO4(aq) + Cu(s)
The heating curve for this is shown on the next slide:
When was the zinc added to the solution?
According to the data points, what is the maximum temperature?
Can we determine a better maximum temperature that accounts for any potential heat loss through the walls of the cup to the surroundings?
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Example: 1.15 g of anhydrous lithium chloride, LiCl, was added to 25.0 g of warm water in a coffee cup calorimeter. A temperature increase of 3.80 K was recorded. Calculate the enthalpy change of solution for 1 mol of lithium chloride.
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How do we experimentally determine the heat of combustion?
IB Value is 1367 kJ/mol
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Homework: Worksheet
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When a molecule changes states of matter (from solid to liquid to gas), the temperature of the molecule REMAINS CONSTANT until the sample has completely changed state. This means we can't use q = mcΔT to calculate the heat content during this process. In these moments we can use:
q = nΔH
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The possible phase changes are:
molar heat of fusion: ΔHfus solid melting to liquid
molar heat of solidification: ΔHsolid = ΔHfus freezing a liquid
molar heat of vaporization: ΔHvap liquid evaporating to gas
molar heat of condensing: ΔHcond = ΔHvap gas to liquid
molar heat of solution: ΔHsoln = dissolving a solid in solvent
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Molar heats of fusion, vaporization, etc are CONSTANT values for specific substances, provided we assume standard conditions
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Example: How much heat is needed to melt 125 g of ice, assuming the ice starts at 0 oC?
How much heat would be released when the same sample of ice froze again?
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Example: A 252 K sample of ice (be careful, this is below zero!) has a mass of 1.85 kg. How much heat is needed to melt this ice, and increase its temperature to 375 K, where it will be a gas?
*Start by drawing a heating curve that includes phase changes!
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Worksheet
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Thermochemical Equations
We can write chemical equations that include the enthalpy change during that reaction. When we combine these two things, a chemical equation becomes a THERMOCHEMICAL EQUATION
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How do we calculate the values of enthalpy change for thermochemical equations?
ΔHfΘ is the standard enthalpy change of formation, which is the energy change that occurs when 1 mol of a particular substance is formed. We use existing information about these enthalpies to calculate energies for particular reactions.
ANY ELEMENT IN ITS NATURAL STATE HAS ΔHfΘ = 0
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We can use the exact same idea for many types of reaction for example, the standard enthalpy change of combustion can be calculated using ΔHcΘ, which is the amount of heat that is released when 1 mol of a particular substance burns in the presence of oxygen. Notice, we calculated this in the previous example.