Oxidation-Reduction Reactions

Redox Reactions Most chemical reactions involve transfer of electrons from one atom to another. These have come to be known as oxidation-reduction reactions, or redox reactions. What are some examples of this? Combustion reactions – oxidation The rusting of metals The way living systems produce and utilize energy The operation of a car battery These chemical changes are all classified as "electron- transfer" or oxidation-reduction reactions.

Oxidation/Reduction are archaic terms The origins of the term oxidation come from the observation that almost all elements reacted with oxygen to form compounds called oxides. A typical example is the corrosion or rusting of iron as described by the chemical equation: 4 Fe + 3 O 2 → 2 Fe 2 O 3

Reduction Reduction, was the term originally used to describe the removal of oxygen from metal ores, which "reduced" the metal ore to pure metal as shown below: 2 Fe 2 O C → 3 CO Fe

Oxidation: adding oxygen Reduction: removing oxygen Based on the two examples above, oxidation can be defined very simply as, the "addition" of oxygen; and reduction, as the "removal" of oxygen. In the 19 th century, before atomic theory was deeply understood, chemists observed simply that during oxidation, the masses and volumes of substances increased, while during reduction, the masses and volumes decreased – hence the term reduction. Oxidation and reduction occur simultaneously. Iron oxide is reduced, losing oxygen and carbon is oxidized, gaining oxygen.

Which statement correctly describes a redox reaction? The oxidation half-reaction and the reduction half-reaction occur simultaneously. The oxidation half-reaction occurs before the reduction half-reaction. The oxidation half-reaction occurs after the reduction half-reaction. The oxidation half-reaction occurs spontaneously but the reduction half-reaction does not.

Oxidation-reduction (redox) no oxidation happens without reduction. no reduction happens without oxidation. Hence the term “redox” reaction is commonly used.

Electron Transfer Before the discovery of the electron, chemists had no idea what was really going on in redox reactions. Today we know that oxidation is the loss of electrons, or gain of oxygen reduction is the gain of electrons, or the loss of oxygen. Taking a simple example: the production of table salt from its elements.

The production of NaCl requires the transfer of an electron from sodium to chlorine. We can divide this transfer into two “half reactions,” as shown below:

The substance in a redox reaction that donates electrons is called a reducing agent. The substance in a redox reaction that accepts electrons is called an oxidizing agent. In this reaction, which substance is the oxidizing agent and which is the reducing agent?

In an oxidation-reduction reaction, reduction is defined as the loss of protons gain of protons loss of electrons gain of electrons

Redox Reactions Involving Nonmetals Only The situation is a bit more complex when considering molecular compounds, which contain nonmetals exclusively. As all nonmetals have similarly high electronegativity values, we cannot assume that there will be a transfer of electrons between them in an oxidation-reduction reaction.

For the valence electrons involved can no longer be thought of as being "lost or gained" between the atoms Electrons are only partially transferred, moving closer to that atom which has the higher electronegativity (and away from the atom of lower electronegativity). This "shift" of electrons results in an unequal distribution of charge The more electronegative atom becomes more "negative" and the atom of lower electronegativity becomes more "positive".

The accurate determination of the distribution of charge resulting from these "electron shifts" is very difficult. However, guidelines have been devised to simplify the process. In general, - the more electronegative atom receives a negative oxidation state - the atom with the lower electronegativity receives a positive oxidation state. These conventions are arbitrary approximations.

Non-Metal Oxidation States The multiple oxidation numbers (other than the top one) are for bonding with other non-metals. They are bookkeeping devices to keep track of electrons in reactions.

Example: SO 2 Oxygen is more electronegative: ox # -2 We assume electrons are “owned” by oxygen So sulfur’s oxidation # is +4, to balance of the four electrons of the two oxygens.

In which compound does chlorine have an oxidation number of +7? HClO 4 HClO 3 HClO 2 HClO

Oxidation Numbers Chemists need to keep track of the movement of electrons, whether they be transferred or shared, in chemical reactions. A number of rules have developed over time to facilitate this; they work for both ionic and covalent compounds.

Each element is assigned an oxidation number – a positive or negative number. The oxidation number of a free element is 0. What is a free element? One that is not combined with another element. O 2, N 2, O 3, He, S 8, Au, are all examples of naturally occurring free elements.

In which substance does hydrogen have an oxidation number of zero? LiH H 2 O H 2 S H 2

Oxidation numbers of ions The oxidation number of a simple ion(not polyatomic) is its charge. For Na + the oxidation number is +1; for O -2 the oxidation number is –2.

Alkali Metals and Alkaline Earth Metals Alkali metals (group 1) have an oxidation number of +1. Alkaline Earths (group 2) have an oxidation number of +2.

Hydrogen Hydrogen in compounds usually has an oxidation number of +1. In hydrides, such as NaH, hydrogen has a charge of –1.

In which species does hydrogen have an oxidation number of -1? H 2 O H 2 NaH NaOH

Oxygen Oxygen in compounds usually has an oxidation number of –2. Perioxide ion, O 2 2- is an exception In oxygen-fluorine compounds, oxygen can have a positive oxidation number.

Compounds and Polyatomic Ions In both molecular and ionic compounds, the sum of the oxidation numbers must equal 0 – compounds are electrically neutral. In polyatomic ions, the sum of the oxidation numbers equals the charge on the ion.

What is the oxidation number of chromium in the chromate ion, CrO 4 2- ?

A general rule to follow when assigning oxidation #s: When you encounter a compound with three or more elements, find the oxidation numbers of the ones on either end first. The middle one is usually the hard one and is the difference between the positive charges of the left one (usually the metal) and the negative charges of the far right element (the non-metal).

Problems In the compound Al 2 O 3, what are the oxidation numbers of the elements? What is the sum of the oxidation numbers? Find the oxidation numbers of the elements in the following compounds and ions: H 2 CO 3 Na 2 Cr 2 O 7 FeCl 3 PO 4 3- OF 2 H 2 O 2

What is the oxidation number assigned to manganese in KMnO 4 ?

F2F2 LiHAlCl 3 CO 2 NO 3 - KClHCO 3 - CaH 2 NO 2 KClOPO 4 3- H2O2H2O2 NO 2 - KClO 2 NaIO 3 OF 2 Cr 2 O 7 2- KClO 3 FeBr 2 NH 3 O3O3 NaClO 4 H 2 SO 4 P2O5P2O5 Cu 2 OCuOPb(OH) 2 Pb(NO 3 ) 4 FeSFe 2 S 3 CaCO 3 NH 3

Base your answer to the question on the information below. Aluminum is one of the most abundant metals in Earth’s crust. The aluminum compound found in bauxite ore is Al 2 O 3. Over one hundred years ago, it was difficult and expensive to isolate aluminum from bauxite ore. In 1886, a brother and sister team, Charles and Julia Hall, found that molten (melted) cryolite, Na 3 AlF 6, would dissolve bauxite ore. Electrolysis of the resulting mixture caused the aluminum ions in the Al 2 O 3 to be reduced to molten aluminum metal. This less expensive process is known as the Hall process. Write the oxidation state for each of the elements in cryolite, Na 3 AlF 6 :

In which substance is the oxidation number of Cl equal to +1? Cl 2 Cl 2 O AlCl 3 HClO 2

Oxidation Numbers in Reactions Let’s now consider what happens to elements that participate in redox reactions. Remember, for any element, oxidation is a loss of electrons reduction is a gain of electrons Let’s take a familiar chemical reaction and break down what is happening: Zn(s) + HCl(aq) → ZnCl 2 (aq) + H 2 (g)

In a single replacement reaction such as this, what exactly is happening to the zinc? Let’s look at each element one at a time and apply the oxidation rules: Zn – according to rule 1, the oxidation number of a pure element is 0. H in HCl – according to rule 4, hydrogen has an oxidation number of +1. Cl in HCl – we can determine the oxidation number of Cl in a number of ways. Cl must be assigned an oxidation number of –1 to balance hydrogen.

Zn(s) + HCl(aq) → ZnCl 2 (aq) + H 2 (g) Zn in ZnCl 2 – Zn appears to have lost electrons here. How many? If the chloride ion remains unchanged, the total charge on the two chloride ions is –2. So Zn becomes Zn 2+. Has zinc been oxidized or reduced? H in H 2 –Hydrogen has an oxidation number of 0 because it exists in its elemental state. What has happened to hydrogen in this reaction?

Let’s break this reaction into half-reactions: Zn 0  Zn e - Zn is oxidized – it loses 2 electrons 2H + + 2e -  H 2 Two hydrogen ions are reduced – they gain two electrons.

Which statement correctly describes what occurs when this reaction takes place in a closed system? Atoms of Zn(s) lose electrons and are oxidized. Atoms of Zn(s) gain electrons and are reduced. There is a net loss of mass. There is a net gain of mass.

Redox Equation Worksheet Assign oxidation numbers to each element Write the half reactions for the oxidized element and reduced element (ignore any element that is not changed in the reaction) Identify the oxidizing agent and reducing agent Balance each equation

P 4 + Cl 2 → PCl 3 Pb + AgNO 3 → Ag + Pb(NO 3 ) 2 Al + Fe 3 O 4 → Al 2 O 3 + Fe NO 2 → NO + O 2

Na 3 N → Na + N 2 Ca + H 2 O → Ca(OH) 2 + H 2 Fe + O 2 → Fe 2 O 3 Cl 2 + KI → KCl + I 2 Ag 2 O → Ag + O 2

Given the reaction: 2Na(s) + 2H 2 O(l) → 2NaOH(aq) + H 2 (g) Which substance undergoes oxidation? Na NaOH H 2 H 2 O

Given the redox reaction: Cr 3+ + Al → Cr + Al 3+ As the reaction takes place, there is a transfer of electrons from Al to Cr 3+ electrons from Cr 3+ to Al protons from Al to Cr 3+ protons from Cr 3+ to Al

When the equation __Pb 2+ + __Au 3+ → __Pb 4+ + __Au is correctly balanced using the smallest whole- number coefficients, the coefficient of Pb 2+ will be

In the reaction: 2Al(s) + 3Cu 2+ (aq) → 2Al 3+ (aq) + 3Cu(s), the Al(s) gains protons loses protons gains electrons loses electrons

Recognizing Redox Reactions we learned that reactions can be classified into four general categories – Combination Decomposition Single Replacement Double Replacement Combustion, because of its importance, is often considered along with the other four.

In the reaction: Cl 2 + H 2 O → HClO + HCl, the hydrogen is oxidized, only reduced, only both oxidized and reduced neither oxidized nor reduced

Redox or not? This classification system is based on the types of reactants involved. If we consider only the electrons of a chemical system, we can classify reactions another way: oxidation-reduction reactions, where electrons are transferred reactions where electrons are not transferred

Synthesis (combination) Reactions How does the the old way of looking at reactions apply to the new way? Take a typical synthesis reaction: H 2 (g) + O 2 (g) → H 2 O(g) Look at the oxidation numbers! Combination reactions are generally redox reactions.

Decomposition What about decomposition? NH 4 NO 2 (s) → N 2 (g) + 2H 2 O(l) Decomposition reactions are also usually redox reactions.

Single Replacement Single replacement reactions? Zn + 2HCl → ZnCl 2 + H 2 Single replacement reactions are always redox reactions.

Double Replacement NaOH + HCl → NaCl + H 2 O Double replacement reactions and acid base reactions are not redox reactions.

Combustion? C 6 H 12 O 6 + 6O 2 → 6CO 2 + 6H 2 O Combustion reactions are redox reactions.

Given the equations A, B, C, and D: AgNO 3 + NaCl → AgCl + NaNO 3 Cl 2 + H 2 O → HClO + HCl CuO + CO → CO 2 + Cu NaOH + HCl → NaCl + H 2 O Which two equations represent redox reactions? (1) A and B (3) C and A (2) B and C (4) D and B