代做CHEM 191 MODULE 4 CHEMICAL REACTIONS 2: ENERGY AND RATE代写数据结构语言程序

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CHEM 191

MODULE 4

CHEMICAL REACTIONS 2: ENERGY AND RATE

Learning Objectives.

By the end of this module, you should be able to:

• Write thermochemical equations and describe reactions as exothermic or endothermic

• Calculate energy changes for reactions using data from thermochemical equations

• Calculate the enthalpy of a reaction from bond dissociation enthalpy data

• Use collision theory to explain how changing reaction conditions affects the rate of the reaction

• Predict the effect of change in concentration, particle size, temperature and the addition of a catalyst on the rate of a given reaction

• Describe the role of a catalyst in a chemical reaction

Reference: ESA Chapters 11, 17 and 18.

CHEMICAL REACTIONS AND ENERGY

In Module 3 we discussed the processes involved in chemical reactions by considering from a particle perspective with atoms being reorganized through the breaking of bonds in reactant and the making of the new bonds in the products. The difference in the energy needed to break then bonds and the energy released when the bonds are made is the basis of our understanding the energy changes that occur in chemical reactions. The study of the interconversion of heat and other forms of energy, a very  important branch of chemistry, is known as thermodynamics. It enables chemists to predict, with some degree of confidence,  whether a particular chemical reaction will proceed in a given direction. Thermochemistry, a sub branch of thermodynamics, is the study of heat changes that  take place during chemical reactions.

We have already observed that energy is involved in physical changes such as changes of state. When a solid is heated, the forces of attraction that hold the particles of the solid in place in the lattice can be broken. This allows the particles more freedom of movement and, as more heat is supplied, their kinetic energy (energy of movement) increases. Temperature is a measure of the average kinetic energy of all the particles in a substance.

Nearly every chemical reaction either produces or absorbs energy. This is   generally in the form. of heat, although light and sound are sometimes also  observed. Human activity has been characterised by the use of combustion reactions involving the burning of carbon-based fuels such as  wood and coal for thousands of years. More recently, industrialisation has  come about through the burning of a wider range of fossil fuels, in  particular oil and natural gas, with the resulting increase in carbon dioxide  in the atmosphere contributing to global climate change. This has led to a  quest for ways to trap and use renewable sources of energy such as solar  power and for the development of technologies for using new energy sources  such as hydrogen gas.

The combustion of natural gas, methane, produces carbon dioxide and water:

CH4(g) + 2O2(g)  →   CO2(g)  +  2H2O(g)  +  energy

The same products are formed during respiration, the process by which glucose is reacted with oxygen  to provide energy for metabolic processes:

C6H12O6(s) + 6O2(g)  →  6 CO2(g)  +  6H2O(g)  +  energy

ENDOTHERMIC AND EXOTHERMIC REACTIONS

When a “gummy bear‟ is placed in some molten potassium chlorate there is a huge explosion as the sugar is oxidised to carbon dioxide and water. The energy of the molecules in the sugar (reactant) is greater than the energy  of the carbon dioxide and water (products) and  the enormous amount of excess energy is released to the surroundings as heat and light.

You can watch this experiment athttps://www.youtube.com/watch?v=7lfnOTNlqtY

When hydrogen and oxygen gases react together, water is produced along with a considerable amount of energy. We could write the equation for the reaction as:

2H2(g)   +   O2(g)        →   2H2O(g)  +  energy

This energy is liberated in the form. of heat and in the form. of sound energy (there is a loud bang).  All  the energy lost by the system is gained by the surroundings and we say it is an exothermic process. Energy released is by definition, negative and in an exothermic process the  total energy of the products must be lower than the total energy ofthe reactants.

When mercury(II) oxide, HgO, decomposes energy must be supplied for the reaction to proceed. The equation for the reaction can be written as:

2HgO(s)   +   energy   →       2 Hg(ℓ) +   O2(g)

This reaction is called an endothermic process and the energy supplied is by definition, positive. In endothermic processes the total energy of the products must be higher than the total energy of the reactants.

ENTHALPY

The energy content of a system is the sum of the kinetic and potential energies of all the system  components. It is not possible to measure the total energy of a system but the energy change during a  reaction is measurable and particularly useful. When a reaction is carried out at a constant pressure the  heat of the reaction (amount of heat transferred during the reaction) is defined as the enthalpy change,  and is given the symbol ΔH. The enthalpy (heat content) of a system has the symbol Hand “Δ” is used  to signify a “change   in” or “difference in” . Hence, ΔH for a reaction will be the difference in enthalpy  between the reactants and   the products.

ΔH = Hfinal - Hinitial   = Hproducts - Hreactants

This means that for an endothermic process Hfinal  > Hinitial  so ΔH > 0. For an exothermic process Hfinal  < Hinitial  so ΔH < 0.

Focusing Questions 1:

1.         What does the symbol ‘ΔH’ stand for?

2.         What is the definition of enthalpy?

3.         Which is greater in an endothermic reaction, energy of reactants or energy of products?

4.          Does a positive enthalpy value indicate energy gained or lost?

Watch the video “Enthalpy” (12 min), which is on Nuku.

Activity 4.1 - Enthalpy and Thermochemical Equations

INFORMATION

Most chemical reactions are accompanied by a change in energy which is usually released to or absorbed from the surroundings in the form. of heat. The result is a change in temperature. The thermodynamic name associated with thermal energy is enthalpy, which has the symbol H.

When a reaction gives off heat to its surroundings, the energy of the system decreases. The reaction is said to be exothermic and the change in enthalpy ΔH is negative.

When a reaction absorbs heat from its surroundings, the energy of the system increases. The reaction is said to be endothermic and the change in enthalpy ΔH is positive.

A chemical equation for which the enthalpy change ΔH is given is called a thermochemical equation.

For example: C(s) + ½O2(g) → CO(g) ΔH = -111 kJ mol-1 (1)

This equation tells us that when 1 mol of solid carbon reacts with ½ mol of gaseous oxygen 1 mol of gaseous carbon monoxide is formed and 111 kJ of energy are released.

If the equation is reversed energy is needed to break CO into C and O2.

Increasing the number of moles increases the energy needed/released. So

2CO(g) → 2C(s) + O2(g) ΔH = +222 kJ mol-1 (2)

It is important that physical states be included in thermodynamic equations. For example, the enthalpy changes for the following reactions are different.

H2(g) + ½O2(g) → H2O(g) ΔH = -242 kJ mol-1 (3)

H2(g) + ½O2(g) → H2O(ℓ) ΔH = -285 kJ mol-1 (4)

Questions

1.          Complete the table below by describing the type of reaction as endothermic or exothermic

Type of reaction	ΔH	Energy
(a)	> 0	is absorbed
(b)	< 0	is released

Questions 2 to 7 below refer to the reactions in the INFORMATION Table above

2.    Consider reaction (1) in the information above. When carbon is burned in oxygen will energy be given or will it be needed for the reaction to occur? Justify your answer.

3.    Explain why ΔHfor reaction (2) is twice as big as for reaction (1) but has the opposite sign.

4.    Write a sentence to explain what equation (3) means.

5.    How does equation (4) differ from equation (3)?

6.    Is more or less energy released when hydrogen and oxygen react to form gaseous or liquid water?

7.    Give a reason for the difference in the ΔH values in equations (3) and (4).

8.    The neutralisation reaction of an acid solution with sodium hydroxide solution is an exothermic  process. If you hold a test tube in which this reaction occurs, will it feel hot or cold? Give reasons for your answer.

9.    Describe the following reactions as exothermic or endothermic

(a)   The evaporation of water                       (b)         The burning of natural gas

(c)  The freezing of water                              (d)         The dissolving of sugar in water

THERMOCHEMICAL EQUATIONS

A thermochemical equation links the enthalpy change for a reaction to the molar amounts in a  balanced equation.

For example: The energy released when hydrogen and oxygen react to produce water is 572 kJ per mol. This can be  written as:

2H2(g)   +   O2(g)     →   2H2O(g)        ΔrH = -572 kJ mol-1

ΔrH is the enthalpy change for the reaction.

•     It is important that physical states be included in thermodynamic equations. For example the   enthalpy changes for producing H2O(ℓ) is different than for H2O(g)

2H2(g)   +   O2(g)      →      2H2O(ℓ)        ΔrH = -584 kJ mol-1

•     When the coefficients in an equation are changed the value of ΔrH is changed accordingly. For  example: H2(g)  +  ½ O2(g)  →  H2O(g)            ΔrH = -286 kJ mol-1

•     Similarly if the reaction is reversed the magnitude of the value of ΔrH does not change but the  sign is reversed.

2H2O(g)   →   2H2(g)   +   O2(g)       ΔrH = +572 kJ mol-1

•     Enthalpy changes for given amounts of reactants can be calculated using ΔrH and the masses of reactants or products used.

For example: To calculate the heat energy released when 200 g of octane, C8H18, is burned. C8H18(ℓ)   +  12.5 O2(g)  → 8CO2(g)  +  9H2O(g)       ΔrH = -5470 kJ mol-1

M (C8H18) = 114 gmol-1

n(C8H18) = m/M = 200 g/114 gmol-1  = 1.75 mol

From the equation: 1 mol C8H18(ℓ)  releases 5470 kJ

So 1.75 mol releases:  5470 kJ mol-1  × 1.75 mol = 9596 kJ

NOTE: The enthalpy change quoted for a reaction is given the units ‘kJ per mole” . This means per “mole of reaction”.

To explain this we could consider a mythical reaction: 2A  +  3B    →  4C

When one mole of reaction has taken place, two moles of A and three moles of B will have reacted to produce four moles of C. So one mole of reaction means “for the reaction as in the given equation” .

If more or less than one mole of reaction is  involved, the energy per mole of reaction is determined first and then this value is scaled according to  how much reaction actually took place.

For example: For the reaction           2H2O(g)   →    2H2(g)   +   O2(g)       ΔrH = +572 kJ mol-1

When 2 mol of water decomposes to give 2 mol of hydrogen gas and 1 mol of oxygen gas, 572 kJ of energy is needed. However, for the reaction:

H2(g)  +  ½ O2(g)  →  H2O(g)          ΔrH = -286 kJ mol-1

the mole of reaction refers to the energy released when 1 mol H2  reacts with ½ mol O2  to produce 1 mol H2O.

Watch the video “Thermochemical Equations” (12 min) which is on Nuku.

Exercise 4.1

(a)     Given:   C(s)  +  ½O2(g)  →  CO(g)   ΔH= -111 kJ mol-1, calculate the enthalpy change for the following reactions:

(i)    6C(s)  +  3O2(g)  →   6CO(g) (ii)   2CO(g) →   2C(s)  +  O2(g)

(iii) Calculate the energy released when 50.0 g of carbon are burned in oxygen to produce CO gas

(b)     (i)    Calculate the energy used to vapourise 10.0 g of hydrazine, N2H4(ℓ), given:

N2H4(ℓ)    →   N2H4(g)   ΔH = +10 kJ mol-1

(ii)  Calculate the mass of hydrazine that needs to be condensed to release 500 kJ of energy

(c)      When Sulfur dioxide is converted to sulfur trioxide in the manufacture of sulfuric acid, energy is released.

2SO2(g)  +  O2(g)     →   2SO3(g) ΔH= -188 kJ mol-1

(i)    Calculate the energy released when 200.0 g of sulfur dioxide are burnt in excess oxygen (M(SO2) = 64.0 gmol-1)

(ii)   Calculate the mass of SO2 needed to produce 3000 kJ of energy

BOND ENERGY AND CHEMICAL CHANGE Bond properties

Bond order, bond length and bond enthalpy are all inter-related. Bond order relates to the number of electron  pairs shared between any two bonded atoms. In single bonds where one electron pair is shared the bond order is one. For a double bond, where two pairs of electrons are shared, the bond order is two and it  is three for a triple bond.

Bond length is the distance between the nuclei of the atoms in a bond at the point of minimum energy. Bond lengths depend on the size of the atoms  involved in the bond and the bond order. As the  bond order increases the nuclei are more strongly attracted to the bonded electrons pulling the atoms closer  together and the bond length will decrease.

Bond enthalpy or bond dissociation enthalpy is the energy needed to break a bond in 1 mole of gaseous molecules. The strength of a bond will depend on how strongly the nuclei are attracted to the bonding electrons. Since energy is always needed to break a bond, enthalpy of bond breaking will   always be positive i.e. ΔHbond breaking  > 0. The same amount of energy needed to break a bond is  released when the  same bond is made.

Table 4.1 Bond length and bond enthalpy

Bond	Bond length / pm	Bond enthalpy / kJ mol-1
C-H	109	414
C-O	143	358
C=O	122	804

H-O	96	463
N-H	101	391
H-H	74	436
C-C	154	346
C=C	134	614
C≡C	120	839
N≡N	110	945
C-N	147	286
O=O	121	498

Activity 4.2 Making and Breaking Bonds

Chemical reactions can be a lot like playing with Lego®—you must take apart part of your last creation before you can replace it with something new. For many chemical reactions, we have to first break bonds in the reactants before we can put the atoms back into a new arrangement to form. the products. Both of these processes involve changes in energy. The net energy change for a reaction is called the heat of reaction or the change in enthalpy (ΔH).

	Reaction	Energy change / kJ mol-1		Reaction	Energy change / kJ mol-1
A	PCl3(g) → P(g) + 3Cl(g)	+966.7	E	P(g) + 3Cl(g) → PCl3(g)	-966.7
B	PCl5(g) → P(g) + 5Cl(g)	+1297.9	F	P(g) + 5Cl(g) → PCl5(g)	-1297.9
C	PF3(g) → P(g) + 3F(g)	+1470.4	G	P(g) + 3F(g) → PF3(g)	-1470.4
D	PF5(g) → P(g) + 5F(g)	+2305.4	H	P(g) + 5F(g) → PF5(g)	-2305

Use the table to answer the following questions:

1.    Identify four reactions where bonds are being broken.

2.    Identify four reactions where bonds are being formed.

3.    Decide on the correct word to complete each sentence below:

(i)      When bonds are (broken/formed) there is a positive energy change.

(ii)     Breaking bonds is (endothermic/exothermic).

(iii)    When bonds are (broken/formed) there is a negative energy change.

(iv)    Forming bonds is (endothermic/exothermic).

4.    Find two reactions in the table that are exact opposites of each other, that is, one reaction is the reverse of the other reaction.

(i)      How do the changes in energy for the reverse reactions compare?

(ii)     Explain your answer to (i) considering what you learned from Questions 1–3 about bond breaking and bond formation.

5.    Consider the data in the table

(i)      What are the units for the energy changes?

(ii)     For Reaction A of how many P–Cl bonds are broken with the 966.7 kJ of energy listed?

(iii)    Calculate the energy needed to break one mole of P–Cl bonds in Reaction A.

6.    Use the data in the table above to answer the following questions.

(i)      Calculate the energy needed to break one mole of P–Cl bonds in Reaction B.

(ii)     Do the P–Cl bonds in different molecules require the same amount of energy to break?

Bond energies can be very useful (as you will soon discover) for calculating the net energy change in a reaction. However, a table listing the bond energies for even the most common substances would be several pages long. For this reason, chemists often approximate energy changes using average bond energy.