REACTIONS IN AQUEOUS SOLUTIONS

WHAT IS AN AQUEOUS SOLUTION?

WHAT HAPPENS WHEN STUFF DISSOLVES?

When some ionic compounds are placed in water, they dissociate, or break apart into their respective cation/anion pairs. In this Aqueous state, they are more readily

available to react chemically with other substances in the solution than if they were solid.

Ionic compounds are made up of two halves. Cation

Comes first. Positive charge. Can be mono- or poly- atomic.

Anion Comes second. Negative charge. Can be mono- or poly- atomic.

PRACTICE: DISSOCIATE THE FOLLOWING IONIC COMPOUNDS…

CaCl2

Pb(NO3)2

Li3P

Na2SO4

(NH4)3PO3

SOLUBILITY When a substance readily dissolves in water

it is said to be soluble in water. Solubility is the mass of substance that will

dissolve in water at a given temperature. Vocabulary

Solute – What is being dissolved. Solvent – What is doing the dissolving. Solution – Solvent with Solute dissolved in it. Soluble – will dissolve. Insoluble – will not dissolve. Slightly Soluble – will only dissolve a little bit. Moderately Soluble – will dissolve more than a

little but not all the way.

WHY DO SOME THINGS DISSOLVE?

Different Substances dissolve to different extents.

Whether or not something will dissolve in water depends on two things:1. Polarity

The generic rule of solubility is like dissolves like. Water is a polar solvent and will only dissolve polar

things. All ionic compounds, acids, and bases are polar but

they do not all dissolve in water.

2. Bond Energy Bond energy holds ions in an ionic compound together. In order to dissolve, the polar nature of water must be

strong enough to overcome the attraction between the ions in a compound.

THE POLARITY OF WATERWhy is water such a good solvent?

•Water is a polar molecule.

•That means that the electrons are unevenly shared.

•The oxygen in water hogs the electrons in its bonds with hydrogen, giving it a partial negative charge.

•The hydrogen side is left with a partial positive charge.

•When ionic compounds are placed into water, they break apart because the partial positive and negative sides of water attack the negative and positive ions respectively.

•In other words, the ionic compounds are pulled apart by the polar water molecules.

WATER BREAKS APART A SOLUBLE SALT

Remember, Salt is another word for Ionic compounds.

HOW TO TELL IF SOMETHING DISSOLVES …

ONLY STRONG ELECTROLYTES DISSOLVE COMPLETELY IN H2O

Strong Electrolytes are completely soluble in water. Strong Acids Strong Bases Soluble Salts

Ionic compounds are sometimes called salts.

Weak electrolytes are only moderately or slightly soluble in water. Weak acids Weak bases Moderately or slightly soluble salts.

STRONG ACIDS & BASES

HClHI

HBrHNO3

HClO3

H2SO4

KOHNaOHLiOHCsOHRbOH

Ba(OH)2

Ca(OH)2

Sr(OH)2

Strong Acids Strong Bases

If it is not on the list of strong acids or bases, it can be considered weak.

TABLE 5-3 SIMPLE RULES FOR THE SOLUBILITY OF SALTS IN WATER (MODIFIED FROM TEXT)

1. Most nitrate (NO3-) salts are soluble.

2. Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+ & Rb+) and the ammonium ion (NH4

+) are soluble.

3. Most chloride, bromide and iodide salts are soluble. Notable exceptions are salts containing the ions Ag+, Pb2+, & Hg2

2+

4. Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4, Hg2SO4, & CaSO4.

5. Most hydroxide salts are only slightly soluble. The important soluble hydroxides are NaOH & KOH. The compounds Ba(OH)2, Sr(OH)2, & Ca(OH)2 are marginally soluble.

6. Most sulfide (S2-), carbonate (CO32-), and phosphate

(PO43-) salts are only slightly soluble.

SIDE BY SIDE COMPARISON

Soluble Salts (aqueous in water)

NO3-

Alkali metal ions Li +, K +, Na +, Cs +, Rb+

NH4+

Most Cl-, Br-, & I-

Most SO42-

NaOH, KOH Moderately = Ba(OH)2,

Ca(OH)2

Slightly or Insoluble Salts (solid in water)

AgCl, PbCl2, Hg2Cl2 AgBr, PbBr2, Hg2Br2

AgI, PbI2, Hg2I2 BaSO4, PbSO4, CaSO4

OH-

S2-

CO32-

PO43-

PRACTICE: DETERMINE SOLUBILITY

Write an equation to show what happens to each substance in water. CaF2

HF

Mg(OH)2

Ca3(PO4)2

HCl

KOH

IONS IN AQUEOUS SOLUTION ARE ABLE TO REACT… THE QUESTION IS,

“WILL THEY?”

WHAT HAPPENS WHEN A REACTION TAKES PLACE?

When a pair of ions comes in contact in an aqueous environment, there is a chance that they will react to form a compound that is NOT soluble in water.

This new compound will form and fall out of solution. If a solid forms, a non-uniform cloud will form in the

solution. If a gas forms, bubbles will appear and escape the

solution. If liquid forms, it can be hard to see easily, unless

there is a color change, or the new liquid is immiscible in water.

DETERMINING THE FORMULA OF A POSSIBLE PRODUCT…

Make a list of the ions in solution.

Pair them up, remember make cation-anion pairs.

Some Rules to help decide on products:1. The products will NOT be the same as the

reactants.2. Solid, liquid, and gas compounds must have a

zero net charge. Substances with ionic charges are ALWAYS aqueous.

3. They must contain both cations and anions. Most compounds contain only ONE OF EACH.

WHAT MAKES CHEMISTRY HAPPEN?

o.O

DRIVING FORCE A chemical reaction will only take place if

there is a force driving it forward. Remember entropy? Everything wants to attain the lowest possible

state of energetic existence. If a reaction leads to a state of lower energy,

then reactants will have a tendency to form products.

Driving Force – a force that leads reactants to react chemically and form products. The Driving Force causes a reaction to go in the direction of the arrow.

4 COMMON DRIVING FORCES1. Formation of a solid

A Precipitation Reaction occurs when two aqueous solutions are combined and a solid precipitates, or forms and falls out of solution.

2. Formation of water An Acid Base Reaction occurs when bases act to

neutralize acids, and water forms.

3. Transfer of electrons In a RedOx Reaction, one substance loses electrons in

an Oxidation Reaction, and another substance gains electrons in a Reduction Reaction.

4. Formation of a gas Similar to formation of a solid, except the product is a

gas.

Understanding the driving force behind a chemical reaction can aid in predicting the products of a chemical reaction.

DRIVING FORCES:#1 FORMATION OF A SOLID

Solubility Rules and Precipitation Reactions

PRECIPITATION REACTIONS Precipitation Reactions occur when two

solutions are mixed and an insoluble solid falls out of solution. The solid product is called a precipitate.

Precipitation Reactions are also classified as Double Displacement Reactions. This is because the reactants contain cation-anion

pairs that swap partners, or displace each other. A reaction will occur between two ionic pairs IF one

of the possible products is insoluble. Insoluble products precipitate driving the reaction

forward. To know if either possible product will be a

precipitate, you must know the Solubility Rules which can be found on Table 7.1 (page 178).

SIDE BY SIDE COMPARISON

Soluble Salts (aqueous in water)

NO3-

Alkali metal ions Li +, K +, Na +, Cs +, Rb+

NH4+

Most Cl-, Br-, & I-

Most SO42-

NaOH, KOH Moderately = Ba(OH)2,

Ca(OH)2

Slightly or Insoluble Salts (solid in water)

AgCl, PbCl2, Hg2Cl2 AgBr, PbBr2, Hg2Br2

AgI, PbI2, Hg2I2 BaSO4, PbSO4, CaSO4

OH-

S2-

CO32-

PO43-

PREDICTING PRECIPITATION REACTIONS When two solutions of aqueous salts are

mixed, the solubility rules can be used to predict whether or not a reaction will occur:1. Write the formulas of the two salts as if they

were reactants.2. Determine the formulas of the possible

products by swapping the anions of the two salts.

3. Check the solubility rules.

If one or both of the possible products is insoluble (or marginally or slightly soluble) than a reaction will occur and a precipitate will form.

USE THE SOLUBILITY RULES TO PREDICT THE PRODUCTS OF A REACTION:

Predict what will happen when the following salts are mixed in aqueous solution. Write the balanced equation for any reaction that occurs.

1. potassium nitrate & barium chloride2. sodium sulfate & lead(II) nitrate3. potassium hydroxide and iron(III) nitrate

KNO3(aq) + BaCl2(aq) no rxn

Na2SO4(aq) + Pb(NO3)2(aq) PbSO4(s) + 2NaNO3(aq)

3KOH(aq) + Fe(NO3)3(aq) Fe(OH)3(s) + 3KNO3(aq)

USE THE SOLUBILITY RULES TO PREDICT THE PRODUCTS OF A REACTION:

Predict what will happen when the following salts are mixed in aqueous solution. Write the balanced equation for any reaction that occurs.

1. barium nitrate & sodium chloride2. sodium sulfide & copper(II) nitrate3. ammonium chloride & lead(II) nitrate

Ba(NO3)2 (aq) + NaCl(aq) no rxn

Na2S (aq) + Cu(NO3)2(aq) CuS (s) + 2NaNO3(aq)

2NH4Cl(aq) + Pb(NO3)2(aq) PbCl2(s) + 2NH4NO3

(aq)

TYPES OF CHEMICAL EQUATIONS

DESCRIBING REACTIONS IN AQUEOUS SOLUTION

There are several different ways that a single chemical reaction can be represented.

Different styles of chemical equations are used to express relative information as needed.

We will discuss three.1. Molecular Equations2. Complete Ionic Equations3. Net Ionic Equations

CONSIDER THE FOLLOWING REACTION An aqueous solution of potassium chromate

is added to an aqueous solution of barium nitrate and a solid precipitate forms.

Can you write a chemical equation to represent the reaction that takes place?

The equation you wrote is a molecular equation. We will now compare this equation to two new equations that show the same chemical reaction in different ways, the complete ionic equation and the net ionic equation.

1. MOLECULAR EQUATIONS

Molecular equations show the complete formulas of all reactants and products.

The molecular equation for the potassium chromate and barium nitrate reaction is:

K2CrO4(aq) + Ba(NO3)2 BaCrO4(s) + 2KNO3(aq)

These are the type of equations studied when first learning how to balance chemical equations. (Chapter 6)

2. COMPLETE IONIC EQUATIONS

In a complete ionic equation, all substances that are strong electrolytes and in the aqueous are represented as ions. Reminder: Strong Electrolytes = Strong Acids, Strong Bases, and Soluble Salts

Which compounds dissociate in our example equation?K2CrO4(aq) + Ba(NO3)2 BaCrO4(s) + 2KNO3(aq)

2K+(aq) + CrO4

2-(aq) + Ba2+

(aq) + 2NO3-(aq) BaCrO4(s) + 2K+

(aq) + 2NO3-

(aq)

This type of equation better represents the physical state of the reactants and products as they occur in the reaction environment. The precipitate is not represented as ions because it falls out

of solution as an insoluble solid.

3. NET IONIC EQUATIONS & SPECTATOR IONS

The Net Ionic Equation shows only those components that are directly involved in the chemical change that occurs.

The Spectator Ions are those components that are present as both reactants and products, and remain unchanged by the reaction.

Consider our Example:2K+(aq) + CrO4

2-(aq) + Ba2+(aq) + 2NO3-(aq) BaCrO4(s) + 2K+(aq) + 2NO3

-(aq

1. What is the net ionic equation? CrO4

2-(aq) + Ba2+(aq) BaCrO4(s)

2. What are the Spectator Ions?K+ & NO3

-

WRITING EQUATIONS FOR REACTIONS

Write the molecular, complete ionic, and net ionic chemical equations for the reaction between aqueous silver nitrate and aqueous sodium chloride.

Molecular Equation

NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq)

Complete Ionic Equation

Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq) AgCl(s) + Na+(aq) + NO3

-(aq)

Net Ionic Equation

Cl-(aq) + Ag+(aq) AgCl (s)

WRITING EQUATIONS FOR REACTIONS

Write the molecular, complete ionic, and net ionic chemical equations for the reaction between aqueous potassium hydroxide and aqueous iron(III) nitrate.

Molecular Equation

3KOH(aq) + Fe(NO3) 3(aq) Fe(OH) 3 (s) + 3KNO3(aq)

Complete Ionic Equation

3K+(aq) + 3OH-(aq) + Fe3+(aq) + 3NO3-(aq) Fe(OH) 3 (s) + 3K+(aq) + 3NO3

-(aq)

Net Ionic Equation

Fe3+(aq) + 3OH-(aq) Fe(OH) 3 (s)

WRITING EQUATIONS FOR REACTIONS

Write the molecular, complete ionic, and net ionic chemical equations for the reaction between aqueous sodium sulfide and aqueous copper(II) nitrate.

Molecular Equation

Na2S(aq) + Cu(NO3)2(aq) CuS(s) + 2NaNO3(aq)

Complete Ionic Equation

2Na+(aq) + S2-(aq) + Cu2+(aq) + 2NO3-(aq) CuS(s) + 2Na+(aq) +2 NO3

-(aq)

Net Ionic Equation

Cu2+(aq) + S2-(aq) CuS(s)

WRITING EQUATIONS FOR REACTIONS

Write the molecular, complete ionic, and net ionic chemical equations for the reaction between aqueous ammonium chloride and aqueous lead(II) nitrate.

Molecular Equation

2NH4Cl(aq) + Pb(NO3)2(aq) PbCl2 (s) + 2NH4NO3(aq)

Complete Ionic Equation

2NH4+(aq) + 2Cl-(aq) + Pb2+(aq) + 2NO3

-(aq) PbCl2 (s) +2 NH4+(aq) + 2NO3

-(aq)

Net Ionic Equation

Pb2+(aq) + 2Cl-(aq) PbCl2 (s)

DRIVING FORCES:#2 PROTON TRANSFER

Acid-Base Reactions

WHAT ARE ACIDS AND BASES?

Arrhenius (Late 1800s) Brǿnsted & Lowry (Early 1900s)

Acid – A substance that produces H+ ions when it dissolves in water. HA H+ + A-

Base – A substance that produces OH- ions when it dissolves in water. BOH B+ + OH-

Acid – A proton donor. HA H+ + A-

Base – A proton acceptor. B- + H+ BH

This new definition was necessary to accommodate substances that behave chemically as bases, but do not contain OH- ions.

STRONG ACIDS Strong Acids are Strong Electrolytes

They dissociate completely in water. 100% of the HA molecules break up into H+ and A- ions.

There are six common strong acids chemists MUST know:1. HCl - Hydrochloric Acid2. HI - Hydroiodic Acid3. HBr - Hydrobromic Acid4. HNO3 - Nitric Acid

5. HClO3 - Chloric Acid

6. H2SO4 - Sulfuric Acid *1st dissociation only

WEAK ACIDS Weak Acids are Weak Electrolytes

They DO NOT dissociate completely in water. 1% of the HA molecules break up into H+ and A- ions.

There are five common weak acids chemists MUST know:1. HF - Hydrofluoric Acid2. H3PO4 - Phosphoric Acid *1st dissociation only

3. HC2H3O2 - Acetic Acid

4. H3BO3 - Boric Acid *1st dissociation only

5. C6H8O7 - Citric Acid

STRONG BASES Strong Bases are Strong Electrolytes

They dissociate completely in water. 100% of the BOH molecules break up into B+ and OH- ions.

There are eight common strong bases chemists MUST know:1. KOH - Potassium Hydroxide2. NaOH - Sodium Hydroxide3. LiOH - Lithium Hydroxide4. CsOH - Cesium Hydroxide5. RbOH - Rubidium Hydroxide6. Ba(OH)2 - Barium Hydroxide

7. Ca(OH)2 - Calcium Hydroxide

8. Sr(OH)2 - Strontium Hydroxide

WEAK BASES Weak Bases are Weak Electrolytes

They DO NOT dissociate completely in water. 1% of the BOH molecules break up into B+ and OH- ions.

There are six common weak bases chemists MUST know:1. NH3 - Ammonia

2. C5H5N - Pyridine

3. CH3NH2 - Methylamine

4. C3H5O2NH2 - Alanine

5. Be(OH)2 - Beryllium Hydroxide

6. Mg(OH)2 - Magnesium Hydroxide

ACID-BASE REACTIONS Proton Transfer

This occurs in every acid-base reaction. When acids and bases react, a proton is

transferred from the acid to the base. This is the driving force of acid-base reactions. Once the proton is transferred, it is in a position

of less chemical potential. Formation of water

If the acid contains H+ and the base contains OH- the acid base reaction will form water.

The remaining cation and anion form a salt. If both the acid and base are weak, no

reaction occurs because no ions are present to initiate the proton transfer.

NEUTRALIZATION Acid-Base Reactions are also called

Neutralization Reactions Acids and Bases are extremely corrosive

materials. This means that they destroy things they come

into contact with, including skin, plant tissue, and even metals when the acid or base is strong enough.

For this reason, chemists should be VERY careful when handling them.

When acid-base reactions occur, corrosive materials become harmless materials (water and salt). Another way to say this is to say that they are

neutralized.

THE EQUATIONS OF A STRONG ACID-STRONG BASE REACTION

Consider the reaction between nitric acid and potassium hydroxide.

Molecular EquationHNO3(aq) + KOH(aq) KNO3(aq) + H2O(l)

Complete Ionic Equation – both acid & base dissociate.

H+(aq) + NO3-(aq) + K+(aq) + OH-(aq) K+(aq) + NO3

-(aq) + H2O(l)

Net Ionic Equation – always shows formation of water.

H+(aq) + OH-(aq) H2O(l)

THE EQUATIONS OF A WEAK ACID-STRONG BASE REACTION

Consider the reaction between acetic acid and sodium hydroxide.

Molecular EquationHC2H3O2(aq) + NaOH(aq) NaC2H3O2 (aq) + H2O(l)

Complete Ionic Equation – only the base dissociates.HC2H3O2(aq) + Na+(aq) + OH-(aq) Na+(aq) + C2H3O2

-(aq) + H2O(l)

Net Ionic EquationHC2H3O2(aq) + OH-(aq) H2O(l) + C2H3O2

-(aq)

THE EQUATIONS OF A STRONG ACID-WEAK BASE REACTION

Consider the reaction between hydrochloric acid and ammonia.

Molecular EquationHCl(aq) + NH3 (aq) NH4

+(aq) + Cl-(aq)

Complete Ionic Equation – only the acid dissociates.

H+(aq) + Cl-(aq) + NH3(aq) NH4+(aq) + Cl-(aq)

Net Ionic EquationH+(aq) + NH3(aq) NH4

+(aq)

DRIVING FORCES:#3 TRANSFER OF ELECTRONS

Oxidation – Reduction Reactions

WHAT ARE REDOX REACTIONS? Oxidation-Reduction Reactions are also called Redox

Reactions. Redox Reactions are reactions in which one or more electrons

is transferred. Electron transfers involve energy transfers. Electrons move from positions of higher to lower energy releasing

energy in the process. Most reactions used for energy production by living things

are Redox Reactions. Combustion Reactions are Redox Reactions. Reactions between metals and nonmetals are also Redox

Reactions.

Transfer of electrons can occur even when neither reactants nor products involve ionic species.

REDOX REACTION BETWEEN A

METAL & NONMETAL

2Al(s) + 3Cu(OH)2(aq) 2Al(OH)3(aq) + 3Cu(s)

Electrons are lost by Aluminum Al solid has a charge of zero Al in Al(OH)3 has a charge of 3+

Electrons are gained by Copper Cu in Cu(OH)2 has a charge of 2+ Cu solid has a charge of zero

In this RedOx Reaction, 6 Electrons are transferred from 2 Aluminum atoms to 3 Copper ions.

WHY DO ELECTRON TRANSFERS OCCUR IN NONIONIC COMPOUNDS?

Electronegativity can be thought of as the ability of an atom to draw electrons towards itself in a covalent bond. This property increases across the periodic table from left to right,

and going up. (F is the most electronegative element) The electron effectively belongs to the more electronegative

element in the bond, but there is no charge because the bond is covalent. The most electronegative atom takes one or more electrons from the

least electronegative atom.

OXIDATION STATES Oxidation States provide a way to keep track

of electrons in oxidation-reduction reactions that do not contain ions.

The oxidation states of the atoms in a covalent compound is the imaginary charges the atoms would have if the shared electrons are divided based on the atoms’ relative abilities to attract electrons.

Its an imaginary charge and is represented as a +/-# superscript. Ionic Charges are represented as #+/-

RULES FOR ASSIGNING OXIDATION STATES

Fluorine’s oxidation state is -1 in all its compounds because it is the most electronegative element.

Oxygen has an oxidation state of -2 in almost all its compounds.

Hydrogen has a +1 oxidation state in its covalent compounds.

Elemental atoms have an oxidation state of zero. Elements that occur as diatomic molecules have an

oxidation state of zero in this state, including O2. The sum of the oxidation states must be zero for an

electrically neutral compound. Monoatomic ions have the same oxidation state as their

ionic charges in aqueous solution and in ionic compounds. The sum of the oxidation states of the atoms in polyatomic

ions is equal to the polyatiomic ion’s charge.

ASSIGNING OXIDATION STATES Assign oxidation states to all atoms in the

following:1. NaCl2. H2O

3. CO2

4. SF6

5. NO3-

1. +1 + -1 = 0

2. 2(+1) + -2 = 0

3. +4 + 2(-2) = 0

4. +6 + 6(-1) = 0

5. +5 + 3(-2) = -1

OXIDATION-REDUCTION REACTIONS

Review: Chemical Reactions that involve electron transfer are called oxidation-reduction reactions.

Oxidation-Reduction reactions actually include two half reactions, an oxidation and a reduction.

Oxidation – the half reaction in which one or more electrons is lost. The element with an increased oxidation number is

oxidized and the substance that contains it is called the reducing agent.

Reduction – the half reaction in which one or more electrons is gained. The element with decreased oxidation number is

reduced and the substance that contains it is called the oxidizing agent.

LEO GOES GEROxidation Reduction

M M+ + e-

2X- X2 + 2e-

Loses Electrons

Is or Is Contained in the Reducing Agent

Electron Donor

X + e- X-

M3+ +3e- M

Gains Electrons

Is or Is Contained in the Oxidizing Agent

Electron Acceptor

REDOX HALF REACTIONS In a reaction between sodium and chlorine, an

electron from sodium is transferred to chlorine. 2Na + Cl2 2NaCl

Think about the ions that Na and Cl form…Which one is going to take the electron? Chlorine will; it forms an anion.

The half reactions show which species loses and which gains the electron(s) being transferred. Remember one loses and become a cation (LEO), and

the other gains and becomes an anion (GER): Sodium is oxidized: Na Na+ + e-

Chlorine is reduced: Cl + e- Cl-

The electron is transferred from the metal (sodium) to the more electronegative nonmetal (chlorine).

MORE EXAMPLES OF REDOX RXNS

Magnesium reacts with Oxygen to form Magnesium Oxide: 2Mg(s) + O2(g) 2MgO(s) The Half Reactions:

Mg Mg2+ + 2e-

O + 2e- O2-

Magnesium is oxidized and is the reducing agent. Oxygen is reduced and is the oxidizing agent.

Aluminum reacts with Iron(III) Oxide to form Iron and Aluminum Oxide: Al(s) + Fe2O3(s) Al2O3(s) + Fe(s) The Half Reactions:

Al Al3+ + 3e-

Fe3+ + 3e- Fe Aluminum is oxidized and is the reducing agent. Iron is reduced and is the oxidizing agent.

BALANCING REDOX EQUATIONS

• Oxidation-Reduction Reactions that take place in aqueous solution can be complex and very difficult to balance by the inspection or the try & try again method.

• To balance Redox reactions it is useful to separate the reaction into two half reactions.

• The halves include a reaction that involves the substance being reduced, and a reaction that involves the substance being oxidized.

• The two halves are balanced individually, following a set of balancing rules the are unique to either acidic or basic aqueous solutions because different ions are available to react in these two types of solutions.

• Once balanced, the halves are put back together to form one balanced net ionic equation for the Redox Reaction in question.

OXIDATION-REDUCTION REACTIONS Identify the atoms that are oxidized and reduced

and specify the oxidizing and reducing agents in the following reaction:

2Al(s) + 3I2(s) 2AlI3(s)

Assign oxidation states

Aluminum is oxidized and is the reducing agent. Iodine is reduced and is the oxidizing agent.

OXIDATION-REDUCTION REACTIONS Identify the atoms that are oxidized and reduced, and specify the

oxidizing and reducing agents in each of the following equations:

2PbS(s) + 3O2(g) 2PbO(s) + 2SO2(g)

PbO(s) + CO(g) Pb(s) + CO2(g)

Determine the change in oxidation state for each atom.

1. Sulfur is oxidized & Oxygen is reduced. O2 is the oxidizing agent & PbS is the reducing agent. Lead’s oxidation state does not change.

2. Lead is reduced & Carbon is oxidized. PbO is the oxidizing agent & CO is the reducing agent. Oxygen’s oxidation state does not change.

IDENTIFYING ELECTRON TRANSFER IN REDOX REACTIONS Write the half reactions for the following RedOx

reactions. Identify which species is being oxidized and which is being reduced. Indicate which is the oxidizing agent and which is the reducing agent.

2Al(s) + 3I2(s) 2AlI3(s)

Half Reactions:Al Al3+ + 3e-

I + e- I-

Aluminum is oxidized and is the reducing agent. Iodine is reduced and is the oxidizing agent.

IDENTIFYING ELECTRON TRANSFER IN REDOX REACTIONS Write the half reactions for the following

RedOx reactions. Identify which species is being oxidized and which is being reduced. Indicate which is the oxidizing agent and which is the reducing agent.

2Cs(s) + F2(g) 2CsF(s)

Half Reactions:Cs Cs2+ + 2e-

F + e- F-

Cesium is oxidized and is the reducing agent. Fluorine is reduced and is the oxidizing agent.