CHEMISTRY 161 Chapter 4

CHEMICAL REACTIONS 2 HgO (s) → 2Hg (l) + O 2(g) aq 1. properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)

1.PROPERTIES OF AQUEOUS SOLUTIONS homogeneous mixture of two or more substances solvent solute substance in a large amount substance in a small amount N 2 gas phase O 2 (air) Ag solid phase Au (alloys) H 2 O liquid phase NaCl (sea water)

EXP1 iodine in ethyl alcohol (C 2 H 5 OH) EXP2 table salt in water (H 2 O) does not conduct electricity (molecular solid) I 2 does conduct electricity (ionic solid) Na + Cl -

AQUEOUS SOLUTION solutes solute water (H 2 O) electrolytesnon-electrolytes solution conducts electricity solution does not conduct electricity EXP3

electrolytesnon-electrolytes solution conducts electricity solution does not conduct electricity

non-electrolyte weak electrolyte strong electrolyte methanol sugar ethanol water darkbright ionic compounds (NaCl, KF) NaOH HCl H 2 SO 4 CH 3 COOH HCOOH HF medium EXP5

SOLUTION concentration

SOLUTION percentage concentration % = g [solute] / g solvent X g of sodium chloride are solved in 150 g of water. Calculate the percentage concentration 8 %

solubility of a solute number of grams of solute that can dissolve in 100 grams of solvent at a given temperature SOLUTION 36.0 g NaCl can be dissolve in 100 g of water at 293 K

GAS PHASE SOLUTION Saturn solvent H 2 /He solute CH 4, PH 3

LIQUID SOLUTION Europa solvent H 2 O solute MgSO 4

SOLID SOLUTION Triton solvent N 2 solute CH 4

methanol sugar ethanol water ionic compounds (NaCl, KF) NaOH HCl H 2 SO 4 CH 3 COOH HCOOH HF ELECTROLYTES

migrating negative and positive charges Kohlrausch NaCl

DISSOCIATION ‘breaking apart’ NaCl (s) → Na + (aq) + Cl - (aq) NaOH (s) → Na + (aq) + OH - (aq) HCl (g) → H + (aq) + Cl - (aq) strong electrolytes are fully dissociated Ca(NO 3 ) 2 (s) → Ca 2+ (aq) + 2 NO 3 - (aq) EXP5 polyatomic ions do NOT dissociate

O HH δ-δ- δ+δ+ δ+δ+

SOLVATION cationsanions

SOLVATION non-electrolyte

CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) weak electrolytes are not fully dissociated reversible reaction (chemical equilibrium) → ← NaCl (s) → Na + (aq) + Cl - (aq) strong electrolytes are fully dissociated

CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)

2.1. PRECIPITATION REACTIONS solution 1solution 2 solution 1 + solution 2

2.1. PRECIPITATION REACTIONS formation of an insoluble product (precipitate) NaCl(aq) + AgNO 3 (aq) AgCl(s) + NaNO 3 (aq) EXP 6

insoluble compounds 1.M + compounds (M = H, Li, Na, K, Rb, Cs, NH 4 ) 2. A - compounds (A = NO 3, HCO 3, ClO 3, Cl, Br, I) (AgX, PbX 2 ) 3. SO 4 2- (Ag, Ca, Sr, Ba, Hg, Pb) 4. CO 3 2-, PO 4 3-, CrO 4 2-, S 2- (Ag, Ca, Sr, Ba, Hg, Pb)

NaCl(aq) + AgNO 3 (aq) → AgCl(s) + NaNO 3 (aq) balanced molecular equation (table to determine which compound precipitates)

balanced ionic equation 1. NaCl(s) → Na + (aq) + Cl - (aq) 2. AgNO 3 (s) → Ag + (aq) + NO 3 - (aq) 3. Na + (aq) + Cl - (aq) + Ag + (aq)+ NO 3 - (aq) → AgCl(s) + Na + (aq) + NO 3 - (aq) spectator ions

Ba(NO 3 ) 2 (aq) + Na 3 PO 4 (aq) 1.which compound falls out? 2. balanced molecular equation 3. balanced ionic equations 4. identify spectator ions Cs 2 CrO 4 (aq) + Pb(NO 3 ) 2 (aq) Ba(NO 3 ) 2 (aq) + Na 2 SO 4 (aq)

CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)

ACIDS AND BASES Arrhenius (1883) ACIDS BASES NaOH (s) → Na + (aq) + OH - (aq) MOH → M + (aq) + OH - (aq) HCl (g) → H + (aq) + Cl - (aq) HAc → H + (aq) + Ac - (aq) ionization

IDENTIFICATION Litmus Paper acid base red blue Säure Base EXP7

ACIDS AND BASES ACIDS BASES HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O HAc (aq) + MOH (aq) → MAc (aq) + H 2 O and NEUTRALIZE EACH OTHER acid + base salt + water

H+H+ ≈ m Na+ ≈ m ACIDS AND BASES

HCl (g) → H + (aq) + Cl - (aq) H + (aq) + H 2 O H 3 O + (aq) HCl (g) + H 2 O → H 3 O + (aq) + Cl - (aq) one step hydronium ion

(aq)(l)(aq) hydronium ion acid base

cationhydronium ion

PROPERTIES OF ACIDS 1.acids have a sour taste vinegar – acetic acid lemons – citric acid 2. acids react with some metals to form hydrogen 2 HCl(aq) + Mg(s) → MgCl 2 (aq) + H 2 (g) 3. acids react with carbonates to water and carbon dioxide 2 HCl(aq) + CaCO 3 (s) → CaCl 2 (aq) + [H 2 CO 3 ] H 2 CO 3 → H 2 O(l) + CO 2 (g) EXP8 EXP9 4. some acids are hygroscopic H 2 SO 4 (conc)

BASES 1.bases have a bitter taste 2. bases feel slippery soap 3. aqueous bases and acids conduct electricity

EXAMPLES KOH(aq) and HF(aq) Mg(OH) 2 (aq) and HCl(aq) Ba(OH) 2 (aq) and H 2 SO 4 (aq) NaOH(aq) and H 3 PO 4 (aq) (stepwise)

Bronsted (1932) ACIDS HAc → H + (aq) + Ac - (aq) proton donors BASES proton acceptor B + H + (aq) → BH + (aq)

weak electrolyte CH 3 COOH(aq) + H 2 O(l) H 3 O + (aq) + CH 3 COO - (aq) NH 3 (aq) + H 2 O(l) NH OH - strong electrolyte HCl(aq) + H 2 O(l) → H 3 O + (aq) + Cl - (aq) HNO 3 (aq) + H 2 O(l) → H 3 O + (aq) + NO 3 - (aq) donor versus acceptor

CH 3 COOH(aq) + H 2 O(l) H 3 O + (aq) + CH 3 COO - (aq) NH 3 (aq) + H 2 O(l) NH 4 + (aq)+ OH - (aq) H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH - (aq) water can be either an acid or a base AUTO DISSOCIATION

monoprotic acids diprotic acid HF, HCl, HBr, HNO 3, CH 3 COOH H 2 SO 4 → H + (aq) + HSO 4 - (aq) HSO 4 - (aq) H + (aq) + SO 4 2- (aq) triprotic acid H 3 PO 4 H + (aq) + H 2 PO 4 - (aq) H 2 PO 4 - (aq) H + (aq) + HPO 4 2- (aq) HPO 4 2- (aq) H+(aq) + PO 4 3- (aq) EXP10

CHEMICAL PROPOERTIES 1. Non-metal oxides react with water to form an acid (acetic anhydrides) Cl 2 O 7, SO 2, Br 2 O 5 + H 2 O

CHEMICAL PROPERTIES 2. Soluble metal oxides react with water to form a base (base anhydrides) MgO, Al 2 O 3 + H 2 O

NAMING ACIDS AND BASES prefix hydro- the suffix –ic to the stem of the nonmetal name followed by the word acid binary acids

NAMING ACIDS AND BASES oxo acids acids contain hydrogen, oxygen, plus another element main group 5 HNO 3 nitric acid HNO 2 nitrous acid H 3 PO 4 phosphoric acid H 3 PO 3 phosphorous acid

H 2 SO 4 sulfuric acid H 2 SO 3 sulfurous acid main group 6 main group 7 HClO 4 perchloric acid HClO 3 chloric acid HClO 2 chlorous acid HClO hypochlorous acid

Acids in the Solar System Venus H 2 SO 4 (g) Europa H 2 SO 4 (s)

Acids in the Interstellar Medium

Orion NH 3, H 2 O, H 2 S CH 3 COOH HCOOH HF, HCl

CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)

1. oxidation KEY CONCEPTS loss of electrons 2. reduction acceptance of electrons NUMBER OF ELECTRONS MUST BE CONSERVED

1. oxidation EXAMPLE 2. reduction !!!balance electrons!!! Na + Cl - Na  Na + + e Cl e  2 Cl - CaO, Al 2 O 3

substance that lost the electrons reduction agent substance that gained the electrons oxidizing agent oxidizing agent is reduced reducing agent is oxidized 2 Na + Cl 2  2 Na + Cl -

EXAMPLE 1 solid state reaction of potassium with sulfur to form potassium sulfide EXAMPLE 2 solid state reaction of iron with oxygen to form iron(III)oxide

OXIDATION NUMBER ionic compounds ↔ molecular compounds NaClHF, H 2 Na + Cl - ? electrons are fully transferred covalent bond charges an atom would have if electrons are transferred completely

HF H + + F - molecular compoundionic compound F - oxidation state -1 H + oxidation state +1 EXAMPLE 1

H2OH2O molecular compound ionic compound 2 H + + O 2- H + oxidation state +1 O 2- oxidation state -2 EXAMPLE 2

H2H2 molecular compound ionic compound H + + H - EXAMPLE 3 OXIDATION NUMBER OF FREE ELEMENTS IS ZERO

RULE 1 OXIDATION NUMBER OF FREE ELEMENTS IS ZERO H 2, O 2, F 2, Cl 2, K, Ca, P 4, S 8

RULE 2 monoatomic ions oxidation number equals the charge of the ion group I M + group II M 2+ group III M 3+ (Tl: also +1) group VII (w/ metal) X -

RULE 3 oxidation number of hydrogen +1 in most compounds (H 2 O, HF, HCl, NH 3 ) -1 binary compounds with metals (hydrides) (LiH, NaH, CaH 2, AlH 3 )

RULE 4 oxidation number of oxygen -2 in most compounds (H 2 O, MgO, Al 2 O 3 ) -1 in peroxide ion (O 2 2- ) (H 2 O 2, K 2 O 2, CaO 2 ) -1/2 in superoxide ion (O 2 - ) (LiO 2 )

RULE 5 oxidation numbers of halogens F: -1 (KF) Cl, Br, I: -1 (halides) (NaCl, KBr) Cl, Br, I: positive oxidation numbers if combined with oxygen (ClO 4 - )

RULE 6 charges of polyatomic molecules must be integers (NO 3 -, SO 4 2- ) oxidation numbers do not have to be integers -1/2 in superoxide ion (O 2 - )

MENUE 1.oxidation states of group I – III metals 2.oxidation state of hydrogen (+1, -1) 3. oxidation states of oxygen (-2, -1, -1/2, +1) 4.oxidation state of halogens 5.remaining atoms

oxidizing agents OCl - Cl - ????? EXP10

reducing agent 2 Na + 2 H 2 O  H NaOH EXP11/12

NO NO 2 NO + NO - NO 2 - NO 3 - PO 4 3- SO 4 2- SO 3 SO 2 KO 2 K2OK2O BrO - KClO 4

1.redox reactions 2. oxidation versus reduction 3. oxidation numbers versus charges 4. calculation of oxidation numbers REVISION

TYPES OF REDOX REACTIONS 1.combination reactions A + B → C 2. decomposition reactions C → A + B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions

1.combination reactions A + B → C two or more compounds combine to form a single product S 8 (s) + O 2 (g) → SO 2 (g) 1.oxidation numbers 2. balancing charges

MENUE 1.oxidation states of group I – III metals 2.oxidation state of hydrogen (+1, -1) 3. oxidation states of oxygen (-2, -1, -1/2, +1) 4.oxidation state of halogens 5.remaining atoms

2. decomposition reactions C → A + B breakdown of one compound into two or more compounds HgO(s) → Hg(l) + O 2 (g) 1.oxidation numbers 2. balancing charges KClO 3 (s) → KCl(s) + O 2 (g)

3. displacement reactions A + BC → AC + B an ion or atom in a compound is replaced by an ion or atom of another element 3.1. Hydrogen displacement 3.2. Metal displacement 3.3. Halogen displacement

3.1. Hydrogen displacement group I and some group II metals (Ca, Sr, Ba) react with water to form hydrogen Na(s) + H 2 O(l) → NaOH + H 2 (g) less reactive metals form hydrogen and the oxide in water (group III, transition metals) Al(s) + H 2 O(l) → Al 2 O 3 (s) + H 2 (g)

3.1. Hydrogen displacement even less reactive metals form hydrogen in acids Zn(s) + HCl(aq) → ZnCl 2 (aq) + H 2 (g) EXP12

Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au activity series of metals displace H from water displace H from steam displace H from acids

Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au likes to donate electrons does not like so much to donate electrons EXP13

3.2. Metal displacement V 2 O 5 (s) + 5 Ca(s) → 2 V(s) + 5 CaO(s) TiCl 4 (g) + 2 Mg (l) → Ti(s) + 2 MgCl 2 (l)

3.3. Halogen displacement F 2 > Cl 2 > Br 2 > I 2 reactivity (‘likes’ electrons) Cl 2 (g) + 2 KBr(aq) → 2 KCl(aq) + Br 2 (l) 00+1 Br 2 (g) + 2 KI(aq) → 2 KBr(aq) + I 2 (s)

4. disproportionation reactions an element in one oxidation state is oxidized and reduced at the same time H 2 O 2 (aq) → 2 H 2 O(l) + O 2 (g) Cl 2 (g) + 2 OH - (aq) → ClO - (aq) + Cl - (aq) + H 2 O(l)

SUMMARY 1.combination reactions A + B → C 2. decomposition reactions C → A + B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions

STOCHIOMETRY (CONCENTRATION) molar concentration Molarity (M) =

How many grams of AgNO 3 are needed to prepare 250 mL of M AgNO 3 solution?

How many mL of M NaOH are required to react completely with 15.4 mL of M H2SO4? 2 NaOH + H 2 SO 4 Na 2 SO 4 + 2H 2 O

How many mL of M NaOH are required to react completely with 20.1 mL of 0.2 M HCl? NaOH + HCl NaCl + H2O

How many grams of iron(II)sulfide have to react with hydrochloric acid to generate 12 g of hydrogen sulfide?

How many moles of BaSO4 will form if 20.0 mL of M BaCl2 is mixed with 30.0 mL of M MgSO4? BaCl 2 + MgSO 4 BaSO 4 + MgCl 2 This is a limiting reagent problem

How many ml of a 1.5 M HCl will be used to neutralize a 0.2 M Ba(OH) 2 solution? How many ml of a 1.5 M HCl will be used to prepare 500 ml of a 0.1 M HCl? XX=

LIMITING REACTANT C 2 H 4 + H 2 O C 2 H 5 OH EXP14

excess reactant limiting reactant

How many grams of NO can form when 30.0 g NH 3 and 40.0 g O 2 react according to 4 NH O 2 4 NO + 6 H 2 O