You have already learned about chemical formulas for elements and compounds. You have also learned how to balance chemical equations. Finally, you have some familiarity with the basic reaction types listed below:
In this activity you will review these reaction types and study some examples. After some instruction into the details of how the reactions work you will practice predicting products of reactions. In this way you will come to a deeper understanding of the classification of chemical reactions.
In a synthesis reaction two reactants combine to form one product:
Some example synthesis reactions involve only non-metals. These reactions are hard to predict. This is because there are often multiple possible combinations. Still, a few synthesis reactions involving non-metals are frequently used and easy to memorize. Consider the following reactions to be some you should know by heart:
Reactions between metals and non-metals are much easier to predict. As you know, the elements in the periodic table are arranged in groups with similar characteristics. Elements in group 1 (Li, Na, K, etc.) are written as pure elements using their atomic symbol. When any element in this group forms an ion by reacting with a non-metal it always has a +1 charge: Li+1, Na+1, K+1. Memorize these predictable charges.
The predictability of the charges adopted by elements when they react is what makes predicting the products of the reaction between metals and non-metals relatively simple. For example, Mg becomes Mg+2 in compounds and O becomes O–2:
For reactions between a metal and a non-metal, here is how it works:
For example, consider a reaction between potassium and chlorine:
Note that ionic compounds are always solids in these reactions.
In a decomposition reaction one reactant forms two or more products:
Some decomposition reactions are easy to predict. They simply take a compound and break it down into the component elements. They are simply the reverse of synthesis reactions as discussed above. Here are some examples:
Predicting the products of the decomposition of simple ionic compounds of two elements requires only that you write the correct formulas of both elements as products before you balance the chemical equation. As you do so, remember that the subscripts in the formulas of the compound being decomposed depend on the charges of the ions in ionic compounds, not on the molecular formula of the component elements. Sodium chloride (NaCl) does not form Na + Cl because the correct formula for chlorine is Cl2, not Cl.
Other decomposition reactions follow a specific pattern based on the polyatomic ion that is part of the compound being broken down. Compounds containing the carbonate ion (CO32–) break down to form an oxide compound and carbon dioxide. Compounds containing the chlorate ion (ClO3–) break down to form a chloride and oxygen gas. For example:
Predicting the products for these decomposition reactions requires that you substitute other metals in for Na or Ca as used in the examples.
In a single replacement reaction one reactant is an element, and the other is a compound. In the reaction the pure element becomes an ion and combines with an ion from the compound. One of the elements in the compound in turn becomes a pure element (as neutral atoms).
One common type of single replacement reaction involves a solid metal (with neutral atoms) that reacts with an aqueous solution of an ionic compound which includes a metal ion. The metal ion gains electrons to become neutral. The electrons come from the neutral metal atoms, which lose electrons to become an ion. Note that there is always another ion which is unchanged by the reaction. In the first reaction below, the nitrate ion (NO3–) switches from being part of a compound with the silver ion to being part of a compound with the copper ion.
To predict the products of a single replacement reaction, several things are required:
For example, consider a reaction between magnesium and aluminum nitrate:
In a double replacement reaction both reactants are aqueous ionic compounds. The ionic compounds exchange ions to form two new compounds.
There are two types of double replacement reactions that you need to know. One of these is called a precipitation reaction. Some combinations of ions are insoluble in water so when these ions meet, a solid forms and falls to the bottom of the liquid, like precipitation from the sky. The second is an acid-base reaction. Note the states of matter for these examples, which are both precipitation reactions:
Some ions mostly have compounds which are insoluble in water while others are always or nearly always soluble, no matter what ion they are with in a compound. Any ion that is Always Soluble will be soluble even if it is combined with an ion labeled as Always Insoluble. Ions soluble with exceptions are only insoluble with the cations shown.
|Cl–1, Br–1, I–1||SO4–2||
|exceptions: Ag+1, Pb+2, Hg2+2||
exceptions: Sr+2, Ba+2,
To predict the products of a double replacement precipitation reaction, several things are required:
For example, consider a reaction between sodium phosphate and calcium nitrate:
Another type of double replacement reaction is the acid-base reaction. An acid, which always contains hydrogen ions (H+), reacts with a base, which always includes hydroxide ions (OH–), to make a salt and water. The word ‘salt’ here means any ionic compound, not just sodium chloride. Here is an example reaction:
To predict the products of an acid-base reaction, several things are required:
In a combustion reaction a hydrocarbon reacts with oxygen to make carbon dioxide and water:
Sometimes the hydrocarbon molecule includes one or more oxygen atoms but the reaction is still the same:
There is really no prediction involved in writing combustion reaction equations. The second reactant and both products are always the same. Just memorize that oxygen is a reactant and that carbon dioxide and water are products. Here are a few examples of balanced combustion reaction equations:
Some fuels, such as natural gas and propane, are gases. Other fuels are usually kept in liquid form, for example, alcohol, gasoline, and diesel fuel. Although we may write them in a chemical equation as a liquid, they all actually burn in the vapor form. Even such fuels as candle wax, which is a solid at room temperature, actually burn in the vapor phase. In the case of a candle, the heat of the flame melts and then vaporizes the wax before the wax is burned in the flame.
Answer the following questions.