Thermochemical Changes

Thermochemical Changes

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Thermochemistry The study of the energy
Evolved or absorbed in chemical reactions Changes of physical transformations (Example: melting, boiling)

Thermochemistry Deals with the energy exchange of transformations (chemical reactions phase transitions, nuclear reactions, dissolving of ionic compounds) Includes calculations of quantities such as the heat capacity, heat of combustion, heat of formation, enthalpy and free energy.

The atoms and molecules in any sample of matter are in constant random motion.
Depending on the phase or state of a sample of matter, the atoms and molecules are capable of vibrating, rotating and translating to varying extents. Heat is the name given to the energy transferred between two substances that have different temperatures.

Temperature Measurement
Celsius and Kelvin temperatures scales have the same size units (Example:1 oC change is equal to 1 K change)

Temperature Conventions
∆t is used for Celsius measurements ∆T is used for Kelvin measurements Increase in temperature - positive Decrease in temperature - negative

Primary Sources of Energy
Chemical - fossil fuels, plants Solar - direct radiant, wind, water Nuclear - fission, fusion Geothermal - geysers, hot springs

Useful forms of Energy Heat Electrical Mechanical Light Sound

Heat as Energy Almost all forms of energy are eventually converted to thermal energy If heat is so spread out that it cannot be used it is called low level heat

An Introduction to Energetics
Kinetic Energy (Ek) is related to the motion of an entity Molecular motion can by translational (straight-line), rotational and vibrational Chemical Potential Energy (Ep) is energy stored in the bonds of a substance and relative intermolecular forces Thermal Energy is the total kinetic energy of all of the particles of a system. Increases with temperature. Symbol (Q), Units (J), Formula used (Q=mcΔT) Temperature is a measure of the average kinetic energy of the particles in a system Heat is a transfer of thermal energy. Heat is not possessed by a system. Heat is energy flowing between systems. When a substances changes temerpature, its molecules undergo a change in kinetic energy. An increase in temperature signifies an increase in kinetic energy. A decrease in temperature signifies a decrease in kinetic energy.

The Law of Conservation of Energy
DO YOU REMEMBER?? The Law of Conservation of Energy During physical and chemical processes, energy may change form, but it may never be created nor destroyed. If a chemical system gains energy, the surroundings lose energy If a chemical system loses energy, the surroundings gain energy Examples: When octane (C3H8, the main component of gasoline) is burned in your car engine, chemical bond energy (potential energy) is converted into mechanical energy (pistons moving in the car engine; kinetic energy) and heat. When we turn on a light switch, electrical energy is converted into light energy and, you guessed it, heat energy.

DO YOU REMEMBER?? Exothermic ENDOTHERMIC A change in a chemical energy where energy/heat EXITS the chemical system Results in a decrease in chemical potential energy A change in chemical energy where energy/heat ENTERS the chemical system Results in an increase in chemical potential energy

Heat as Energy Exothermic change - reaction releases E (surroundings increase in temp.) Endothermic change - reaction absorbs E (surroundings decrease in temp.)

THE LAW OF CONSERVATION OF ENERGY
Why does the graph plateau? Remember, a phase change is primarily changing a substances potential energy. Bonds are being broken or formed, changing the Ep, during the plateau of the graph. Remember it takes energy to break bonds, and energy is released when bonds are formed When the graph is climbing, the kinetic energy of the particles is increasing because additional thermal energy is being added.

An Introduction to Energetics
Which has more thermal energy, a hot cup of coffee or an iceberg? An iceberg! Put very roughly, thermal energy is related to the amount of something you have multiplied by the temperature. Let's assume your iceberg is at the freezing point of water - 0 degrees Celsius (~273 Kelvin). Now your cup of coffee might be 75 degrees Celsius (~350 Kelvin). 350 isn't a whole lot more than 270, but an iceberg is thousands of times larger than a cup of coffee. Even though the iceberg is at a lower temperature, it contains more thermal energy because the particles are moving and it's much larger than the cup of coffee.

Thermal Energy Calculations
There are three factors that affect thermal energy (Q = mcΔt): Mass (m) Type of substance (c) c is the specific heat capacity, the quantity of energy required to raise the temperature of one gram of a substance by one degree Celsius Temperature change (Δt) Example: Consider a bathtub and a teacup of water! All water has the same specific heat capacity which is 4.19J/g°C. However, the bathtub would take considerably more energy to heat up! The quantity of kinetic energy, Q, lost or gained when a sample of matter with mass “m” undergoes a change in Temperature Remember that substances with large heat capacities can absorb large qauntites if heat energy with only small increases in temperature Substances with small heat capacities can absorb small quanities of heat energy with significant increases in temperature Water has a high heat capacity

Example: Determine the change in thermal energy when 115 mL of water is heated from 19.6oC to 98.8oC? Mass = density x volume Show how L = kg and mL = g The density of a dilute aqueous solution is the same as that of water; that is, 1.00g/mL or 1.00kg/L c water = 4.19J/g °C or kJ/kg °C or kJ/L °C

Example: A sample of ethanol absorbs 23.4 kJ of energy when its temperature increases by 14.25°C . The specific heat capacity of ethanol is 2.44 J/g°C . What is the mass of the ethanol sample? Q = 23.4 kJ Δt = °C c = 2.44 J/g°C m = ? Q = mcΔt m = Q/cΔt m = kJ (2.44J/g°C)(+14.25°C) m = kg = kg

Example How much heat energy is absorbed when 10.0kg of aluminum initially at a temperature of 20 degrees Celsius, warms up to its melting point , 660 degrees Celsius, but does not melt?

Example A 2.50kg sample of water has a temperature of 85 degrees Celsius. What is the final temperature of the water if it loses 220 kJ of heat energy?

Example What mass of water, at 25 degrees Celsius, requires 2.25 MJ of heat energy to warm up 100 degrees Celsius? (Assume the water being heated is below sea level in order to allow for heating to 100 degrees Celsius without changing states.)

Example: Day 2 A 55.0g block of aluminum cools from 53.2°C to 25.0°C. What amount of energy is lost by the aluminum block?

Example: When 50 mL of 1.0 mol/L hydrochloric acid is neutralized completely by 75 mL of 1.0 mol/L sodium hydroxide in a polystyrene cup calorimeter, the temperature of the total solution changes from 20.2°C to 25.6°C. Determine the enthalpy change that occurs in the chemical system.

Endothermic/Exothermic Reactions
Recall that energy is required to break bonds & is released when bonds are formed.

Endothermic/Exothermic Reactions
A net absorption of energy results in an endothermic reactions while a net release of energy gives an exothermic reaction.

Energy of Chemical Systems
Kinetic Energy Moving electron w/in atoms Vibration of atoms connected by chemical bonds Rotation & translation of molecules Potential Energy Covalent and /or ionic bonds between atoms Intermolecular forces

How do we measure Q? With a simple laboratory calorimeter, which consists of an insulated container made of three nested polystyrene cups, a measured quantity of water, and a thermometer. The chemical is placed in or dissolved in the water of the calorimeter. Energy transfers between the chemical system and the surrounding water is monitored by measuring changes in the water temperature. “Calorimetry is the technological process of measuring energy changes of an isolated system called a calorimeter” More to come Includes: Thermometer, stirring rod, stopper or inverted cup, two Styrofoam cups nested together containing reactants in solution

Comparing Q’s Negative Q value Positive Q value
An exothermic change Heat is lost by the system The temperature of the surroundings increases and the temperature of the system decreases Example: Hot Pack Question Tips: “How much energy is released?” Positive Q value An endothermic change Heat is gained by the system The temperature of the system increases and the temperature of the surroundings decreases Example: Cold Pack Question Tips: “What heat is required?”

Enthalpy The total of the kinetic and potential energy within a chemical system is called its enthalpy. (Energy possessed by the system) Enthalpy is communicated as a difference in enthalpy between reactants and products, an enthalpy change, ΔrH . Units (usually kJ) Enthalpy: H (aka Heat Content) - Enthalpy is the change in the potential energy between the reactants and products at constant pressure and is most often equal to the heat absorbed or released by the system.

Enthalpy Change (ΔH) ΔH = HP – HR (IUPAC defn.)
It is not possible to determine the absolute enthalpy of a system (H) but changes in enthalpy can be measured (∆H) ΔH is a difference in enthalpy between reactants and products of a particular chemical system when the pressure or volume of a system is held constant ΔH = HP – HR (IUPAC defn.) Enthalpy of reaction is measured in kJ

Enthalpy Change Enthalpy change can be communicated in the following four ways The molar enthalpy of a process, such as formation or combustion, (units kJ/mol) The enthalpy change of a reaction (units kJ) The enthalpy of a reaction written as a reactant or product term in the equation (units kJ) An energy diagram representing a reaction (units kJ)

Enthalpy Change First it is important to be familiar with some common definitions The standard molar enthalpy change of formation of a compound, is the heat energy absorbed or released when 1 mol of compound is formed from its elements in their standard states. Values of can be found on the table of standard molar enthalpies of formation. The variable can also be represented as The standard molar enthalpy of formation of elements in their standard states is defined as 0 kJ/mol.

Signs Negative enthalpy values indicate exothermic reactions (energy is released to the surroundings thus is lost by the system) Positive enthalpy values indicate endothermic reactions (energy is gained by the system from its surroundings)

Types of Enthalpy As noted earlier enthalpy changes occur during other processes (eg. phase changes, nuclear changes, dissolving of ionic compunds) so your text uses the symbol ∆rH to indicate the enthalpy of reaction

IUPAC Symbols If the reaction is:
combustion the symbol used can be ∆c H formation the symbol used can be ∆f H decomposition the symbol used can be ∆d H dissolving to form a solution the symbol used can be ∆sol H dilution of a solution the symbol used can be ∆dil H

Standard Enthalpy of Reaction (∆ r H o)
Since enthalpy changes differ depending upon pressure and temperature standard enthalpy of reaction values are for SATP (1 mol/L, 25 oC & 100 kPa) This means that the initial and the final conditions of the reaction system must be in a standard state.

Communicating Enthalpy Changes of Chemical Reactions
Enthalpy changes can be communicated by four methods: Stating the molar enthalpy of a reaction Thermochemical equations ΔH notation Potential energy diagrams

ENTHALPY CHANGES We will use ΔcH = - mcΔt
In a simple calorimetry experiment involving a burning candle and a can of water, the temperature of 100 mL of water increases from 16.4°C to 25.2°C when the candle is burned for several minutes. What is the enthalpy change of this combustion reaction? Assuming: ΔcH = Q (The energy lost by the chemical system, (burning candle), is equal to the energy gained by the surroundings (calorimeter water) Assuming: Q = mcΔt then ΔcH = mcΔt Is the value of ΔcH going to be positive or negative?? If the surroundings (water) gained energy, then the system (burning candle) lost it. So based on the evidence, the enthalpy change of combustion for this reaction is J We will use ΔcH = - mcΔt

Is this an Endothermic or Exothermic reaction??
ENTHALPY CHANGES When 50 mL of 1.0 mol/L hydrochloric acid is neutralized completely by 75 mL of 1.0 mol/L sodium hydroxide in a polystyrene cup calorimeter, the temperature of the total solution changes from 20.2°C to 25.6°C. Determine the enthalpy change that occurs in the chemical system. Based upon the evidence available, the enthalpy change for the neutralization of hydrochloric acid in this context is recorded as kJ. Is this an Endothermic or Exothermic reaction??

DO YOU REMEMBER?? Exothermic ENDOTHERMIC
A change in a chemical energy where energy/heat EXITS the chemical system Results in a decrease in chemical potential energy ΔH is negative A change in chemical energy where energy/heat ENTERS the chemical system Results in an increase in chemical potential energy ΔH is positive

Molar Enthalpy Molar enthalpy: ΔrHm the change in enthalpy expressed per mole of a substance undergoing a specified reaction (kJ/mol) Have we had other quantities expressed per mole? YES! How will we calculate this? the change in enthalpy per mole of a specified chemical undergoing change in a system is important b/c it allows chemists to do calculations using balanced chemical reactions enthalpy change and molar enthalpy use the same symbol (ΔH) so pay attention to the context and units of the question Enthalpy is measured in kJ while molar enthalpy is measured in kJ/mol IUPAC symbol ∆rHm both the substance and the reaction type must be specified Well known formations reaction & combustion enthalpies are tabulated (data book & p 347, 377) If a reaction is not well known or obvious then the chemical equation must be stated.

Molar Enthalpy= kJ/mol IUPAC Symbol = ΔrHm
Molar Enthalpy: change in enthalpy per mol of a specified chemical undergoing a change. Important because it lets us do calculations using chemical reactions Enthalpy change and Molar enthalpy use the same symbol ΔH, so pay attention to the context and units of the question. Enthalpy = kJ Molar Enthalpy= kJ/mol IUPAC Symbol = ΔrHm

ΔH Notation Describing the enthalpy change (ΔH) for a balanced equation To calculate an enthalpy change a molar enthalpy value and a balanced equation is required.

Example Standard molar enthalpy of formation of CuO (s) is
Hof (CuO) = kJ/mol (Data Booklet) Example: kJ is released to the surroundings during the formation of CuO (s) (exothermic) under SATP conditions

Molar enthalpy of a reaction
Note: if the balanced chemical reaction indicates that the specified compound has a coefficient other than 1 the molar enthalpy must be multiplied by the number of moles to obtain the total enthalpy of reaction. When determining a molar enthalpy change, it is critical to specify which substance a molar enthalpy is for.

Try This Formation of water

Try This **When hydrogen is combusted the product is H2O (g) .
Hydrogen has been called the fuel of the 21st century and may overcome fossil fuels in your lifetime. Calculate the enthalpy change for the combustion of hydrogen gas. **When hydrogen is combusted the product is H2O (g) . ∆fHo (H2O) = kJ/mol (Table 10.1 p 377)

Method 1 Molar Enthalpy Answer:
If 1 mol of water formed produces – kJ then the 2 mol produced in this reaction will produce 2 x kJ = kJ

Method 2 Enthalpy Change
Note: Standard Molar Enthalpies of Combustion & Formation Table H2(g) + ½ O2(g) –> H2O(l) ΔH = kJ Note: Liquid water’s value differs from gaseous water.

Molar Enthalpy Predict the change in enthalpy due to the combustion of 10.0 g of propane used in a camp stove. The molar enthalpy of combustion of propane is kJ/mol. 2. Predict the enthalpy change due to the combustion of 10.0 g of butane in a camp heater. The molar enthalpy of combustion of butane is kJ/mol.

Reaction for the combustion of propanone:
CH3COCH3 (l) + 4O2(g)  3CO2(g) + 3H2O(g) Molar Enthalpy of propanone combustion: HCH3COCH3(l) = kJ/mol What is the molar enthalpy of the combustion of propanone per mole of carbon dioxide formed?

Molar Enthalpy and Calorimetry
Can we measure the molar enthalpy of reaction using calorimetry? Yes, but indirectly. We can measure a change in temperature, we can then calculate the change in thermal energy (Q=mct). Then, using the law of conservation of energy we can infer the molar enthalpy. In doing so, we must assume that the change in enthalpy of the chemicals involved in a reaction is equal to the change in thermal energy of the surroundings. From this equation, any one of the five variables can be determined as an unknown.

Molar Enthalpy In a research laboratory, the combustion of 3.50 g of ethanol in a sophisticated calorimeter causes the temperature of 3.63 L of water to increase from 19.88°C to °C. Use this evidence to determine the molar enthalpy of combustion of ethanol. ** You don’t have to equate the two formulas to solve this. Instead, you can calculate Q, then use that value as ΔrH, and solve for either the chemical amount or the molar enthalpy of reaction. ΔcHm = 1.26 x 103 kJ/mol = 1.26 MJ/mol Q = 95.8 kJ = ΔH

Heat Capacity Q = CΔt

Thermochemical Equations Example
Combustion of propane

The formation of dinitrogen tetraoxide

Decomposition Reactions
The values in the data booklet apply for decomposition reactions if you reverse the enthalpy sign. Example: decomposition of water H2O(l) –> H2(g) + ½ O2(g) ΔH = kJ

If the equation 2NaOH + H2SO4  Na2SO4 + 2H2O kJ was rewritten such that the energy term had a value of kJ the respective balancing coefficients for the equation would be what?

If the equation 2Ag + I2  2AgI kJ was rewritten to produce 1 mol of AgI, the magnitude of the energy term would be

Try This If the combustion of 1.28 mol of benzene
(C6H6 (l)), produces an enthalpy change of kJ, what is the enthalpy change for this combustion as shown for a balanced chemical equation? 2 C6H6(l) + 15 O2(g) –> 12 CO2(g) + 6 H2O(g)

Step 1 Find the molar enthalpy of combustion of benzene

Step 2 Find the enthalpy change for 2 moles of benzene

Example How much energy would be released by the combustion of propane if 30.0g of carbon dioxide was formed?

Example How much energy is evolved from the combustion of 25.0g of ethane if ΔcH° = kJ/mol?

Example Commercial drain cleaner typically contain sodium hydroxide and aluminum. When the solid cleaner is poured down the drain water is added, a reaction occurs. The given equation represents this reaction If the given reaction produces 6.75 moles of NaAlO2, then the heat released is what?

Thermochemical equations
Chemical equations which show an energy change as part of the balanced chemical equation

Example: Combustion of benzene
2 C6H6(l) O2(g) –> 12 CO2(g) H2O(g) ΔcH = MJ As a thermochemical equation becomes: 2 C6H6(l) + 15 O2(g) –> 12 CO2(g) + 6 H2O(g) MJ Note: Endothermic energy goes on reactant side (ΔH = + ) Exothermic energy goes on product side (ΔH = - )

Communicating Enthalpy
We will be learning how to communicate enthalpy changes in four ways: By stating the molar enthalpy of a specific reactant in a reaction By stating the enthalpy change for a balanced reaction equation By including an energy value as a term in a balanced reaction equation By drawing a chemical potential energy diagram

COMMUNICATING ENTHALPY #1
By stating the molar enthalpy of a specific reactant in a reaction Why do we use standard conditions in chemistry (i.e. SATP)? We use a standard set of conditions so that scientists can create tables of precise, standard values and can compare other values easily Do we have standard conditions for enthalpy?? Yes, we will be using SATP (but liquid and solid compounds must only have the same initial and final temperature – most often 25°C)

How do we communicate that standard conditions are used for reactants and products?
With a ° superscript, such as ΔfHm° or ΔcHm° (See data booklet pg. 4 and 5) *For well-known reactions such as formation and combustion, no chemical equation is necessary, since they refer to specific reactions with the Δf or Δm ** Would the sign for ΔfHm° be the opposite of the sign for ΔdHm° (decomposition)? YES! *For equations that are not well known or obvious, then the chemical equation must be stated along with the molar enthalpy.

By stating the molar enthalpy of a specific reactant in a reaction Example #1: This means that the complete combustion of 1 mol of methanol releases kJ of energy according to the following balanced equation Example #2: This does not specify a reaction, so a chemical equation must be stated along with the molar enthalpy. This is not a formation reaction, since not all of the reactants are elements, so this could not have been communicated with Δf

By stating the enthalpy change beside a balanced reaction equation Do we know how to calculate enthalpy change?? The enthalpy change for a reaction can be determined by multiplying the chemical amount (from the coefficient in the equation) by the molar enthalpy of reaction (for a specific chemical)

Start with a balanced chemical equation.
Example: Sulfur dioxide and oxygen react to form sulfur trioxide. The standard molar enthalpy of combustion of sulfur dioxide, in this reaction, is kJ/mol. What is the enthalpy change for this reaction? Start with a balanced chemical equation.

2)Then determine the chemical amount of SO2 from the equation =
2 mol (this is an exact #, don’t use for sig digs) 3) Then use to determine the enthalpy change for the whole reaction. 4) Then report the enthalpy change by writing it next to the balanced equation.

By stating the enthalpy change beside a balanced reaction equation THE ENTHALPY CHANGE DEPENDS ON THE ACTUAL CHEMICAL AMOUNT OF REACTANTS AND PRODUCTS IN THE CHEMICAL REACTION. THEREFORE, IF THE BALANCED EQUATION IS WRITTEN DIFFERENTLY, THE ENTHALPY CHANGE SHOULD BE REPORTED DIFFERENTLY Both chemical reactions agree with the empirically determined molar enthalpy of combustion for sulfur dioxide

By stating the enthalpy change beside a balanced reaction equation THE ENTHALPY CHANGE DEPENDS ON THE ACTUAL CHEMICAL AMOUNT OF REACTANTS AND PRODUCTS IN THE CHEMICAL REACTION. THEREFORE, IF THE BALANCED EQUATION IS WRITTEN DIFFERENTLY, THE ENTHALPY CHANGE SHOULD BE REPORTED DIFFERENTLY Example 2: 2Al(s) + 3Cl2(g)  2AlCl3(s) ΔfH° = kJ What is the molar enthalpy of formation of aluminum chloride? ΔfHm° = kJ = kJ/mol AlCl3 2 mol

By stating the enthalpy change beside a balanced reaction equation EXAMPLE: The standard molar enthalpy of combustion of hydrogen sulfide is kJ/mol. Express this value as a standard enthalpy change for the following reaction equation: SOLUTION:

By including an energy value as a term in a balanced reaction equation If a reaction is endothermic, it requires additional energy to react, so is listed along with the reactants If a reaction is exothermic, energy is released as the reaction proceeds, and is listed along with the products In order to specify the initial and final conditions for measuring the enthalpy change of the reaction, the temperature and pressure may be specified at the end of the equation

By including an energy value as a term in a balanced reaction equation EXAMPLE: Ethane is cracked into ethene in world-scale quantities in Alberta. Communicate the enthalpy of reaction as a term in the equation representing the cracking reaction. Does the kJ mean exothermic or endothermic?

By including an energy value as a term in a balanced reaction equation EXAMPLE: Write the thermochemical equation for the formation of 2 moles of methanol from its elements if the molar enthalpy of formation is kJ/mol 2 C(s) H2(g) + O2(g)  2 CH3OH(l) ___?_____ ΔfH = 2 mol ( kJ/mol) = kJ (Exothermic) 2 C(s) H2(g) + O2(g)  2 CH3OH(l) kJ

By drawing a chemical potential energy diagram During a chemical reaction, observed energy changes are due to changes in chemical potential energy that occur during a reaction. This energy is a stored form of energy that is related to the relative positions of particles and the strengths of the bonds between them. As bonds break and re-form and the positions of atoms are altered, changes in potential energy occur. Evidence of a change in enthalpy of a chemical system is provided by a temperature change of the surroundings. A chemical potential energy diagram shows the potential energy of both the reactants and products of a chemical reaction. The difference is the enthalpy change (obtained from calorimetry) Guidelines: The vertical axis represents Ep. The reactants are written on the left, products on the right, and the horizontal axis is called the reaction coordinate or reaction progress. energy changes in chemical reactions are due to a change in potential energy as bonds break & reform the relative positions of atoms and the strengths of the bonds is altered Evidence of a change in enthalpy is provided by a change in temperature of the surroundings

Chemical Potential Energy Diagrams
Show the Ep of the reactants before the reaction and the products after the reaction As in chemical equations, the reactants are written on the left and the products go on the right

Exothermic Reactions Chemical potential energy diagrams can illustrate enthalpy changes Diagrams show the reactants storing a lot of energy then turning into products that cannot store the same amount of energy. The reactants are higher on the graph

Exothermic Reactions An exothermic reaction releases heat to its surroundings thus losing energy Temperature of surroundings increases

Endothermic Reactions
An endothermic reaction gains heat from its surroundings thus the temp. of the surroundings decreases Diagrams show using energy. The products use and store more energy than the reactants so are higher on the graph.

COMMUNICATING ENTHALPY
During an exothermic reaction, the enthalpy of the system decreases and heat flows into the surroundings. We observe a temperature increase in the surroundings. During an endothermic reaction, heat flows from the surroundings into the chemical system. We observe a temperature decrease in the surroundings.

Summary In an exothermic reaction the products have less Ep than the reactants & energy is released to the surroundings as the products form In an endothermic reaction the products have more Ep than the reactants & E is absorbed from the surroundings

EXAMPLE: Communicate the following enthalpies of reaction as a chemical potential energy diagram. The burning of magnesium to produce a very bright emergency flare. The decomposition of water by electrical energy from a solar cell.

Due at the start of class tomorrow
Your Task Practice Questions #9-13 page 493 Practice Questions #3,4,5 page 494 Due at the start of class tomorrow

Energy Diagrams Insert diagrams from page 21 KEY

Let’s Review Quickly… Endothermic Reaction
Activation Energy (Ea) - The energy level that the reactant molecules must overcome before a reaction can occur More to come later Endothermic Reaction In an endothermic reaction in which the change in enthalpy between the products and the reactants is positive, there must be an extra input of energy above the energy level of the products in order for a reaction to occur.

Exothermic Reaction Even in an exothermic reaction in which the change in enthalpy between the products and the reactants is negative, there must be an input of energy to start the reaction.

Enthalpy (H) - The sum of the internal energy of the system plus the product of the pressure of the gas in the system and its volume: Exothermic - Reaction in which a system RELEASES heat to its surroundings. H is negative ( H < 0) Ea is the activation energy

Endothermic - Reaction in which a system ABSORBS heat from its surroundings.
H is positive ( H > 0)

Review Example Write the equation representing the reaction for the formation for magnesium carbonate, and determine the standard molar enthalpy of formation for magnesium carbonate

Review Example Continued…
Determine the reaction enthalpy change for the equation representing the formation of magnesium carbonate balanced using whole number coefficients. Include the value of enthalpy change as an energy term in the equation.

Review Example Continued…
Draw the potential energy diagram that represents the formation of magnesium carbonate

Review Example 2 The standard molar enthalpy of neutralization for sodium hydroxide with any monoprotic strong acid is kJ/mol. What quantity of heat energy would be released if 750 mL of mol/L sodium hydroxide were neutralized with an appropriate quantity of strong acid?

Review Example 3 Heptane burns in a camping stove. What quantity of energy is released for every 1.00 kg of carbon dioxide produced?

Review Example 4 Heptane burns in a camping stove. What quantity of energy is released for every 1.00 kg of heptane is produced?

Predicting Enthalpy Changes
The enthalpy change for any reaction is the difference between the sum of the formation enthalpies of the products and the sum of the formation enthalpies of the reactants. The following equation is used for predicting enthalpy change of a reaction

Example Determine the enthalpy change for the complete combustion of 1 mol of propane gas

Example Determine the enthalpy change for the reaction of ethane with water

Example Determine the enthalpy change when 1 mol of ammonium nitrate decomposes to produce water vapour and dinitrogen monoxide during an explosion

Example Butane is added to gasoline to make it perform better under rigorous conditions of the Canadian winter. Approximately how much energy is produced when one mole of butane burns in an automobile engine to give gaseous products?

For Your Viewing Pleasure…

HESS’ LAW Do you think it is convenient or possible to use a calorimeter to test all chemical reactions? NO! Sometimes two products are created simultaneously, sometimes a reaction is too small to be able to measure accurately. So what do scientists do? Theoretically, we assume that the enthalpy change of a physical or chemical process depends only on the initial and final conditions. It is independent of the pathway, process or number of intermediate steps required. Illustration: Bricks are being moved from the ground up to the second floor. But there are two pathways to do this: Move from the 1st to 2nd floor Move to third floor and then carry down one flight In both cases, the overall change in position is the same. Hess’ Law involves the idea that a reaction could be shown as a series of steps rather than one step. The law can be done the slow way with reaction manipulation or the fast way with the formula.

Germain Henri Hess ( ) ● Is important primarily for his thermochemical studies Hess's Law states that the heat transferred, or change in enthalpy (H), in a reaction is the same regardless whether the reaction occurs in a single step or in several steps. The method of calculating the enthalpy of reaction developed by Hess is called Hess's Law of Heat Summation. If a series of reactions are added together, the net change in the heat of the reaction is the sum of the enthalpy changes for each step.

Hess’ Law of Constant Heat Summation aka Law of Additivity of Enthalpies of Reaction
G. W. Hess suggested in 1840 that: - The change in enthalpy for any reaction depends only on the nature of the reactants and products and is independent of the number of steps or the pathway taken between them. Hess’s Law can be written as an equation: The uppercase Greek Letter, Σ (sigma) means “the sum of” The enthalpy change of a net chemical reaction is the arithmetic sum of the enthalpies of the individual reactions combined to make the net reaction. In mathematical terms, this can be expressed as follows: For example, add together the following four equations (get from page 26) Anything that is both a product and a reactant cancels out when the sums of the coefficients are equal. For example, adding the following reactions: Also, anything that appears on the left-hand side of the reactions can be cancelled out or reduced, with like chemicals on the right hand side. For examples, adding the following reactions: Equations must often be manipulated before they are added together to obtain the desired net equation. When multiplying an equation by a certain factor, each coefficient must be multiplied by that factor. The reaction enthalpy must also be multiplied by the same number. If the reaction is reversed multiply the reaction’s enthalpy by -1. When constructing the net reaction, first take care of the compounds in the final reaction and the rest will usually fall into place.

Hess’ Law Hess’ discovery allows us to determine enthalpy change without direct calorimetry, using two rules that you already know: If a chemical equation is reversed, then the sign of ΔrH changes 2) If the coefficients of a chemical equation are altered by multiplying or dividing by a constant factor, then the ΔrH is altered by the same factor Rules for using Hess's LawIf the reaction is multiplied (or divided) by some factor, H must also be multiplied (or divided) by that same factor. If the reaction is reversed (flipped), the sign of H must also be reversed.

Hess’ Law Hess's Law The enthalpy change for any reaction depends only on the energy states of the initial reactants and final products and is independent of the pathway or the number of steps between the reactant and product.

Predicting Enthalpy Changes
Whether a product is formed from a one step reaction or a series of reactions the enthalpy change will be the same provided the initial reactants and final products are the same and the same initial and final conditions are present.

Using Hess’s Law Write a balanced net chemical reaction Manipulate given formation equations to yield the net equation (x, ÷, and/or reverse ∆rH) Cancel and add equations to yield the net equations Add component enthalpy changes to obtain the net enthalpy change Determine molar enthalpy if required

Hess’s Law Example - Combustion of Methane
ΔrH = kJ CH4(g) + 2 O2(g) CO2 (g) + 2 H2O (g) Reactants Products Alternate Path ΔHr = kJ CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) Hess said sum these two energies ΔHformation = kJ + ( kJ) ΔHdecomp = kJ Reactants break into elements C(s) + 2 H2(g) + 2 O2(g) Products form from these elements

Hess’ Law #1 Example: Use Hess’ Law to determine the enthalpy change for the formation of carbon monoxide. This reaction can not be studied calorimetrically but we are given the following information to help solve this equation Our job now, is to manipulate the equations so they will add to yield the net equation We need 1 mol of C(s) to start the equation, so leave (1) unaltered However, we want 1 mol of CO as a product, so reverse equation (2) and divide all terms by 2 ** Remember whatever you do to the equation, affects the ΔH the same way ΔcH = kJ (original equation) 1) Reversed equation; ΔH = kJ 2) Divide equation by 2; Divide ΔH by 2 = kJ

Hess’ Law #1 Example: Use Hess’ Law to determine the enthalpy change for the formation of carbon monoxide. Now cancel and add the remaining reactants and products to yield the net equation. Add the component enthalpy changes to obtain the net enthalpy change. The process of using Hess’ Law is a combination of being systematic and using trial and error. Do what needs to be done to the given equations so they add to get the equation you want.

Hess’ Law #1 Example: Use Hess’ Law to determine the enthalpy change for the formation of carbon monoxide. Sketching a potential energy diagram might help you ensure that you have made the appropriate additions and subtractions

Example Calculations 1) Nitrogen and oxygen gas combine to form nitrogen dioxide according to the following reaction: Calculate the change in enthalpy for the above overall reaction, given: This problem is very simple. If we simply add up the two reactions keeping all the reactants on the left and all the products on the right, we end up with the overall equation that we are given. Since we didn't make any changes to the individual reactions, we don't make any changes toH. If we add upH as well, we find the change in enthalpy:

From the following enthalpy changes:
Calculate the value of H for the reaction:

Example The formation of liquid hexane is represented by the following overall net equation: 6C(s) + 7H2(g)  C6H14(l) C6H14(l) /2 O2(g)  6CO2(g) + 7H2O(l) C(s) + O2(g)  CO2(g) H2(g) + ½ O2(g)  H2O(l) Using Hess’ Law, determine the molar enthalpy of the formation for hexane

Hess’ Law Example Use the following equations to answer the next question Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(aq) kJ MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(g) kJ H2(g) + ½ O2(g)  H2O(g) kJ What is the experimental determination of the equation, Mg(s) + ½ O2(g)  MgO(s)

Hess’s Law Example - Combustion of Methane
Potential Energy Diagram

Hess’ Law #2 Example: One of the methods the steel industry uses to obtain metallic iron is to react iron(III) oxide with carbon monoxide Fe2O3(s) + 3CO(g)  3CO2(g) + 2Fe(s) ΔrH = ?? 1) CO(g) + ½ O2(g)  CO2(g) ΔfH = kJ 2) 2Fe(s) + 3/2O2(g)  Fe2O3(s) ΔfH = kJ 3( CO(g) + ½ O2(g)  CO2(g)) ΔfH =3( kJ) = kJ reverse Fe2O3(s)  2Fe(s) + 3/2O2(g) ΔfH = kJ = kJ Fe2O3(s) + 3CO(g)  3CO2(g) + 2Fe(s) ΔrH = kJ

Hess’ Law #3 Example: What is the standard enthalpy of formation of butane? ΔfHm° = ??? First, we need to be able to write this balanced formation equation. 4C(s) + 5H2(g)  C4H10(g) The following values were determined by calorimetry: What will we need to do to get our net equation? ΔfHm° = kJ/1 mol = kJ/mol C4H10 Reverse equation (1) and change the ΔH sign Multiply equation (2) and its ΔH by 4 Multiply equation (3) and its ΔH by 5/2

Synthesis of Sulfur Trioxide
S8 (s) + 8 O2 (g) -> 8 SO2 (g) 2 SO2 (g) + O2 (g) -> 2 SO3 (g) ∆Ho – from formn tables ⅛ S8 (s) + O2 (g) -> SO2 (g) ∆Ho = SO2 (g) + ½ O2 (g) -> SO3 (g) ∆Ho = To obtain the overall eqn: ⅛ S8 (s) + 3/2 O2 (g) -> SO3 (g) ∆Ho = kJ/mol

Hess’ Law If you can add equations then you can add ΔH’s ΔHnet = Σ ΔrH
Note: Σ - means “sum of”

Recall these rules of enthalpy change:
If a chemical equation is reversed then the sign for ΔHr changes If the coefficients are changed by a constant factor then ΔHr changes in the same way

Why does Hess’s Law work?
There are only a limited number or formation reactions (one for each compound) But there are a seemingly infinite number of pathways for each reaction…. Therefore, if we always go through the elements pathway, it will limit the amount of information required to calculate the theoretical energy of reaction

Enthalpy change for a reaction can be obtained by:
ΔrHo ↓ ↓ ↓ Calorimetry Hess’s Law (Experimental value) (Predicted value) nH = mcΔt ΔrHo = Σn∆fpHo - Σn∆frHo - manipulating intermediate steps

Hess’ Law Example 20kJ of heat is involved when 4.2 mol of Y is used in the exothermic reaction above. Which of the following reactions is correct? 2X(s) + 3Y(g)  X2Y3(g) H = +20 kJ 2X(s) + 3Y(g)  X2Y3(g) H = -20 kJ 2X(s) + 3Y(g) kJ  X2Y3(g) 2X(s) + 3Y(g)  X2Y3(g) + 14kJ

Example Determine the enthalpy change for the hydrogenation reaction of benzene to make cyclohexane.

Molar Enthalpy of Formation
Molar enthalpies of formation are defined as the enthalpy change when one mole of a compound forms from its elements NOTE: The enthalpy of formation for an element is 0 kJ Examples: Using your data booklet, find the following: Δf Hm ° CH4(g) = kJ/mol Δf Hm ° O2(g) = 0 kJ/mol (Δf H elements = 0) Δf Hm ° CO2(g) = kJ/mol Δf Hm ° H2O(g) = kJ/mol

Reference energy state
A convention which describes elements as the reference point at which potential energy is zero Allows chemists to compare enthalpy changes to the reference state of zero and tabulate these values

Thermal stability Tendency of a compound to resist decomposition when heated. The standard molar enthalpies of formation give an indication of thermal stability. The lower the value, the greater the thermal stanility

Thermal Stability Example
Example: List the following compounds in decreasing order of thermal stability; A. carbon dioxide, B. ethyne, C. magnesium chloride, D. hydrogen bromide Answer (C, A, D, B )

Standard Molar Enthalpies of Formation
Hess’s Law can be stated mathematically as: ΔrHo = Σ(n∆fHo products) - Σ(n∆fHo reactants) [Enthalpy change = the sum of the molar enthalpy of formation of products minus the sum of the molar enthalpy of formation of reactants] [ Note: When using this equation do NOT change any signs of values, the equation accounts for whether the enthalpy values are for products or reactants. ]

Predicting the Enthalpy Change of a Reaction Component
The previous equation can be rearranged to calculate the total formation enthalpy of the reactants (or of the products) as follows:

When a compound reacts with an element, the sum of the reactant formation enthalpies is a simple multiple of that compound’s formation enthalpy. Example: Methylhydrazine burns in a bomb calorimeter according to the following balanced equation: Determine the molar enthalpy of formation of methylhydrazine.

The standard molar enthalpy of combustion of cyclobutane to gaseous products is kJ/mol. Determine the standard molar enthalpy of formation of cyclobutane.

When 1.00g of diphemylmethanone, burns in a closed system to produce carbon dioxide and water, kJ of heat energy is released. Determine the molar enthalpy of formation of diphenylmethanone.

Why do we care about the standard molar enthalpies of formation, ΔfH° ??? Because we are going to use them to predict standard enthalpy changes for chemical reactions. How? Using this crazy formula!! What does it mean? The net enthalpy change for a chemical reaction, ΔrH°, is equal to the sum of the chemical amounts times the molar enthalpies of formation of the products, ΣnΔf pHm °, minus the chemical amounts times the molar enthalpies of formation of the reactants, ΣnΔf RHm ° Clear as mud?? Basically, the equation says that the change in enthalpy is the total chemical potential energy of the products minus the reactants. Epproducts – Epreactants We will need to use an example to figure this out.

Calculate the molar enthalpy of formation for two moles of carbon monoxide from its elements. 2C(s) + O2(g)  2CO(g) ΔfHm = 2 mol( kJ) mol(0 kJ) + 1 mol(0 kJ) mol mol mol = -221.o kJ 2 mol = kJ/mol

Methane is burned in furnaces and in some power plants. What is the standard molar enthalpy of combustion of methane? Assume that water vapor is a product. Need a balanced chemical equation: CH4(g) + O2(g)  CO2(g) + 2H2O(g) Use the formula and the data booklet to calculate the ΔcH° We found all of the Δf Hm for the compounds Are we finished with kJ?? NO!

Methane is burned in furnaces and in some power plants. What is the standard molar enthalpy of combustion of methane? Assume that water vapour is a product. This can also be communicated as an enthalpy change diagram. Note that the labeling of the y-axis is different from that in a chemical potential energy diagram. Epproducts – Epreactants

Practice How can molar enthalpies of formation be used to calculate enthalpies of a reaction? Consider the slacking of lime, calcium oxide, represented by the following chemical reaction equation

Practice

Your Task Questions #3abc, 4ab, 5ab,6ab,8abcd on page 514-15
Questions #24,25 page 520

Enthalpy Changes Three types of Change:
Phase - intermolecular bonds break/form - 100 – 102 kJ/mol Chemical - intramolecular bonds break/form - 102 – 105 kJ/mol Nuclear - nuclear binding energy break/form - > 1010 kJ/mol

Enthalpy Changes Compared

Phase Changes When a substance changes phase there is no change in temperature. Thus there is no change in the average Ek of the substance Since heat (energy) has been added the Ep must have increased Example:

Chemical Changes Chemical changes involve the rearranging of atoms which changes the total bond energy of the system If the temperature is held constant only Ep will change Example:

Nuclear Change Nuclear reactions involve the rearrangement of nucleons resulting in a change in Ep Example:

% Efficiency = E output x 100% E input
Energy Efficiency % Efficiency = E output x 100% E input Input – theoretical (Hess calculation) Output – actual (calorimetry)

Energy Efficiency

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Reaction Rate Some reactions occur very quickly while others occur more slowly. Example: combustion of butane vs formation of rust

Kinetics = Rates of Reaction
Collision-Reaction Theory A chemical sample consists of entities (ions, atoms, molecules) that are in constant, random motion at various speeds, rebounding elastically from collisions with each other (kinetic energy is conserved during elastic collisions) For a reaction to proceed, reactants must collide An effective collision requires sufficient energy to react and the correct orientation, so that bonds can be broken and new bonds formed The more collisions there are, the greater the potential for effective collision. based on the idea that reactant particles must collide for a reaction to occur Recall CRT: Reactant s must collide Proper orientation –> collision geometry Minimum energy –> activation energy Collision Geometry: Reacting particle must collide in the proper orientation Ineffective collisions involve entities that rebound elastically from the collision

Collision-Reaction Theory Ineffective collisions involve entities that rebound and do not rearrange and form new substances.

Factors affecting Reaction Rate: Concentration: more reactant particles in a given volume increases the number of collisions per second Surface Area: more opportunity for collisions, the more collisions there will be Temperature: the faster the particles are moving, the more energy they have to create an effective collision

Bond Energy and Enthalpy Changes
Bond energy is the energy required to break a chemical bond; it is also the energy released when a bond is formed. - bonded particles + energy  separated particles - separated particles  bonded particles + energy The change in enthalpy represents the net effect from breaking and making bonds. ΔrH = energy released from bond making – energy required for bond breaking Exothermic reaction: making > breaking (ΔrH is negative) Endothermic reaction: breaking > making (ΔrH is positive) In an energy diagram, all parts of the graph are labelled, including the reactants and products, each axis, and the step up or step down in terms of enthalpy change

In general, the following rules regarding the energy of reactions apply:
If the total energy input is greater than the total energy output, ΔrH° is positive, and the reaction is endothermic (the reaction absorbs energy from the surrounding) If the total energy output is greater that the total energy input, ΔrH° is negative, and the reaction is exothermic (the reaction releases energy to the surroundings)

Example Consider the formation of hydrogen iodide gas from its gaseous elements This reaction is exothermic, so the products posses more total potential energy than the reactant. This reaction absorbs energy from its surroundings. The energy absorbed to break H-H bonds and the I-I bonds in the reactants (Ein) is greater than the energy released to form the H-I bond in the product (Eout)

Potential Energy Diagram for the Formation of Hydrogen Iodide

LET’S SEE IF YOU GET IT Draw energy pathway diagrams for general endothermic and a general exothermic reaction. Label the reactants, products, enthalpy change, activation energy, and activated complex.

ACTIVATION ENERGY OF A REACTION
Activation Energy – (EA) The minimum collision energy required for effective collision Dependent on the kinetic energy of the particles (depend on T) Analogy: If the ball does not have enough kinetic energy to make it over the hill – the trip will not happen. Same idea, if molecules collide without enough energy to rearrange their bonds, the reaction will not occur. (ineffective collision) Reactants must collide with sufficient energy to break bonds in reactants

The rate at which a reaction proceeds is determined in part by the reaction’s activation energy. The higher a reaction’s value of Ea, the slower the reaction proceeds.

The activated complex occurs at the at the maximum potential energy point in the change along the energy pathway. Is this an exothermic or endothermic change? Exothermic. This means the initial energy absorbed to break the nitrogen-oxygen bond is less than the energy released when a new carbon-oxygen bond forms. The energy pathway is the relative potential energy of the chemicals system as it moves from reactants through activated complex to products. The activated complex is the chemical entity containing the collided reactants. The flat line region to the left, the molecules are moving towards each other but are still distant from each other AS the molecules approach more closely, they are affected by repulsion forces and begin to slow down. Some of their kinetic energy is changed to chemical potential energy and is stored as a repelling electric field between them If the molecules have enough kinetic energy- which means more energy than is required to get to the energy level of the activated complex- get close enough for the bond structure to rearrange The activated complex occurs at the maximum potential energy point in the change along the energy pathway

If a large quantity of energy is needed to start a reaction and if the reaction progresses relatively slowly, then the activation energy is large A spontaneous reaction at room temperature and a higher rate of reaction is interpreted as a relatively small activation energy In general, the greater the EA, the slower the reaction. It takes longer for more particles to achieve kinetic energy necessary for effective collision.

What does this diagram indicate? At Temperature 2, a greater number of particles will have the activation energy required Will this increase the rate of the reaction? Yes Never assume that a reaction with a large enthalpy value is likely to be faster than an endothermic or less exothermic reaction. The rate of a reaction is mostly determined by the activation energy. Another misconception is that an endothermic reaction proceeds when energy is added. An endothermic reaction absorbs energy as it proceeds. Adding or removing any amount of energy to or from a reaction can never change it from an endothermic to an exothermic reaction or vice versa.

Is this an exothermic or endothermic change? Endothermic. A continuous input of energy, usually heat, would be needed to keep the reaction going, and the enthalpy change would be positive. Exothermic: the potential energy of the products is less than that of the reactants Endothermic: the potential energy of the products is greater than that of the reactants

Energy Exchanges in Chemical Reactions
Chemical reactions can produce energy, in exothermic reactions, or absorb energy, in endothermic reactions The energy released or absorbed is linked with the formation or breaking of bonds. In fact, all chemical reactions involve bond breaking (endothermic) and bond making (exothermic) steps.

Bond Energy and Reactions

Bond Energy and Enthalpy Changes
Bond energy is the energy required to break a chemical bond; it is also the energy released when a bond is formed. The change is enthalpy represents the net effect from breaking and making bonds Exothermic reactions: Endothermic reactions: Page 533

Your Task Question # 1-6, 8abc

CATALYSTS AND REACTION RATE
A catalyst is a substance that increases the rate of a chemical reaction without being consumed itself in the overall process. A catalyst reduces the quantity of energy required to start the reaction, and results in a catalyzed reaction producing a greater yield in the same period of time than an uncatalyzed reaction. It does not alter the net enthalpy change for a chemical reaction Catalysts lower the activation energy, so a larger portion of particles have the necessary energy to react = greater yield A substance that increases a reaction rate w/o being consumed itself in the process Reduce the amount of energy required to start a reaction resulting in a greater yield in the same time How a catalyst works: It is thought that catalysts reduce the activation energy required for a reaction by providing an alternate pathway for the reaction What effect does a catalyst have on the enthalpy change of a reaction? Hess’ Law suggests that the enthalpy of a reaction depends only on the enthalpies of the initial and final states, no matter what series of step exist between the products and reactants. A catalyst has no effect on the enthalpy change of a reaction.

How do catalysts work?? Scientists do not really understand the actual mechanism. Catalysts are also usually discovered through trial and error. What they do know is that they provide an alternative, lower energy pathway from reactants to products. Most of the catalysts (enzymes) for biological reactions work by shape and orientation. They fit substrate proteins into locations on the enzyme as a key fits into a lock, enabling only specific molecules to link or detach on the enzyme. Almost all enzymes catalyze only one specific reaction Enzymes: Biological catalysts Very complex molecules (usually proteins) Catalysts speed up a reaction by providing an alternate reaction pathway that has lower activation energy than the uncatalyzed reaction.

Reaction Mechanisms Steps making up the overall reaction Each step = elementary reaction Reaction intermediates: substances formed in one elementary reaction and consumed in another The rate-determining step of a reaction is the step with the highest activation energy. It is called the rate-determining step because it is the slowest step. Reaction mechanism describes the individual reaction steps and the intermediates formed during the reaction, starting with the reactants and finishing with the products. Intermediates are chemical entities that form with varying stability at the end of a step in a reaction mechanism. They are more stable (have lower potential energy) than the activated complexes, but less stable than the reactants and the products. The intermediate then reacts in a subsequent step and does not appear in the final reaction mixture. For example: Step 1: reactant(s)  activated complex 1 intermediate 1 Step 2: intermediate 1  activated complex 2  intermediate 2 Step 3: intermidate 2  activated complex 3  product(s)

Potential Energy Diagram for a Catalyzed and Uncatalyzed Reaction

Practice Molecules A-A and B-B combine to form 2 A-B. The bond in A-B is stronger than the bond in A-A or B-B. Draw an energy diagram to represent the given reaction.

Uses of Catalysts The oil industry uses catalysts in the cracking and reforming of crude oil and bitumen to produce consumer products (gasoline) Catalysts allow oil companies to increase the reaction rate while decreasing the energy (often decreasing temperature) required.

Your Task These questions will be due at the start of class on Monday, NO EXCEPTIONS. You will have today’s class and whatever is left after we go through your exam on Friday to work through them. Practice Questions #1-3ab page 526 Questions # 4,5abcd,7 page 531 Questions #2, 3, 5ab, 7ab,8abc,9 page 534 Read and take notes on page 540 Questions #2abcd, 3abcde,4,5 page 542 Question 16a-h page 546 Question 10 page 547 Question 13,14,16, 21abc, 22,23 page 548/49 Question 31 page 550

The Technology of Energy Measurement

“Calorimetry is the technological process of measuring energy changes of an isolated system called a calorimeter” Includes: Thermometer, stirring rod, stopper or inverted cup, two Styrofoam cups nested together containing reactants in solution

Assumptions of Simple Calorimeter
A polystyrene foam cup is assumed to be perfect insulation  small amounts of energy transfer is ignored Only the water releases, or absorbs, heat energy to or from the reaction. It is assumed that no heat is gained or lost to the surroundings unless otherwise indicated. The specific heat capacity of solution= the specific heat capacity of liquid water VSolution= VH2O= mH2O(l). When a solution’s concentration is less than 2mol/L, water constitutes the majority of the solution. Liquid water has a density very close to 1.00g/mL. This is assumed to be the density of aqueous solutions if no extra information is given.

Equations The given assumptions allow you to calculate the kinetic energy change of the calorimeter Qcal= (𝑚𝑐∆𝑇) A combination of the first and second laws of Thermodynamics, where energy is transferred rather than created or destroyed, allows for the assumption that the kinetic energy change of the calorimeter is equal to the enthalpy change of the calorimeter is equal to the enthalpy change of the reaction ∆𝐻, given the following equation: ∆𝐻=𝑛∆𝐻°

Combining ∆𝐻=𝑛∆𝐻° with the idea that energy gained by the water in a calorimeter equals the energy lost by the reaction, the following equation may be written as: ∆𝐻=−𝑄= 𝑚𝑐∆𝑇 ⇒𝑛∆𝐻°=−(𝑚𝑐∆𝑇)

Example When 50.0 mL of 1.00 mol/L NaOH, with an initial temperature of 17.4 degrees Celsius is mixed with 50.0mL of 1.00 mol/L HCl, with an initial temperature of degrees Celsius, the final reaction mixture reaches a temperature of degrees Celsius. What is the molar enthalpy change for the neutralization of the base?

Example When excess zinc is added to 50.0mL of mol/L aqueous copper (II) sulfate, the calorimeter warms by 12.5 degrees Celsius. What is the molar enthalpy of reduction of aqueous copper(II) ions?

Energy - Recall Science 10:
A heating curve of water shows the types of heat energy that exist. Ek – related to motion and directly related to temp. Ep – related to bonds and bonding forces

Kinetic heat energy is very easy to calculate
Energy Measurement Kinetic heat energy is very easy to calculate Ek = mc ∆ t Potential energy, can not be measured by may be deduced! There is no way that we can measure the energy stored between two bonds.

BUT - If a potential energy change occurs in an insulated container...
AND – we consider the Laws of Thermodynamics… First Law of Thermodynamics: Heat gained (q) = Heat lost (q) Second Law of Thermodynamics: heat flows from a warmer body to a cooler body We can use a kinetic energy change to determine the energy stored in bonds… indirectly!

Calorimetry Information about potential energy changes comes from calorimetry experiments We will consider two types of calorimetry measurements: Solution calorimetry using a simple calorimeter Bomb calorimetry

Limitation of Calorimetry
Not every reaction of interest to chemists and engineers can be studied by calorimetry so a method of predicting enthalpies of reaction is an important part of thermochemistry

Thermochemical Changes

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