Presentation on theme: "1 Ch. 4 Patterns of Chemical Reactivity lDemo: Pour together two clear colorless liquids lDid a chemical reaction occur? How do you know? lDemo: AlkaSeltzer."— Presentation transcript:
1 Ch. 4 Patterns of Chemical Reactivity lDemo: Pour together two clear colorless liquids lDid a chemical reaction occur? How do you know? lDemo: AlkaSeltzer in water, or calcium in water lDid a chemical reaction occur? 04m13vd1
2 Observing and Predicting Reactions lHow do we know whether a reaction occurs? What clues does nature offer? Make a list. lReview photos of reactions lGo to next topic topic
3 What clues does nature offer that a chemical reaction occurred? l lpptppt l lcrystalcrystal l lcolorcolor l lgasgas l lfumesfumes l lsmokesmoke l ltemperaturetemperature l lflamesflames l lmagneticmagnetic l lsoundsound l llightlight l lsolid decompsolid decomp l lexplosionexplosion l lsolid dissol.solid dissol. l lodorodor l lelect. cond.elect. cond. l lpH changepH change l ldensitydensity l lelectrictyyelectrictyy Go to next topictopic
16 Explosion Dynamite Building Demolition Whale Removal return
17 Odor Certain molecules, especially those containing sulfur or nitrogen, have distinctive odors. return
18 Electrical Conductivity Ba(OH) 2 + H 2 SO 4 return
19 Density/Volume Sugar + H 2 SO 4 return
20 pH Change return
21 Magnetic Properties Fe + S 8
22 Generate Electricity lChemical reaction in the battery return
Observing and Predicting Reaction Patterns lPredictions: l do an experiment l use periodicity l use classifications of reactions lexample: combustion reactions involve the reaction of an element or a compound with oxygen, usually with the evolution of heat
24 In the following particulate representations, a circle represents an atom and different circles represent different elements. Using these representations, draw pictures of all the different types of atomic/molecular changes these substances could undergo. In the following particulate representations, a circle represents an atom and different circles represent different elements. Using these representations, draw pictures of all the different types of atomic/molecular changes these substances could undergo. Reaction Classifications A BB C D E F
25 Combination (or Synthesis) B B C D E F AA AA C D C D C D F E BB
26 Decomposition + C D C D E F E F +
27 Single Replacement ++A C D A D C A E F A E F ++
28 Double Replacement + + 1:30 C D E F C F E D
29 Reactions in Solution Precipitation Reactions lcompound 1 + compound 2 compound 3 + compound 4 lAlso called double replacement or metathesis reactions. lexchange of ionic partners AB + CD AD + CB lPb(NO 3 ) 2 (aq) + K 2 CrO 4 (aq) PbCrO 4 (s) + 2KNO 3 (aq) lAnd other related reactions 04m10vd1
31 Precipitation Reactions lPrecipitation reactions: (An example of a "double replacement" or "metathesis" reaction). lPrecipitation Reactions form a solid when two solutions are combined. lAn example is the combining aqueous potassium chromate with aqueous lead nitrate to form the precipitate lead chromate (still used in school bus paint!!)
32 Describing Reactions in Solution lTo identify the precipitate or predict the formation of a precipitate the solubilities of compounds can be used. These rules should already be memorized! lTable 4.1, pg 144
33 Solubility Principles lMost nitrate and acetate salts are soluble. lMost salts containing the alkali metal ions (Li +, Na +, K +, Cs +, Rb + ) and the ammonium ion (NH 4 + ) are soluble. lMost chloride, bromide and iodide salts are soluble. Notable exceptions are salts containing the ions Ag +, Pb 2+ and Hg 2 2+.
34 Solubility Principles lMost sulfate salts are soluble. Notable exceptions are BaSO 4, PbSO 4, HgSO 4 and CaSO 4. lMost hydroxide salts are only slightly soluble. Important soluble hydroxides are NaOH and KOH. Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2 are somewhat soluble*. lMost sulfide, carbonate, chromate, and phosphate salts are only slightly soluble**. * Note Group 2 trends : As you go down the group sulfate solubility decreases and hydroxide solubility increases. ** Slightly soluble compounds will form precipitates using "normal" concentrations.
35 Describing Reactions in Solution lFor reactions involving ionic compounds, we can write the reaction as a molecular equation (or formula equation). This shows the normal (complete) formulas of all compounds: lExample: K 2 CrO 4 (aq) + Pb(NO 3 ) 2 (aq) PbCrO 4 (s) + 2 KNO 3 (aq)
36 Describing Reactions in Solution lWe can rewrite the same reaction as a complete ionic equation - Shows a picture of all that actually occurs in solution lstrong electrolytes represented as ions in solution lweak and non- electrolytes still written in molecular (non-ionized) aqeuous state. lExample: 2K + (aq) + CrO 4 2- (aq) + Pb 2+ (aq) + 2NO 3 - (aq) 2K + (aq) + CrO 4 2- (aq) + Pb 2+ (aq) + 2NO 3 - (aq) PbCrO 4 (s) + 2 K + (aq) + 2 NO 3 - (aq) PbCrO 4 (s) + 2 K + (aq) + 2 NO 3 - (aq)
37 Describing Reactions in Solution lA net ionic equation includes only the solution components involved in the reaction (spectator ions, which do not undergo change, are omitted) Pb 2+ (aq) + CrO 4 2- (aq) PbCrO 4 (s)
38 Stoichiometry lStoichiometry of Precipitation Reactions – based on Chapter 3 stoichiometry concepts, but using molarity (concentration) relationships. lPractice with Chapter 3 & Molarity! lSample: How many grams of lead(II) hydroxide can be formed when 22.5 mL of M Pb(NO 3 ) 2 solution reacts with excess sodium hydroxide? (Hint: Use a BCA table).
39 Acid-Base Reactions lDefinitions: lArrhenius: lAcid - forms H + ions in solution (e.g HCl) lBase - forms OH - ions in solution (e.g. NaOH) lBrønsted-Lowry : lAcid - proton (H + ) donor (e.g. HCl) lBase - proton acceptor e.g. NH 3 :NH 3 + H + NH 4 +
40 Acid-Base Reactions lGeneral reaction : Acid + base(metallic hydroxide) salt + water l(neutralization reaction) le.g. HCl and NaOH lmolecular equation : HCl(aq) + NaOH(aq) H 2 O(l) + NaCl(aq) lcomplete ionic equation. : H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) H 2 O(l) + Na + (aq) + Cl - (aq) lnet ionic equation: H + (aq) + OH - (aq) H 2 O (l)
41 Acid-Base Titrations lAcid-base titrations (volumetric analysis) – determine an unknown quantity through titration. lTitration involves adding a precisely measured volume of a solution of known concentration (the titrant) into a solution containing the substance being analyzed (the analyte). lThe titrant reacts with the analyte in a known manner, such as an acid-base reaction.
42 Acid-Base Titrations lAn indicator marks the equivalence point (or stoichiometric point) where just the right amount of titrant has been added to completely react with the analyte. lThe endpoint is where the indicator actually changes color, which hopefully occurs near the equivalence point.
43 Acid-Base Reactions lNote the similarities to precipitation reactions. lAcid-Base reactions are another variation of a double replacement reaction. The key is the production of water. lOther common double replacement reactions produce gases.
44 Acid-Base Stoichiometry lThere are numerous variations on the acid- base reaction. Be sure to read through the many examples in Section 4.8. We will consider these examples now from a modeling perspective.
45 Acid-Base Reactions lYou first want to examine the acid-base reaction (similar to predicting a precipitation reaction). Here are some general steps (they can and should vary depending on the problem): 1. List the major species present in solution before the reaction occurs. Decide what reaction will occur (look for formation of water or gases) 2. Write a balanced equation. (leave space for a BCA table) 3. Calculate the moles of reactants. For solutions, use the volumes of the original solutions and their molarities (before mixing). Input into a BCA table.
46 Acid-Base Reactions 4. Determine the limiting reactant if appropriate. 5. Analyze the problem and find the moles of reactant or product asked for. 6. Convert to grams or volume of solution if asked for *All problems are different. Dont force a problem into a particular solution method.
47 Reaction Classes Combination Reactions lelement + element compound lmetal + nonmetal ionic compound l 2Na(s) + Cl 2 (g) 2NaCl(s) lnonmetal + nonmetal covalent cmpd l 2H 2 (g) + O 2 (g) 2H 2 O(l) lDraw a molecular diagram of this type of reaction 03m10an1 01m11vd1
48 Reactants Product Combination Reactions
49 Combination: K + Cl 2
50 Reaction Classes Addition Reactions lelement + compound compound l Cl 2 + 2TiCl 3 2TiCl 4 l Cl 2 + C 2 H 4 C 2 H 4 Cl 2 lDraw a molecular diagram of this type of reaction
51 Addition Reactions Reactants Product
52 Reaction Classes Decomposition Reactions lCompound 2 elements or element + compound or 2 compounds lOxides, peroxides O 2 Nitrates NO 2 or NO 2 - Nitrates NO 2 or NO 2 - lCarbonates CO 2 lHydrates H 2 O lAmmonium salts NH 3 lDraw a molecular diagram of this type of reaction 04m03an1 CHMVID06
53 Decomposition of HgO
54 Decomposition Reactions lReactantProducts
55 ReactantProducts Decomposition Reactions
56 Reaction Classes Single-Displacement Reactions lelement + cmpd cmpd + element (The more metallic element in the compound is displaced.) lcarbon + metal oxides l 3C + Fe 2 O 3 3CO + 2Fe lmetals + water l Ca(s) + 2H 2 O(aq) Ca(OH) 2 (aq) + H 2 (g)
62 Reaction Classes Single-Displacement Reactions lnonmetals + salts l Cl 2 (aq) + 2KI(aq) 2KCl(aq) + I 2 (aq) lWhat do all these types of reactions have in common???
63 Oxidation-Reduction Reactions lreactions in which electrons are transferred lcauses a change in the charge of an ion or of oxidation state of an element in a molecule lOxidation states - numbers assigned to elements lused to keep track of electrons (not the same as charge, but related)
64 Rules for Assigning Oxidation States (Table 4.2) lThe oxidation state of an uncombined element is zero (includes diatomic elements H 2,N 2, O 2, F 2, Cl 2, Br 2 and I 2 ). lThe oxidation state of a monatomic ion is the same as its charge (e.g. the sulfide ion, S 2-, has an oxidation state of -2).
65 Rules for Assigning Oxidation States (Table 4.2) lOxygen has an oxidation state of -2 in covalent compounds (except in peroxides (O 2 2- ) where each oxygen is assigned an oxidation state of -1). lIn covalent compounds hydrogen is assigned an oxidation state of +1. (Hydrogen has a -1 charge in hydrides such as lithium hydride (LiH) or sodium hydride (NaH).
66 Rules for Assigning Oxidation States (Table 4.2) lIn compounds, fluorine always has an oxidation state of -1. lThe sum of the oxidation states of the elements in a neutral compound must equal zero. lThe sum of the oxidation states of the elements in a polyatomic ion must equal the charge on the polyatomic ion.
67 Rules for Assigning Oxidation States (Table 4.2) lOxidation states may be non-integers. For example in iron (III) oxide (Fe 3 O 4 ), the iron has an oxidation state of 8/3 (eight-thirds).
68 Rules for assigning oxidation states lPractice: Identify the oxidation state of each atom in the following compounds: lMagnesium nitrate lLithium nitride lSodium nitrite
69 Characteristics of Redox Reactions lOxidation la loss of electrons lan increase in oxidation state lthe substance oxidized is the reducing agent (gives electrons to another substance) lReduction la gaining of electrons la decrease in oxidation state lthe substance reduced is the oxidizing agent (takes electrons away from another substance)
70 Balancing Redox Reactions lBy the half-reaction method : lIn acidic solution l1.Write separate oxidation and reduction reactions for the reaction. l2.For each half reaction : lbalance all the elements except hydrogen and oxygen lbalance oxygen atoms using H 2 O lbalance hydrogen atoms using H + lbalance the charge using electrons
71 Balancing Redox Reactions l3.If necessary, balance electrons lost and gained in each half reaction by multiplying one or both half reactions by an integer. l4.Add the half-reactions and cancel out like species. l5.Check to make sure charges and elements are balanced.
72 Example 4.19 lPotassium dichromate is a bright orange compound that can be reduced to a blue- violet solution of chromium(III) ions. In acidic conditions, potassium dichromate reacts with ethyl alcohol as follows: Cr 2 O 7 2- (aq) + C 2 H 5 OH(l) Cr 3+ (aq) + CO 2 (g) + H 2 O(l) lBalance this equation using the half reaction method.
73 Balancing Redox Reactions lIn Basic solution (see example 4.20) : l1.Balance as in an acidic solution (see above). l2.Add a number of OH- ions equal to the H+ ions present to both sides of each half reaction to for H 2 O. l3.Eliminate the number of H 2 O molecules that appear on both sides of the equation. l4.Check to make sure charges and elements are balanced.
74 Group (Partner) Quiz 1.Give the oxidation state of each element in sodium chlorate (NaClO 3 ) 2.In the following reaction, identify the oxidizing agent, the reducing agent, the substance being oxidized, and the substance being reduced Br – (aq) + MnO 4 - (aq) Br 2 (l) + Mn 2+ (aq) 3.Balance the above RedOx reaction that occurs in acidic solution.
75 Classify the following reactions, based on the changes happening at an atomic/molecular level. 1. AlF 3 (aq) + 3H 2 O(l) Al(OH) 3 (s) + 3HF(aq) 2. BaCl 2 (aq) + Na 2 SO 4 (aq) BaSO 4 (s) + 2NaCl(aq) 3. Ca(OH) 2 (s) CaO(s) + H 2 O(g) 4. Ca(s) + 2H 2 O(l) Ca(OH) 2 (aq) + H 2 (g) 5. CaO(s) + CO 2 (g) CaCO 3 (s) 6. Cl 2 (aq) + 2NaI(aq) 2NaCl(aq) + I 2 (aq) 7. Cu(s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag(s) 8. Fe(s) + 2HCl(aq) FeCl 2 (aq) + H 2 (g) 9. H 2 SO 3 (aq) H 2 O(l) + SO 2 (g) 10. 2HgO(s) 2Hg(l) + O 2 (g) 11. KOH(aq) + HNO 3 (aq) KNO 3 (aq) + H 2 O(l) 12. 4Li(s) + O 2 (g) 2Li 2 O(s) 13. Na 2 S(aq) + 2HCl(aq) 2NaCl(aq) + H 2 S(g) 14. NH 3 (g) + HCl(g) NH 4 Cl(s) 15. NiCO 3 (s) NiO(s) + CO 2 (g) 16. P 4 (s) + 10F 2 (g) 4PF 5 (g) double displacement decomposition single displacement combination decompositiondecomposition double displacement combination combinationdecompositioncombination