9.3 REACTIVITY
Reactive metals lose their outer electrons more readily
 form the cation rather than the elemental form
 more easily oxidized
 stronger reducing agents
 react faster (↑reaction rate) with acids or water
Reactive non-metals gain outer electrons more readily
 form the anion rather than the elemental form
 more easily reduced
 stronger oxidizing agent
TRENDS
METALS
 Representative Groups (1A, 2A, 3A)
Reactivity increases down a metal group  Li < Na < K
and
Be < Mg < Ca
 increased energy level  less electrostatic attraction
Elemental forms of potassium, sodium, and calcium are very rare and unstable
Reactivity decreases across a period

Na > Mg > Aℓ
1A > 2A > 3A
 related to ionization energies
 Transition metals (B) – have lower reactivities than Group A elements
Base metals (Zn, Fe, Pb) are more reactive than precious metals (Cu, Ag, Au, Pt)
 Hydrogen − less reactive than Group A and base transition metals
− more reactive than precious metals
 acids (H+1) oxidize base metals easier and get reduced to form hydrogen gas (H0)2
NON-METALS
 Halogens
Reactivity decreases down the halogens

F > Cℓ >Br > I
 increased energy level  less electrostatic attraction
Elemental fluorine is rare and displaces all other halogens to form compounds.
Elemental bromine and iodine are relatively stable.
 Molecular oxygen (O0)2 reacts with many compounds and elements to form oxides (O2−).
O2 is an oxidizing agent that is reduced
 Powerful Oxidizing Agents
Permanganate ion
MnO4−
Dichromate ion
Cr2O72−
Chromate ion
CrO42−
Nitrate
NO3−
Mn+7 manganese (VII)
Cr6+ chromium (VI)
Cr6+ chromium (VI)
N5+ nitrogen (V) – stronger oxidizing agent than H+
10.2.2. Deduce the feasibility of a redox reaction from a given reactivity series.
 Inclass Assignment /Homework
p.485 12.37, 12.39, 12.40, 12.41, 12.43
Electrochemistry: Voltaic Cells and Electrolytic Cells
Addison – Wesley: Ch 12 (p.503)
Brady – Holum: Ch 17 (p. 695)
Electrochemical processes  conversions between chemical energy and electrical energy.
Voltaic Cell = Galvanic Cell = wet cell

an electrochemical cell in which a spontaneous redox reaction produces electricity.

two half cells connected by conductors and a salt bridge
Spontaneous anode = most reactive metal (most easily oxidized)
cathode = least reactive metal (least easily oxidized)
Electrodes
– solid conductors (metal bars)
 identified by reaction not charge
Anode  oxidation location − voltaic
Cathode
 reduction location+ voltaic
+ electrolyic
− electrolytic
Electrolyte
 liquid containing ions (salt solution)
Half-cell
 the electrode and electrolye where either the oxidation or reduction occurs
Salt bridge
 a tube filled with an electrolyte (KNO3 or KCℓ ) that allows ions to be transferred between the halfcells and complete the circuit.
ex. Zn – Cu Voltaic Cell
Oxidation Half cell
Zn0  Zn+2 + 2 e−
Anode: Zinc
Electrolyte: zinc sulfate solution
electron acceptor  (−) charge
Zn(s) | ZnSO4 (aq)
Anode salt bridge
Salt Bridge
K2SO4
 SO42− 
|| CuSO4(aq) | Cu (s)
Cathode
Lead Storage Battery  rechargeable car battery
Oxidation Half cell
Pb0  Pb+2 + 2 e−
Anode: lead
Electrolyte: sulfuric acid solution
electron acceptor  (−) charge
OX
RED
NET
Reduction Half-cell
Cu+2 + 2 e−  Cu0
Cathode: copper
Electrolyte: copper (II) sulfate solution
electron donor  (+) charge
Reduction Half-cell
Pb+4 + 2 e−  Pb+2
Cathode: lead (IV) oxide PbO2
Electrolyte: sulfuric acid solution
electron donor  (+) charge
Pb(s) + SO42−(aq)  PbSO4(s) + 2e−
E0ox = + 0.36 V
+
2−
−
0
PbO2(s) + 4H (aq) + SO4 (aq) + 2e  PbSO4(s) + 2 H2O (ℓ) E red = + 1.69 V
Pb(s) + PbO2(s) + H2SO4 (aq)  2 PbSO4(s) + 2 H2O(ℓ)
E0net = + 2.05 V
Reverse Reaction  Recharge
supply 2.0 V of electricity = Electrolysis
2 PbSO4(s) + 2 H2O(ℓ) + 2.05V  Pb(s) + PbO2(s) + H2SO4 (aq)
Practice:
A-W p. 521 #6, 25, 39
Electrolytic Cells
electrolyisis = splitting using electricity
Electrolytic cells
 use electric (DC) currect
 cause nonspontaneous redox reactions form products
Anode = oxidation

Cathode = reduction
(same as voltaic)
Describe how current is conducted in an electrolytic cell.
 Electrons flow from the anode to the cathode in the external circuit.
 The DC cell pushes electrons to the cathode making it negatively charged.
 The negative cathode donates electrons for reduction.
 The DC cell attracts electrons from the anode making it positively charged.
 The positive cathode accepts electrons from oxidation.
ex. electrolysis of water
reaction 2 H2O  2 H2 + O2
Anode: water is oxidized
Cathode: water is reduced
2 H2O + 2e−  H2 + 2OH−
E0ox = − 0.42 V
+
−
RED 2 H2O  O2 + 4H + 4e
E0red = − 0.82 V
NET 2 H2O  2 H2 + O2
E0net = − 1.66 V
 Hydrogen gas is evolved at the anode
 Oxygen gas is eveolved at the cathode

OX
Deduce the products for the electrolysis of a molten salt.
ex. Electrolysis of molten sodium chloride (Down’s cell)
reaction 2 NaCℓ(ℓ)  2 Naℓ) + Cℓ2(g)
Anode: molten chloride ions are oxidized
OX
Cathode:
molten sodium ions are reduced
 Chlorine gas is evolved at the anode
 Molten sodium metal collects at the cathode
Voltaic Cell
Electrolytic Cell
2 Cℓ−(ℓ)  Cℓ2 (g) + 2e-RED 2 Na+(ℓ) + 2e−  2Na(ℓ)
Anode = oxidation
Cathode = reduction
electrode gains electrons
electrode loses electrons
negative electrode
positive electrode
positive electrode
negative electrode
accepts electrons for oxidation
provides electrons for reduction
Electroplating
Describe and explain the use of electrolysis in electroplating.
the application by electrolysis of a thin coating of one metal oven another
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Practice:
B-H
p. 730 # 17.16 – 17.18