Chapter 3 — Reactions

1 2

Important – Read Before Using Slides in Class Instructor: This PowerLecture contains two PowerPoint presentations: one with lecture content and art from the text, the other with videos and animations. For animations and videos to run properly, we recommend that you run this PowerPoint presentation from the PowerLecture disc inserted in your computer. Also, for the mathematical symbols to display properly, you must install the supplied font called “Symb_chm,” supplied as a cross-platform TrueType font in the “Font_for_Lectures” folder in the "Media" folder on this disc. For instructions on inserting video slides into the lecture presentation, please next slide. If you prefer to customize the presentation or run it without the PowerLecture disc inserted, the animations and videos will only run properly if you also copy the associated animation and video files for each chapter onto your computer. Follow these steps: 1. Go to the disc drive directory containing the PowerLecture disc, and then to the “Media” folder, and then to the “PowerPoint_Lectures” folder. 2. In the “PowerPoint_Lectures” folder, copy the entire chapter folder to your computer. Chapter folders are named “chapter1”, “chapter2”, etc. Each chapter folder contains the PowerPoint Lecture file as well as the animation and video files. For assistance with installing the fonts or copying the animations and video files, please visit our Technical Support at http://cengage.com/support or call (800) 423-0563. Thank you.

Chapter 3 Chemical Reactions

Jeffrey Mack California State University, Sacramento

3

Chemical Reactions

Instructions for Inserting Video Slides into Lectures Throughout the lecture presentation, references the video slides are given in the “Notes” section at the bottom of your screen. For Office 2003 or earlier users: 1) Copy the slide from Video Presentation 2) Paste to the Lecture Presentation For Office 2007 or later users: 1) Under the Home tab, select the New Slide drop down menu. 2) On that drop down menu, select Reuse Slides. 3) In the search panel that opens up on the right, select browse and find the location of this chapter’s video presentation. 4) From the same search panel, select the slide you wish to insert. PLEASE NOTE: Slide insertion using the video presentation only works if you follow the instructions on the previous for copying the folder over to your computer.

Chemical Reactions Evidence of a chemical reaction: • Gas Evolution • Temperature Change • Color Change • Precipitation (insoluble species forms) In general, a reaction involves a rearrangement or change in oxidation state of atoms from reactants to products.

Reactants: Zn + I2

Product: ZnI2

Chemical Equations Chemical Equations show: • the reactants and products in a reaction. • the relative amounts in a reaction. Example: 4 Al(s) + 3 O2(g)  2 Al2O3(s) • The numbers in the front are called stoichiometric coefficients • The letters (s), (g), (l) and (aq) are the physical states of compounds.

Chapter 3 — Reactions

2

Reaction of Phosphorus with Cl2

Notice the stoichiometric coefficients and the physical states of the reactants and products.

Chemical Equations 4 Al(s) + 3 O2(g)  2 Al2O3(s)

Notice the stoichiometric coefficients and the physical states of the reactants and products.

Chemical Equations Law of the Conservation of Matter • Because the same number of atoms are present in a reaction at the beginning and at the end, the amount of matter in a system does not change.

This equation states that: 4 Al atoms + 3 O2 molecules react to form 2 formula units of Al2O3 or... 4 moles of Al + 3 moles of O2 react to form 2 moles of Al2O3

Chemical Equations • Since matter is conserved in a chemical reaction, chemical equations must be balanced for mass! • In other words, there must be same number of atoms of the each kind on both sides of the equatoin.

Reaction of Iron with Cl2

Lavoisier, 1788

2HgO(s)  2 Hg(l) + O2(g)

Balancing Chemical Reactions Steps in balancing a chemical reaction using coefficients: 1. Write the equation using the formulas of the reactants and products. Include the physical states (s, l, g, aq etc…) 2. Balance the compound with the most elements in the formula first using integers as coefficients. 3. Balance elements on their own last. 4. Check to see that the sum of each individual elements are equal on each side of the equation. 5. If the coefficients can be simplified by dividing though with a whole number, do so.

Chapter 3 — Reactions

3

Balancing Chemical Equations: Example

Balancing Chemical Equations: Example

balance last

C2H6 + 2 C’s & 6 H’s

balance last



O2

CO2

2 O’s

+

H2O

1 C & 2 O’s

C2H6 +

2 H’s & 1 O



O2

2 C’s & 6 H’s

CO2

2 O’s

+

H2O

1 C & 2 O’s

2 H’s & 1 O

balance H first

___C2H6

+



O2

CO2

This side will always have an even # of O-atoms

Balancing Chemical Equations: Example

2 C’s & 6 H’s

balance last



O2

CO2

2 O’s

+

H2O

1 C & 2 O’s

C2H6 +

2 H’s & 1 O

balance H first

2 2H6 ___C

This side has an odd # of O-atoms

Balancing Chemical Equations: Example

balance last

C2H6 +

3 H2O ___

+



O2

2 C’s & 6 H’s

CO2

2 O’s

+

H2O

1 C & 2 O’s

2 H’s & 1 O

balance H first

+



O2

CO2

3 H2O ___

+

2 2H6 ___C

+



O2

CO2

3 H2O ___

+

balance C next

2C2H6 +

Balancing Chemical Equations: Example

2 C’s & 6 H’s

CO2

2 O’s

+

H2O

1 C & 2 O’s

C2H6 +

2 H’s & 1 O

balance H first

2 2H6 ___C



O2

CO2

3 H2O ___

+

CO2

2 O’s

+

H2O

1 C & 2 O’s

2 H’s & 1 O

2 2H6 ___C

+



O2

CO2

3 H2O ___

+

balance C next

O2



4 CO2 ___

+

6H2O

balance O

2C2H6 +



O2

2 C’s & 6 H’s balance H first

+

balance C next

2C2H6 +

6H2O

balance last



O2

+

Balancing Chemical Equations: Example

balance last

C2H6 +

4 CO2 ___



O2

2C2H6 +

O2



4 CO2 ___

+

6H2O

balance O

7 O2 ____



4CO2 +

6H2O

2C2H6 +

7 O2 ____

4 C’s 12 H’s 14 O’s



4CO2 + 4 C’s 12 H’s 14 O’s

6H2O

Chapter 3 — Reactions

Balancing Equations

4 Balancing Equations: Practice

___ Al(s) + ___ Br2(l)  ___ Al2Br6(s) ___C3H8(g) + ___ O2(g)



___ CO2(g) + _____ H2O(g) ___B4H10(g) + ___ O2(g)  ___ B2O3(g)

+ ___ H2O(g)

Balancing Equations: Practice

Balancing Equations: Practice

• Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. • Write the balanced chemical equation for this reaction.

• Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. • Write the balanced chemical equation for this reaction. _ Mg(OH)2(s) + _ HCl(aq) 

Balancing Equations: Practice

Balancing Equations: Practice

• Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. • Write the balanced chemical equation for this reaction.

• Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. • Write the balanced chemical equation for this reaction.

_ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l)

_ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l) Balance with a coefficient of “2” in front of both HCl and water.

Chapter 3 — Reactions

Balancing Equations: Practice • Solid magnesium hydroxide reacts with hydrochloric acid to form aqueous magnesium chloride and water. • Write the balanced chemical equation for this reaction. _ Mg(OH)2(s) + _ HCl(aq)  _ MgCl2(aq) + _ H2O(l) Balance with a coefficient of “2” in front of both HCl and water. Mg, Cl, O and H are now balanced.

Chemical Equilibrium When writing chemical reactions one starts with: Reactants products N2(g) + 3H2(g)

2NH3(g)

5 Chemical Equations: Review • What Scientific Principles are used in the process of balancing chemical equations? • What symbols are used in chemical equations: gasses: _____ liquids: _____ solids: _____ aqueous species in solution: _____ • What is the difference between P4 and 4P in an eq.? • In balancing a chemical equation, why are the reactant and product subscripts not changed?

Chemical Equilibrium

Once equilibrium is achieved, reaction continues, but there is no net change in amounts of products or reactants.

Some reactions can also run in reverse: 2NH3(g)

N2(g) + 3H2(g)

Under these conditions, the reaction can be written:

3H2 (g)  N2 (g)

2NH3 (g)

Double arrows indicate “Equilibrium”.

Classifying Compounds • Salts (ionic compounds): Composed of a metal and non metal element(s). • Acids: Arrhenius definition Produce H+(aq) in water Examples: HCl, HNO3, HC2H3O2 • Bases: Arrhenius definition Produce OH(aq) in water Examples: NaOH, Ba(OH)2, NH3

Classifying Compounds • Molecular Compounds: • Covalently bonded atoms, not acids, bases or salts. • Compounds like alcohols (C2H5OH) or table sugar (C6H12O6) • These never break up into ions.

Chapter 3 — Reactions

6

Classifying Compounds

Classifying Compounds

• Classify the following as ionic, molecular, acid or base.

• Classify the following as ionic, molecular, acid or base.

Compound Na2SO4 Ba(OH)2 H3PO4 CH4 P2O5 NH3 HCN

Type

Reactions in Aqueous Solutions Aqueous Solutions: Water as the solvent Solution =

solute

That which is dissolved (lesser amount)

+

solvent That which is dissolves (greater amount)

Compound Na2SO4 Ba(OH)2 H3PO4 CH4 P2O5 NH3 HCN

Type ionic base acid molecular molecular base acid

Reactions in Aqueous Solutions Many reactions involve ionic compounds, especially reactions in water — aqueous solutions. KMnO4 in water

K+(aq) + MnO4-(aq)

There are three types of aqueous solutions: Those with Strong Electrolytes Those with Weak Electrolytes & those with non-Electrolytes

Ionic Compounds (CuCl2) in Water

Strong Electrolyte When ions are present in water, the solutions conduct electricity! Ions in solution are called ELECTROLYTES Examples of Strong Electrolytes: HCl (aq), CuCl2(aq) and NaCl (aq) are strong electrolytes. These dissociate completely (or nearly so) into ions. Strong Electrolytes conduct electricity well.

Chapter 3 — Reactions

Strong Electrolytes HCl(aq), CuCl2(aq) and NaCl(aq) are strong electrolytes. These dissociate completely (or nearly so) into ions.

7 Weak Electrolytes Acetic acid ionizes only to a small extent, it is a weak electrolyte. CH3CO2H(aq)

CH3CO2 (aq)  H (aq)

Weak electrolytes exist in solution under equilibrium conditions. The small concentration of ions conducts electricity poorly. Weak electrolytes exit primarily in their molecular form in water.

Weak Electrolytes Weak electrolytic solutions are characterized by equilibrium conditions in solution: When acetic acid dissociates, it only partially ionizes.

HC2H3O2 (aq)

Weak Electrolytes Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq)  CH3CO2-(aq) + H+(aq)

H+ (aq) + C2H3O2 (aq)

95%

5%

The majority species in solution is acetic acid in its molecular form. When writing a weak electrolyte in solution, one NEVER breaks it up into the corresponding ions!

HC2H3O2 (aq)

H ×+ (aq) + C2H3O2 (aq)

Non-Electrolytes Some compounds dissolve in water but do not conduct electricity. They are non-electrolytes. Examples include: • sugar • ethanol • ethylene glycol Non-electrolytes do not dissociate into ions!

Species in Solution: Electrolytes Strong electrolytes: Characterized by ions only (cations & anions) in solution (water).

Conduct electricity well Weak electrolytes:

Characterized by ions (cations & anions) & molecules in solution.

Conduct electricity poorly Non-electrolytes:

Characterized by molecules in solution.

Do not conduct electricity

Chapter 3 — Reactions

8

Solutes in Aqueous Solutions

Solubility Rules How do we know if a compound will be soluble in water? For molecular compounds, the molecule must be polar. We will discuss polarity later, for now I will tell you whether or not a molecular compound is polar… For ionic compounds, the compound solubility is governed by a set of SOLUBILITY RULES! You must learn the basic rules on your own!!!

Water Solubility of Ionic Compounds

Types of Reactions in a Solution

If one ion from the “Soluble Compound” list is present in a compound, then the compound is water soluble.

Precipitation Reactions: A reaction where an insoluble solid (precipitate) forms and drops out of the solution. Acid–base Neutralization: A reaction in which an acid reacts with a base to yield water plus a salt. Gas forming Reactions: A reaction where an insoluble gas is formed. Reduction and Oxidation Reactions (RedOx): A reaction where electrons are transferred from one reactant to another.

Chemical Reactions in Water EXCHANGE: Precipitation Reactions

EXCHANGE REACTIONS Pb(NO3) 2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3 (aq)

EXCHANGE Gas-Forming Reactions

REACTIONS

REDOX REACTIONS

EXCHANGE Acid-Base Reactions

The anions exchange places between cations. A precipitate forms if one of the products in insoluble.

Chapter 3 — Reactions

9

Precipitation Reactions The “driving force” is the formation of an insoluble solid called a precipitate.

Precipitation Reactions Which species is the precipitate? Pb(NO3)2(?) + 2KI(?)  2KNO3(?) + PbI2(?)

Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s) BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2 NaCl(aq) Precipitates are determined from the solubility rules.

Precipitation reactions Which species is the precipitate?

Precipitation Reactions Which species is the precipitate?

Pb(NO3)2(?) + 2KI(?)  2KNO3(?) + PbI2(?) From the solubility rules:

Pb(NO3)2(?) + 2KI(?)  2KNO3(?) + PbI2(?) From the solubility rules:

All nitrate salts are soluble, therefore: Pb(NO3)2(aq) + 2KI(?)  2KNO3(aq) + PbI2(?)

All nitrate salts are soluble, therefore: Pb(NO3)2(aq) + 2KI(?)  2KNO3(aq) + PbI2(?) All potassium salts are soluble, therefore: Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(?)

Precipitation Reactions Which species is the precipitate? Pb(NO3)2(?) + 2KI(?)  2KNO3(?) + PbI2(?) From the solubility rules: All nitrate salts are soluble, therefore: Pb(NO3)2(aq) + 2KI(?)  2KNO3(aq) + PbI2(?) All potassium salts are soluble, therefore: Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(?) By the solubility rules:

PbI2 is the ppt.

Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s)

Net Ionic Equations Molecular Equation: all species listed as formula units or in molecular form. reactants  products • Note all states of each reactant or product by: (s), (l), (g) or (aq) Ionic Equation: All soluble (aq) species present are listed as ions. • Leave all (s), (l) or (g) species as is. They do not dissociate into ions Net Ionic Equation: • From the ionic equation, cancel out any species that appear on either side of the equation. • These are known as the “spectator ions” and they are never part of a net ionic equation!

Chapter 3 — Reactions

10

Writing Net Ionic Equations

Writing Net Ionic Equations

Molecular Equation:

Molecular Equation:

Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s)

Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s) Total Ionic Equation: Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)  2K+(aq) + 2NO3– (aq) +

Writing Net Ionic Equations

PbI2(s)

Writing Net Ionic Equations

Molecular Equation:

Molecular Equation:

Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s) Total Ionic Equation: Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq) 

Never break up any (s), (l) or (g) or molecular (aq) species!

2K+(aq) + 2NO3– (aq) +

PbI2(s)

Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s) Total Ionic Equation: Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq) 

Never break up any (s), (l) or (g) or molecular (aq) species!

2K+(aq) + 2NO3– (aq) +

PbI2(s)

Cancel out the spectator ions to yield the net ionic equation: Pb2+ (aq) + 2I–(aq)  PbI2(s)

Acids & Bases Arrhenius Definition: • An acid is any substance that increases the H+(aq) concentration in an aqueous solution. HX(aq)  H+(aq) + X–(aq) • A base is any substance that increases the OH–(aq) concentration in an aqueous solution. MOH(aq)  M+(aq) + OH–(aq)

Acids and Bases Brönsted-Lowry: • An acid is any substance that donates H+(aq) to another species in an aqueous solution. HX(aq) + H2O(l) H3O+(aq) + X–(aq) H3O+(aq) = H+(aq)

• A base is any substance that accepts an H+(aq) in an aqueous solution. H+(aq) + NH3(aq) NH4+(aq)

Chapter 3 — Reactions

11

Acids

Strong Acids Examples: Strong acids are almost completely ionized in water. (strong electrolytes) HX (aq) (X = Cl, Br & I)

hydro ___ ic acid

HNO3 (aq)

nitric acid

HClO4 (aq)

perchloric acid

H2SO4 (aq)*

sulfuric acid

* Only the

1st

H is strong, sulfuric acid dissociates via: H2SO4 (aq)  H+ (aq) + HSO4– (aq)

Acids An acid: H3O+ in water

Weak Acids Examples: Weak Acids are incompletely ionized in water. (weak electrolytes) Weak acids are governed by dynamic equilibrium. HC2H3O2 (aq)

acetic acid (vinegar)

nitrous acid

HNO2 (aq)

hydrosulfuric acid

H2S (aq)

hydrogen sulfate ion

HSO4–(aq)

Weak acids are always written in their molecular form. See you text and home work for more examples.

Strong Bases Bases: A base is a substance that produces OH– (aq) ions in water by dissociation in water: H O( ) NaOH(s)  Na (aq)  ΟΗ aq 2

Bases Base: OH- in water NaOH(aq)  Na+(aq) + OH-(aq)

Strong bases are almost completely ionized in aqueous solution. (Strong electrolytes) Examples: Hydroxides of Group 1 (MOH(aq) where M = Li, Na, K ect…) and Ca, Sr, Ba.* *Ca(OH)2, Sr(OH)2 & Ba(OH)2 are slightly soluble, but that

which dissolves is present as ions only.

NaOH is a strong base

Chapter 3 — Reactions

12

Weak Bases

Ammonia, NH3

Weak Bases: NH3 acts as a base by reacting with water: 



NH3(aq) + H2O(l)

NH4+(aq) + OH –(aq)

Ammonia can also accept H+ from an acid: 



NH3(aq) + H+(aq)

NH4+(aq)

Reactions of Acids & Bases: Acid-Base Neutralization Acid + Base  Salt + Water (usually) HA (aq) + MOH(aq)  MA(aq) + HOH(l) Strong acid - Strong base neutralization: HBr(aq)/KOH(aq) Molecular Equation: HBr(aq) + KOH(aq)  KBr (aq) + H2O(l) Total Ionic Equation:

/

H+ (aq) + Br– (aq)+ K+(aq) + OH– (aq) 

/

Acid-Base Reactions • The “driving force” is the formation of water. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(liq) • Net ionic equation OH-(aq) + H3O+(aq)  2 H2O(l) • This applies to ALL reactions of STRONG acids and bases.

K+(aq) + Br– (aq) + H2O(l)

/

/

Net Ionic equation: H+ (aq) + OH– (aq)  H2O (l)

Reactions of Acids & Bases: Acid-Base Neutralization Reactions of weak acids and strong bases: Molecular Equation: HC2H3O2(aq) + NaOH(aq)  NaC2H3O2(aq) + H2O(l) Total Ionic Equation:

/

/

HC2H3O2(aq) + Na+(aq) + OH–(aq)  Na+(aq) + C2H3O2–(aq) + H2O(l) Leave in molecular form

Net Ionic: HC2H3O2(aq) + OH–(aq)  C2H3O2–(aq) + H2O(l)

Chapter 3 — Reactions

13

Non-Metal Acids

Bases Metal oxides form bases in aqueous solution

Nonmetal oxides can form acids in aqueous solutions: Examples: CO2(aq) + H2O(s)  H2CO3(aq)

CaO(s) + H2O(l)  Ca(OH)2(aq)

SO3(aq) + H2O(s)  H2SO4(aq) Both gases come from the burning of fossil fuels. CaO in water. Indicator shows solution is basic.

Gas-Forming Reactions

Gas-Forming Reactions Metal carbonate salts react with acids to the corresponding metal salt, water and carbon dioxide gas. 2HCl(aq) + CaCO3(s)  CaCl2(aq) + H2CO3(aq) decomposes Similarly:

H2O(l) + CO2(g)

HCl(aq) + NaHCO3(s)  NaCl(aq) + H2O(l) + CO2(g) acid base salt water Neutralization!!!

Gas-Forming Reactions

Gas-Forming Reactions

Group I metals: Na, K, Cs etc.. react vigorously with water 2K(s) + 2H2O(l)  2KOH(aq)+ H2(g)

Metals & acid:

Some metals react vigorously with acid solutions: Zn(s) +

2H+(aq)



Zn2+(aq)

+ H2(g)

CaCO3(s) + H2SO4(aq)  2 CaSO4(s) + H2CO3(aq)

Carbonic acid is unstable and forms CO 2 & H2O H2CO3(aq)  CO2 + water (The antacid tablet contains citric acid + NaHCO3)

Chapter 3 — Reactions

Oxidation-Reduction Reactions Thermite reaction:

14 Oxidation-Reduction Reactions REDOX = reduction & oxidation

Fe2O3(s) + 2Al(s) 

O2(g) + 2 H2(g)  2 H2O(l)

2Fe(s) + Al2O3(s)

Oxidation-Reduction Reactions • Oxidation involves a reactant atom or compound losing electrons. • Reduction involves a reactant atom or substance gaining electrons. • Neither process can occur alone… that is, there must be an exchange of electrons in the process. • The substance that is oxidized is the reducing agent • The substance that is reduced is the oxidizing agent oxidized

reduced

Mg(s) + 2H+(aq)  reducing agent

Mg2+(aq) + H2(g)

oxidizing agent

2. If an atom is charged, then the charge is the oxidation numbers . Ion

Oxidation Number

Mg2+(aq)

+2

Cl(aq)

1

Sn4+(s)

+4

2 2

Hg (aq)

• Chemists use oxidation numbers to account for the transfer of electrons in a RedOx reaction. • Oxidation numbers are the actual or apparent charge on atom when alone or combined in a compound. 1. The atoms of pure elements always have an oxidation number of zero. Examples: Mg(s) Hg(l) I2(s)

All have an oxidation number of zero (0)

O2(g)

Oxidation Numbers

Examples:

Oxidation Numbers

+2/2 = +1 for each Hg atom

Oxidation Numbers 3. In a compound, fluorine always has an oxidation numbers of 1. 4. Oxygen most often has an oxidation number of 2. » *When combined with fluorine, oxygen has a positive O.N. » *In peroxide, the O.N. is 1.

5. In compounds, Cl, Br & I are 1 (Except with F and O present) 6. In compounds, H is +1, except as a hydride (H: 1)

Chapter 3 — Reactions

15

Oxidation Numbers Examples:

Oxidation Numbers Most common oxidation numbers:

compound

Oxidation Numbers

HF(g)

H = +1

F = 1

H2O(l) OF2(g)

H = +1 O = +2

O = 2 F = 1

Na2O2(s)

Na = +1

O = 1

HCl(g)

H = +1

Cl = 1

NaH(l)

Na = +1

H = 1

Oxidation Numbers 7. For neutral compounds, the sum of the oxidation numbers equals zero. For a poly atomic ion, the sum equals the charge. Examples: +2 + 2 × (−1) =0

Oxidation Numbers Determine the oxidation number of iron in the following compound: ? + 3(1) = 0

MgCl2

Fe(OH)3

3 + 4 × (+1) = +1

 4

Iron must have an oxidation number of +3!

NH

Recognizing a Redox Reaction In a RedOx reaction, the species oxidized and the species reduced are identified by the changes in oxidation numbers :

3CH4 (g)  Cr2O72 (aq)  8H (aq) 

Oxidation numbers:

+1

Practice: Identify the species that is Oxidized and Reduced by assigning oxidation numbers in the following reaction.

0

3CH3OH(l)  2Cr 3 (aq)  4H2O(l)

2Ag+ (aq) + Cu(s) ® 2Ag(s) + Cu2+ (aq) 0

+2

Oxidation numbers:

Since silver goes from +1 to zero, it is reduced. Since copper goes from zero to +2, it is oxidized. The reaction is balanced for both mass and charge.

Answer:

Chapter 3 — Reactions

Practice: Identify the species that is Oxidized and Reduced by assigning oxidation numbers in the following reaction.

3CH4 (g)  Cr2O72 (aq)  8H (aq) 

16 Practice: Identify the species that is Oxidized and Reduced by assigning oxidation numbers in the following reaction.

3CH4 (g)  Cr2O72 (aq)  8H (aq) 

3CH3OH(l)  2Cr 3 (aq)  4H2O(l)

3CH3OH(l)  2Cr 3 (aq)  4H2O(l)

Answer: • The carbon in methane (CH 4) is oxidized (4 to 2)

Answer: • The carbon in methane (CH 4) is oxidized (4 to 2) • Chromium in dichromate is reduced (+6 to +3)

Redox Reactions

Oxidation-Reduction Reactions • Iron gains 3 electrons (+3 to 0) oxidation number change. It is Reduced. • Carbon loses 2 electrons (+2 to +4) it is Oxidized.

Redox Reactions

Redox Reactions

REDOX = reduction & oxidation Corrosion of aluminum

Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) In all reactions if a species is oxidized then another species must also been reduced

2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)

Chapter 3 — Reactions

17 Electron Transfer in a Redox Reaction

Redox Reactions Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

e

e

2Ag+(aq) + Cu(s)  Cu2+(aq) + 2Ag(s)

• • • •

Redox Reactions in Our World

Two electrons leave copper. The silver ions accept them. The copper metal is oxidized to copper (II) ion. The silver ion is reduced to solid silver metal.

Examples of Redox Reactions

Batteries Corrosion

Fuels

Metal + halogen 2 Al + 3 Br2  Al2Br6

Manufacturing metals

Examples of Redox Reactions

Examples of Redox Reactions Metal + acid Mg + HCl Mg = reducing agent H+ = oxidizing agent

Nonmetal (P) + Oxygen  P4O10

Metal (Mg) + Oxygen  MgO

Metal + acid Cu + HNO3 Cu = reducing agent HNO3 = oxidizing agent

Chapter 3 — Reactions Reviewing What You’ve Learned • You have the following items available to you: Deionized water, pH paper, test tubes various metal nitrate salts, common acid and base solutions. • Suggest a simple test or set of tests for identifying the unknown substances. Use proper terminology and write balanced chemical equations where applicable. • Justify your answers thoroughly.

18 Reviewing What You’ve Learned • How would you determine whether or not a test tube containing a clear colorless solution is water or sulfuric acid? • Given a white powder that my be silver chloride or sodium chloride. • Whether a compound is silver nitrate or sodium nitrate.3