Transcript Slide 1

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
solute
solvent
substance in a large amount
substance in a small amount
N2
gas phase
(air)
O2
Ag
solid phase
(alloys)
Au
H2O
liquid phase
(sea water)
NaCl
EXP1
iodine in ethyl alcohol (C2H5OH)
does not conduct electricity
(molecular solid)
I2
EXP2
table salt in water (H2O)
does conduct electricity
(ionic solid)
Na+Cl-
AQUEOUS
SOLUTION
solute
water (H2O)
solutes
solution conducts
electricity
solution does not
conduct electricity
EXP3
electrolytes
non-electrolytes
electrolytes
non-electrolytes
solution conducts
electricity
solution does not
conduct electricity
non-electrolyte
methanol
sugar
ethanol
water
weak electrolyte
CH3COOH
HCOOH
HF
EXP5
strong electrolyte
ionic compounds
(NaCl, KF)
NaOH
HCl
H2SO4
dark
medium
bright
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%
SOLUTION
solubility of a solute
number of grams of solute that can dissolve
in 100 grams of solvent at a given temperature
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
H 2O
solute
MgSO4
SOLID SOLUTION
Triton
solvent
N2
solute
CH4
ELECTROLYTES
methanol
sugar
ethanol
water
ionic compounds
CH3COOH
(NaCl, KF)
HCOOH
NaOH
HF
HCl
H2SO4
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)
Ca(NO3)2 (s) → Ca2+(aq) + 2 NO3- (aq)
strong electrolytes are fully dissociated
polyatomic ions do NOT dissociate
EXP5
δO
δ+
H
H
δ+
SOLVATION
cations
anions
SOLVATION
non-electrolyte
NaCl (s) → Na+ (aq) + Cl- (aq)
strong electrolytes are fully dissociated
+ (aq) + CH COO- (aq)
CH3COOH (aq) →
H
←
3
weak electrolytes are not fully dissociated
reversible reaction
(chemical equilibrium)
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-, PO43-, CrO42-, S2(Ag, Ca, Sr, Ba, Hg, Pb)
balanced molecular equation
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
(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) + Na2SO4 (aq)
Ba(NO3)2(aq) + Na3PO4(aq)
Cs2CrO4(aq) + Pb(NO3)2(aq)
1. which compound falls out?
2. balanced molecular equation
3. balanced ionic equations
4. identify spectator ions
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
ACIDS
HAc → H+ (aq) + Ac- (aq)
ionization
HCl (g) → H+ (aq) + Cl- (aq)
BASES
MOH → M+ (aq) + OH- (aq)
Arrhenius (1883)
NaOH (s) → Na+ (aq) + OH- (aq)
IDENTIFICATION
Litmus Paper
acid
red
Säure
base
blue
Base
EXP7
ACIDS AND BASES
ACIDS
BASES
and
NEUTRALIZE
EACH OTHER
HAc (aq) + MOH (aq) → MAc (aq) + H2O
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O
acid + base
salt + water
ACIDS AND BASES
Na+
H+
≈ 10-10 m
≈ 10-15 m
ACIDS AND BASES
HCl (g) → H+ (aq) + Cl- (aq)
H+(aq) + H2O
H3O+(aq)
HCl (g) + H2O → H3O+ (aq) + Cl- (aq)
one step
hydronium ion
(aq)
acid
(l)
(aq)
base
hydronium ion
(aq)
cation
hydronium 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) → MgCl2(aq) + H2(g)
EXP8
3. acids react with carbonates to water and carbon dioxide
2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3]
H2CO3 → H2O(l) + CO2(g)
4. some acids are hygroscopic
H2SO4 (conc)
EXP9
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)
ACIDS
proton donors
HAc → H+ (aq) + Ac- (aq)
BASES
proton acceptor
Bronsted (1932)
B + H+ (aq) → BH+ (aq)
strong electrolyte
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
HNO3(aq) + H2O(l) → H3O+(aq) + NO3-(aq)
weak electrolyte
CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3COO-(aq)
NH3(aq) + H2O(l)
NH4+ + OH-
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
HF, HCl, HBr, HNO3, CH3COOH
diprotic acid
H2SO4 → H+(aq) + HSO4-(aq)
HSO4-(aq)
H+(aq) + SO42-(aq)
triprotic acid
H3PO4
H+(aq) + H2PO4-(aq)
H2PO4-(aq)
H+(aq) + HPO42-(aq)
HPO42-(aq)
H+(aq) + PO43-(aq)
EXP10
CHEMICAL PROPOERTIES
1. Non-metal oxides react with water to form an acid
(acetic anhydrides)
SO3 (g ) + H22O
O  H 2SO 4 (aq)
sulfuric acid
H22O  2HNO3 (aq) nitric acid
N 2 O5 (g ) + H
CO 2 (g ) +
 H22O  H 2CO3 (aq)
Cl2O7, SO2, Br2O5
carbonic acid
CHEMICAL PROPERTIES
2. Soluble metal oxides react with water to form a base
(base anhydrides)
CaO(s) +HH22OO Ca(OH) 2 (aq) calcium hydroxide
Na 2 O(s) +HH2O
2O 2NaOH(aq )
MgO, Al2O3
sodium hydroxide
NAMING ACIDS AND BASES
binary acids
prefix hydrothe suffix –ic
to the stem of the nonmetal name followed by the word acid
HCl(g ) hydrogen chloride
HCl(aq) hydrochloric acid
H 2S(g ) hydrogen sulfide
H 2S(aq) hydrosulfuric acid
NAMING ACIDS AND BASES
oxo acids acids
contain hydrogen, oxygen, plus another element
main group 5
HNO3
HNO2
nitric acid
nitrous acid
H3PO4
H3PO3
phosphoric acid
phosphorous acid
main group 6
H2SO4
H2SO3
sulfuric acid
sulfurous acid
main group 7
HClO4
HClO3
HClO2
HClO
perchloric acid
chloric acid
chlorous acid
hypochlorous acid
Acids in the Solar System
Europa
Venus
H2SO4(s)
H2SO4(g)
Acids in the Interstellar Medium
NH3, H2O, H2S
CH3COOH
HCOOH
HF, HCl
Orion
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)
KEY CONCEPTS
1. oxidation
loss of electrons
2. reduction
acceptance of electrons
NUMBER OF ELECTRONS MUST BE CONSERVED
EXAMPLE
Na+Cl-
1. oxidation
Na  Na+ + e
2. reduction
Cl2 + 2 e  2 Cl-
!!!balance electrons!!!
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
EXAMPLE 1
H+ + F-
HF
molecular compound
ionic compound
H+
oxidation state +1
F-
oxidation state -1
EXAMPLE 2
H2O
2 H+ + O2-
molecular compound
ionic compound
H+
oxidation state +1
O2-
oxidation state -2
EXAMPLE 3
H2
molecular compound
H+ + Hionic compound
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-, SO42-)
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
K2O
PO43-
NO+
SO42-
KO2
SO3
NO3NO2
NO
KClO4
BrO-
SO2
NO2-
NO-
REVISION
1.redox reactions
2. oxidation versus reduction
3. oxidation numbers versus charges
4. calculation of oxidation numbers
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)
KClO3(s) → KCl(s) + O2(g)
1. oxidation numbers
2. balancing charges
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
activity series of metals
Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au
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
does not like so much to donate electrons
likes 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)
0
+1 -1
+1 -1
0
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
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)
molarity(M) 
molesof solute
= litersof solution
How many grams of AgNO3 are needed to prepare
250 mL of 0.0125 M AgNO3 solution?
0.531 g AgNO3
How many mL of 0.124 M NaOH are required
to react completely with 15.4 mL of 0.108 M H2SO4?
2 NaOH + H2SO4
Na2SO4 + 2H2O
26.8 mL NaOH
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
0.0120 mol BaSO 4
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?
Vdil X M dil
=
Vconcd X M concd
LIMITING REACTANT
EXP14
C2H4 + H2O
C2H5OH
limiting reactant
excess 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