Succeed in chemistry using this paperback edition of CHEMISTRY & CHEMICAL REACTIVITY, Hybrid with OWL, Eighth Edition, which includes access to OWL Online Web Learning and its built-in interactive eBook. Packed with clear explanations, easy-to-follow problem-solving strategies, and dynamic study tools, the book combines thorough instruction with the powerful multimedia tools you need to develop a deeper understanding of general chemistry concepts. With OWL, you can learn at your own pace to ensure you've mastered each concept before you move on. The authors emphasize the visual nature of chemistry, illustrating the close interrelationship of the macroscopic, symbolic, and particulate levels of chemistry. The book's built-in access to the OWL online learning system helps you maximize your study time and improve your success in the course, while the interactive and customizable Cengage YouBook (interactive eBook) enhances your understanding through videoso and animations and gives you the ability to highlight, add notes, and more-including to option to download GO CHEMISTRY mini video lectures on to the key topics in the text for quick, on-the-go review on your iTunes, video iPods/iPhones, other personal video players, and QuickTime. Table of Contents.
Comprehension Checkpoint first theorized that while substances changed form during a chemical reaction, the mass of the system did not change. a.Lavoisier. b.Dalton Types of chemical reactions There is a staggering array of. Chemical occur constantly within our bodies, within plants and animals, in the air that circulates around us, in the lakes and oceans that we swim in, and even in the where we grow crops and build our homes.
In fact, there are so many chemical reactions that occur that it would be difficult, if not impossible, to understand them all. However, one that helps us to understand them is to categorize chemical reactions into a few, general types.
While not a perfect, placing reactions together according to their similarities helps us to identify patterns, which in turn allows predictions to be made about as yet unstudied reactions. In this module, we will consider and provide some context for a few categories of reactions, specifically:, decomposition, single replacement, double replacement, REDOX (including combustion), and acid-base reactions. No matter the type of, one universal truth applies to all. For a to be classified as a chemical reaction, i.e., one where a chemical change takes place, a new substance must be produced. The formation of a new substance is nearly always accompanied by an change, and often with some kind of physical or observable change. The physical change can be of different types, such as the formation of bubbles of a, a, or a color change.
These changes are clues to the existence of a chemical reaction and are important triggers for further by chemists. Synthesis reactions Prior to Lavoisier’s work, it was poorly understood that there were different gases made up of different. Instead, various gases were commonly mischaracterized as types of 'air' or air missing parts – for example, terms commonly used were 'inflammable air,' or 'dephlogisticated air.' Lavoisier thought differently and was convinced that these were different substances. He conducted where he mixed inflammable air with dephlogisticated air and a spark, and he found that the substances combined to produce water.
Songr 1 9 17 setup. In response, he renamed inflammable air 'hydrogen' from the Greek hydro for 'water' and genes for 'creator.' In so doing, Lavoisier was identifying a. In general, a synthesis reaction is one in which simpler substances combine to form another more complex one. Hydrogen and oxygen (which Lavoisier also renamed dephlogisticated air) combine in the presence of a spark to form water, summarized by the chemical equation shown below (for more on chemical equations see the section called Anatomy of a chemical equation), it represents a simple synthesis reaction. Equation 1 2H 2(g) + O 2(g) → 2H 2O (l) Decomposition reactions In 1774, the scientist turned his curiosity to a called cinnabar – a brick red mineral. When he placed the mineral under sunlight amplified by a powerful magnifying glass, he found that a was produced which he described as having an “exalted nature” because a candle burned in the gas brightly (Priestley, 1775).
Without realizing it, Priestley had discovered oxygen as a result of a decomposition. Decomposition reactions are often thought of as the opposite of reactions since they involve a being broken down into simpler compounds or even. In the case of Priestley’s oxygen, he had broken down mercury (II) oxide (cinnabar) with into its individual elements. The reaction can be summarized in the following equation.
Equation 2 2HgO (s) → 2Hg (l) + O 2(g) Single replacement reactions The British chemist and meteorologist John Daniell, invented one of the very first practical batteries in 1836 (Figure 6). In his cell, Daniell utilized a very common single replacement.
His early cells were complicated affairs, with ungainly parts and complicated constructs, but by contrast, the chemistry behind them was really quite simple. Figure 6: Daniell cell batteries. In certain, a single constituent can substitute for another one already joined in a chemical. The Daniell cell works because zinc can substitute for copper in a of copper sulfate, and in so doing exchange that are used in the battery cell. The can be summarized as follows.
Equation 3 Zn (s) + CuSO 4(aq) → ZnSO 4(aq) + Cu (s) This particular single displacement is called a metal displacement since it involves one metal replacing another metal, and many types of batteries are based on metal replacement. However, several other types of single replacement reactions exist, such as when a metal can replace hydrogen from an or from water, or a halogen can replace another halogen in certain. Combustion reactions The controlled use of fire was a crucial development for early civilization. While it’s difficult to pin down the exact time that humans first tamed the that produce fire, recent suggests it may have occurred at least a million years ago in a South African cave (Berna et al. Chemically, is no more than the of a fuel (wood, oil, gasoline, etc.) with oxygen. For combustion to take place there must be a fuel and oxygen. However, these reactions often require (discussed in more detail in the module ), which can be provided by a ‘spark’ or source of for ignition.
Fuel, oxygen, and energy are the three things make up what is known as the fire triangle (Figure 7), and any one of them being absent means that combustion will not take place. Figure 7: The fire triangle is made up of three things - fuel, oxygen, and energy. Image © Gustavb In the modern world, many of the fuels that are typically burned for, are – substances that contain both hydrogen and carbon (as discussed in more detail in our module).
Plants produce hydrocarbons when they grow, and thus make an excellent fuel source, and other hydrocarbons are produced when plants or animals over time (such as natural, oil, and other substances). When these fuels combust, the hydrogen and carbon within them combine with oxygen to produce two very familiar, water, and carbon dioxide. One simple example is the of natural gas, or methane, CH 4. Equation 4 CH 4(g) + 2O 2(g) → CO 2(g) + 2H 2O (l) As with the of all fuels, and are, too, and it is these products that are used to cook our food or to heat our homes. Reduction-oxidation reactions Each of the four types of above are sub-categories of a single type of known as redox reactions. A redox reaction is one where reduction and oxidation take place together, hence the name. The individual processes of oxidation and reduction can be defined in more than one way, but whatever the definition, the two processes are symbiotic, i.e., they must take place together.
In one definition, oxidation is described as the in which a loses, and reduction is a process where a species gains electrons. In this way, we can see how the pair must take place together. If a chemical substance is to lose electrons (and therefore be oxidized), then it must have another, interdependent chemical substance that it can give those electrons to.
In the process, the second substance (the one gaining electrons) is said to be reduced. Without such an electron acceptor, the original species can never lose the electrons and no oxidation can take place. When the electron acceptor is present, it gets reduced and the redox combination process is complete. Redox of this type can be summarized by a pair of equations – one to show the loss of electrons (the oxidation), and the other to show the gain of electrons (the reduction). Using the example of the Daniell cell above. Equation 6 Zn + Cu 2+ → Zn 2+ + Cu Other definitions of oxidation and reduction also exist, but in every case, the two halves of the redox remain symbiotic – one loses and the other gains.
The loss from one cannot happen without the other species gaining. Double displacement reactions When soap won’t easily produce a lather in water, the water is said to be ‘hard’. Hard water causes all kinds of problems that go beyond just making it difficult to form a lather. The buildup of in water pipes (known as ‘scale’), can block the flow of water and can cause problems in industrial processes. Textile manufacturing and the beverage industry rely heavily on water. In those situations, the quality of the water can make a difference to the end, so controlling the water composition is crucial.
Hard water contains magnesium or calcium in the form of a dissolved such as magnesium chloride or calcium chloride. When soap (sodium stearate) comes into contact with either of those salts, it enters into a double displacement that forms the known as ‘soap scum’. A double displacement (also known as a double replacement reaction) occurs when two ionic substances come together and both substances swap partners. Equation 8 CaCl 2(aq) + 2NA(C 17H 35COO) (aq) → 2NaCl (aq) + Ca(C 17H 35COO) 2(s) The calcium stearate is what we call soap scum, which is formed by the of the sodium stearate (the soap) in a double replacement reaction with calcium chloride. Acid-Base reactions Acid-base happen around, and even inside of us, all the time.
From the classic elementary school baking soda volcano to the of digestion, we encounter and on a daily basis. When a hydrogen loses its only, it forms a positive, H +. This hydrogen ion is the essential component of all, and indeed one definition of an acid is that of a hydrogen ion donor. Such as the citric acid in lemon juice, the ethanoic acid in vinegar, or a typical laboratory acid like hydrochloric acid, all give their hydrogen away in known as acid-base. The chemical opposites of acids are known as, and bases can be defined as hydrogen ion acceptors. Whenever an acid donates a hydrogen ion to a base, an acid-base reaction has taken place, for example, when hydrochloric acid donates a hydrogen ion to a base such as sodium hydroxide.
Comprehension Checkpoint The type of chemical reaction where a single constituent can substitute for another one already joined in a chemical compound is:. a.redox. b.single replacement Anatomy of a chemical equation Chemical equations are always linked to since they are the shorthand by which chemical are described. That fact alone makes equations incredibly important, but equations also have a crucial role to play in describing the quantitative aspect of chemistry, something that we formally call stoichiometry. All take on the same, basic format. The starting substances, or, are listed using their chemical to the left-hand side of an arrow, with multiple reactants separated with plus signs.
In the case of a between carbon and oxygen. Equation 10c C (s) + O 2(g) → CO 2(g) Finally, in order to ensure that this representation abides by the of of, the equation may need to be balanced by the addition of numbers in front of each that create equal numbers of of each on each side of the equation. In the case of the formation of carbon dioxide from carbon and oxygen, there is no need for the addition of such numbers (called the stoichiometric coefficients), since 1 carbon atom and 2 oxygen atoms appear on each side of the equation. Energy changes In nature, are often driven by exchanges in. In this respect, are generally separated into two categories – those that release energy and those that energy. Exothermic reactions are those that release to the surroundings (Figure 8, right). Are an obvious example because the energy released by the reaction is converted into the and seen in the immediate surroundings.
By contrast, endothermic reactions are those that from the surroundings (Figure 8, left). In this situation, one may have to up the or add some other form of energy to the before seeing the reaction proceed.
Figure 8: On the left is an endothermic reaction, where energy is absorbed from the surroundings. In contrast, on the right is an exothermic reaction, which releases energy into the surroundings. In both cases it is important to note that is neither created nor destroyed, rather it is transferred from one type of energy to another, for example from chemical energy to that of. The energy that goes into the formation of chemical is exchanged for other types of energy with the around that. A classic example is the reaction, in which plants light energy from the sun in order to create bonds between that make up, which are stored as chemical energy for later use by the plant. The of is essentially the reverse of photosynthesis, where the bonds in sugar are broken and the released energy is then used by the plant.
Physical change - a change in a substance that does not involve a change in the identity, including a change in state between solid, liquid and gas phases. Chemical change - a change which one or more substances are converted into different substances, some indications are burning, rusting, color changes etc. Physical change can be changed but in chemical change chemicals are mixed together for example if we are sick that is a physical change we have to change healthy as well when we mix salt in a dish,can we take out that again? Never this is chemical change A physical change is when the object undergoes a change that can be reversed, such as mixing water with salt, melting ice, or freezing or vaporizing water. A chemical change is a change that changes a substance into something else entirely, and therefore cannot be reversed.
Examples would be mixing vinegar and baking soda or turning batter into bread. Chemical change is when during a chemical reaction a new substance will be formed while in physical change no new substances are formed.
Chemical Reactivity Hazards
Chemical Reactivity Reaction Examples Examples of Organic Reactions 1. Ionic Reactions The principles and terms introduced in the previous sections can now be summarized and illustrated by the following three examples. Reactions such as these are called ionic or polar reactions, because they often involve charged species and the bonding together of. Ionic reactions normally take place in liquid solutions, where solvent molecules assist the formation of charged intermediates. The substitution reaction shown on the left can be viewed as taking place in three steps.
The first is an acid-base equilibrium, in which HCl protonates the oxygen atom of the alcohol. The resulting conjugate acid then loses water in a second step to give a carbocation intermediate.
Finally, this electrophile combines with the chloride anion nucleophile to give the final product. The addition reaction shown on the left can be viewed as taking place in two steps.
The first step can again be considered an acid-base equilibrium, with the pi-electrons of the carbon-carbon double bond functioning as a base. The resulting conjugate acid is a carbocation, and this electrophile combines with the nucleophilic bromide anion. The elimination reaction shown on the left takes place in one step. The bond breaking and making operations that take place in this step are described by the curved arrows. The initial stage may also be viewed as an acid-base interaction, with hydroxide ion serving as the base and a hydrogen atom component of the alkyl chloride as an acid.
There are many kinds of molecular rearrangements. The examples shown on the left are from an important class called tautomerization or, more specifically, keto-enol tautomerization. Tautomers are rapidly interconverted constitutional isomers, usually distinguished by a different bonding location for a labile hydrogen atom (colored red here) and a differently located double bond. The equilibrium between tautomers is not only rapid under normal conditions, but it often strongly favors one of the isomers (acetone, for example, is 99.999% keto tautomer). Even in such one-sided equilibria, evidence for the presence of the minor tautomer comes from the chemical behavior of the compound.
Tautomeric equilibria are catalyzed by traces of acids or bases that are generally present in most chemical samples. Since many ionic reactions proceed by bonding interactions between electrophiles and nucleophiles, it is important to understand how these qualities vary from compound to compound, and how they may be enhanced by acid or base catalysts. To explore this matter further.
Radical Reactions If methane gas is mixed with chlorine gas and exposed to sunlight an explosive reaction takes place in which chlorinated methane products are produced along with hydrogen chloride. An unbalanced equation illustrating this reaction is shown below; the relative amounts of the various products depends on the proportion of the two reactants that are used. CH 4 + Cl 2 + energy CH 3Cl + CH 2Cl 2 + CHCl 3 + CCl 4 + HCl How does this reaction take place? Gas phase reactions, such as the chlorination of methane, do not normally proceed via ionic intermediates.
Strong evidence indicates that neutral radical intermediates, sometimes called free radicals, play a role in this and many other similar transformations. A radical is an atomic or molecular species having an unpaired, or odd, electron. Some radicals, such as nitrogen dioxide (NO 2) and nitric oxide (NO) are relatively stable, but most are so reactive that isolation and long-term study under normal conditions is not possible. A set of radical reactions called a chain reaction can account for all the facts observed for this process. The reaction is initiated by the input of energy (heat or light). The weak chlorine-chlorine bond is broken homolytically to give chlorine atoms. In these two reactions radical intermediates abstract an atom from one of the reactant molecules.
If a chlorine atom abstracts a hydrogen from methane in the first step, the resulting methyl radical abstracts a chlorine atom from chlorine in the second step, regenerating a chlorine atom. This is therefore a chain reaction. In principle a chain reaction should continue until one or both of the reactants are consumed. In practice, however, such reactions stop before completion and have to be reinitiated. This happens whenever two radical intermediates meet and combine to give a stable molecule, thus terminating the chain of reactions. Since radical intermediates are extremely reactive and are present in very low concentration, the probability that two such intermediates will collide is small. Consequently, the chain reaction will proceed through many cycles before termination occurs.
More About Free Radicals This treatment of free radicals just scratches the surface of this important class of reactions. To learn more about free radical chemistry.
Bond Energy Bond Energy Since reactions of organic compounds involve the making and breaking of bonds, the strength of bonds, or their resistance to breaking, becomes an important consideration. For example, the chlorination of methane, discussed previously, was induced by breaking a relatively weak Cl-Cl covalent bond. Bond energy is the energy required to break a covalent bond homolytically (into neutral fragments). Bond energies are commonly given in units of kcal/mol or kJ/mol, and are generally called bond dissociation energies when given for specific bonds, or average bond energies when summarized for a given type of bond over many kinds of compounds.
Tables of bond energies may be found in most text books and handbooks. The following table is a collection of average bond energies for a variety of common bonds.
Such average values are often referred to as standard bond energies, and are given here in units of kcal/mole. The SI unit of energy is the joule, symbol J. To convert kilocalories into kilojoules multiply by 4.184. A useful site for unit conversions may be reached. Standard Bond Energies Single Bonds ΔHº. Single Bonds ΔHº. Multiple Bonds ΔHº.
H–H 104.2 B–F 154 C=C 146 C–C 83 B–O 123 N=N 109 N–N 38.4 C–N 73 O=O 119 O–O 35 N–CO 86 C=N 147 F–F 36.6 C–O 85.5 C=O ( CO 2) 192 Si–Si 52 O–CO 110 C=O (aldehyde) 177 P–P 51 C–S 65 C=O (ketone) 178 S–S 54 C–F 116 C=O (ester) 179 Cl–Cl 58 C–Cl 81 C=O (amide) 179 Br–Br 46 C–Br 68 C=O (halide) 177 I–I 36. C–I 51 C=S ( CS 2) 138 H–C 99 C–B 94 N=O ( HONO 2) 143 H–N 93 C–Si 83 P=O ( POCl 3) 110 H–O 111 C–P 73 P=S ( PSCl 3) 70 H–F 135 N–O 55 S=O ( SO 2) 128 H–Cl 103 S–O 87 S=O ( DMSO) 93 H–Br 87.5 Si–F 132 P=P 84 H–I 71 Si–Cl 86 P≡P 117 H–B 90 Si–O 110 C≡O 258 H–S 81 P–Cl 79 C≡C 200 H–Si 70 P–Br 65 N≡N 226 H–P 77 P–O 96 C≡N 213. Average Bond Dissociation Enthalpies in kcal mole -1 Some useful and interesting conclusions may be drawn from this table. First, a single bond between two given atoms is weaker than a double bond, which in turn is weaker than a triple bond.
Second, hydrogen forms relatively strong bonds (90 to 110 kcal) to the common elements found in organic compounds (C, N & O). Third, with the exception of carbon and hydrogen, single bonds between atoms of the same element are relatively weak (35 to 64 kcal). Indeed, the fact that carbon forms relatively strong bonds to itself as well as to nitrogen, oxygen and hydrogen is a primary factor accounting for the very large number of stable organic compounds.
Energetics Reaction Energetics Chemical reactions involve breaking and making some (or even all) of the bonds that hold together the atoms of reactant and product molecules. Energy is required to break bonds, and since the strengths of different kinds of bonds differ, there is often a significant overall energy change in the course of a reaction. In the combustion of methane, for example, all six bonds in the reactant molecules are broken, and six new bonds are formed in the product molecules (equation 1). Reactants Products The sum of the product bond strengths in this case is greater than the sum of reactant bond strengths; consequently, the products are energetically (or thermodynamically) more stable than the reactants, and energy is released in the form of heat.
Such reactions are called exothermic. It is helpful to think of exothermic reactions as proceeding from a higher energy (less stable) reactant state to a lower energy (more stable) product state, as shown in the diagram on the right. Reactions in which the products are higher energy than the reactants require an energy input to occur, and are called endothermic.
Photosynthesis (equation 2) is an important example of an endothermic process. Energy in the form of photons (sunlight) drives the reaction, which requires chlorophyll as a catalyst.
Reactants Products Common sense suggests that molecules in which the bonds are all strong will be more stable than molecules having weaker bonds. Previously we defined bond strengths as the energy required to break a bond into neutral fragments (radicals or atoms).
The sum of all the bond energies of a molecule can therefore be considered its atomization energy, i.e. The energy required to break the molecule completely into its component atoms. If this concept is applied to a group of isomers, it should be clear that all the isomers will have a common atomization state, and that the total bond energy of each isomer is inversely related to that isomer's potential energy. Thus, that isomer having the greatest total bond energy has the lowest potential energy and is thermodynamically most stable. To summarize, bond energy is energy that must be introduced to break a bond, and is not a component of a molecule's potential energy. The three C 6H 12 isomers on the right illustrate this relationship.
Cyclohexane is made up of six C-C sigma bonds and twelve C-H sigma bonds configured in a strain-free six-membered ring. The isomer having a double bond, 1-hexene, on the other hand, has four C-C single bonds (all sigma) and one C-C double bond (one sigma and one pi bond).
Since the pi bond is weaker than a sigma bond, cyclohexane has a larger total bond energy (by nearly 20 kcal/mol) and is thermodynamically more stable than 1-hexene. The four-membered ring compound, ethylcyclobutane, has the same kinds of bonds as cyclohexane, but they are weakened by ring strain to such a degree that this isomer is even less stable (thermodynamically) than 1-hexene. The Nature of Energy The term energy is used here in a general and rather imprecise way. For a more comprehensive treatment. Activation Energy Since exothermic reactions are energetically (thermodynamically) favored, a careless thinker might conclude that all such reactions will proceed spontaneously to their products. Were this true, no life would exist on Earth, because the numerous carbon compounds that are present in and essential to all living organisms would spontaneously combust in the presence of oxygen to give carbon dioxide-a more stable carbon compound. The combustion of methane (eq.1), for example, does not occur spontaneously, but requires an initiating energy in the form of a spark or flame.
The flaw in this careless reasoning is that we have focused only on the initial (reactant) and final (product) states of reactions. To understand why some reactions occur readily (almost spontaneously), whereas other reactions are slow, even to the point of being unobservable, we need to consider the intermediate stages of reactions.
Chemical Reactivity Chart
Exothermic Single Step Reaction Endothermic Single Step Reaction Exothermic Two Step Reaction Every reaction in which bonds are broken will have a high energy transition state that must be reached before products can form. In order for the reactants to reach this transition state, energy must be supplied and reactant molecules must orient themselves in a suitable fashion. The energy needed to raise the reactants to the transition state energy level is called the activation energy, ΔE . An example of a single-step exothermic reaction profile is shown on the left above, and a similar single-step profile for an endothermic reaction is in the center. The activation energy is drawn in red in each case, and the overall energy change (ΔE) is in green. The profile becomes more complex when a multi-step reaction path is described. An example of a two-step reaction proceeding by way of a high energy intermediate is shown on the right above.
Here there are two transition states, each with its own activation energy. The overall activation energy is the difference in energy between the reactant state and the highest energy transition state. We see now why the rate of a reaction may not correlate with its overall energy change. In the exothermic diagram on the left, a significant activation energy must be provided to initiate the reaction. Since the reaction is strongly exothermic, it will probably generate enough heat to keep going as long as reactants remain.
The endothermic reaction in the center has a similar activation energy, but this will have to be supplied continuously for the reaction to proceed to completion. What is the source of the activation energy that enables a chemical reaction to occur?
Chemical Reactivity Of An Atom
Often it is heat, as noted above in reference to the flame or spark that initiates methane combustion. At room temperature, indeed at any temperature above absolute zero, the molecules of a compound have a total energy that is a combination of translational (kinetic) energy, internal vibrational and rotational energies, as well as electronic and nuclear energies. The temperature of a system is a measure of the average kinetic energy of all the atoms and molecules present in the system. As shown in the following diagram, the average kinetic energy increases and the distribution of energies broadens as the temperature is raised from T 1 to T 2. Portions of this thermal or kinetic energy provide the activation energy for many reactions, the concentration of suitably activated reactant molecules increasing with temperature, e.g. Orange area for T 1 and yellow plus orange for T 2.
(Note that the area under a curve or a part of a curve is proportional to the number of molecules represented.) Distribution of Molecular Kinetic Energy at Two Different Temperatures, T 1 & T 2 2. Reaction Rates and Kinetics Chemical reactivity is the focus of chemistry, and the study of reaction rates provides essential information about this subject. Some reactions proceed so rapidly they seem to be instantaneous, whereas other reactions are so slow they are nearly unobservable. Most of the reactions described in this text take place in from 0.2 to 12 hours at 25 ºC. Temperature is important, since fast reactions may be slowed or stopped by cooling, and slow reactions are accelerated by heating.
When a reaction occurs between two reactant species, it proceeds faster at higher concentrations of the reactants. These facts lead to the following general analysis of reaction rates. Reaction Rate = Number of Collisions between Reactant Molecules per Unit of Time Fraction of Collisions with Sufficient Energy to pass the Transition State an Orientational or Probability Factor Since reacting molecules must collide to interact, and the necessary activation energy must come from the kinetic energy of the colliding molecules, the first two factors are obvious. The third (probability) factor incorporates the orientational requirements of the reaction. For example, the addition of bromine to a double bond at the end of a six-carbon chain (1-hexene) could only occur if the colliding molecules came together in a way that allowed the bromine molecule to interact with the pi-electrons of the double bond.
The collision frequency of reactant molecules will be proportional to their concentration in the reaction system. This aspect of a reaction rate may be incorporated in a rate equation, which may take several forms depending on the number of reactants. Three general examples are presented in the following table. Reaction Type Rate Equation Reaction Order A —— B Reaction Rate = kA First Order Reaction (no collision needed) A + B —— C + D Reaction Rate = kA.B Second Order Reaction A + A —— D Reaction Rate = kA 2 Second Order Reaction These rate equations take the form Reaction Rate = kX nY m, where the proportionality constant k reflects the unique characteristics of a specific reaction, and is called the rate constant. The concentrations of reactants X and Y are X and Y respectively, and n & m are exponential numbers used to fit the rate equation to the experimental data. The sum n + m is termed the kinetic order of a reaction.
The first example is a simple first order process. The next two examples are second order reactions, since n + m = 2.
The kinetic order of a reaction is usually used to determine its. In writing a rate equation we have disconnected the collision frequency term from the activation energy and probability factors defined above, which are necessarily incorporated in the rate constant k. This is demonstrated by the following equation. The complex parameter A incorporates the probability factor. Because of the exponential relationship of k and the activation energy small changes in ΔE will cause relatively large changes in reaction rate. An increase in temperature clearly acts to increase k, but of greater importance is the increase in average molecular kinetic energy such an increase produces. This was illustrated in a, increase in temperature from T 1 to T 2 producing a larger proportion of reactant molecules having energies equal or greater than the activation energy (designated by the red line.
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