Rules for Assigning Oxidation #’s

Elements by themselves are neutral Elements by themselves are neutral Ox # = 0 Ex: Na 0, O 2 0, Mg 0 Ex: Na 0, O 2 0, Mg 0 Monoatomic Ions show their oxidation # as their written charge Monoatomic Ions show their oxidation # as their written charge Ex: H +1, Cl -1, Al +3 Ex: H +1, Cl -1, Al +3

Group #1 Elements in compounds: Group #1 Elements in compounds: Ox # = +1Ex: LiBr, NaCl Group #2 Elements in compounds: Group #2 Elements in compounds: Ox # = +2Ex: MgBr 2, CaCl 2

Aluminum in compounds: Aluminum in compounds: Ox # = +3Ex: AlCl 3 Fluorine in compounds Fluorine in compounds Ox # = -1Ex: LiF, MgF 2

Hydrogen: Hydrogen: Ox # = +1 (most common) Ox # = +1 (most common) when combined with nonmetals when combined with nonmetals Ex: HCl, H 2 SO 4, H 2 S Ex: HCl, H 2 SO 4, H 2 S Ox # = -1 Ox # = -1 When combined with metals (metal hydrides) When combined with metals (metal hydrides) Ex: LiH, MgH 2 Ex: LiH, MgH 2

Oxygen: Oxidation # = -2 (usually) Oxygen: Oxidation # = -2 (usually) Ex: CO 2, H 2 O Ex: CO 2, H 2 O Special Cases: Special Cases: Peroxides: Oxygen is a diatomic ion (O 2 ) -2 with total charge of -2, each oxygen has -1 charge Peroxides: Oxygen is a diatomic ion (O 2 ) -2 with total charge of -2, each oxygen has -1 charge Ex: H 2 O 2, BaO 2 Ex: H 2 O 2, BaO 2 When bonded with Fluorine: Ox # = +2 When bonded with Fluorine: Ox # = +2 Ex: OF 2 Ex: OF 2

How to Find Charges If No “Rule”? Sum of Individual Ox # of all the elements in a compound = 0 Sum of Individual Ox # of all the elements in a compound = 0 All compounds are NEUTRAL All compounds are NEUTRAL Ex: NaClO 3 Ex: NaClO 3 What is the ox # for each element? What is the ox # for each element? Draw your Lines…. Draw your Lines….

How to Find Charges If No “Rule”? Sum of Ox # of all atoms in a polyatomic ion is equal to the charge of that ion. Sum of Ox # of all atoms in a polyatomic ion is equal to the charge of that ion. Ex: PO 4 -3 Ex: PO 4 -3 What is the charge of each element? What is the charge of each element?

Let’s Practice What is the oxidation number of sulfur in Na 2 S 2 O 3 ? What is the oxidation number of sulfur in Na 2 S 2 O 3 ? (1) -1 (1) -1 (2) +6 (3) +2 (3) +2 (4) +4

What is the oxidation number of chromium in the chromate ion, CrO 4 2– ? What is the oxidation number of chromium in the chromate ion, CrO 4 2– ? (1) +6 (1) +6 (2) +3 (3) +2 (3) +2 (4) +8

In which compound does chlorine have the highest oxidation number? In which compound does chlorine have the highest oxidation number? (1) NaClO (1) NaClO (2) NaClO 3 (3) NaClO 2 (3) NaClO 2 (4) NaClO 4

What is the oxidation number assigned to manganese in KMnO 4 ? What is the oxidation number assigned to manganese in KMnO 4 ? (1) +7 (1) +7 (2) +3 (3) +2 (3) +2 (4) +4

Mahjong Oxidation # Game: Mahjong Oxidation # Game:

Oxidation and Reduction Reactions

Oxidation (Read only) Original definition: When substances combined with oxygen. Ex: All combustion (burning) reactions CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) All “rusting” reactions 4Fe(s) + 3O 2 (g)2Fe 2 O 3 (s)

Reduction (Read Only) Original Definition: Reaction where a substance “gave up” oxygen. Called “reductions” because they produced products that were “reduced” in mass because gas escaped. Ex: 2Fe 2 O 3 (l) + 3C(s) 4Fe(l) + 3CO 2 (g)

Oxidation/Reduction Reactions Oxidation: LOSS of one or more electrons. Reduction: GAIN of one or more electrons Oxidation & reduction always occur together.

Deal with a movement of ELECTRONS Deal with a movement of ELECTRONS between atoms during a reaction. Electrons travel from what is oxidized Electrons travel from what is oxidized towards what is reduced.

One atom loses e-, the other gains e- This is called “electron transfer”

Remember!!

Or…Remember

Redox Reactions: ALWAYS involve changes in charge If charges don’t change it’s NOT a Redox rxn.

Conservation of “Charge” Total electrons lost = Total electrons gained Redox Reactions must balance electrons as well as atoms!

Let’s Practice Which changes occur when Pt 2+ is reduced? (1) Pt 2+ gains electrons and its oxidation number increases. (2) Pt 2+ gains electrons and its oxidation number decreases. (3) Pt 2+ loses electrons and its oxidation number increases. (4) Pt 2+ loses electrons and its oxidation number decreases.

Answer: 1 Answer: 1

Given the balanced equation representing a redox reaction: 2Al + 3Cu 2+ → 2Al Cu Which statement is true about this reaction? (1) Each Al loses 2e- and each Cu 2+ gains 3e- (2) Each Al loses 3e- and each Cu 2+ gains 2e- (3) Each Al 3+ gains 2e- and each Cu loses 3e- (4) Each Al 3+ gains 3e- and each Cu loses 2e-

Answer: 2 Answer: 2

In an oxidation-reduction reaction, reduction is defined as the In an oxidation-reduction reaction, reduction is defined as the (1) loss of protons (1) loss of protons (2) loss of electrons (3) gain of protons (3) gain of protons (4) gain of electrons

Which change in oxidation number indicates oxidation? Which change in oxidation number indicates oxidation? (1) –1 to +2 (2) +2 to –3 (3) –1 to –2 (3) –1 to –2 (4) +3 to +2

When a neutral atom undergoes oxidation, the atom’s oxidation state When a neutral atom undergoes oxidation, the atom’s oxidation state (1) decreases as it gains electrons (2) decreases as it loses electrons (3) increases as it gains electrons (4) increases as it loses electrons

Oxidizing/Reducing Agents Oxidizing Agent: substance reduced Gains electrons Gains electrons Reducing Agent: substance oxidized Loses electrons Loses electrons The “Agent” is the “opposite”

Identify What is Changing in Charge What is oxidized and reduced? What are the oxidizing and reducing agents? Ex: 3Br 2 + 2AlI 3 2AlBr 3 + 3I 2

Br 2 + 2AlI 3 2AlBr 3 + 3I 2 Br 2 is reduced and is the oxidizing agent I -1 is oxidized and is the reducing agent

What is oxidized and reduced? What are the oxidizing and reducing agents? Mg + CuSO 4 MgSO 4 + Cu 2K + Br 2 2KBr Cu + 2AgNO 3 Cu(NO 3 ) 2 + 2Ag NOTE: Atoms in a polyatomic ion DO NOT change in charge!

Mg + CuSO 4 MgSO 4 + Cu Mg oxidized (reducing agent) Cu +2 reduced (oxidizing agent) K + Br 2 2KBr K oxidized (reducing agent) Br 2 reduced (oxidizing agent) Cu + 2AgNO 3 Cu(NO 3 ) 2 + 2Ag Cu oxidized (reducing agent) Ag +1 reduced (oxidizing agent)

Redox or Not Redox (that is the question…) Redox Reactions: must have atoms changing in charge. Not all reactions are redox. Easy way to spot a redox reaction!!! Look for elements entering and leaving compounds.

Is it Redox? Look for Changes in Charge! Are elements entering and leaving compounds? Synthesis: Ex:2H 2 + O 2 2H 2 O Decomposition: Ex:2KClO 3 2KCl + 3O 2

Is it Redox? Synthesis: YES Ex:2H 2 + O 2 2H 2 O Decomposition: YES Ex:2KClO 3 2KCl + 3O 2

Is it Redox? Combustion: CH 4 + 2O 2 CO 2 + 2H 2 0 Single Replacement: Zn + CuCl 2 ZnCl 2 + Cu

Is it Redox? Combustion:YES CH 4 + 2O 2 CO 2 + 2H 2 0 Single Replacement:YES Zn + CuCl 2 ZnCl 2 + Cu

Is it Redox? Double Replacement: AgNO 3 + LiCl AgCl + LiNO 3

Is it Redox? Double Replacement: NO!!!! Ions switch partners, but don’t change in charge AgNO 3 + LiCl AgCl + LiNO 3 Remember charges of atoms inside polyatomic Ions do not change!

Let’s Practice Which equation represents an oxidation reduction reaction? (1) CH 4 + 2O 2 → CO 2 + 2H 2 O (2) H 2 SO 4 + Ca(OH) 2 → CaSO 4 + 2H 2 O (3) MgCrO 4 + BaCl 2 → MgCl 2 + BaCrO 4 (4) Zn(NO 3 ) 2 + Na 2 CO 3 → 2NaNO 3 + ZnCO 3

Which reaction is an example of an oxidation reduction reaction? (1) AgNO 3 + KI →AgI + KNO 3 (2) Cu + 2 AgNO 3 → Cu(NO 3 ) Ag (3) 2 KOH + H 2 SO 4 → K 2 SO 4 + 2H 2 O (4) Ba(OH) 2 + 2HCl → BaCl 2 + 2H 2 O

Writing Half Reactions Redox Reactions are composed of two parts or half reactions. Half Reactions Show: Element being oxidized or reduced. Change in charge # of moles of e - being lost or gained

Writing Half Reactions Na + F 2 2NaF Oxidation: Na Na e- or2Na 2Na e- Note: e- are “lost” (on the right of arrow) Reduction:F + 1e- F -1 orF 2 + 2e- 2F -1 Note: e- are “gained” (on the left of arrow)

Ox’s Have Tails!! Oxidation Half reactions always have “tails” of electrons Oxidation Half reactions always have “tails” of electrons Na Na e-

Zn + CuCl 2 ZnCl 2 + Cu Ox:ZnZn e- Red:Cu e-Cu Tutorial on Writing Half Reactions

Balancing Simple Redox Rxns Must be: Balanced for Mass ATOMS balance Balanced for Charge Total e- Lost = Total e- Gained

Applications of Redox Reactions Corrosion of Metals Metals gets oxidized by oxygen in the air forming metal oxides (rust) 4Fe(s) + 3O 2 (g) → 2Fe 2 O 3 (s) Prevention: Use paint, oil, “Plating” or attach to negative terminal of a battery. Gold doesn’t rust…Why?

Ships or pipes made of Fe will “rust away” into Fe +2 ions when oxidized. Ships or pipes made of Fe will “rust away” into Fe +2 ions when oxidized. Blocks of more reactive metals (see Table J) like Mg or Zn are attached to ships or underground pipes. Blocks of more reactive metals (see Table J) like Mg or Zn are attached to ships or underground pipes. Electrons travel from more reactive metal to less reactive iron preventing it from oxidizing. Electrons travel from more reactive metal to less reactive iron preventing it from oxidizing.

Aluminum forms an oxide coating that protects the metal from further corrosion. Aluminum forms an oxide coating that protects the metal from further corrosion.

Photograph Development involves oxidation and reduction of silver atoms and ions

Bleach acts on stains by oxidizing them, getting reduced in the process Explosives form gases like N 2 from nitrogen compounds!

Car Batteries Lead/Acid Battery Lead/Acid Battery Lead is both oxidized and reduced Lead is both oxidized and reduced Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq) → 2PbSO 4 (s) + 2H 2 O(l) W_KkEhCy68&safe=active

Reactivity of Metals Reference Table J Metals Higher on Table J are more ‘active” “Active” metals are easily oxidized (lose electrons), and good “reducing agents”

Copper replaces silver! Single Replacement Reactions: Elements higher on Table J “replace” elements that are lower. Cu 0 (s) + AgNO 3 (aq)Ag 0 (s) + CuNO 3 (aq) Ag 0 (s) + CuNO 3 (aq) wouldn’t happen!!!

Reactivity of Nonmetals Reference Table J Nonmetals higher on Table J are more “active” “Active” nonmetals are more easily reduced (gain electrons), and good “oxidizing agents”

Balancing Harder Redox Reactions (Honors)

Oxidation Number Method (Balancing in Acid Solution) Find ox #’s and use brackets to connect elements Find ox #’s and use brackets to connect elements changing in charge. Balance atoms changing in charge Balance atoms changing in charge Find total e- involved in each change Find total e- involved in each change If necessary balance e- by multiplication If necessary balance e- by multiplication Balance all other atoms except H and O Balance all other atoms except H and O Balance oxygen by adding H 2 O to side deficient Balance oxygen by adding H 2 O to side deficient Balance hydrogen by adding H +1 to side deficient Balance hydrogen by adding H +1 to side deficient Check for balance with respect to atoms and charge. Check for balance with respect to atoms and charge.

Half Reaction Method (Ion/Electron Method) (In acid solution) Separate equation into two “basic” half reactions Separate equation into two “basic” half reactions Balance all atoms except H and O Balance all atoms except H and O Balance oxygen by adding H 2 O Balance oxygen by adding H 2 O Balance hydrogen by adding H +1 Balance hydrogen by adding H +1 Balance charge by adding electrons to more positive side Balance charge by adding electrons to more positive side If necessary balance e- by multiplication If necessary balance e- by multiplication Add together half reactions and simplify Add together half reactions and simplify Check for balance of atoms and charge Check for balance of atoms and charge

Crash Course Chemistry: Redox Crash Course Chemistry: Redox A1HM3s&index=11&list=PL8dPuuaLjXtPH zzYuWy6fYEaX9mQQ8oGr A1HM3s&index=11&list=PL8dPuuaLjXtPH zzYuWy6fYEaX9mQQ8oGr A1HM3s&index=11&list=PL8dPuuaLjXtPH zzYuWy6fYEaX9mQQ8oGr A1HM3s&index=11&list=PL8dPuuaLjXtPH zzYuWy6fYEaX9mQQ8oGr