The first law of thermodynamics is applied to a system that undergoes a chemical change.
There is no work done in Chapter 7 of thermochemistry.
Bomb calorimeters do not usually carry chemical reactions at constant volume.
As the system expands or contracts, a small amount of pressure-volume work is done.
The change in the quantity is on the left side of the expression.
For a constant-pressure process, such as a reaction occurring in a container open to the atmosphere, the heat transferred is equal to the enthalpy change, C/H, of the system.
The first law for a constant-pressure process is Equation (7.14).
There is a relationship between C/H and C/U.
For a constant pressure process, we can write q, qp, and w.
The work associated with the energy change at constant change in volume of the system under a constant external pressure is the last term in this expression.
To see how much pressure there is.
A 40 kJ quantity of heat is added and the system does 15 kJ of work, then the system is returned to its original state by cooling and compression.
The system shrinks into a smaller volume when 5 0 work on it.
If the heat transferred in this process is measured under constant-pressure conditions at a constant temperature, we get -566.0 kJ, indicating that 566.0 kJ of energy has left the system as heat.
This alternative expression can be written using the ideal gas equation.
The number of moles of gas in 12 mol CO22 and ni,gas is the number of moles of gas in the reactants.
The calculation shows that the PC/V term is small compared to C/U and C. As a result of the work done on the system by the surroundings, the volume of the system decreases.
The change in the number of moles of gas is called C/ngas.
qP - qV is about 2.5 kJ in a process for which C/ngas is +2 mol.
We have considered internal energy and enthalpy changes for a system in which specified amounts of reactants are converted into specified amounts of products.
Consider the enthalpy of reaction for the fire.
C/rH is the enthalpy of the reaction for the burning of C H O.
To calculate the quantity of heat produced, the first step is to determine the number of moles in 1.00 kg of sucrose, and then use that value and the C/rH value for the reaction to calculate the quantity of heat produced.
Use the C/rH value to convert frommol C12H22O11 to kJ of heat.
The negative sign indicates that the heat is dissipated.
The heat produced by a combustion reaction is not immediately transferred to the surroundings.
A sample of 0.1045 M HCl(aq) was neutralized by NaOH.
The heat evolved in this neutralization.
In the preceding discussion, we used the first law of thermodynamics, C/U, to show that C/H was the same as qP.
C/H is a form of the first law that is convenient for constantpressure processes.
The answer is simple: C/H is the heat transferred under constant-pressure conditions.
Two different symbols, C/ H and qP, are used to represent the same thing.
The answer to this question is disappointing.
There is no simple physical meaning or interpretation of the combination.
For convenience, this combination of quantities has been introduced.
Because most processes are carried out at constant pressure, we most often measure the heat transferred as qP, a quantity that may also be represented as C/H.
When a liquid is in contact with the atmosphere, energetic molecules at the surface of the liquid can overcome forces of attraction to their neighbors and enter the gaseous state.
The liquid needs to absorb heat from its surroundings to replace the energy carried off by the vaporizing molecule.
The enthalpy of vaporization is the amount of heat required to evaporate a fixed quantity of liquid.
We described the melting of a solid.
The enthalpy of fusion is called the energy and Applied Chemistry requirement.
H2O1l2 fusH is 6.01 kJ mol, which is 1 at 273.15 K on the C/ symbol.
For the process in which 50.0 g of water is converted from liquid to vapor at 10.0 degC, you can calculate C/H by taking the Enthalpy Changes Accompanying Changes in States of Matter.
The process begins with raising the temperature of liquid water from 10.0 to 25.0 degC, and then completely vaporizing the liquid at 25.0 degC.
The sum of the changes in the two steps is the total enthalpy change.
The quantity of water must be expressed in moles so that we can use the molar enthalpy of vaporization.
The system gains energy when the enthalpy change is positive.
The reverse would be true for condensation of water at 28.0 degC and cooling it to 10.0 degC.
The enthalpy of fusion of 6.01 kJ>mol is used for ice.
The ideal gas at a pressure of 1 bar and a temperature of interest is the pure gas.
Although temperature is not part of the definition of a standard state, it still has to be specified in the values of C/rHdeg.
Unless otherwise stated, the values given in this text are for 298.15 K 125 degC2.
The standard-state pressure should be changed from 1 atm to 1 bar more than 30 years ago, but some data tables are still based on the 1 atm standard.
The differences in values resulting from the change in standard-state pressure are small enough to be ignored.
We will mostly use standard enthalpy changes in the rest of the chapter.
Chapter 13 contains the details of nonstandard conditions.
D products are lower than the reactants.
An exothermic reaction is the burning of sugar.
The heat is absorbed from the surroundings.
The absolute values of enthalpy are represented by horizontal lines.
The higher the horizontal line, the more H it represents.
The amount of heat involved in changing reactants and products from one temperature to another is what determines the difference in pointing up.
Calcu reactions are caused by these quantities of heat.
We write an expression of this type for each reactant and product with measured reactions.
To get the value of C/rHdeg at another temperature, you have to use one temperature.
The enthalpy concept is useful because it can be calculated from a small number of measurements.
Consider the stan dard enthalpy change in the formation of NO1g2 from its elements.
2 must be one.
If a process is reversed, the function of state reverses sign will change.
The enthalpy change for the overall process is the sum of the individual steps' enthalpy changes.
The state function property of enthalpy is what leads to Hess's law.
C/rH is the same regardless of the path taken from the initial state to the final state.
A slight reaction will occur if we try to get hydrogen and graphite to react.
Several other hydrocarbons will form as well, and the product will not be limited to propane 1C3H82.
We can't directly measure C/rHdeg for reaction.
The greatest value of the law is here.
We can calculate enthalpy changes that we can't measure directly.
The values can be used to calculate C/rHdeg.
Combine the appropriate chemical equations to determine an enthalpy change.
The products of the burning of carbon-hydrogen-oxygen compounds are CO21g2 and H2O1l2.
C(graphite) and H21g2 are reactants.
3 mol C(graphite) and 4 mol H21g2 have been consumed, and 1 mol C3H 81g2 has been produced.
This is what is required in the equation.
The modified equations can now be combined.
Assess Hess's law can be used to determine the enthalpy of reaction by using a series of unrelated reactions.
We were able to determine the enthalpy of the reaction of another reaction by taking three unrelated combustion reactions.
C3H61g2 has a standard heat of -2058 kJ>mol.
The data from this example can be used to determine C/rHdeg for the hydrogenation of propene to propane.
Determine the standard enthalpy of combustion of one mole of CH3 CH1 OH2 CH31l2 from the data in Practice example 7-9A and the following equation.
The heat of reaction between carbon and hydrogen is 226.7, 52.3, and 84.7 kJ mol-1, respectively.
The enthalpy diagram is shown in the margin.
We did not write numerical values on the enthalpy axis in the diagrams we drew.
These changes can be dealt with.
It is still useful to have a starting point of zero.
We mean the vertical distance between the mountaintop and the sea level.
Everest is +8848 m, and Badwater is 86 m. The zero is related to the enthalpies of elements and other substances relative to this zero.
The most stable forms of elements at one bar and the given temperature are the reference forms.
The degree symbol shows that the enthalpy change is a standard enthalpy change, and the subscript f shows that the reaction is one in which a substance is formed from its elements.
The expression formation of the most stable form of an element is the same as before.
The standard enthalpy of formation is 0 for a pure element.
The most stable forms of several elements are listed here.
There is an interesting situation with carbon.
Carbon can be found in the form of diamond.
There is a measurable enthalpy difference between them.
We assign C/fHdeg1graphite2 and C/ Diamond and Graphite to you.
The most stable form of bromine is Br21l2.
If obtained at 298.15 K and 1 bar pressure, the condenses to Br21g2.
The enthalpies of formation are 30.91 kJ>mol.
The reference form is not usually the most stable form.
Solid white phosphorus has been chosen as the reference form despite it being converted to red phosphorus over time.
The enthalpies of formation are C/fHdeg3P1s, white 24 and red 24.
Table 7.2 contains the standard enthalpies of formation of some common substances.
It suggests that the standard enthalpies of formation are related to the structure.
The standard enthalpies of formation will be used.
There are values for reactions in which one mole of substance is formed.
The data has been rounded off to four significant figures.
The enthalpy of formation of formaldehyde is 108.6 kJ>mol.
Write the equation to which it applies.
We need a fractional coefficient in this equation.
We must remember to use the HCHO(g) ments in their most stable form when answering these types of problems.
In this example, the conditions were stated to be 298.15 K and 1 bar.
The enthalpy for the formation of leucine, C formaldehyde, HCHO(g) 6H13O2N1s2 is 637.3 kJ>mol.
Write the equation to which it applies.
The exothermic reaction is what it is.
A compound with a positive value of C/fHdeg is formed from elements.
If the reaction is reversed, the compound breaks down into its elements.
We sometimes say that the compound is not stable.
This doesn't mean that the compound can't be made, but it does suggest a tendency for the compound to enter into chemical reactions yielding products with lower enthalpies of formation.
When no other criteria are available, chemists sometimes use enthalpy change as a rough indicator of the likelihood of a chemical reaction occurring.
Later in the text, we'll present better criteria.
Standard enthalpies of reaction are one of the primary uses of forma tion.
When baking soda is used in baking, the standard enthalpy of reaction for the decomposition of sodium bicarbonate is calculated.
The following equations yield an equation when added together.
Imagine a place where the decomposition of sodium bicarbonate takes place.
The Na2CO3(s) 2 mol Na(s), 2 mol C(graphite), 1 mol H21g2, and 3 mol O21g2 should be combined in the second step.
enthalpy is a state function and the overall change of any state function is independent of the path chosen.
The standard enthalpy changes of 2 NaHCO3 are the individual steps.
Na2CO3(s) + CO2(g) + C/ rHdeg
The stances A, B, C, D, and so on are represented by the following general equation.
The standard enthalpy of reaction, C/ rHdeg, is obtained using the following equation.
The equation can be expressed as a single weighted sum, which was introduced in Chapter 4--see page 138.
A + sign for a product and a - sign for a reactant are included in the stoichiometric number for a given reactant or product.
If we substitute these values into equation (7.23), we get equation (7.22), which proves that equation (7.23) is just another way of expressing equation (7.22).
The terms are formed by combining the standard enthalpies of formation by the corresponding coefficients or numbers, both of which are simply numbers.
There are units of kJ mol-1 C/ rHdeg.
The process involves the conversion of pure, unmixed reactants in their standard states to pure, unmixed products in their standard states.
When we learn about other quantities, this concept will be further explored.
The standard enthalpy of combustion of ethane, C2H61g2, a component of natural gas is calculated using equation (7.22).
This type of problem can be solved with an equation.
The standard enthalpy of formation for a number of compounds can be found in Appendix D.
The equation is the relationship we need.
Table 7.2 contains the data we substitute into the relationship.
We must subtract the sum of the products' standard enthalpies of formation from the sum of the reactants' standard enthalpies of formation in these types of problems.
The standard enthalpy of formation of an element in its reference form is zero.
The term can be dropped at any time in the calculation.
Data from Table 7.2 can be used to calculate the standard enthalpy of combustion of ethanol, C2H5OH1l2.
The essential step is to rearrange the equation to separate the unknown C/fHdeg from one side of the equation.
A way of organizing the data is shown.
We know the standard enthalpy of reaction with a chemical equation.
We are asked to come up with a standard enthalpy of formation.
The standard enthalpy of reaction to formations for reactants and products is related to the equation.
We organize the data needed in the calculation by writing the chemical equation for the reaction with C/ data listed under the chemical formulas.
We can use known data and rearrange the equation to get a single term on the left.
The problem only involves numerical calculations.
We were able to see how to use equation (7.22) by organizing the data as shown.
We needed to use the correct states for the compounds to get the correct answer.
The water in the product is always liquid.
We would have gotten the wrong answer if we had used the standard enthalpy.
Determine the standard enthalpy of formation of C6H12O61s2.
The standard enthalpy of burning 1CH322O1g2 is -31.70 kJ>g.
Net ionic equations are used to represent many chemical reactions in a solution.
A strong acid can be neutralized by a strong base.
We should be able to calculate neutralization by using formation data in expression, but we have to have formation data for individual ion.
There is a problem with getting these.
A single type of ion cannot be created in a chemical reaction.
We compare the enthalpies of formation of other ion to the reference ion.
The zero ion we choose is H+1aq2.
Let's see how we can use expression and data from equation to determine the enthalpy of formation of OH-1aq2.
C/fHdeg3H2O(l 24) has a -55.8 kJ mol-1.
The relevant data should be introduced after you write the net ionic equation.
Then use the equation.
The data should be in a table.
The heat given off by the system is the standard enthalpy of reaction determined here.
The precipitation of Ag2CO31s2 has a standard enthalpy change of -38 kJ per mole.
Explain how you would do this, and give any additional data you might need.
One of the most important uses of thermochemical measurements and calcula Cm1H2O2n is to assess materials as energy sources.
There are no H rials, called fuels, in these mate "hydrate" of carbon.
Fossil fuels comprise the majority of current energy needs.
The fuels are derived from millions of years ago.
Solar energy is the original source of the energy locked into these fuels.
Store energy from one source C61H2O26 if the sugar m is n.
The evolution of heat occurs when the reaction is reversed.
The principal structural material of plants is the complex cellulose carbohydrate.
It may take about 300 million years for this process to progress.
Coal is an organic rock consisting of carbon, hydrogen, and oxygen, together with small quantities of nitrogen, sulfur, and mineral matter.
Natural gas was formed in a different way.
The ocean floor was covered with sand and mud when the remains of plants and animals fell there.
Over time, the sand and mud were converted to sandstone by the weight of overlying layers of sand and mud.
The sandstone rock formation's high pressures and temperatures transformed the organic matter into oil and gas.
The deposits range in age from 250 million to 500 million years.
A typical natural gas consists of 85% methane 1CH42, 10% ethane 1C2H62, 3% propane 1C3H82, and small quantities of other gases.
A typical petroleum consists of several hundred different hydrocarbons that range in complexity from C1 to C40 or higher.
The table lists the approximate heats of the fossil fuels.
Fossil fuels have an environmental effect.
oxides of sulfur are produced by sulfur impurities in fuels.
The high temperatures asso cause the reaction of N2 and O2 in air to form oxides of gasoline nitrogen.
The environmental problem known as acid rain is caused by natural gas.
Environmental issues don't usually think of CO as an air pollutant because it is a natural and necessary component of air and is not toxic.
Its effect on the environment could be very significant.
The graphs show the history of energy consumption.
Coal and natural gas are seen as the main sources of energy for the foreseeable future, followed by petroleum.
Nuclear power and renewable power are included.
Earth's surface is warmed by this radiation.
Some of the absorbed energy is reradiated.
Some atmospheric gases, such as CO2, methane, and water vapor, absorb some of the radiation, and the energy retained in the atmosphere causes a warming effect.
Earth would be covered with ice if it weren't for it.
The atmospheric carbon dioxide concentration has varied from 180 to 300 parts per million over the past 400,000 years.
By the year 2011, the level was reflected back into space by the atmosphere and is still rising.
Increasing atmospheric carbon dioxide and some, such as certain concentrations, are caused by the burning of carbon-containing fuels such as UV light, wood, coal, natural gas, and gasoline, and the destruction of tropical stratospheric ozone.
Plants consume CO2 from the atmos the radiation from sunlight.
There are estimates of a doubling of the surface.
Predicting the probable effects of a CO on the Earth's surface is done largely through computer models, and it is very difficult to know all the factors by CO2 and other greenhouse gases.
CO2 is transparent to visible and some UV light, but it absorbs IR radiation.
The bulk flow of warm air out of the warms atmosphere is prevented by the glass in a greenhouse.
The concentration of CO2 in the atmosphere increased from 2003 to 2011.
Increased cloud formation and increased water evaporation could be a result of global warming.
Increased cloud cover could reduce the amount of solar radiation reaching Earth's surface.
Over the past 50 years, the average annual temperature for Alaska and Northern Canada has increased.
Over this time period, Alaskan winter temperatures have increased by an average of 3.5 degrees.
Tens of millions of people in Bangladesh would be displaced by a sea level increase of up to 1 m.
Simon Fraser/Science Photo Librar meters closer to the poles could show the characteristic of certain areas of the globe.
An ice core from the ice endemic could also expand.
There is a growing body of room.
Evidence shows that the ice core supports the likelihood of global warming.
The correlation between atmospheric CO2 con gases and trace elements it centration and temperature for the past 160,000 years is strong.
The data show periods of low CO2 levels and higher temperatures.
CO2 isn't the only greenhouse gas.
methane 1CH42, ozone 1O32, nitrous oxide (N2O), and trends in the pollution of the chlorofluorocarbons are some of the stronger gases.
Strategies beyond curtailing the use of chlorofluorocarbons and fossil fuels have not been developed to counter a possible global warming.
Climate change is similar to several other major environmental issues in that research, debate, and action are all likely to occur at the same time.
Coal and other energy sources have more reserves in the United States than oil and gas.
The use of coal has not increased in recent years despite the relative abundance.
The costs and dangers of deep mining of coal are considerable.
Deep mining is more harmful to the environment than surface mining.
Coal can be converted to gaseous or liquid fuels, either in surface installations or while the coal is still underground.
Before cheap natural gas became available in the 1940s, gas produced from coal was widely used in the United States.
The principal gasification reaction is very cold.
The heat requirements are met by the partial burning of coal.
A typical producer gas consists of 23% CO, 18% H2, 8% CO2, and 1% CH4 by volume.
Air is used in its production.
The heat value of natural gas is only 10% to 15% because the N2 and CO2 are noncombustible.