Chapter 4 Reactions in Aqueous Solution

Observing and Predicting Reactions Chapter 4 Explosion!! Reactions in Aqueous Solution Fire! Precipitate Smoke How do we know when a reaction is taking place? Bubbles Change in Temperature 1 Properties of Aqueous Solutions 2 Electrolytes in Aqueous Solutions Most reactions in general chemistry take place in an aqueous (aq) environment. What does that mean? – Solution: homogeneous mixture of two or more substances Electrolytes conduct electricity in aqueous solutions by moving toward the oppositely charged electrode NaCl – Solute: substance present in smaller amount Nonelectrolytes, however, do not have charges in aqueous solutions and are not attracted to the electrodes – Solvent: substance present in greater amount – Aqueous solution: solvent is water Electrolytes Properties of Aqueous Solutions Electrolytic Properties Substances behave differently when they are placed in water; specifically ionic versus covalent compounds. NaCl CH3OH NaCl “breaks apart” or ionizes Na+ ClNa+ ClNa+ ClElectrolyte Non-electrolytes Strong electrolyte: Weak electrolyte: substance that, when dissolved in water, results in a solution that can conduct electricity substance that is a poor conductor of electricity when dissolved in water Non-electrolyte: substance that doesn’t conduct electricity when dissolved in water (CH3OH – methanol) CH3OH CH3OH CH3OH does not “break apart”. It does not ionize CH3OH NonElectrolyte Ionize 100% Ionize <100% Ionize 0% Six General Types of Reactions Double Replacement / Precipitation Reactions • Double Replacement (Precipitation) Soluble ionic compound AB dissolved in H2O • Single Replacement Soluble ionic compound CD dissolved in H2O Most of the time, products of these reactions can be easily predicted • Acid – Base Neutralization Insoluble precipitate (BC) formed from mixing AB + CD • Combustion Will not dissolve in H2O • Combination Sometimes products of these reactions cannot be easily predicted • Decomposition In general: AB(aq) + CD(aq) → AD(aq) + BC(s) Pb(NO3)2 Example: (aq) + 2KI (aq) →? How do we know what is soluble? Predicting Products of Double Replacement We have rules… MgCl2 MgCl2 (aq) (aq) + 2AgNO3 + AgNO3 (aq) To predict products, switch partners with cations (aq) → ??? → Mg(NO3)2 (aq) + 2AgCl (s) To determine states of mater, check the solubility rules of the products If the ionic compound is soluble → (aq) If the ionic compound is insoluble → (s) If both products are soluble (aq), there is NO REACTION Pb(NO3)2 (aq) + 2KI (aq) → 2KNO3 (aq) + PbI2 (s) Solubility Rules Chemistry Humor…tell all your friends! We’re a clever Lot! • Determine if the following ionic compounds will be soluble (aq) or insoluble (s) in water: • K2CO3 • BaSO4 • PbCl2 • NaClO4 • Ag2S • (NH4)3PO4 • Cu(OH)2 Molecular, Ionic, and Net Ionic Equations Solubility Rules Answers MgCl2 • Determine if the following ionic compounds will be soluble (aq) or insoluble (s) in water: soluble (aq) • K2CO3 • BaSO4 insoluble (s) • PbCl2 insoluble (s) • NaClO4 soluble (aq) • Ag2S insoluble (s) • (NH4)3PO4 soluble (aq) • Cu(OH)2 insoluble (s) (aq) + 2AgNO3 (aq) → Mg(NO3)2 (aq) + 2AgCl (s) Molecular Equation Any ionic compound that is aqueous (aq) is separated into ions Mg2+(aq) + 2Cl-(aq) + 2Ag+(aq) + 2NO3-(aq) → Mg2+(aq) + 2NO3-(aq) + 2AgCl(s) Ionic Equation Any ion found on both sides of the equation is removed from the net ionic equation. Solids liquids and gases are not removed 2Cl- (aq) + 2Ag+(aq) → 2AgCl (s) Net Ionic Equation 14 Complete/Balance These Equations Complete/Balance These Equations Give the net ionic equation for each • Identify any precipitate(s) that form(s) and identify the spectator ions. • Na2S(aq) + CuCl2(aq)
• MgSO4(aq) + BaCl2(aq)
• KNO3(aq) + CaCl2(aq)
• CuSO4(aq) + NaOH(aq)
• Na2S(aq) + CuCl2(aq)
2 NaCl(aq) + CuS(s) – Cu2+(aq) + S2-(aq)
CuS(s) • MgSO4(aq) + BaCl2(aq)
MgCl2(aq) + BaSO4(s) – Ba2+(aq) + SO42-(aq)
BaSO4(s) • KNO3(aq) + CaCl2(aq)
KCl(aq) + Ca(NO3)2(aq) – No reaction 15 Acid-Base Reactions Acid-Base Reaction Acid: a substance, when dissolved in water, increases the H3O+ concentration. HCl(aq) + H2O(l) → H3O+ (aq) + Cl-(aq) NaOH(s) + H2O(l) → H2O(l) + OH-(aq) + Na+(aq) H2O(l) NaOH(s) → OH-(aq) + Na+(aq) HCl Sometimes represented as H+ Base: a substance, when dissolved in water, increases the OH- concentration. -or- • CuSO4(aq) + 2 NaOH(aq)
Cu(OH)2(s) + Na2SO4(aq) – Cu2+(aq) + 2 OH-(aq)
Cu(OH)2(s) Mg(OH)2 Acid + Base → H2O + Salt Water is created from H+ and OH- 2HCl(aq) + Mg(OH)2(s) → 2H2O(l) + MgCl2(aq) The salt is created from spectator ions 16 Single-Replacement (Displacement) Acid-Base Neutralization A single-replacement is somewhat similar to a double replacement in that one chemical “replaces” another. • What does the word “neutralization” mean to you? • How might these two beakers react? Except, instead of two ionic compounds reacting, one of the reactants is a pure metal (not an ion). And there is no precipitate. A(s) + BC(aq) → AC(aq) + B(s or g) In general: Na+ OH- + Cl- Na+ To predict products, “switch places” Solid Cu metal (no charge) becomes a Cu2+ ion Cl- H+ Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s) Molecular Equation H+ Net Ionic OH- Equation Single-Displacement Reactions The Ag+ ion becomes solid metal (no charge) Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s) Single Displacement: Li + H2O Some metals can also displace hydrogen from water and acids. In this case, the solid metal and hydrogen switch places. Products are predicted by “switching places” also. What’s actually happening? Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s) 2Li(s) + 2H2O(l) → H2(g) + 2LiOH(aq) Notice that the NO3-’s aren’t really doing anything. Ag+(aq) NO3-(aq) NO3-(aq) Cu2+(aq) → In both reactions, hydrogen is displaced to give hydrogen gas, H2. Ag(s) Mg(s) + 2HCl(aq) → H2(g) + MgCl2(aq) Cu(s) Cu(s) 21 22 Will this reaction take place? The Activity Series If we mix any metal with any aqueous ionic solution (or acid), will a reaction take place? Ummm….no. Once again, we have rules. For a single replacement reaction to occur, the solid metal must be higher on the list than the metal it is displacing The first five will react with H2 O Anything above H2 will react with an acid Represents Acids Predict whether a reaction occurs Li K Ba Ca Na Mg Al Mn Zn Cr Fe Co Ni Sn H2 Cu Ag Hg Pt Au • Na + H2O • Fe + HCl • Ni + Pb(NO3)2 • Ag + Mg(NO3)2 • Zn + Co(NO3)2 24 Combustion Reactions What does your local (non-nuclear) power plant, your running car, and wildfires all have in common? Recognizing a Combustion Reaction In each case, a carbon containing compound is being burned in oxygen. Coal Gasoline Plant Matter CxHyOz + O2(g) → CO2(g) + H2O(g) CO2(g) + H2O(g) In each case, a combustion reaction is taking place. During the reaction, carbon dioxide and water are produced (in a perfect world). In reality, side-reactions also take place. We’ll ignore these for now. Combination and Decomposition Reactions Combination A combustion reaction usually will contain: A reactant containing C and H or C, H, and O + O2 as the other reactant CO2 and H2O as products → Combination Reactions Combination reactions that produce ionic compounds can be easily predicted Simpler compounds and/or elements → More complex compound(s) Predict: A+B→C Cl2(g) + Na(s) → ?? Cl2(g) + 2Na(s) → 2NaCl(s) 2H2(g) + O2(g) → 2H2O(g) Decomposition More complex compound(s) → Simpler compounds and/or elements However, when covalent compounds are formed, more than one compound is possible Predict: C(s) + O(g) → ?? C→A+B 2H2O(g) → 2H2(g) + O2(g) Decomposition Reactions CO2(g) CO(g) Even if all the products of combination reactions cannot be predicted, you should be able to recognize this type of reaction Decomposition Reactions Many decomposition reactions produce gases • Oxides, peroxides (O22-) → O2 2C7H5(NO2)3(s) → 7CO(g) + 7C(s) + 5H2O(g) + 3N2(g) TNT Gases produced Explosions are caused by decomposition of unstable compounds. In the reaction, gases are rapidly produced and expand from heat formed in the reaction. Think about how much more space gases occupy than solids and why this might be destructive!! • Nitrates (NO3-) → NO2 • Carbonates (CO32-) → CO2 • Ammonium (NH4+) salts → NH3 32 Types of Reactions Summary Predict Products and Types of Reactions • Precipitation: use Solubility Rules • (AB + CD
AD + CB) • Acid-Base Neutralization: acid + base
salt + water • (AB + CD
AD + H2O) • Combination: start with elements • (A + B
AB) • Decomposition: often produces gas • (AB
A + B) • Single Displacement: use Activity Series to predict if a reaction occurs • (A + BC
B + AC) • Combustion: – Hydrocarbon + O2(g)
CO2(g) + H2O(g) Predict Products and Types of Reactions Cl Cl ClH+ H+ Cl- + H+ ClH+ ClCa Ca Ca Ca Ca Ca Ca Predict Products and Types of Reactions Ag+ NO3- + Cl- Na+ + Cl Cl Cl Ag+ Cl Na Na Na Na Na Na NO3- NO3Ag+ Na+ Cl- ClNa+ Classify (and balance) these rxns Predict Products and Types of Reactions H+ + Cl- K+ OH- K+ H+ Cl- OH- 1. AlF3(aq) + H2O(l)
Al(OH)3(s) + HF(aq) 2. Fe(s) + HCl(aq)
3. Ca(OH)2(s)
CaO(s) + H2O(g) 4. Ca(s) + H2O(l)
5. CaO(s) + CO2(g)
6. Cu(s) + AgNO3(aq)
7. BaCl2(aq) + Na2SO4(aq)
40 Classify these reactions by type Classify these reactions by type 8. H2SO3(aq)
H2O(l) + SO2(g) 9. HgO(s)
Hg(l) + O2(g) 10. KOH(aq) + HNO3(aq)
11. Li(s) + O2(g)
12. Mg(OH)2(aq) + 2HCl(aq)
13. NH3(g) + HCl(g)
NH4Cl(g) 14. NiCO3(s)
15. P4(s) + F2(g)
PF5(g) 1. AlF3(aq) + 3H2O(l)
Al(OH)3(s) + 3HF(aq) 2. Fe(s) + 2HCl(aq)
FeCl2(aq) + H2(g) 3. Ca(OH)2(s)
CaO(s) + H2O(g) 4. Ca(s) + 2H2O(l)
Ca(OH)2(aq) + H2(g) 5. CaO(s) + CO2(g)
CaCO3(s) 6. Cu(s) + 2AgNO3(aq)
Cu(NO3)2(aq) + 2Ag(s) 7. BaCl2(aq) + Na2SO4(aq)
BaSO4(s) + 2NaCl(aq) 8. H2SO3(aq)
H2O(l) + SO2(g) 9. 2HgO(s)
2Hg(l) + O2(g) 10. KOH(aq) + HNO3(aq)
KNO3(aq) + H2O(l) 11. 4Li(s) + O2(g)
2Li2O(s) 12. Mg(OH)2(aq) + 2HCl(aq)
2MgCl2(aq) + H2O(l) 13. NH3(g) + HCl(g)
NH4Cl(s) 14. NiCO3(s)
NiO(s) + CO2(g) 15. P4(s) + 10F2(g)
4PF5(g) Oxidation-Reduction Reactions precipitation single disp./H2 decomposition single disp./H2 combination single disp./metal precipitation decomposition decomposition neutralization combination neutralization combination decomposition combination Assigning Oxidation Numbers Oxidation-reduction reactions are reactions in which electrons are transferred to/from atoms. While seemingly a different “type” of reaction than the six previously mentioned, it another way of describing those reactions. In order to determine if a redox reaction takes place, we first have to assign oxidation numbers to all of the atoms in a reaction Single-Replacement Assigning Oxidation Numbers Combination Double-replacement Decomposition Acid-Base Combustion Usually redox reactions Assigning Oxidation Numbers Assigning Oxidation Numbers #3) Oxygen usually has an oxidation number of –2 in its compounds. Exceptions: Oxygen has an oxidation number of –1 in peroxides, which contain 2the O2 ion (e.g., K2O2 and H2O2). #1) The oxidation number of atoms of an uncombined element (in elemental state) is 0. Fe(s) O2(g) N2(l) KNO3(aq) CH4(g) +1 H2O2(aq) Exception to the rule -1 -2 #2) Hydrogen has an oxidation number of +1 (unless it is combined with alkali metals, in which case it has an oxidation number of –1). H2O(l) H2O(g) #4) For binary ionic compounds, the oxidation number is the same as the charge of the ion. NaCl(aq) NaH Exception to the rule +1 -1 MgS(s) +2 -2 AlP(s) +3 -3 -1 #5) For covalent compounds, the oxidation number of the halogens is -1. 45 Assigning Oxidation Numbers Assigning Oxidation Numbers #6) The sum of the oxidation numbers of all atoms in a substance must equal the total charge on the species NaCl +1 For many covalent compounds, you will need to use rule #6 to determine oxidation numbers for elements not covered by the other rules Since NaCl is not a charged species, the oxidation numbers must add up to 0. + -1 = 0 CO2 ? + 2(-2) = 0 Carbon is +4 CO ? + 1(-2) = 0 Carbon is +2 CCl4 ? + 4(-1) = 0 Carbon is +4 H3O+ 3(+1) + -2 = +1 Since H3O+ is a charged species, the oxidation numbers must add up to the charge. Some elements can only be determined using the last rule CH4 ? + +1(4) = 0 Hydrogen is +1. But what about carbon? Carbon has to be -4 Example Assigning Oxidation Numbers • H2SO4 – H = +1 – O = –2 – S unknown, so leave this for last. • The overall charge on this compound is 0. • Use algebra to solve for S: • 2(+1) + 1(x) + 4(-2) = 0 • Determine values of the oxidation number of each element in these compounds or ions: H2O SO2 CCl4 H2O2 Fe3(PO4)2 MnO4NaNO3 KClO4 • Solve for each element: MgCr2O7 49 Assigning Oxidation Numbers 50 Oxidation-Reduction Reactions Now that you have assigned oxidation numbers, let’s determine if something is oxidized or reduced. • Determine values of the oxidation number of each element in these compounds or ions: H2O H: +1, O: -2 SO2 S: +4, O: -2 CCl4 C: +4, Cl: -1 H2O2 H: +1, O: -1 Fe: +2, P: +5, O: -2 Fe3(PO4)2 Mn: +7, O: -2 MnO4Na:+1, N:+5, O: -2 NaNO3 K: +1, Cl: +7, O: -2 KClO4 2H2(g) + O2(g) → 2H2O(g) 0 0 +1 -2 H2 is oxidized because the oxidation number increased. It lost electrons. O2 is reduced because the oxidation number decreased. It gains electrons. 51 Oxidation-Reduction Reactions • The substance which is oxidized is called the reducing agent. • The substance which is reduced is called the oxidizing agent. • 4 Fe(s) + 3O2(g)
2Fe2O3(s) Oxidation-Reduction Reactions • Identify the element oxidized/reduced. Also identify the oxidizing agent and the reducing agent. • Cu(s) + 2AgNO3(aq)
2Ag(s) + Cu(NO3)2(aq) • Zn(s) + 2HCl(aq)
ZnCl2(aq) + H2(g) • H2O(l) + Na(s)
H2(g) + NaOH(aq)