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CHEMISTRY 161
Chapter 4
CHEMICAL REACTIONS
2 HgO(s) → 2Hg(l) + O2(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
N2 gas phase O2
(air)
Ag solid phase Au(alloys)
H2O liquid phase NaCl(sea water)
EXP1
iodine in ethyl alcohol (C2H5OH)
EXP2
table salt in water (H2O)
does not conduct electricity(molecular solid)
I2
does conduct electricity(ionic solid)
Na+Cl-
AQUEOUS SOLUTION
solutes
solute
water (H2O)
electrolytes non-electrolytes
solution conducts electricity
solution does not conduct electricity
EXP3
electrolytes non-electrolytes
solution conducts electricity
solution does not conduct electricity
non-electrolyte weak electrolyte strong electrolyte
methanol
sugar
ethanol
water
dark bright
ionic compounds
(NaCl, KF)
NaOH
HCl
H2SO4
CH3COOH
HCOOH
HF
medium
EXP5
SOLUTION
concentration
SOLUTION
percentage concentration
% = g [solute] / g solvent X 100
12 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
H2/He
solute
CH4, PH3
LIQUID SOLUTION
Europa
solvent
H2O
solute
MgSO4
SOLID SOLUTION
Triton
solvent
N2
solute
CH4
methanol
sugar
ethanol
water
ionic compounds
(NaCl, KF)
NaOH
HCl
H2SO4
CH3COOH
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(NO3)2 (s) → Ca2+(aq) + 2 NO3- (aq)
EXP5
polyatomic ions do NOT dissociate
O
H H
δ-
δ+ δ+
SOLVATION
cations anions
SOLVATION
non-electrolyte
CH3COOH (aq) H+ (aq) + CH3COO- (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 1 solution 2 solution 1 + solution 2
2.1. PRECIPITATION REACTIONS
formation of an insoluble product
(precipitate)
NaCl(aq) + AgNO3(aq)
AgCl(s) + NaNO3(aq) EXP 6
insoluble compounds
1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4)
2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I)(AgX, PbX2)
3. SO42-
(Ag, Ca, Sr, Ba, Hg, Pb)
4. CO32-, PO4
3-, CrO42-, S2-
(Ag, Ca, Sr, Ba, Hg, Pb)
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
balanced molecular equation
(table to determine which compound precipitates)
balanced ionic equation
1. NaCl(s) → Na+(aq) + Cl-(aq)
2. AgNO3(s) → Ag+(aq) + NO3-(aq)
3. Na+(aq) + Cl-(aq) + Ag+(aq)+ NO3-(aq) →
AgCl(s) + Na+ (aq) + NO3-(aq)
spectator ions
Ba(NO3)2(aq) + Na3PO4(aq)
1. which compound falls out? 2. balanced molecular equation
3. balanced ionic equations4. identify spectator ions
Cs2CrO4(aq) + Pb(NO3)2(aq)
Ba(NO3)2 (aq) + Na2SO4 (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) + H2O
HAc (aq) + MOH (aq) → MAc (aq) + H2O
and
NEUTRALIZE EACH OTHER
acid + base salt + water
H+ ≈ 10-15 m
Na+≈ 10-10 m
ACIDS AND BASES
ACIDS AND BASES
HCl (g) → H+ (aq) + Cl- (aq)
H+(aq) + H2O H3O+(aq)
HCl (g) + H2O → H3O+ (aq) + Cl- (aq)
one stephydronium ion
(aq) (l) (aq) (aq)
hydronium ion
acid base
cation hydronium ion
PROPERTIES OF ACIDS
1. acids have a sour taste
vinegar – acetic acidlemons – citric acid
2. acids react with some metals to form hydrogen
2 HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
3. acids react with carbonates to water and carbon dioxide
2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3]H2CO3 → H2O(l) + CO2(g)
EXP8
EXP9
4. some acids are hygroscopic
H2SO4 (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 H2SO4(aq)
NaOH(aq) and H3PO4(aq)
(stepwise)
Bronsted (1932)
ACIDS
HAc → H+ (aq) + Ac- (aq)
proton donors
BASES
proton acceptor
B + H+ (aq) → BH+ (aq)
weak electrolyte
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
NH3(aq) + H2O(l) NH4+ + OH-
strong electrolyte
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
HNO3(aq) + H2O(l) → H3O+(aq) + NO3-(aq)
donor versus acceptor
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
NH3(aq) + H2O(l) NH4+(aq)+ OH-(aq)
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
water can be either an acid or a base
AUTO DISSOCIATION
monoprotic acids
diprotic acid
HF, HCl, HBr, HNO3, CH3COOH
H2SO4 → H+(aq) + HSO4-(aq)
HSO4-(aq) H+(aq) + SO4
2-(aq)
triprotic acid
H3PO4 H+(aq) + H2PO4-(aq)
H2PO4-(aq) H+(aq) + HPO4
2-(aq)
HPO42-(aq) H+(aq) + PO4
3-(aq)
EXP10
CHEMICAL PROPOERTIES
1. Non-metal oxides react with water to form an acid
(acetic anhydrides)
3 2 2 4
2 5 2 3
2 2 2 3
SO ( ) H O H SO ( ) sulfuric acid
N O ( ) H O 2HNO ( ) nitric acid
CO ( ) H O H CO ( ) carbonic acid
g aq
g aq
g aq
Cl2O7, SO2, Br2O5
+ H2O
+ H2O
+ H2O
CHEMICAL PROPERTIES
2. Soluble metal oxides react with water to form a base
(base anhydrides)
MgO, Al2O3
2 2
2 2
CaO( ) H O Ca(OH) ( ) calcium hydroxide
Na O( ) H O 2NaOH( ) sodium hydroxide
s aq
s aq
+ H2O
+ H2O
NAMING ACIDS AND BASES
2 2
HCl( ) hydrogen chloride HCl( ) chlor
H S( ) hydrogen sulfide H S( ) sulfur
g aq hydro ic acid
g aq hydro ic acid
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
HNO3 nitric acidHNO2 nitrous acid
H3PO4 phosphoric acidH3PO3 phosphorous acid
H2SO4 sulfuric acidH2SO3 sulfurous acid
main group 6
main group 7
HClO4 perchloric acidHClO3 chloric acidHClO2 chlorous acidHClO hypochlorous acid
Acids in the Solar System
Venus
H2SO4(g)
Europa
H2SO4(s)
Acids in the Interstellar Medium
Orion
NH3, H2O, H2S
CH3COOH
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
Cl2 + 2 e 2 Cl-
CaO, Al2O3
substance that lost the electrons reduction agent
substance that gained the electrons oxidizing agent
oxidizing agent is reduced
reducing agent is oxidized
2 Na + Cl2 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
NaCl HF, H2
Na+Cl- ?electrons are fully transferred covalent bond
charges an atom would have if electrons are transferred completely
HF H+ + F-
molecular compound ionic compound
F- oxidation state -1
H+ oxidation state +1
EXAMPLE 1
H2O
molecular compound ionic compound
2 H+ + O2-
H+ oxidation state +1
O2- oxidation state -2
EXAMPLE 2
H2
molecular compound ionic compound
H+ + H-
EXAMPLE 3
OXIDATION NUMBER OF FREE ELEMENTS IS ZERO
RULE 1
OXIDATION NUMBER OF FREE ELEMENTS IS ZERO
H2, O2, F2, Cl2, K, Ca, P4, S8
RULE 2
monoatomic ions
oxidation number equals the charge of the ion
group I M+
group II M2+
group III M3+ (Tl: also +1)
group VII (w/ metal) X-
RULE 3
oxidation number of hydrogen
+1 in most compounds
(H2O, HF, HCl, NH3)
-1 binary compounds with metals (hydrides)
(LiH, NaH, CaH2, AlH3)
RULE 4
oxidation number of oxygen
-2 in most compounds
(H2O, MgO, Al2O3)
-1 in peroxide ion (O22-) (H2O2, K2O2, CaO2)
-1/2 in superoxide ion (O2-) (LiO2)
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 (ClO4
-)
RULE 6
charges of polyatomic molecules must be integers
(NO3-, SO4
2-)
oxidation numbers do not have to be integers
-1/2 in superoxide ion (O2-)
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 H2O H2 + 2 NaOH
EXP11/12
NONO2
NO+
NO-NO2-
NO3-
PO43- SO4
2-
SO3
SO2
KO2
K2O
BrO-KClO4
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
S8(s) + O2(g) → SO2(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) + O2(g)
1. oxidation numbers
2. balancing charges
KClO3(s) → KCl(s) + O2(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) + H2O(l) → NaOH + H2(g)
less reactive metals form hydrogen and the oxide in water (group III, transition metals)
Al(s) + H2O(l) → Al2O3(s) + H2(g)
3.1. Hydrogen displacement
even less reactive metals form hydrogen in acids
Zn(s) + HCl(aq) → ZnCl2(aq) + H2(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
V2O5(s) + 5 Ca(s) → 2 V(s) + 5 CaO(s)
TiCl4(g) + 2 Mg (l) → Ti(s) + 2 MgCl2(l)
3.3. Halogen displacement
F2 > Cl2 > Br2 > I2
reactivity (‘likes’ electrons)
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)0 0+1+1 -1 -1
Br2(g) + 2 KI(aq) → 2 KBr(aq) + I2(s)
4. disproportionation reactions
an element in one oxidation state is oxidized and reduced
at the same time
H2O2(aq) → 2 H2O(l) + O2(g)
Cl2(g) + 2 OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(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)
solution of literssolute of moles (M)molarity =
How many grams of AgNO3 are needed to prepare250 mL of 0.0125 M AgNO3 solution?
30.531 g AgNO
How many mL of 0.124 M NaOH are required
to react completely with 15.4 mL of 0.108 M H2SO4?
26.8 mL NaOH
2 NaOH + H2SO4 Na2SO4 + 2H2O
How many mL of 0.124 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
0.600 M BaCl2 is mixed with 30.0 mL of 0.500 M MgSO4?
BaCl2 + MgSO4 BaSO4 + MgCl2
This is a limiting reagent problem
40.0120 mol BaSO
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?
dil dil concd concdV M V MX X=
LIMITING REACTANT
C2H4 + H2O C2H5OHEXP14
excess reactantlimiting reactant
How many grams of NO can form when 30.0 g NH3 and 40.0 g O2 react according to
4 NH3 + 5 O2 4 NO + 6 H2O
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