Stoichiometry coefficient

Stoichiometry

( /ˌstɔɪkiˈɒmɨtri/) is a branch of chemistry that deals with the relative quantities of reactants and products in chemical reactions. In a balanced chemical reaction, the relations among quantities of reactants and products typically form a ratio of whole numbers. For example, in a reaction that forms ammonia (NH₃), exactly one molecule of nitrogen (N₂) reacts with three molecules of hydrogen (H₂) to produce two molecules of NH₃:

N₂ + 3H₂ → 2NH₃

Stoichiometry can be used to calculate quantities such as the amount of products (in mass, moles, volume, etc.) that can be produced with given reactants and percent yield (the percentage of the given reactant that is made into the product). Stoichiometry calculations can predict how elements and components diluted in a standard solution react in experimental conditions. Stoichiometry is founded on the law of conservation of mass: the mass of the reactants equals the mass of the products.
Reaction stoichiometry describes the quantitative relationships among substances as they participate in chemical reactions. In the example above, reaction stoichiometry describes the 1:3:2 ratio of molecules of nitrogen, hydrogen, and ammonia.
Composition stoichiometry describes the quantitative (mass) relationships among elements in compounds. For example, composition stoichiometry describes the nitrogen to hydrogen (mass) relationship in the compound ammonia: i.e., one mole of nitrogren and three moles of hydrogen are in every mole of ammonia.
A stoichiometric amount or stoichiometric ratio of a reagent is the amount or ratio where, assuming that the reaction proceeds to completion:

1. all reagent is consumed,
2. there is no shortfall of reagent, and
3. no residues remain.

A nonstoichiometric mixture, where reactions have gone to completion, will have only the limiting reagent consumed completely.
While almost all reactions have integer-ratio stoichiometry in amount of matter units (moles, number of particles), some nonstoichiometric compounds are known that cannot be represented by a ratio of well-defined natural numbers. These materials therefore violate the law of definite proportions that forms the basis of stoichiometry along with the law of multiple proportions.
Gas stoichiometry deals with reactions involving gases, where the gases are at a known temperature, pressure, and volume, and can be assumed to be ideal gases. For gases, the volume ratio is ideally the same by the ideal gas law, but the mass ratio of a single reaction has to be calculated from the molecular masses of the reactants and products. In practice, due to the existence of isotopes, molar masses are used instead when calculating the mass ratio.

Contents  [hide]  1 Etymology 2 Definition 3 Different stoichiometries in competing reactions 4 Stoichiometric coefficient 5 Stoichiometry matrix 6 Gas stoichiometry 7 Stoichiometric air-fuel ratios of common fuels 8 References 9 External links

[edit] Etymology

"Stoichiometry" is derived from the Greek words στοιχεῖον (stoicheion, meaning element]) and μέτρον (metron, meaning measure.) In patristic Greek, the word Stoichiometria was used by Nicephorus to refer to the number of line counts of the canonical New Testament and some of the Apocrypha.

[edit] Definition

Stoichiometry rests upon the very basic laws which help to understand it better i.e law of conservation of mass, the law of definite proportions (i.e., the law of constant composition) and the law of multiple proportions. In general, chemical reactions combine in definite ratios of chemicals. Since chemical reactions can neither create nor destroy matter, nor transmute one element into another, the amount of each element must be the same throughout the overall reaction. For example, the amount of element X on the reactant side must equal the amount of element X on the product side.
Stoichiometry is often used to balance chemical equations (reaction stoichiometry). For example, the two diatomic gases, hydrogen and oxygen, can combine to form a liquid, water, in an exothermic reaction, as described by the following equation:

Reaction stoichiometry describes the 2:1:2 ratio of hydrogen, oxygen, and water molecules in the above equation.
The term stoichiometry is also often used for the molar proportions of elements in stoichiometric compounds (composition stoichiometry). For example, the stoichiometry of hydrogen and oxygen in H₂O is 2:1. In stoichiometric compounds, the molar proportions are whole numbers.
Stoichiometry is not only used to balance chemical equations but also used in conversions, i.e., converting from grams to moles, or from grams to millilitres. For example, to find the number of moles in 2.00 g of NaCl, one would do the following:

In the above example, when written out in fraction form, the units of grams form a multiplicative identity, which is equivalent to one (g/g=1), with the resulting amount of moles (the unit that was needed), is shown in the following equation,

Stoichiometry is also used to find the right amount of reactants to use in a chemical reaction (stoichiometric amounts). An example is shown below using the thermite reaction,

This equation shows that 1 mole of aluminium oxide and 2 moles of iron will be produced with 1 mole of iron(III) oxide and 2 moles of aluminium. So, to completely react with 85.0 g of iron(III) oxide (0.532 mol), 28.7 g (1.06 mol) of aluminium are needed.

[edit] Different stoichiometries in competing reactions

Often, more than one reaction is possible given the same starting materials. The reactions may differ in their stoichiometry. For example, the methylation of benzene (C₆H₆), through a Friedel-Crafts reaction using AlCl₃ as catalyst, may produce singly methylated (C₆H₅CH₃), doubly methylated (C₆H₄(CH₃)₂), or still more highly methylated products, as shown in the following example,

In this example, which reaction takes place is controlled in part by the relative concentrations of the reactants.

[edit] Stoichiometric coefficient

In layman's terms, the stoichiometric coefficient (or stoichiometric number in the IUPAC nomenclature^[1]) of any given component is the number of molecules which participate in the reaction as written.
For example, in the reaction CH₄ + 2 O₂ → CO₂ + 2 H₂O, the stoichiometric coefficient of CH₄ would be 1 and the stoichiometric coefficient of O₂ would be 2.
In more technically-precise terms, the stoichiometric coefficient in a chemical reaction system of the i–th component is defined as

or

where N_i is the number of molecules of i, and ξ is the progress variable or extent of reaction (Prigogine & Defay, p. 18; Prigogine, pp. 4–7; Guggenheim, p. 37 & 62).

The extent of reaction ξ can be regarded as a real (or hypothetical) product, one molecule of which is produced each time the reaction event occurs. It is the extensive quantity describing the progress of a chemical reaction equal to the number of chemical transformations, as indicated by the reaction equation on a molecular scale, divided by the Avogadro constant (it is essentially the amount of chemical transformations). The change in the extent of reaction is given by dξ = dn_B/n_B, where n_B is the stoichiometric number of any reaction entity B (reactant or product) an dn_B is the corresponding amount.^[2]

The stoichiometric coefficient ν_i represents the degree to which a chemical species participates in a reaction. The convention is to assign negative coefficients to reactants (which are consumed) and positive ones to products. However, any reaction may be viewed as "going" in the reverse direction, and all the coefficients then change sign (as does the free energy). Whether a reaction actually will go in the arbitrarily-selected forward direction or not depends on the amounts of the substances present at any given time, which determines the kinetics and thermodynamics, i.e., whether equilibrium lies to the right or the left.
If one contemplates actual reaction mechanisms, stoichiometric coefficients will always be integers, since elementary reactions always involve whole molecules. If one uses a composite representation of an "overall" reaction, some may be rational fractions. There are often chemical species present that do not participate in a reaction; their stoichiometric coefficients are therefore zero. Any chemical species that is regenerated, such as a catalyst, also has a stoichiometric coefficient of zero.
The simplest possible case is an isomerism

in which ν_B = 1 since one molecule of B is produced each time the reaction occurs, while ν_A = −1 since one molecule of A is necessarily consumed. In any chemical reaction, not only is the total mass conserved, but also the numbers of atoms of each kind are conserved, and this imposes corresponding constraints on possible values for the stoichiometric coefficients.
There are usually multiple reactions proceeding simultaneously in any natural reaction system, including those in biology. Since any chemical component can participate in several reactions simultaneously, the stoichiometric coefficient of the i–th component in the k–th reaction is defined as

so that the total (differential) change in the amount of the i–th component is

.

Extents of reaction provide the clearest and most explicit way of representing compositional change, although they are not yet widely used.
With complex reaction systems, it is often useful to consider both the representation of a reaction system in terms of the amounts of the chemicals present { N_i } (state variables), and the representation in terms of the actual compositional degrees of freedom, as expressed by the extents of reaction { ξ_k }. The transformation from a vector expressing the extents to a vector expressing the amounts uses a rectangular matrix whose elements are the stoichiometric coefficients [ ν_i k ].
The maximum and minimum for any ξ_k occur whenever the first of the reactants is depleted for the forward reaction; or the first of the "products" is depleted if the reaction as viewed as being pushed in the reverse direction. This is a purely kinematic restriction on the reaction simplex, a hyperplane in composition space, or N‑space, whose dimensionality equals the number of linearly-independent chemical reactions. This is necessarily less than the number of chemical components, since each reaction manifests a relation between at least two chemicals. The accessible region of the hyperplane depends on the amounts of each chemical species actually present, a contingent fact. Different such amounts can even generate different hyperplanes, all of which share the same algebraic stoichiometry.
In accord with the principles of chemical kinetics and thermodynamic equilibrium, every chemical reaction is reversible, at least to some degree, so that each equilibrium point must be an interior point of the simplex. As a consequence, extrema for the ξ's will not occur unless an experimental system is prepared with zero initial amounts of some products.
The number of physically-independent reactions can be even greater than the number of chemical components, and depends on the various reaction mechanisms. For example, there may be two (or more) reaction paths for the isomerism above. The reaction may occur by itself, but faster and with different intermediates, in the presence of a catalyst.
The (dimensionless) "units" may be taken to be molecules or moles. Moles are most commonly used, but it is more suggestive to picture incremental chemical reactions in terms of molecules. The N's and ξ's are reduced to molar units by dividing by Avogadro's number. While dimensional mass units may be used, the comments about integers are then no longer applicable.

[edit] Stoichiometry matrix

In complex reactions, stoichiometries are often represented in a more compact form called the stoichiometry matrix. The stoichiometry matrix is denoted by the symbol, .
If a reaction network has n reactions and m participating molecular species then the stoichiometry matrix will have corresponding m rows and n columns.
For example, consider the system of reactions shown below:

S₁ → S₂

5S₃ + S₂ → 4S₃ + 2S₂

S₃ → S₄

S₄ → S₅.

This systems comprises four reactions and five different molecular species. The stoichiometry matrix for this system can be written as:

where the rows correspond to S₁, S₂, S₃, S₄ and S₅, respectively. Note that the process of converting a reaction scheme into a stoichiometry matrix can be a lossy transformation, for example, the stoichiometries in the second reaction simplify when included in the matrix. This means that it is not always possible to recover the original reaction scheme from a stoichiometry matrix.
Often the stoichiometry matrix is combined with the rate vector, v to form a compact equation describing the rates of change of the molecular species:

[edit] Gas stoichiometry

Gas stoichiometry is the quantitative relationship (ratio) between reactants and products in a chemical reaction with reactions that produce gases. Gas stoichiometry applies when the gases produced are assumed to be ideal, and the temperature, pressure, and volume of the gases are all known. The ideal gas law is used for these calculations. Often, but not always, the standard temperature and pressure (STP) are taken as 0°C and 1 bar and used as the conditions for gas stoichiometric calculations.
Gas stoichiometry calculations solve for the unknown volume or mass of a gaseous product or reactant. For example, if we wanted to calculate the volume of gaseous NO₂ produced from the combustion of 100 g of NH₃, by the reaction:

4NH₃ (g) + 7O₂ (g) → 4NO₂ (g) + 6H₂O (l)

we would carry out the following calculations:

There is a 1:1 molar ratio of NH₃ to NO₂ in the above balanced combustion reaction, so 5.871 mol of NO₂ will be formed. We will employ the ideal gas law to solve for the volume at 0 °C (273.15 K) and 1 atmosphere using the gas law constant of R = 0.08206 L · atm · K⁻¹ · mol⁻¹ :

PV	= nRT
V

Gas stoichiometry often involves having to know the molar mass of a gas, given the density of that gas. The ideal gas law can be re-arranged to obtain a relation between the density and the molar mass of an ideal gas:

    and    

and thus:

where:
P	= absolute gas pressure
V	= gas volume
n	= number of moles
R	= universal ideal gas law constant
T	= absolute gas temperature
ρ	= gas density at T and P
m	= mass of gas
M	= molar mass of gas

[edit] Stoichiometric air-fuel ratios of common fuels

Fuel	By mass [3]	By volume [4]	Percent fuel by mass
Gasoline	14.6 : 1	—	6.8%
Natural gas	14.5 : 1	9.7  : 1	5.8%
Propane (LP)	15.67 : 1	23.9 : 1	6.45%
Ethanol	9 : 1	—	11.1%
Methanol	6.47 : 1	—	15.6%
Hydrogen	34.3 : 1	2.39 : 1	2.9%
Diesel	14.5 : 1	0.094 : 1	6.8%

Gasoline engines can run at stoichiometric air-to-fuel ratio, because gasoline is quite volatile and is mixed (sprayed or carburetted) with the air prior to ignition. Diesel engines, in contrast, run lean, with more air available than simple stoichiometry would require. Diesel fuel is less volatile and is effectively burned as it is injected, leaving less time for evaporation and mixing. Thus, it would form soot (black smoke) at stoichiometric ratio.


question :

1. Why do we care about Stoichiometry ?
2. How stoichiometric reaction between mercapto silica hybrids with hydrogen peroxide?

Measurement of energy changes in chemical reactions

Energy changes in chemical reactions can always be made as hot, because it is more appropriate when the term is called the heat of reaction. Tool used to measure the heat of reaction is called a calorimeter (actually calorie meter, although the heat is now known more commonly expressed in joules rather than calories). There are several different forms of these tools, one called the Bomb Calorimeter are shown in the image above. Such calorimeters are usually used to study the exothermic reaction, which would not be running when not heated, for example, the combustion reaction of CH 4 with O 2 or the reaction between H 2 and O 2. Tool consists of a container made of strong steel (bombnya) where the reagent is placed. Bomb is inserted in the insulated tub and given a stirrer and a thermometer. Initial temperature of the bath was measured and then the reaction is run by a small turning on the heater wire inside the bomb. Heat released by reaction is absorbed by the bombs and tanks causing temperature rises tool. From changes in temperature and heat capacity tool that has measured the amount of heat supplied by the reaction can be calculated.Heat Capacity and Specific Heat
The properties of water which gives the definition of the origin of the calorie is the amount of temperature change experienced by water or release time to take some heat. The general term for these properties is called the heat capacity is defined as the amount of heat required to change the temperature of an object by 1 0 C.
Heat capacity is extensive, which means that the amount depending on the sample size. For example to raise the temperature of 1 g of water by 1 0 C required 4.18 J (1 cal), but to raise the temperature of 100 g of water by 1 0 C takes 100 times more energy is 418 J. So that 1 g of the sample had a heat capacity of 4.18 J / 0 C while 100 g samples of 418J / 0 C.
Intensive nature of the heat capacity is related to the type of heat (specific heat) which is defined as the amount of heat required to raise the temperature of 1 g of substance by 1 0 C. For water, its specific heat is 4.18 JG-1C-1. Most substances have a smaller specific heat than water. For example iron, only the specific heat 0.452 J g -1 0 C -1. Means less heat is needed to heat 1 g of iron by 1 0 C than water or can also mean that the amount of heat that will raise the temperature of 1 g of iron is greater than on raising the temperature of 1 g of water.
The amount of the specific heat for water due to a little influence from the sea to the weather. In winter the sea water to cool slower than the mainland so that air moves from sea to land is hotter than the air from land to sea. Likewise in the summer, the sea water is slower to heat than the mainland.
The formula:
q = m.c. Δ't
Description:
q = amount of heat (Joule)
m = mass of substance (g)
Δt = change in temperature t final - t initial)
c = heat typeCalorimetry
Calorimetry calorimetry

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 Enthalpy changes
Enthalpy = H = Heat of reaction at constant pressure = Qp
Enthalpy change is the change in energy that accompanies the event of chemical changes at a constant pressure.
a.

Termination of bonds requires energy (= endothermic)
Example: H 2 → 2H - a kJ; DH = + akJ
b.

Bond formation to give the energy (= exothermic)
Example: 2H → H 2 + a kJ; DH =-a kJ
Term used in the enthalpy change:
1.

Standard enthalpy Pembentakan (DHF):
DH animal lays forming 1 mole of compound directly from its elements are measured at 298 K and a pressure of 1 atm. Example: H 2 (g) + 1 / 2 O 2 (g) → H 2 0 (l); DHF = -285.85 kJ
2.

Enthalpy of decomposition:
DH from the decomposition of 1 mol of the compound directly into its elements (= opposite of DH formation). Example: H 2 O (l) → H 2 (g) + 1 / 2 O 2 (g) DH = +285.85 kJ
3.

Standard Enthalpy of Combustion (DHC):
DH to burn 1 mole of compounds with O 2 from the air measured at 298 K and a pressure of 1 atm. Example: CH 4 (g) + 2o 2 (g) → CO 2 (g) + 2H 2 O (l); DHC = -802 kJ
4.

Reaction enthalpy:
DH equation of a reaction in which substances contained in the equation is expressed in units of moles and the coefficients of the equation is simple round. Example: 2Al + 3H 2 SO 4 → Al 2 (SO 4) 3 + 3H 2; DH = -1468 kJ
5.

Enthalpy of Neutralization:
DH produced (always exothermic) in the reaction of neutralization of acid or base. Example: NaOH (aq) + HCl (aq) → NaCl (aq) + H 2 O (l) DH = -890.4 kJ / mol
6.

Lavoisier-Laplace law
"The amount of heat that is released to the formation of 1 mole of substance from the elements-unsurya = amount of heat required to decompose the substance into its constituent elements."
Meaning: If the reaction is reversed then the sign of heat that is formed is also reversed from positive to negative or vice versa Example:
N 2 (g) + 3H 2 (g) → 2NH 3 (g) DH = - 112 kJ
2NH 3 (g) → N 2 (g) + 3H 2 (g) DH = + 112 kJEnthalpy of formation, combustion and decomposition
Thermochemical data are generally set at a temperature of 25 0 C and a pressure of 1 atm, hereinafter referred to standard conditions. Enthalpy changes are measured at a temperature of 25 0 C and a pressure of 1 atm is called the standard enthalpy change and is expressed with the symbol Δ H 0 or ΔH298. While the changes in enthalpy measurement does not refer to the conditions of measurement represented by the symbol ΔH alone.
Is the molar enthalpy of the reaction enthalpy change associated with the quantity of substance involved in the reaction. In the known thermochemical various molar enthalpy, such as the enthalpy of formation, decomposition enthalpy, and enthalpy of combustion.The formation enthalpy
There is a range of important thermochemical equation associated with the formation of one mole of a compound of unsurunsurnya. Enthalpy changes associated with this reaction is called the heat of formation or enthalpy of formation is given the symbol ΔH f. For example thermochemical equation for the formation of water and steam at 100 0 C and 1 atm respectively.
RM1
How can we use this equation to obtain the heat of vaporization of water? Clearly equation (1) should be behind us, and then summed with equation (2). Do not forget to change the sign of ΔH. (If the formation of H 2 O (l) exothermic, as reflected by a negative ΔH f, the reverse process must be endothermic) which means a positive exothermic which means to be endothermic.Exothermic
Exothermic (heat producing) exothermic (heat producing)
RM2Endothermic
rm311
When we add the equation (1) and (2), we can
rm410
And heat the reaction =
rm56
Note that the heat of reaction for all the changes together with the heat of formation of reaction products minus the heat of formation of reactants. Generally it can be written:
rm65
Price changes in reaction enthalpy can be influenced by the temperature and pressure conditions during the measurement. Therefore, the necessary conditions of temperature and pressure are to be given to any thermochemical data.Enthalpy of Combustion

The reaction of a substance with oxygen is called combustion reactions. Flammable substances that are the elements carbon, hydrogen, sulfur, and various compounds of these elements. Said to be perfect if the combustion of carbon (c) burned into CO2, hydrogen (H) burned into H2O, sulfur (S) burned to SO2.
Enthalpy changes in the complete combustion of 1 mol of a substance measured at 298 K, 1 atm is called the standard enthalpy of combustion (standard enthalpy of combustion), which is expressed by Δ Hc 0. Enthalpy of combustion is also expressed in kJ mol -1.
Price enthalpy of combustion of various substances at 298 K, 1 atm are given in Table 3 below.
Table 3. Enthalpy of combustion of various substances at 298 K, 1 atm
gb18
The burning of gasoline is an exothermic process. If gasoline is considered consisting of isooktana, C8H18 (one component of gasoline) determine the amount of heat released on combustion of 1 liter of gasoline. Known enthalpy of combustion isooktana = -5460 kJ mol -1 and a density of isooktan = 0.7 kg L -1 (H = 1; C = 12).
Answer:
Isooktana enthalpy of combustion is - 5460 kJ mol -1. Mass of 1 liter of gasoline = 1 x 0.7 kg liter L-1 = 0.7 kg = 700 grams. Mol gram/114 isooktana = 700 grams = 6.14 mol mol -1. So the heat released on combustion of 1 liter of gasoline is: 6.14 x 5460 kJ mol -1 = 33524.4 kJ mol.Perfect combustion and Not Perfect

Fuel combustion in vehicle engines or in the industry do not burn completely. Complete combustion of hydro carbon compounds (fossil fuels) to form carbon dioxide and water vapor. While imperfect combustion to form carbon monoxide and water vapor. For example:
a. Complete combustion isooktana:
C8H18 (l) +12 ½ O2 (g) -> 8 CO2 (g) + 9 H2O (g) ΔH = -5460 kJ
b. Imperfect combustion isooktana:
C8H18 (l) + 8 ½ O2 (g) -> 8 CO (g) + 9 H2O (g) ΔH = -2924.4 kJ
Arson not Perfect Impact
As shown in the example above, imperfect combustion produces less heat. Thus, imperfect combustion reduces fuel efficiency. Another disadvantage of incomplete combustion of carbon monoxide gas is produced (CO), which are toxic. Therefore, imperfect combustion will pollute the air.The decomposition enthalpy
Decomposition reaction is the reverse of reaction formation. Therefore, in accordance with the principle of conservation of energy, equal to the value of the enthalpy of the decomposition enthalpy of formation, but the opposite sign.
Example:
Known Δ Hf 0 H2O (l) = -286 kJ mol -1, the enthalpy of the decomposition of H2O (l) into hydrogen gas and oxygen gas is + 286 kJ mol -1
H2O (l) -> H2 (g) + ½ O2 (g) ΔH = + 286 kJ




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 Reaction heat and thermochemical
Relations system with the environment Relations with environmental systems
The study of heat is called a thermochemical reaction that is part of the branch of a larger science of thermodynamics. Before talking about the principle of this thermochemical we proceed, will be made once the definition of some terms. One of the terms that will be used is the system. The system is part of the universe that we are studying. It may be for example a chemical reaction that occurs in a beaker. Outside the system is the environment. In explaining a system, we must analyze its properties appropriately. Given temperature, pressure, number of moles of each substance and a liquid, solid or gas. After all these variables are determined means that all the properties of the system is certain, we have described the state of the system.
When changes occur in a system it is said that the system moves from one state to another state. When the system is isolated from the environment so that no heat can flow then changes that occur in the system is adiabatic change. During the adiabatic change, the temperature of the system will shift, when the exothermic reaction will go up while going down when the endothermic reaction. When the system was isolated from its environment, then the heat will flow between the two, so if there is a reaction, the temperature of the system can be made permanent. Changes that occur at constant temperature is called the change isotermik. It has been said, if the reaction is exothermic or endothermic then on the chemicals involved will be a change of potential energy. We measure the heat of reaction will be equal to the change of this potential energy. From now on we will use this change in some quantity that need to be enforced some rules to declare changes in general.
The symbol Δ (Greek letter for delta) is generally used to indicate a change in quantity. For example changes in temperature can be written by ΔT, where T represents temperature. In practice, usually in the show change is by reducing the final temperature with the temperature at first.
ΔT = T final - T initial
Likewise, changes in potential energy
(Ep) Δ (E.P) = EP late - early EP
From this definition obtained an agreement in algebraic sign for exothermic and endothermic changes. In the exothermic change, the potential energy of the reaction is lower than the potential energy EP reagents means lower end of the first EP. So the price ÷ EP has a negative price. Opposite to the endothermic reaction, where the price ÷ EP is positive.Exothermic and endothermic reactions
Endothermic event (right) and exothermic (left) endothermic event (right) and exothermic (left)Exothermic reaction
Exothermic reaction occurs at the heat transfer from the system into the environment or on the reaction heat is released. In the exothermic reaction ΔH = negative prices (-)
Example:
C (s) + O 2 (g) → CO 2 (g) + 393.5 kJ;
ΔH = -393.5 kJEndothermic reaction
In reaction to the heat transfer occurs from the environment to the system or to the reaction heat is needed. At the price endothermic reaction ΔH = positive (+)
Example:
CaCO 3 (s) → CaO (s) + CO 2 (g) - 178.5 kJ; ΔH = +178.5 kJ
The process of exothermic and endothermic processes The process of exothermic and endothermic process

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 Definition of thermochemical
thermochemical thermochemical can be defined as part of chemistry that studies the dynamics of chemical reactions or changes by observing thermal / thermal only. One of applied science in everyday life is a chemical reaction in our bodies where the production of energy-the energy required or issued for all the tasks that we do. Combustion of fuels like oil and coal used for electricity generation. Gasoline burned in a car engine will produce power that causes the car running. If we have a gas stove means we burn methane (the main component of natural gas) that produce heat for m gold ak. And through a sequence of reactions called metabolism, food you eat will produce the energy we need for the body to function.
Almost all chemical reactions there is always the energy is taken or removed. Let us examine the occurrence of this and how do we know of any changes in energy.
Thermochemical thermochemical events Events
Suppose we will perform a chemical reaction in an enclosed area so that no heat can escape or enter into the reaction mixture. Or the reaction is carried out in such a way that the total energy remains the same. Also suppose that the potential energy of the reaction is lower than the potential energy of reagents so that reactions occur when there is a decrease of potential energy. But this energy can not just disappear because the total energy (kinetic and potential) should remain constant. Therefore, when its potential energy falls, then the mean kinetic energy must increase the potential energy turns into kinetic energy. The addition amount of kinetic energy will cause the price of the average kinetic energy of molekulmolekul ride, which we see as a rise in temperature of the reaction mixture. The reaction mixture became hot.
Most chemical reactions are not sealed from the outside world. When the reaction mixture to heat as described below, the heat can flow around. Any changes that can release energy to the surroundings are called exothermic changes. Note that if there is an exothermic reaction, the temperature of the reaction mixture will rise and the potential energy of chemical substances in question will come down.
Sometimes chemical changes occur when there is increased potential energy of the substances concerned. When this happens, then the kinetic energy would go down so that its temperature is also down. When the system is not closed in around him, the heat can flow into the reaction mixture and the change is called endothermic change. Note that if there is an endothermic reaction, the temperature of the reaction mixture will decrease and the potential energy of the substances involved in the reaction will go up.
Event of a fire producing heat the fire produces heat eventsMeasurement of Energy in Chemical Reactions
Standard international unit for energy is Joule (J) derived from the kinetic energy. One joule = 1 kgm 2 / s 2. Equivalent to the amount of energy that belongs to an object with a mass of 2 kg and a speed of 1 m / sec (if in units of English, the object with a mass of 4.4 lb and the speed of 197 ft / min or 2.2 mile / hour).
1 J = 1 kg m 2 / s 2
Smaller unit of energy used in physics called the ERG that cost = 1 × 10 -7 J. In referring to the energy involved in the reaction between the reactant molecule size is usually replaced with larger units of kilojoules (kJ). One kilojoules = 1000 joules (1 kJ = 1000J).
All forms of energy can be converted entirely into heat and when a chemist measures the energy, usually in the form of heat. The usual way is used to express heat is called calorie (cal abbreviation). The definition is derived from the effect of heat on the temperature of the object. At first calorie is defined as the amount of heat required to raise the temperature of 1 gram of water with temperatures from 15 0 C for 1 0 C. Kilocalories (kcal) as well as kilojoules are units that are better suited to express the energy changes in chemical reactions. The unit is also used to indicate kilocalories of energy contained in food.
Upon the acceptance of the SI, is now also joules (or kilojoules) and calories preferably redefined in SI units. Now calories and kilocalories are defined exactly as follows:
1 cal = 4.184 J
1 kcal = 4.184 kJ

Chemical Reaction

What is a Chemical Reaction?

A chemical reaction is a process in which the identity of at least one substance changes. A chemical equation represents the total chemical change that occurs in a chemical reaction using symbols and chemical formulas for the substances involved. Reactants are the substances that are changed and products are the substances that are produced in a chemical reaction.

The general format for writing a chemical equation is

reactant₁ + reactant₂ + … → product₁ + product₂ + …

With the exception of nuclear reactions, the Law of Conservation of Mass–matter is neither created nor destroyed during a chemical reaction– is obeyed in “ordinary” chemical reactions. For this reason a chemical equation must be balanced–the number of atoms of each element must be the same on the reactants side of the reaction arrow as on the products side. Details on balancing chemical equations are found in the units on Stoichiometry and Redox Reactions.

The general format for writing a chemical equation can be written in a short-hand version as

a A + b B + … → c C + d D + …

where the lower case letters are the stoichiometric coefficients needed to balance a specific equation.

The units on Stoichiometry, Redox Reactions, and Acid-Base Chemistry contain additional background reading, example problems, and information on the topics covered in this unit.

Types of Chemical Reactions

Chemists classify chemical reactions in various ways. Often a major classification is based on whether or not the reaction involves oxidation-reduction. A reaction may be classified as redox in which oxidation and reduction occur or nonredox in which there is no oxidation and reduction occurring.

Redox Reactions

A redox reaction can be recognized by observing whether or not the oxidation numbers of any of the elements change during the reaction.

Example Problem: Classify the reactions as either redox or nonredox.

(1) 4 Fe(s) + 3 O₂(g) → 2 Fe₂O₃(s)

(2) NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

(3) Cl₂(g) + H₂O(l) → HCl(aq) + HClO(aq)

Answer: In equation (1), the iron changes oxidation numbers from 0 to +3 and oxygen changes from 0 to -2. Equation (1) represents a redox reaction. In equation (2), there is no change in oxidation numbers for the elements involved: sodium is +1, oxygen is -2, hydrogen is +1, and chlorine is -1 on both the reactants and products sides. Equation (2) represents a nonredox reaction. In equation (3), the chlorine changes from 0 to -1 in HCl and to +1 in HClO. There is no change in the oxidation numbers of hydrogen (+1 in H₂O, HCl, and HClO) and oxygen (-2 in H₂O and HClO). Because chlorine is oxidized and reduced, equation (3) represents a redox reaction.

The reaction described in equation (3) is interesting in that an element in one oxidation state undergoes both oxidation and reduction. Such a redox process is known as a disproportionation reaction. The element undergoing disproportionation must have at least three different oxidation states–the initial one in the reactant and one higher plus one lower in the products.

Most simple redox reactions may be classified as combination, decomposition, or single displacement reactions. In a combinations reaction two reactants react to give a single product. The general format of the chemical equation is

a A + b B + … → c C

A special case of a combination reaction in which the reactants are only elements in their naturally occurring forms and physical states at the temperature and pressure of the reaction is known as a formation reaction.

Example Problem: Identify which reactions are redox combination reactions.

(4) 6 Li₂O(s) + P₄O₁₀(g) → 4 Li₃PO₄(s)

(5) CaO(s) + H₂O(l) → Ca(OH)₂(s)

(6) S(s) + 3 F₂(g) → SF₆(g)

(7) ZnS(s) + 2 O₂(g) → ZnSO₄(s)

(8) SO₂(g) + Cl₂(g) → SO₂Cl₂(g)

Answer: All of the reaction are classified as combination reactions because they involved two or more reactants producing a single product. However, redox is occurring only in equations (6), (7), and (8). In equation (6), S is oxidized from 0 to +6 and F is reduced from 0 to -1; in equation (7), S is oxidized from -2 to +6 and O is reduced from 0 to -2; and in equation (8), S is oxidized from +4 to +6 and Cl is reduced from 0 to -1.

In a decomposition reaction a single reactant breaks down to give two or more substances. The general format of the chemical equation is

a A → b B + c C + …

If the decomposition reaction involves oxidation-reduction, the reaction is often called an internal redox reaction because the oxidized and reduced elements originate in the same compound.

Example Problem: Identify which reactions are redox decomposition reactions.

(9) CuSO₄⋅5H₂O(s) → CuSO₄(s) + 5 H₂O(g)

(10) SnCl₄⋅6H₂O(s) → SnO₂(s) + 4 HCl(g) + 4 H₂O(g)

(11) NH₄NO₂(s) → N₂(g) + 2 H₂O(g)

(12) (NH₄)₂Cr₂O₇(s) → N₂(g) + Cr₂O₃(s) + 4 H₂O(g)

Answer: All of the reactions are classified as decomposition reactions because they involve a single reactant producing two or more substances. However, redox is occurring only in equations (11) and (12). In equation (11), the N in NH₄⁺ is oxidized from -3 to 0 and the N in NO₂^- is reduced from +3 to 0. In equation (12), the N is oxidized from -3 to 0 and the Cr is reduced from +6 to +3.

In a single displacement reaction the atoms or ions of one reactant replace the atoms or ions in another reactant. Single displacement reactions are also known as displacement, single replacement, and replacement reactions. The general format of the chemical equation is

a A + b BC → c AC + d B

Whether or not a redox single displacement reaction occurs will depend on the relative reducing strengths of A and B.

Example Problem: Identify which reactions are redox single displacement reactions.

(13) 2 Al(s) + Fe₂O₃(s) → 2 Fe(s) + Al₂O₃(s)

(14) 2 NaI(aq) + Br₂(aq) → 2 NaBr(aq) + I₂(aq)

(15) Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Answer: All three reactions are redox. Both equations (13) and (14) fit the general format of the single displacement reaction by assigning A as Al, B as Fe, and C as O in equation (13) and A as Br, B as I, and C as Na in equation (14). To classify equation (15) is a little more difficult. The reaction has been represented by a net ionic equation in which the anion has been omitted. If an anion X is added to generate the overall equation, Zn(s) + CuX(aq) → ZnX(aq) + Cu(s), then assigning A as Zn, B as Cu, and C as X shows that this is also a redox single displacement reaction.

In addition to the single redox reactions described above, a redox reaction may be classified as a simple redox electron transfer reaction in which the oxidation numbers of ionic reactants are changed by the direct transfer of electrons from one ion to the other–typically in aqueous solutions. For example

2 Fe³⁺(aq) + Sn²⁺(aq) → 2 Fe²⁺(aq) + Sn⁴⁺(aq)

Many redox reactions do not fit into the classifications described above. For example, redox reactions involving oxygen-containing reactants in aqueous acidic or basic solutions such as

3 Cu(s) + 8 HNO₃(aq, dil) → 3 Cu(NO₃)₂(aq) + 2 NO(g) + 4 H₂O(l)

Cu(s) + 4 HNO₃(aq, conc) → Cu(NO₃)₂(aq) + 2 NO₂(g) + 2 H₂O(l)

or the combustion of oxygen with more than one element in a reactant

2 CH₃OH(g) + 3 O₂(g) → 2 CO₂(g) + 4 H₂O(l)

These types of reactions are classified as complex redox reactions.

Nonredox Reactions

There are several classifications of nonredox reactions–including combination, decomposition, single displacement, and double displacement reactions.

The general format of the chemical equation for a nonredox combination reaction is the same as for a redox combination reaction

a A + b B + … → c C

However, all reactants and the product must be compounds and no changes in oxidation numbers of the elements occur. Usually these reactions involve reactants that are acidic and basic anhydrides.

Example Problem: Identify which reactions are nonredox combination reactions.

(16) 2 Na(s) + Cl₂(g) → 2 NaCl(s)

(17) SO₃(g) + CaO(s) → CaSO₄(s)

(18) SO₂(g) + H₂O(l) → H₂SO₃(aq)

Answer: All three equations are combination reactions, but only equations (17) and (18) are nonredox.

The general format of the chemical equation for a nonredox decomposition reaction is the same as for a redox decomposition reaction

a A → b B + c C + …

However, the reactant and all products must be compounds and no changes in oxidation numbers occur. Quite often one of the products formed will be a gas.

Example Problem: Identify which reactions are nonredox decomposition reactions.

(19) NH₄HCO₃(s) → NH₃(g) + CO₂(g) + H₂O(g)

(20) NH₄NO₂(s) → N₂(g) + 2 H₂O(g)

(21) CaCO₃(s) → CaO(s) + CO₂(g)

Answer: All three equations are decomposition reactions, but only equations (19) and (21) are nonredox.

The general format of the chemical equation for a nonredox single displacement reaction is the same as for a redox single displacement reaction

a A + b BC → c AC + d B

However, there are no changes in the oxidation numbers of the elements during the reaction. Common nonredox single displacement reactions include ligand substitution in complexes and formation of more stable oxygen-containing compounds from less stable oxygen-containing compounds.

Example Problem: Identify which reactions are nonredox single displacement reactions.

(22) [PtCl₄]^2-(aq) + 2 NH₃(aq) → [Pt(NH₃)Cl₂](s) + 2 Cl^-(aq)

(23) Na₂CO₃(s) + SiO₂(s) → Na₂SiO₃(l) + CO₂(g)

(24) 2 AgNO₃(aq) + Cu(s) → Cu(NO₃)₂(aq) + 2 Ag(s)

Answer: All three equations are single displacement reactions, but only equations (22) and (23) are nonredox.

Finally, the last classification of nonredox reactions is that of nonredox double displacement reactions. The general format of the chemical equation is

a AC + b BD → c AD + d BC

with no oxidation or reduction of A, B, C, or D occurring. These reactions are also known as double replacement, “partner” exchange, and metathesis reactions. Usually one or more of the products will be a gas, a precipitate, a weak electrolyte, or water. An important example of a nonredox double displacement reaction is the reaction of an acid with a base under aqueous conditions.

Example Problem: Identify the nonredox double displacement reactions.

(25) CaCO₃(s) + 2 HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

(26) HCl(aq) + KOH(aq) → KCl(aq) + H₂O(l)

(27) AgNO₃(aq) + KCl(aq) → AgCl(s) + KNO₃(aq)

Answer: All three reactions are nonredox double displacement reactions.

Classifying Reactions

Most experienced chemists can classify a given chemical reaction rather easily and quickly by “inspection” of the formulas of the reactants and products in the chemical equation. The first decision most chemists make is to determine whether or not the reaction involves redox. Based on this decision, the answers to a few more specific questions will readily lead to the reaction classification. These specific questions are based on the general formats of the chemical equations for the different classifications of the reactions described above.

To begin learning what these questions are, you might consider using the online analysis program that is available at http://www.xxx.yyy to classify the reactions given in the Example Problem. The basis of this analysis program is outlined by the flow charts given in Figures (1) and (2). Please do not memorize this flow chart–it is simply a tool to help you learn what to look for and what questions should be asked as you classify a given reaction.

Figure 1. Flow chart of questions to classify nonredox reactions. (Used by permission ...)

Figure 2. Flow chart of questions to classify redox reactions. (Used by permission...)

Questions:

1. Why are chemical reactions that occur in the thermodynamic sometimes does not happen without the help of an enzyme?

     2.   what are the benefits to the health of a redox reaction?