Chemical equations examples

Examples 1,1, and 2 in H2 + Br2 - 2 HBr. stoichiometric point The stage in a titration when exactly the right volume of solution needed to complete the reaction has been added, stoichiometric proportions Reactants in the same proportions as their coefficients in the chemical equation. Example equal amounts of H2 and Br2 in the formation of HBr. 

The first term on the right is the total standard molar entropy of the products and the second term is that of the reactants n denotes the various stoichiometric coefficients in the chemical equation. Example 7.7 illustrates how this expression is used. 

Given a mechanism for a reaction, obtain the overall chemical equation. (EXAMPLE 14.8)... 

We make the conversions between solution volumes and amounts of solute in moles using the molarities of the solutions. We make the conversions between amounts in moles of A and B using the stoichiometric coefficients from the balanced chemical equation. Example 4.8 demonstrates solution stoichiometry. 

Stability, Bifurcations, Limit Cycles Some aspects of this subject involve the solution of nonlinear equations other aspects involve the integration of ordinaiy differential equations apphcations include chaos and fractals as well as unusual operation of some chemical engineering eqmpment. Ref. 176 gives an excellent introduction to the subject and the details needed to apply the methods. Ref. 66 gives more details of the algorithms. A concise survey with some chemical engineering examples is given in Ref. 91. Bifurcation results are closely connected with stabihty of the steady states, which is essentially a transient phenomenon. 

In liquid metal solutions Z is normally of the order of 10, and so this equation gives values of Ks(a+B) which are close to that predicted by the random solution equation. But if it is assumed that the solute atom, for example oxygen, has a significantly lower co-ordination number of metallic atoms than is found in the bulk of die alloy, dieii Z in the ratio of the activity coefficients of die solutes in the quasi-chemical equation above must be correspondingly decreased to the appropriate value. For example, Jacobs and Alcock (1972) showed that much of the experimental data for oxygen solutions in biiiaty liquid metal alloys could be accounted for by the assumption that die oxygen atom is four co-ordinated in diese solutions. 

This rule allows you to find AH corresponding to any desired amount of reactant or product. To do this, you follow the conversion-factor approach used in Chapter 3 with ordinary chemical equations. Consider, for example,... 

We are so used to free air that we do not think of oxygen as an important chemical. For example, we buy natural gas (mostly methane) as a fuel and burn it in air to furnish heat to us. If methane were free in the air and oxygen were scarce so that we had to purchase it we would consider oxygen to be the fuel. In either case, the amount of heat released would be that for the reaction represented by the following equation. 

Because atoms are neither created nor destroyed, chemists regard each elemental symbol as representing one atom of the element (with the subscripts giving the number of each type of atom in a formula) and then multiply formulas by factors to show the same numbers of atoms of each element on both sides of the arrow. The resulting expression is said to be balanced and is called a chemical equation. For example, there are two H atoms on the left of the preceding skeletal equation but three H atoms on the right. So, we rewrite the expression as... 

Now there are four H atoms, two Na atoms, and two O atoms on each side, and the equation conforms to the law of conservation of mass. The number multiplying an entire chemical formula in a chemical equation (for example, the 2 multiplying H20) is called the stoichiometric coefficient of the substance. A coefficient of 1 (as for H2) is not written explicitly. 

Although normally the coefficients in a balanced chemical equation are the smallest possible whole numbers, a chemical equation can be multiplied through by a factor and still be a valid equation. At times it is convenient to use fractional coefficients for example, we could write... 

Sometimes we need to construct a balanced chemical equation from the description of a reaction. For example, methane, CH4, is the principal ingredient of natural gas (Fig. H.3). It burns in oxygen to form carbon dioxide and water, both formed initially as gases. To write the balanced equation for the reaction, we first write the skeletal equation ... 

Write, balance, and label a chemical equation on the basis of information given in sentence form (Example H.l). 

Some redox reactions, particularly those involving oxoanions, have complex chemical equations that require special balancing procedures. We meet examples and see how to balance them in Chapter 12. 

STRATEGY We write the chemical equation for the formation of HI(g) and calculate the standard Gibbs free energy of reaction from AG° = AH° — TAS°. It is best to write the equation with a stoichiometric coefficient of 1 for the compound of interest, because then AG° = AGf°. The standard enthalpy of formation is found in Appendix 2A. The standard reaction entropy is found as shown in Example 7.9, by using the data from Table 7.3 or Appendix 2A. 

A note on good practice As these examples have shown, it is important to specify the chemical equation to which the equilibrium constant applies. 

If a chemical equation can be expressed as the sum of two or more chemical equations, the equilibrium constant for the overall reaction is the product of the equilibrium constants for the component reactions. For example, consider the three gas-phase reactions... 

However, it is best to use the full chemical equation when working with titrations to ensure the correct stoichiometry. For example, if hydrochloric acid is used to neutralize Ca(OH)2, we must take into account the fact that each formula unit of Ca(OH)2 provides two OH ions ... 

We consider oxidation first. To show the removal of electrons from a species that is being oxidized in a redox reaction, we write the chemical equation for an oxidation half-reaction. A half-reaction is the oxidation or reduction part of the reaction considered alone. For example, one battery that Volta built used silver and zinc plates to carry out the reaction... 

Balancing the chemical equation for a redox reaction by inspection can be a real challenge, especially for one taking place in aqueous solution, when water may participate and we must include HzO and either H+ or OH. In such cases, it is easier to simplify the equation by separating it into its reduction and oxidation half-reactions, balance the half-reactions separately, and then add them together to obtain the balanced equation for the overall reaction. When adding the equations for half-reactions, we match the number of electrons released by oxidation with the number used in reduction, because electrons are neither created nor destroyed in chemical reactions. The procedure is outlined in Toolbox 12.1 and illustrated in Examples 12.1 and 12.2. 

Balance chemical equations for redox reactions by the halfreaction method (Toolbox 12.1 and Examples 12.1 and 12.2). 

Write the chemical equation for a cell reaction, given the cell diagram (Toolbox 12.2 and Example 12.4). 

The interhalogen IFT can be made only by indirect routes. For example, xenon difluoride gas can react with iodine gas to produce 1FV and xenon gas. In one experiment xenon difluoride is introduced into a rigid container until a pressure of 3.6 atm is reached. Iodine vapor is then introduced until the total pressure is 7.2 atm. Reaction is then allowed to proceed at constant temperature until completion by solidifying the IF as it is produced. The final pressure in the flask due to the xenon and excess iodine vapor is 6.0 atm. (a) What is the formula of the mterhalogen (b) Write the chemical equation for its formation. 

A reaction mechanism is a series of simple molecular processes, such as the Zeldovich mechanism, that lead to the formation of the product. As with the empirical rate law, the reaction mechanism must be determined experimentally. The process of assembling individual molecular steps to describe complex reactions has probably enjoyed its greatest success for gas phase reactions in the atmosphere. In the condensed phase, molecules spend a substantial fraction of the time in association with other molecules and it has proved difficult to characterize these associations. Once the mecharrism is known, however, the rate law can be determined directly from the chemical equations for the individual molecular steps. Several examples are given below. 

These examples suggest that the learner is not always aware which aspects of our use of symbolic representation in chemistiy are intended to be significant. This is an area where further work would be useful, as clearly teachers need to do more to induct learners into the intended symbolism we use in teaching the subject. In the next section, these issues will be explored further in the particular context of learning about chemical equations. 

Consider the examples of some of the forms of chemical equations (and related representations) met in school and college (i.e. middle and senior high school) science and chemistiy classes that are shown in Table 4.1. For the purposes of this chapter half-equations (Example 11) and symbolic representations of processes such as ionisation (Example 10) will be included under the generic heading of chemical equations . Table 4.1 does not include examples of chemical reactions and reaction schemes that include stmctural formulae, as are commonly nsed in organic chemistiy. 

Table 4.1 Some examples of types of chemical equations met in learning chemistry 2Mg + O2 — 2MgO...

Chemical eqrrations are used to represent chemical processes such as chemical reactions. A key feature then of a chemical equation is that it has two parts, representing before and after the process, separated by an arrow or other signifier of the process itself. Each of the examples in Table 4.1 has this stractiue. 

Chemical equations may be presented both as words and formulae. Consider the second example in Table 4.1 ... 

Other examples in Table 4.1 offer additional complications. Example 11 provides an electrode potential, which is related to, but not the same as, an energy change. Example 10 represents an endothemtic process (AH > 0). Both these examples include terms that should not be found in chemical equations representing reactions. 

Just as mass and energy must be conserved, so also must electrical charge. Yet free electrons are not found stable in nature under the conditions of chentistry on earth, so caimot appear as reactants or products in representations of chemical reactions. Example 11 is a half-equation , something that represents a common pattern in chemical reactions, but only occms when coupled to another suitable half-equation (i.e., this reduction process must be paired with an oxidation process that releases electrons), e.g. 

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