Use a different order? 10-4 Decomposition, 10-1 Combination, 10-2 Aqueous ions, 10-5 hydrates, 10-3 acids and bases?

Chapter 10 Chemical Reactivity Now that we understand how molecules are made, let s move on to the real fun part of chemistry, how do the molecules react with each other Use a different order? 10-4 Decomposition, 10-1 Combination, 10-2 Aqueous ions, 10-5 hydrates, 10-3 acids and bases? 10-1 Combination reactions A combination reaction occurs when two reactants combine to form a single product. Here are a couple of different combination reactions where we have two reactants and one product: 4Al(s) + 3O 2(g) 2Al2O 3 (s) 4Fe(s) + 3O (g) 2Fe O (s) 2 2 3 Both of these reactions are examples of combustion reactions, where one substance (a fuel) is burned in the presence of some oxidizer (Often O ) Anybody recognize the second reaction? Iron rusting? We have to coat irons surface is some way to prevent iron from rusting and the metal from flaking away as Fe2O 3. Anybody know why we don t do the same for aluminum? The Fe2O 3 flakes away from the Fe surface, exposing allowing the Fe underneath to also rust. Al2O 3 does not flake away, In fact I tightly adheres to the Al underneath so the process stops! In the above cases the product is a metal and a nonmetal, so it is an ionic compound. What do you remember about the physical properties of ionic compounds? Solids with high melting points We can do a similar combination combustion reaction with nonmetals: C(s) + O 2(g) CO 2(g) S(s) + O (g) SO (g) 2 2 Here our products are covalent compounds. Note that these covalent compounds are gases instead of solids. We have a lot more variability in the physical states of covalent compounds they can be solids or liquids are gases at RT. So far all of our reactants have been elements. We can also have combination reactions with molecules: Na2O(s) + CO 2(g) Na2CO 3(s) NH (g) + HCl(g) NH Cl(s) 3 4 2

2 10-2 Nomenclature of Polyatomic Ions So our reactions have involved only solids, liquids or gases. Many reactions involve chemicals dissolved in water or the aqueous (aq) state. We need to discuss what the aqueous state really is before we look at any reaction that occurs in water. What happens when compounds dissolve in water First of all, there is a big difference in what happens to ionic compounds and covalent compounds when they dissolve in water. Remember the demonstration I did back at the beginning of chapter 6, when I used light bulb to demonstrate that Ionic compounds dissolved in water carry current, but covalent compounds don t? Lets explain that on a molecular level. First question. Will all covalent compounds dissolve in water? If not, which one do and which ones don t? (Polar do, non-polar don t) Why do polar compounds dissolve? Both water and polar covalents have a dipole, and the dipole of one molecule interacts with the dipole of the other, so the molecules spread out to maximize their interactions Figure 10.4. ( We will study this more in chapter 16) What ions do is different. Remember that an ionic solid is held together by alternating + and - charges. When exposed to water the water dipole interacts with these charges. The negative ends of the dipoles interact with the + ions, and the positive ends of the dipoles interact with the - ions as shown in figure 10.5 This explains why many ionic solids that are so strong can completely dissolve in water at RT. And why the water can now carry a current. This is actually a chemical reaction: + + The Na (aq) represents the Na ion surrounded by water molecules. This - complex is called a solvated ion. Ditto for Cl (aq) You can also write the reaction as: But I prefer the first way because it emphasizes the true chemical species in solution.

Again we can ask the same question we did with the covalent compounds. Does this happen for all ionic compounds? And again the answer is no. Some ionic compounds have such a strong attraction between the ions that they won t break apart in water and, instead, stay as a solid. This is a very important point. The ionic compounds that dissolve to give you a clear solution can only do this when water interacts so strongly with the ions that it can literally tear the ionic solid apart into millions and millions of aqueous ions interacting with the water molecules around them. On the other hand there are lots of ionic compounds that have such a strong ionic lattice that water cannot break the compound into individual ions. If an ionic compound doesn t turn into aqueous ions, then it remains as a solid and does not dissolve. Covalent ions We have now talked about covalent compounds and ionic compounds, What about covalent ions? When we were doing Lewis formula s for covalent 2- compounds, I occasionally threw in some covalent ions like NH 4+ or SO 4. What happens to these molecules in water? When covalent ions interact with water ALL COVALENT BONDS REMAIN INTACT. The covalent ions does NOT decompose into individual atomic ions 3 This is shown graphically in figure 10.6 with MgSO 4 There are lots of these polyatomic ions in chemistry, and you need to know the names and formula s for many of them. Table 10.1 You must memorize the formula and name of the following polyatomic ions: hydroxide, cyanide, acetate, nitrite, nitrate, chlorate, perchlorate, peroxide, carbonate, sulfite, sulfate, oxalate, phosphate, ammonium and mercury(i) The book goes into the naming system that is used to come up with some of these names. As far as I am concerned, just memorize the names and don t worry about where they came from. (Unless this is something you want to know) Practice Problems: What is the correct name for NH4CH3COO? (Ammonium acetate) What is the correct formula for sodium sulfite? (Na SO ) 2 3

4 Clicker questions: Name and formulas of some salts made with complex ions 10-3 Bases and Acids Bases One definition for a base is that it is a compound that yields hydroxide ions (OH - ) when dissolved in water Many (but not all) metal oxides make bases when they undergo a combination reaction with water, Thus: Metal oxides are generally basic when dissolve in water. Here are some examples: Na2O(s) + H2O(l) 2NaOH(aq) BaO(s) + H2O(l) Ba(OH) 2(aq) -or- + - Na2O(s) + H2O(l) 2Na (aq) + 2 OH (aq) 2+ - BaO(s) + H O(l) Ba (aq) + 2 OH (aq) 2 The above reaction works best with the reactive metals on the left hand of the periodic table. The non-reactive metals to the right on the periodic table don t do so well. Al2O 3(s) + H2O(l) No reactions Acids One definition of an acid is that it is a compound that produces hydrogen ions + (H ) when dissolved in water + If you think about it, what is the H ion? It is a single proton with no electrons + around it. A free proton like this cannot exist in water, so the H ion is actually a + shortcut method of writing a family of related water complexes like H3O (aq), + + + H5O 2 (aq) and H9O 4 (aq). The hydronium ion, H3O (aq) is the main species in + + this mixture, so the Finicky chemist will use H3O (aq), instead of H (aq) to refer to this acid species Many nonmetal oxides undergo a combination reaction with water to form acids.

5 Examples include: CO 2(g) + H2O(l) H2CO 3(aq) SO (g) + H O(l) H SO 3 2 2 There are many compounds with H in them that are not acids, things like H 2(g) or CH 4(g) ore CH3CH2OH. How do you know when you have an acid, and when you have a plain old covalent compound? Usually when your write the chemical formula of an acid, you write the acid H as the first atom in the molecular formula. If there are other non-acidic hydrogens they are given later in the formula. Examples: HCN, HCl, HC2H3O 2 are hydrocyanic, hydrochloric and acetic acid respectively. CH CH CH and CH OH are not acids 3 2 3 3 Some compounds have more than one acidic proton. These compounds are called polyprotic acids Examples: H2SO 4 - Sulfuric Acid, a diprotic acid H PO - Phosphoric acids, a triprotic acid 3 4 Oxyacids Table 10.3 Any acid that contains an oxygen in the formula is called an oxyacid. The names of the oxyacids are tied to the names of the ions that you obtain when an acid deprotonates and goes into it s anion form. If the anion name ends in -ite, the name of the corresponding acid is -ous acid If the anion name ends in -ate, the name of the corresponding acid is -ic acid Examples: HNO HNO (NO anion nitrate hence nitric acid) (NO anion, nitrite, hence nitrous acid) - 3 3-2 2 Clicker questions: H2SO3 H2SO4 Perchloric acid Binary Acids Table 10.5 Binary acids consist of two atoms, one of which has to be the acidic hydrogen.

6 To name a binary acid you add the prefix hydro- to the anion and tack on the suffix -ic acid to the end of the anion name. Example: HCl(aq) Hydrochloric acid Clicker questions HI hydrosulfuric acid One final note about th physical form of an acid. To be an acid you have to be in a state that let s you ionize. Thus HCl(g) or HNO 3(g) are technically not acids. It is only when they are dissolved in a solvent to become HCl(aq) or HNO 3 (aq) that these compound become true acids. Some of the common properties of acids and bases are given in table 10.6 10-4 Decomposition Reactions Back to chemical reactions A decomposition reaction is just the opposite of a combination reaction In a decomposition reaction a single reactant breaks down into two or more products. Some examples of decomposition reactions are: The airbag on your car uses the decomposition reaction 2NaN (s) 2Na(s) + 3N (g) to inflate your airbag 3 2 IF you want a little more force, and a little more hair-trigger on the reaction, you can use some much less stable compounds in the following reactions Pb(N 3) 2(s) Pb(s) + 3 N 2(g) Hg(N 3) 2(s) Hg(l) + 3N 2(g) to set off a blasting cap!

7 10-5 Hydrates Many salts can combine with water to form interesting compounds called hydrates. A hydrate is a salt that has incorporated a specific number of water molecules into it s solid crystal structure to make a new molecular structure. For instance CuSO 4(s) is a while powder but it can react with 5 water molecules to form CuSO4 5H2O, a bright blue solid Figure 10.14 No jpg file? Maybe do a demo? Often the hydrated compound can be returned to the anhydrous (without water) form by gentle heating. Note how I wrote the waters of hydration in the formula using a dot convention. We do this because these waters are not that strongly bound to the structure can be added or removed fairly easily. 10-6 Single replacement (Substitution) Reactions In a single replacement or substitution reaction an element in one compound is replaced with a different element Examples: 2Mg(l) + TiCl 4(g) 2MgCl 2(s) + Ti(s) Fe(s) + H SO (aq) FeSO (aq) + H (g) 2 4 4 2 We will see that many metals undergo a single replacement reaction with acids to produce H gas. 2 10-7 Relative Activities of Metals We can use the Single replacement reaction (previous section) to rank metals in terms of relative reactivity For instance Cu is considered more active than Ag because it can displace Ag in the following reaction: Cu(s) + 2AgNO 3(aq) Cu(NO 3) 2(aq) + 2Ag(s) Figure 10.17 or demo

By comparing one metal with another we can come up with the following ranking of metals by activity. Table 10.8 It would be nice if that activity table matched exactly with the periodic table, but it doesn t. The general trend however is that the more reactive metals are on the left and the less reactive metals are on the right. Also in a single group the more reactive metals tend to be lower on the periodic table. If you are interested in metallurgy this activity ranking explains how smelting works. See the text if you are interested. 10-8 Relative Activities of Halogens We can try to rank the activities of nometals based on single displacement reactions in a similar manner, but the reactivities of the nonmetals are too diverse to make a nice table. However, if we focus on just the halides, then you can get something to work. For instance Br 2 is considered to be more active than I 2 because it will displace I2 in the reaction : Br 2(l) + 2NaI(aq) 2NaBr(aq) + I 2(s) (I won t try to demonstrate this because both Br 2 and I 2 give off some pretty nasty gases!) The ranking you get is F 2>Cl 2> Br 2> I2 In general the nonmetals in a single group become more reactive as you go up the periodic table 8

9 10-9 Double-Replacement Reactions In a double-replacement reaction the cations and anions of two compounds are exchanged We will look at two major types of double-replacement reactions. In this section we will examine precipitation reactions, and in the next section we will look at acid-base reactions Let s try to react NaCl(aq) with LiNO 3 (aq) In a double-replacement reaction you exchange the + of on ionic compound with the + of the other. This would give you the products LiCl(aq) and NaNO (aq) The reaction equation would be: NaCl(aq) + LiNO (aq) LiCl (aq) + NaNO (aq) 3 3 Since all reactants and products ionize completely in water, a more appropriate way to write this reaction would be: + - + - + - + - Na (aq) + Cl (aq) + Li (aq) + NO 3 (aq) Li (aq) + Cl (aq) + Na (aq) + NO 3 (aq) 3 The clever chemist quickly realizes that NOTHING HAS HAPPENED! We have all the same ions in the solution that we started with. Nothing has changed, so this is actually a NO REACTION So both NaCl(aq) + LiNO 3(aq) N.R. + - + - Na (aq) + Cl (aq) + Li (aq) + NO (aq) N.R. So simply switching cations around will not give you a reaction if both reactants and products are soluble ionic compounds! 3 To have a chemical reaction some kind of chemistry must occur. The chemistry that we will look at now is the chemistry of precipitation. Back in section 10-2 I told you that some, but not all, ionic compounds dissociate into individual ions in aqueous solution. Other ionic compounds have such a strong bond between the anions and the cations that water cannot break this interaction down, and the compound remains a solid. What happens if we mix two solutions together, one that has the dissolved aqueous cation and another that has a dissolved aqueous anion?

What happens is that due to the random motions of the ions in the mixed solutions a cation will get close enough to an anion that they can interact and form the strong ionic bond and a tiny solid will start to form. As more and more ions randomly bump into this tiny seed, the solid gets larger and larger, until is get so large that the molecules drop out of solution and form a solid on the bottom of the container. This process is called precipitation. Precipitation occurs when aqueous anions and cations coalesce for form a solid that cannot remain in solution. This is how we get our double-replacement reaction to actually do some chemistry and produce a product. We start with two ionic compounds that are completely dissolved into their respective aqueous ions. We then allow the aqueous ions to mix together to form a new solid ionic precipitate. Precipitation reactions Our next task, then, is to identify which ions are soluble and which ones like to form precipitates so we can identify when we have a precipitation reaction It is not possible to predict the solubility (or insolubility) of all compounds, but there are 6 simple rules you can use that will work about 95% of the time. Key task: Memorize and learn how to used Table 10.9 to predict the solubility of ionic compounds. Table 10.9 Note, as the table says, apply in order Practice: Predict the solubility of (NH 4) 2SO 4(s) Soluble - rule 1 CaCO3 Insoluble - rule 5 Al2O3 Insoluble - Rule 5 Pb(NO ) Soluble - rule 2 3 2 Clicker question - Solubility of a couple of salts 10

11 Now let s integrate solubility with precipitation reactions Practice problems: Use the solubility to predict the products of the following reactions BaCl 2(aq) + Na2SO 4(aq) Products: BaSO 4 and NaCl Solubility: BaSO 4 (s) exception to rule 6 ; NaCl (aq) rule 1 Final balanced reaction: BaCl (aq) + Na SO (aq) BaSO (s) + 2NaCl(aq) 2 2 4 4 LiOH(aq) + Pb(NO 3) 2(aq) Products: Pb(OH) 2 and LiNO3 Solubility: Pb(OH) 2 (s) rule 5 and LiNO 3(aq) rule 1 Final Balanced Reaction: 2LiOH(aq) + Pb(NO ) (aq) Pb(OH) (s) +2 LiNO (aq) 3 2 2 3 NaCH3COO(aq) + CaBr 2 (aq) Products: NaBr and Ca(CH3COO) 2 Solubility: NaBr (aq) rule 1; Ca(CH3COO) 2 (aq) rule 2 N.R. Clicker Question: one or two more solubility reactions Writing Reaction Equations So far we have been writing what are called Molecular equations because we wrote all of our compounds as complete molecules. To show what is going on in solution it is more appropriate to write complete ionic equations Key concept: In writing a complete ionic equation all soluble aqueous ionic compounds broken down their component ions. However all covalent compounds, and all solids, liquids, and gases remain unchanged. Practice problems: Let s go back to the previous reactions and write the reactions as complete ionic equations BaCl (aq) + Na SO (aq) BaSO (s) + 2NaCl(aq) Ba (aq) + 2Cl (aq) + 2 Na (aq) + SO (aq) BaSO (s) + 2Na (aq) + 2Cl (aq) 2 2 4 4 2+ - + 2- + - 4 4 - Notice that the Cl 2 in BaCl 2 breaks into 2 independent Cl ions + - and the 2NaCl becomes 2Na and 2 Cl 2LiOH(aq) + Pb(NO ) (aq) Pb(OH) (s) +2 LiNO (aq) 2Li (aq) + 2OH (aq) + Pb (aq) + 2NO (aq) Pb(OH) (s) + 2Li (aq) + 2 Cl (aq) 3 2 2 3 + - 2+ - + - 3 2

12 Clicker question Write a single molecular equation and 3 out of 4 incorrect complete ionic equations and have them pick out the correct equation Notice in the above equations were pretty long and cumbersome to write and included some ions that were unchanged throughout the reaction. We call these ions spectator ions because all they do is watch, they don t do any real chemistry. There is a third way to write a chemical equation that drops these spectator ions out of the written equation because they don t take part in in chemistry. We call this third type of equation a net ionic equation A net ionic equation is written just like a complete ionic equation, however all spectator ions are removed from both sides of the equation Practice: Continuing on with our examples Molecular: BaCl 2(aq) + Na2SO 4(aq) BaSO 4 (s) + 2NaCl(aq) Complete Ionic: 2+ - + 2- + - Ba (aq) + 2Cl (aq) + 2 Na (aq) + SO 4 (aq) BaSO 4(s) + 2Na (aq) + 2Cl (aq) 2+ 2- Net Ionic: Ba (aq) + SO (aq) BaSO (s) 4 4 Molecular: 2LiOH(aq) + Pb(NO 3) 2(aq) Pb(OH) 2(s) +2 LiNO 3 (aq) Complete Ionic: + - 2+ - + - 2Li (aq) + 2OH (aq) + Pb (aq) 2NO 3 (aq) Pb(OH) 2(s) + 2Li (aq) + 2 Cl (aq) - 2+ Net Ionic: 2OH (aq) + Pb (aq) Pb(OH) (s) Notice how much shorter the net ionic equations are to write 10-10 Acid-Base Reactions Our second example of a double replacement reaction is the acid-base reaction. HCl(aq) + NaOH(aq) H2O(l) + NaCl (aq) + - + - + - H (aq) + Cl (aq) + Na (aq) + OH (aq) H2O(l) + Na (aq) + Cl (aq) + - H (aq) + OH (aq) H2O(l) 2 + - Here the aqueous ions H and OH combine to form the covalent water molecule In this reaction a strong, corrosive acid is combined with a strong, corrosive,

base to make neutral water with a little dissolved salt. We call this a neutralizing reaction, and, when the amount of acid is exactly equal to the amount of base we can say that the acid and base have neutralized each other because the acid and basic characters of the two solutions have been nullified. You have heard the term salt. On definition of a salt is that it is the ionic compound that results (along with water) from the reaction of an acid and a base. Practice problems: Predict the product and write the net ionic equation for the following reactions: KOH(aq) + H2SO 4(aq) 2KOH(aq) + H2SO 4(aq) 2H2O(l) + K2SO 4(aq) - + 2OH (aq) + 2H 2H O(l) 2 HC2H3O 2(aq) + Ba(OH) 2(aq) HC2H3O 2(aq) + Ba(OH) 2(aq) 2H2O(l) + Ba(C2H3O 2)(aq) Note Ba(OH) slightly soluble rule 5. Ba(C H O )(aq) soluble rule 2 Clicker Questions Another acid base reaction 10-11 Oxidation-Reduction Reactions 2 2 3 2 Oxidation-Reduction reactions are reactions in which electrons are transferred from one reactant to the other Many of our reaction between metals and nonmetals are oxidation reduction reactions. For instance: 2Na(s) + S(s) Na2S(s) Can you see the electrons? I can t! But if you go to the electron configurations: + 2-2Na + :S: 2Na + ::S:: (Solid broken into ions for clarity) 1 2 4 2 6 [Ne]3s [Ne]3s 3p [Ne] [Ne]3s 3p =[Ar] Now you can see eletrons leaving one atom to go to another. Key Concepts: 1.) When an atom gives up its electron, it gains + charge, and we call this 13

14 oxidation 2.) When an atom gains electrons it loses + charge and we call this reduction 3.) Electrons, like all mass, must be conserved, so in any oxidation reduction chemical reaction the number of electrons lost by some of the reactant atoms must equal the number of electrons gained by other reactant atoms Many of the reactions we have looked at so far are oxidation-reduction reactions One final bit of somewhat funky and confusing nomenclature: In the above reaction: 2Na(s) + S(s) Na2S(s) Na was being oxidized because it lost electrons so its charge increased S was being reduced because it was gaining electrons so its charge decreases Na is also called a reducing agent because it allowed the reduction of S to occur S is called an oxidizing agent because the oxidation of Na could not occur without it Key Concepts: In an oxidation-reduction reaction: The material that loses electrons, gain + charge, is being oxidized, and is considered the reducing agent in the reaction The material that gains electron...loses + charge...is being reduced...is considered to be the oxidizing reagent in the reaction. Practice problems: First identify the atom or molecules being oxidized and reduced in the following reactions, then identify the oxidizing and reducing agents 2Na(s) + Cl 2(g) 2NaCl(s) Na is oxidized and is the reducing agent Cl is being reduced and is the oxidizing agent Clicker question: First identify the atom or molecules being oxidized and reduced in the following reactions, then identify the oxidizing and reducing agents Zn(s) + HgCl (aq) ZnCl (aq) + Hg(l) 2 2

15 Two finals notes before we leave this chapter. 1.Classifying reactions as decomposition, single replacement and double replacement is based on simply looking at the numbers of reactant molecules vs the number of product molecules. This is good for bookkeeping, but does not really tell you much about the underlying chemistry. Later in the chapter we started talking about precipitation reactions, acid-base reactions and oxidation reduction reactions. This viewpoint clearly focuses on the underlying chemistry of the reaction, and I find to be the better way to classify reactions. 2. This chapter very briefly introduced oxidation-reduction reactions. Your text doesn t return to this important class of reactions until chapter 24 which we wouldn t get to until the end of next semester (if we were lucky). Since I don t want you to miss out on this important type of reaction, we will now skip ahead to chapter 24!