For this, we use solubility rules , which are general statements that predict which ionic compounds dissolve are soluble and which do not are not soluble or insoluble. Table 4. We need to consider each ionic compound both the reactants and the possible products in light of the solubility rules in Table 4. If a compound is soluble, we use the aq label with it, indicating it dissolves. If a compound is not soluble, we use the s label with it and assume that it will precipitate out of solution.
If everything is soluble, then no reaction will be expected. Are these soluble? NaCl is by the same rule we just quoted , but what about SrSO 4?
Therefore, we expect a reaction to occur, and the balanced chemical equation would be. You would expect to see a visual change corresponding to SrSO 4 precipitating out of solution Figure 4. Some double-replacement reactions are obvious because you can see a solid precipitate coming out of solution.
Will a double-replacement reaction occur? If we assume that a double-replacement reaction may occur, we need to consider the possible products, which would be NaCl and Fe OH 2.
NaCl is soluble, but, according to the solubility rules, Fe OH 2 is not. Therefore, a reaction would occur, and Fe OH 2 s would precipitate out of solution.
The balanced chemical equation is. Will a double-replacement equation occur? What are the general characteristics that help you recognize single-replacement reactions? What are the general characteristics that help you recognize double-replacement reactions? Assuming that each single-replacement reaction occurs, predict the products and write each balanced chemical equation. Use the periodic table or the activity series to predict if each single-replacement reaction will occur and, if so, write a balanced chemical equation.
Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation. Use the solubility rules to predict if each double-replacement reaction will occur and, if so, write a balanced chemical equation.
Then write a balanced equation for each reaction. Be sure to include the states of all compounds in your equations solid, liquid, aqueous, or gas. If no reaction occurs write the words "no reaction" or NR instead of the products in your balanced equation and indicate why your think there was no reaction. Unless otherwise indicated dispose of all waste in the waste container provided.
Do not put metal strips in the sink. Use 1 mL of each solution unless otherwise specified. For reactions involving metals, use just one piece of metal.
Do not put the metal pieces in the sink. If no discernable initial change is noted, let the reaction mixture stand for at least five to ten minutes before observing again. Not all of the combinations will yield observable reactions. Repeat the reaction if there is any doubt about whether a reaction occurred or not. Be sure to mix the solutions well. Now use the above results to write products for the reactions below. Write NR if no reaction is expected.
Record your observations on these data pages as you perform each reaction. Write a balanced formula equation with state labels for each reaction. If no reaction occurs, follow the instructions in the Procedure. In addition to providing observations and an equation for each reaction, use your results to determine the relative activities of the two elements involved in each reaction. Arrange copper, silver, calcium, zinc, and hydrogen in an activity series from most active to least active on the basis of the results from the displacement reactions that you performed.
Recall that a more active metal displaces a less active metal, a more active metal to is needed to displace hydrogen from water than to displace it from an acid, and that a metal that displaces hydrogen from acid is ranked as more active than hydrogen. Reaction Types Please note that some reactions may be classified as more than one type of reaction; for example, combination and decomposition reactions that involve elemental substances are also oxidation-reduction reactions.
None Nitrates and Acetates are soluble. None Chlorides, Bromides and Iodides are soluble. Sulfates are soluble. Carbonates, Chromates and Phosphates are insoluble. Sulfides are insoluble. Hydroxides are insoluble. Do not touch metals with your hands. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases.
Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners Figure 4.
When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:. The chemical reactions described in which acids and bases dissolved in water produce hydronium and hydroxide ions, respectively, are, by definition, acid-base reactions.
In these reactions, water serves as both a solvent and a reactant. A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base, the products are often a salt and water, and neither reactant is the water itself:. To illustrate a neutralization reaction, consider what happens when a typical antacid such as milk of magnesia an aqueous suspension of solid Mg OH 2 is ingested to ease symptoms associated with excess stomach acid HCl :.
Note that in addition to water, this reaction produces a salt, magnesium chloride. Writing Equations for Acid-Base Reactions Write balanced chemical equations for the acid-base reactions described here:. A double-arrow is appropriate in this equation because it indicates the HOCl is a weak acid that has not reacted completely. Check Your Learning Write the net ionic equation representing the neutralization of any strong acid with an ionic hydroxide. Hint: Consider the ions produced when a strong acid is dissolved in water.
Explore the microscopic view of strong and weak acids and bases. The term oxidation was originally used to describe chemical reactions involving O 2 , but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction redox reactions. A few examples of such reactions will be used to develop a clear picture of this classification. Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride:.
It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction :. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur:. In this reaction, then, sodium is oxidized and chlorine undergoes reduction. Viewed from a more active perspective, sodium functions as a reducing agent reductant , since it provides electrons to or reduces chlorine.
Likewise, chlorine functions as an oxidizing agent oxidant , as it effectively removes electrons from oxidizes sodium. Some redox processes, however, do not involve the transfer of electrons.
Consider, for example, a reaction similar to the one yielding NaCl:. Note: The proper convention for reporting charge is to write the number first, followed by the sign e.
This convention aims to emphasize the distinction between these two related properties. Assigning Oxidation Numbers Follow the guidelines in this section of the text to assign oxidation numbers to all the elements in the following species:.
Assuming the usual oxidation number for oxygen -2 per guideline 3 , the oxidation number for sulfur is calculated as directed by guideline Check Your Learning Assign oxidation states to the elements whose atoms are underlined in each of the following compounds or ions:.
Using the oxidation number concept, an all-inclusive definition of redox reaction has been established. Oxidation-reduction redox reactions are those in which one or more elements involved undergo a change in oxidation number. While the vast majority of redox reactions involve changes in oxidation number for two or more elements, a few interesting exceptions to this rule do exist Example 4.
Definitions for the complementary processes of this reaction class are correspondingly revised as shown here:. Returning to the reactions used to introduce this topic, they may now both be identified as redox processes. Several subclasses of redox reactions are recognized, including combustion reactions in which the reductant also called a fuel and oxidant often, but not necessarily, molecular oxygen react vigorously and produce significant amounts of heat, and often light, in the form of a flame.
Solid rocket-fuel reactions such as the one depicted in Figure 1 in Chapter 4 Introduction are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation:.
Watch a brief video showing the test firing of a small-scale, prototype, hybrid rocket engine planned for use in the new Space Launch System being developed by NASA. The first engines firing at. Single-displacement replacement reactions are redox reactions in which an ion in solution is displaced or replaced via the oxidation of a metallic element.
Two hydrogen atoms must therefore be reduced a decrease in oxidation number and the two electrons required for the reduction must come from the iron. The final product is therefore most likely iron II sulfate.
Decomposition reactions are the most difficult to predict, but there are some general trends that are useful. For example, most metal carbonates will decompose on heating to yield the metal oxide and carbon dioxide. Metal hydrogen carbonates also decompose on heating to give the metal carbonate , carbon dioxide and water.
Identifying these compounds and building an understanding of why and how they are formed is one of the challenges of chemistry. Some examples:. The potential products in double-replacement reactions are simple to predict; the anions and cations simply exchange.
Remember, however, that one of the products must precipitate , otherwise no chemical reaction has occurred.
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