Chemical Reactions in Aqueous Solutions

Chemical Reactions in Aqueous Solutions
Chapter Four Chemical Reactions in Aqueous Solutions

Today… Turn in: Nothing Our Plan: Notes – Synthesis & Decomposition
Begin Worksheet #1 Homework (Write in Planner): WS#1 Due Wednesday

Unit 1 Test Results A 1 B 6 C D 3 F Average 78.47% High Score 96%

Quick Review How do we know a chemical reaction has occurred? What do we observe?

Evidence of a Chemical Rxn
Color Change Precipitate (solid) forms Gas Evolved Heat/Light Given Off Endothermic – heat absorbed (gets cold) Exothermic – heat released (gets hot) More about this in Unit 5 - Thermodynamics

Review of Basics When writing out a balanced equation, it is necessary to indicate what state the substance is in (s, l, g, or aq). For elements, you can look at the PT to determine their standard state. Examples: Mercury Fluorine Iron

There are 7 diatomic elements.
Review of Basics There are 7 diatomic elements. They are only diatomic when they are alone. They are all gases, except Br2, which is a liquid, as indicated on the PT. Remember the Super 7? H2, N2, O2, F2, Cl2, Br2, I2

Review of Basics Soluble means the substance dissolves in solution (water). Label soluble substances aqueous (aq) Insoluble means the substance does not dissolve in solution. Label insoluble substances solid (s) To determine if a substance is soluble or insoluble, use a solubility table.

Solubility Chart

Acids and bases are aqueous Water is almost always a liquid
Other Basic Tips Acids and bases are aqueous Water is almost always a liquid Write water as HOH when it is a reactant Some common gases that aren’t diatomic are CO2, CO, SO3, SO2, H2S, and NH3

Other Basic Tips If a reaction indicates that a catalyst was present, it is not part of the actual reaction. Instead, you write the catalyst above the arrow. Catalysts are not used up in the reaction. They are just there to speed it up. If a reaction is heated, you draw a triangle above the arrow. Heat is used as a catalyst.

Occur when two or more reactants combine to form a single product.
Synthesis Reactions Occur when two or more reactants combine to form a single product. In Chem 1 we called this type “Dating” There are several common types of synthesis reactions.

Synthesis Reactions Type 1: A metal combines with a nonmetal to form a binary salt. Example: A piece of lithium metal is dropped into a container of nitrogen gas. 6Li (s) + N2 (g) → 2Li3N (aq)

Type 2: Metallic oxides and water form bases (metallic hydroxides).
Synthesis Reactions Type 2: Metallic oxides and water form bases (metallic hydroxides). Example: Solid sodium oxide is added to water Na2O (s) + HOH (l) → 2NaOH (aq) Example: Solid magnesium oxide is added to water. MgO (s) + HOH (l) → Mg(OH)2 (aq)

Synthesis Reactions Type 3: Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number. Example: Carbon dioxide is bubbled in water. CO2 (g) + H2O (l) → H2CO3 (aq) Example: Dinitrogen pentoxide is bubbled in water. N2O5 (g) + H2O (l) → 2HNO3 (aq)

Type 4: Metallic oxides and nonmetallic oxides form salts.
Synthesis Reactions Type 4: Metallic oxides and nonmetallic oxides form salts. Example: Solid sodium oxide is added to carbon dioxide. Na2O (s) + CO2 (g) → Na2CO3 (aq) Example: Solid calcium oxide is added to sulfur trioxide. CaO (s) + SO3 (g) → CaSO4 (aq)

Try It Out 1 Solid barium oxide is added to distilled water (p. 142)

Try It Out 2 Sulfur trioxide gas is added to excess water (p. 146)

Decomposition Reactions
Occur when a single reactant is broken down into two or more products. In Chem 1 we called this “The Break Up” There are several common types of decomposition reactions.

Type 1: Metallic carbonates decompose into metallic oxides and carbon dioxide Example: A sample of magnesium carbonate is heated. MgCO3 (s) → MgO (s) + CO2 (g)

Type 2: Metallic chlorates decompose into metallic chlorides and oxygen. Example: A sample of magnesium chlorate is heated. Mg(ClO3)2 (aq) → MgCl2 (aq) + 3O2 (g)

Type 3: Ammonium carbonate decomposes into ammonia, water, and carbon dioxide. Example: A sample of ammonium carbonate is heated. (NH4)2CO3 (aq)→ 2NH3 (g) + H2O (l) + CO2 (g)

Type 4: Sulfurous acid decomposes into sulfur dioxide and water. Example: A sample of ammonium carbonate is heated. H2SO3 (aq) → H2O (l) + SO2 (g)

Type 5: Carbonic acid decomposes into carbon dioxide and water. Example: A sample of carbonic acid is heated. H2CO3 (aq) → H2O (l) + CO2 (g)

Type 6: A binary compound may break down to produce two elements. Example: Molten sodium chloride is electrolyzed. 2NaCl (l) → 2Na (s) + Cl2 (g)

Type 7: Hydrogen peroxide decomposes into water and oxygen. Example: 2H2O2 (aq) → 2H2O (l) + O2 (g)

Type 8: Ammonium hydroxide decomposes into ammonia and water. Example: NH4OH (aq) → NH3 (g) + HOH (l)

Try It Out 1 Solid calcium chlorate is heated in the presence of a magnesium dioxide catalyst

Try It Out 2 A sample of lithium carbonate is heated strongly

STOP Complete Worksheet #1 by next class.

Today… Turn in: Get Out WS#1 Our Plan: Questions on WS#1 Quick Review
Notes – Single & Double Replacement Begin WS#2 Homework (Write in Planner): WS#2 Due Friday

Quick Review Write formulas for the reactants and predicted products for the chemical reactions that follow: Solid calcium carbonate is strongly heated in a test tube. Solid lithium chlorate is heated in the presence of a manganese dioxide catalyst. Excess chlorine gas is passed over hot iron filings.

Reactions that involve an element replacing one part of a compound.
Single Replacement Reactions that involve an element replacing one part of a compound. The products include the displaced element and a new compound. In Chem 1 we called this “Cheating” An element can only replace another element that is less active than itself.

General Activity Series for Metals:
Single Replacement General Activity Series for Metals: Most Active Li Ca Na Mg Al Zn Fe Pb H2 Cu Ag Pt Au Least Active

Single Replacement General Activity Series for Nonmetals: Most Active F2 Cl2 Br2 I2 Least Active

Remember the Trend? Or you can just look at the Periodic Table
The most reactive metals are on the bottom left (Francium) The most reactive nonmetals are on the top right (Fluorine) Hydrogen is the tricky one. Use the pink sheet for electronegativity potentials otherwise.

For metals, the more negative the reduction potential, the more reactive. For nonmetals, the more positive the reduction potential, the more reactive.

3Mg (s) + 2FeCl3 (aq)→ 2Fe (s) + 3MgCl2 (aq)
Single Replacement Type 1: Active metals replace less active metals from their compounds in aqueous solution. Example: Magnesium turnings are added to a solution of iron (III) chloride. 3Mg (s) + 2FeCl3 (aq)→ 2Fe (s) + 3MgCl2 (aq)

2Na (s) + 2HOH (l)→ 2NaOH (aq) + H2 (g)
Single Replacement Type 2: Active metals replace hydrogen in water. Example: Sodium is added to water. 2Na (s) + 2HOH (l)→ 2NaOH (aq) + H2 (g)

2Li (s) + 2HCl (aq) → 2LiCl (aq) + H2 (g)
Single Replacement Type 3: Active metals replace hydrogen in acids. Example: Lithium is added to hydrochloric acid. 2Li (s) + 2HCl (aq) → 2LiCl (aq) + H2 (g)

Cl2 (g) + 2KI (aq) → I2 (g) + 2KCl (aq)
Single Replacement Type 4: Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + 2KI (aq) → I2 (g) + 2KCl (aq)

Single Replacement Type 5: If a less reactive element is combined with a more reactive element in compound form, their will be no resulting reaction. Example: Chlorine gas is bubbled into a solution of potassium iodide. Cl2 (g) + KF (aq) → No Reaction Example: Zinc is added to a solution of sodium chloride. Zn (s) + NaCl (aq) → No Reaction

Try It Out 1 A piece of aluminum metal is added to a solution of gold nitrate. (p. 146)

Try It Out 2 Liquid bromine is shaken with a potassium iodide solution. (p. 152)

Try It Out 3 Small chunks of potassium are added to water. (p. 164)

Try It Out 4 Small strips of platinum are placed in hydrochloric acid.

In Chem 1 we called this “Swapping”
Double Replacement Reactions between two compounds in aqueous solution where the cations and anions appear to “switch partners”. AX + BY → AY + BX In Chem 1 we called this “Swapping”

Double Replacement All double replacement reactions (aka metathesis) must have a “driving force” or reason why the reaction will occur or “go to completion”.

“Driving Force” for reactions: Formation of a precipitate
Double Replacement “Driving Force” for reactions: Formation of a precipitate Formation of a gas Formation of primarily molecular species (nonelectrolytes, water, weak acids)

Double Replacement If one of these “driving forces” is NOT present, then the reaction does not go to completion. This type of reaction is indicated by a double arrow

Reactions that Form Precipitates
There are limits to the amount of a solute that will dissolve in a given amount of water. If the maximum concentration of solute is less than about 0.01 M, we refer to the solute as insoluble in water. When a chemical reaction forms such a solute, the insoluble solute comes out of solution and is called a precipitate.

Silver Iodide Precipitation
A solution containing silver ions and nitrate ions, when added to … … a precipitate of silver iodide. … a solution containing potassium ions and iodide ions, forms … What is the net ionic equation for the reaction that has occurred here? (Hint: what species actually reacted?)

Double Replacement (Precipitate)
In order to predict double replacement reactions yielding precipitates, one must memorize the solubility rules. I will let you use a solubility chart on your test!

Solubility Rules

Example 1 Predict and balance the following double replacement reactions based on the solubility of the products. Use the abbreviations (aq) and (s) for the reactant and products. All reactants are aqueous. Solutions of manganese (II) sulfate and ammonium sulfide are mixed.

Example 2 Solutions of sodium iodide and lead (II) nitrate are combined.

Try it Out 1 Solutions of sodium carbonate and iron (III) nitrate combine.

Try it Out 2 Solutions of lead (II) acetate and calcium chloride are combined.

Double Replacement (Gas)
Common Gases formed: H2S, CO2, SO2, NH3 Reactions that produce three of the gases (CO2, SO2, and NH3) involve the initial formation of a substance that breaks down to give the gas and HOH.

Common Gases H2S Any sulfide (salt of S2-) plus any acid form H2S (g) and a salt. CO2 Any carbonate (salt of CO32-) plus any acid form CO2 (g), HOH, and a salt SO2 Any sulfite (salt of SO32-) plus any acid form SO2 (g), HOH, and a salt. NH3 Any ammonium salt (salt of NH4+) plus any soluble strong hydroxide react upon heating to form NH3 (g), HOH, and a salt.

Example 1: The reaction of Na2SO3 and HCl produces H2SO3: Na2SO3 (aq) + 2HCl (aq) → H2SO3 (aq) + 2NaCl (aq) Bubbling is observed in this reaction because the H2SO3 (sulfurous acid) is unstable and immediately decomposes to give HOH and SO2 gas: H2SO3 (aq) → HOH (l) + SO2 (g) The molecular equation for the overall or complete reaction, therefore, is: Na2SO3 (aq) + 2HCl (aq) → HOH (l) + SO2 (g) + 2NaCl (aq)

Example 2: A typical reaction of a carbonate and an acid is: K2CO3 (aq) + 2HNO3 (aq) → HOH (l) + CO2 (g) + 2KNO3 (aq) Bubbling is also observed in this reaction. Theoretically H2CO3, carbonic acid, is formed, but the acid is unstable and immediately decomposes to form carbon dioxide gas and water according to the following equation: H2CO3 (aq) → HOH (l) + CO2 (g)

Example 3: Ammonium salts and soluble bases react as follows (particularly when the solution is warmed): NH4Cl (aq) + NaOH (aq) → NH3 (g) + HOH (l) + NaCl (aq) The odor of ammonia gas is noted and moist blue litmus paper held near the mouth of the container will turn blue. Theoretically NH4OH, ammonium hydroxide, is produced (also known as ammonia water). The compound is unstable and decomposes into ammonia gas and water: NH4OH (aq) → NH3 (g) + HOH (l)

Example 4: The odor of rotten eggs and bubbling are noted when an acid is added to a sulfide. A typical reaction producing hydrogen sulfide gas is: FeS (s) + 2HCl (aq) → FeCl2 (aq) + H2S (g)

Try It Out 1 Dilute hydrochloric acid is added to a solution of potassium sulfite. (p. 152)

Try It Out 2 Concentrated hydrochloric acid is added to solid manganese (II) sulfide. (p. 150)

Today… Turn in: Get Out WS#2 Our Plan: Questions on WS#2
Quick Review/Video Notes – Aqueous Reactions, Net Ionic Equations, & Math with Ion Concentration Begin WS#3 Homework (Write in Planner): WS#3 Due Monday

A Note about phases A product can only be aqueous if there is an aqueous or liquid reactant. That’s why some of the ones on the first WS that you labeled aqueous were solid. Examples:

Quick Discussion of Weak Electrolytes
One of the driving forces for a reaction is the production of weak electrolytes (water or a weak acid). For a list of strong acids see page 22 in booklet. Everything else is weak. The only thing to note is that if you form a weak acid, it is aqueous but you would have a forward arrow, not a double arrow. See examples on the board.

Quick Review Solutions of sodium fluoride and dilute hydrochloric acid are combined. Dilute acetic acid solution is added to solid magnesium carbonate. A saturated solution of calcium hydroxide was added to a solution of magnesium chloride.

Crash Course in Precipitation

Arrhenius’s Theory of Electrolytic Dissociation
Why do some solutions conduct electricity? An early hypothesis was that electricity produced ions in solution, and those ions allowed the electricity to flow. Arrhenius’s theory: Certain substances dissociate into cations and anions when dissolved in water. The ions already present in solution allow electricity to flow.

Electrolytic Properties of Aqueous Solutions
Electrolytes dissociate to produce ions. The more the electrolyte dissociates, the more ions it produces.

Types of Electrolytes A strong electrolyte dissociates completely.
A strong electrolyte is present in solution almost exclusively as ions. Strong electrolyte solutions are good conductors. A nonelectrolyte does not dissociate. A nonelectrolyte is present in solution almost exclusively as molecules. Nonelectrolyte solutions do not conduct electricity. A weak electrolyte dissociates partially. Weak electrolyte solutions are poor conductors. Different weak electrolytes dissociate to different extents.

Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte?
Strong electrolytes include: Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) Strong bases (IA and IIA hydroxides) Most water-soluble ionic compounds Weak electrolytes include: Weak acids and weak bases A few ionic compounds How do we tell whether an acid (or base) is weak? Nonelectrolytes include: Most molecular compounds Most organic compounds (most of them are molecular)

Aqueous Solutions & Ionic Equations
On the AP Exam, you do not have to write out complete molecular equations like we have been doing. Instead, you have to write ionic equations.

Overall (Total) Ionic Equation
Formulas of the reactants and products are written to show the predominant form of each substance as it exists in aqueous solution. Soluble salts, strong acids, and strong bases are written as separated ions. Everything else is written as a molecule. We will memorize strong acids and strong bases later in the year. For now, a list will be provided.

Memorize Later… Strong Acids Strong Bases HClO4 LiOH RbOH H2SO4 NaOH
CsOH HI KOH Mg(OH)2 HBr Ca(OH)2 HCl Sr(OH)2 HNO3 Ba(OH)2

Overall Ionic Equation Example 1
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)

Overall Ionic Equation Example 2
CaCO3 (aq) + 2HCl (aq) → CaCl2 (aq) + HOH (l) + CO2 (g)

Try it Out (NH4)2S (aq) + 2LiOH (aq) → Li2S (aq) + 2NH3 (g) + 2HOH (l)

Net Ionic Equation Net ionic equations are written to show only the species that react or undergo change in aqueous solutions. The net ionic equation is obtained by eliminating the spectator ions from the overall equation. All that is left are the ions that have changed chemically.

Net Ionic Equation Example 1
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)

Example 2 For each of these word equations, predict the product and write an overall and net ionic equation. Copper (I) nitrate is combined with silver chloride.

Example 3 Magnesium hydroxide is combined with lead (II) sulfate.

Try It Out Mercury (I) nitrate is combined with sodium sulfide.

Ion Concentrations in Solution
“Trick” question … What is the concentration of Na2SO4 in a solution prepared by diluting mol Na2SO4 to 1.00 L? The answer is: … zero … WHY?? And … how do we describe the concentration of this solution?

Calculating Ion Concentrations in Solution
In M Na2SO4: two moles of Na+ ions are formed for each mole of Na2SO4 in solution, so [Na+] = M. one mole of SO42– ion is formed for each mole of Na2SO4 in solution, so [SO42–] = M. An ion can have only one concentration in a solution, even if the ion has two or more sources.

Example 4.1 Calculate the molarity of each ion in an aqueous solution that is M Na2SO4 and M NaCl. In addition, calculate the total ion concentration of the solution.

Try It Out! Exercise 4.1A: Seawater is essentially M NaCl and M MgCl2, together with several other minor solutes. What are the molarities of Na+1, Mg+2, and Cl-1 in seawater?

Example 4.6 One cup (about 240 g) of a certain clear chicken broth yields g AgCl when excess AgNO3(aq) is added to it. Assuming that all the Cl– is derived from NaCl, what is the mass of NaCl in the sample of broth?

Complete Worksheet #4 by Monday.
Stop! Complete Worksheet #4 by Monday. Let’s look at it together.

Today… Turn in: Get Out WS#3 Our Plan: Questions on WS#3 Quick Review
Notes – Acid/Base Reactions & Titrations Begin WS#4 Homework (Write in Planner): WS#4 Due Wednesday

Write the molecular and net ionic equation for the following:
Quick Review Write the molecular and net ionic equation for the following: Chlorine gas is bubbled into a solution of potassium iodide. Aqueous solutions of potassium chromate and silver nitrate react.

Acid-Base Reactions We will cover acid/base reactions in great detail at the end of the school year. For now, we’re just going to cover very simple neutralization reactions. Neutralization is the transfer of PROTONS from an acid to a base. In neutralization, an acid and a base form a salt and water.

Acid-Base Reactions When writing net ionic equations, keep in mind which acids are strong (written in ionic form) and which are weak (written in molecular form). Check the solubility rules of the salt produced. If it is soluble, it is written in ionic form; if it is insoluble it is written in molecular form. The salt produced always consists of a cation from the base and an anion from the acid.

Remember, these are in your notes…
Strong Acids Strong Bases HClO4 LiOH RbOH H2SO4 NaOH CsOH HI KOH Mg(OH)2 HBr Ca(OH)2 HCl Sr(OH)2 HNO3 Ba(OH)2

H2S (g) + 2KOH (aq) → K2S (aq) + 2HOH (l)
Acid-Base Reactions Example 1: Hydrogen sulfide gas is bubbled through excess potassium hydroxide solution. H2S (g) + 2KOH (aq) → K2S (aq) + 2HOH (l)

Acid-Base Reactions Polyprotic acids (more than one hydrogen like H2SO4) can be tricky. If the base is in excess, all hydrogen ions will react with strong base to produce water. See Example 2. If however, this same reaction were described in terms of mixing equal numbers of moles, then the coefficients for both reactants would be one (the same number of H and OH must be given away). See Example 3.

H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)
Acid-Base Reactions Example 2: Dilute sulfuric acid is reacted with excess sodium hydroxide. H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)

H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)
Acid-Base Reactions Example 3: Equal number of moles of sulfuric acid and sodium hydroxide react. H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)

H3PO4 (aq) + 2NaOH (aq) → Na2HPO4 (aq) + 2HOH (l)
Acid-Base Reactions As the following example demonstrates, it is important to take into account the quantity (concentration and amount) of each reactant. Example 4: Equal volumes of 0.1 M phosphoric acid and 0.2 M sodium hydroxide are reacted together. H3PO4 (aq) + 2NaOH (aq) → Na2HPO4 (aq) + 2HOH (l)

Try It Out 1 Equal molar volumes of sodium hydroxide and hydrobromic acid are mixed together. Write the molecular equation and the net ionic equation.

Try It Out 2 Equal volumes of 0.1 M phosphoric acid and potassium hydroxide solutions are mixed.

Titrations A titration is a procedure where one reactant is carefully added to another reactant until the two have combined in their exact stoichiometric proportions. There are different types of titrations, but no matter the type, both reactants are fully consumed at the end of the titration.

Titrations The purpose of a titration is to find the number of moles, grams, concentration, or percentage of an unknown. The unknown is called the ANALYTE. This is done by measuring the precise volume and concentration of a known solution (TITRANT) needed to react completely with the analyte. Stoichiometry is then used to determine whatever quantity you’re looking for.

Acid-Base Titrations Acid-base titrations are the most common type
You conducted this type of titration in Chem 1 (remember the antacid lab?) There are 3 things that you need for a titration: A way to accurately measure the titrant (BURET) A way to know your reaction is complete (INDICATOR or you can look at a TITRATION CURVE) A solution whose concentration you know (STANDARD SOLUTION)

Acid-Base Titrations

Acid-Base Titrations Equivalence point – in an acid-base titration it is the point when the titrant completely neutralizes the analyte. If you’re using an indicator, you choose one that changes color close to the neutralization point. If you are looking at a titration curve, it is the straight line on a graph of volume versus pH.

Acid-Base Titrations When the indicator changes color in a titration, you have reached the endpoint. At the endpoint the titration is stopped and the volume of titrant is recorded. You want the equivalence point and endpoint to be the same!

Titration Curve

Calculation Example 1 Example 4.9: What volume (mL) of M NaOH is required to neutralize mL of M HCl in an acid-base titration?

Calculation Example 2 Example 4.10: A mL sample of an aqueous solution of calcium hydroxide is neutralized by mL of M HNO3 (aq). What is the molarity of the calcium hydroxide solution?

Stop Complete Worksheet #4 by next class and complete the pre-lab for Investigation 4.

Today… Turn in: Get Out WS#4 Our Plan: Questions on WS#4
Investigation 4 Homework (Write in Planner): Lab Report Due Monday Post results (Average concentration of H3PO4) to Edmodo by Friday.

Quick Review Equal volumes of 0.1 M sulfuric acid and 0.2 M potassium hydroxide are mixed. Write the balanced molecular and net ionic equation. (p. 148)

Today… Turn in: Questions on Lab Report so far? Our Plan:
Notes – Oxidation Numbers & Redox Titrations Begin WS#5 (10 minutes) Sample Test (Pep Assembly Today) Homework (Write in Planner): WS#5 Due Monday & Lab Report due Mon

Mixed Review Write net ionic equations for each of the following:
A small piece of sodium metal is add to distilled water. (p. 142) Hydrogen sulfide gas is bubbled into a solution of mercury (II) chloride. (p. 140) Dilute hydrochloric acid is added to a solution of potassium carbonate. (p. 144)

Redox Reactions Balancing Redox Equations is complicated. We will spend an entire week at the end of the year working on it. For now, you just have to identify substances that are oxidized and those that are reduced.

Reactions Involving Oxidation and Reduction
Oxidation: Loss of electrons Reduction: Gain of electrons Both oxidation and reduction must occur simultaneously. A species that loses electrons must lose them to something else (something that gains them). A species that gains electrons must gain them from something else (something that loses them). To remember, use the phrase LEO GER or OIL RIG.

A Redox Reaction: Mg + Cu2+ → Mg2+ + Cu
Electrons are transferred from Mg metal to Cu2+ ions and … … the products are Cu metal and Mg2+ ions.

Oxidation Numbers An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion. Examples: in NaCl, the oxidation number of Na is +1, that of Cl is –1 (the actual charge). In CO2 (a molecular compound, no ions) the oxidation number of oxygen is –2, because oxygen as an ion would be expected to have a -2 charge. The carbon in CO2 has an oxidation number of +4 (Why?)

Rules for Assigning Oxidation Numbers
For the atoms in a neutral species—an isolated atom, a molecule, or a formula unit—the sum of all the oxidation numbers is 0. For the atoms in an ion, the sum of the oxidation numbers is equal to the charge on the ion. In compounds, the group 1A metals all have an oxidation number of +1 and the group 2A metals all have an oxidation number of +2. In compounds, the oxidation number of fluorine is –1.

Rules for Assigning Oxidation Numbers
In compounds, hydrogen has an oxidation number of +1. In most compounds, oxygen has an oxidation number of –2. In binary compounds with metals, group 7A elements have an oxidation number of –1, group 6A elements have an oxidation number of –2, and group 5A elements have an oxidation number of –3. Elements in their standard state (solid, liquid, or gas – look at PT) have an oxidation number of 0.

Identifying Oxidation–Reduction Reactions
In a redox reaction, the oxidation number of a species changes during the reaction. Oxidation occurs when the oxidation number increases (species loses electrons). Reduction occurs when the oxidation number decreases (species gains electrons). If any species is oxidized or reduced in a reaction, that reaction is a redox reaction. Examples of redox reactions: displacement of an element by another element; combustion; incorporation of an element into a compound, etc.

LEO - GER

Example 4.7 What are the oxidation numbers assigned to the atoms of each element in (a) KClO4 (b) Cr2O72– (c) CaH (d) Na2O2 (e) Fe3O4

Fe2O3 (s) + 3CO (g) → 2Fe (s) + 3CO2 (g)
Example 1 Which is oxidized and which is reduced in the following equation: Fe2O3 (s) + 3CO (g) → 2Fe (s) + 3CO2 (g)

MnO2 + 4H+1 + 2Cl-1 → Mn+2 + 2H2O + Cl2
Example 2: Which is reduced/oxidized? MnO2 + 4H Cl-1 → Mn+2 + 2H2O + Cl2

Oxidation Numbers of Nonmetals
The maximum oxidation number of a nonmetal is equal to the group number. For nitrogen, +5. For sulfur, +6. For chlorine, +7. The minimum oxidation number is equal to the group number – 8. ClO4-: what is the ox number of chlorine? Is this an ox agent or a red agent? Cl-? NO3-?

Activity Series of Some Metals
In the activity series, any metal above another can displace that other metal. Mg metal can react with … Will lead metal react with Fe3+ ions? Will iron metal dissolve in an acid to produce H2 gas? … Cu2+ ions to form Cu metal. If a silver coin falls against a nail on a sunken ship, the coin will remain bright for a long time, while the nail will rust badly. Explain.

Applications of Oxidation and Reduction
Everyday life: to clean (bleach) our clothes, sanitize our swimming pools (“chlorine”), and to whiten teeth (peroxide). In foods and nutrition: redox reactions “burn” the foods we eat; antioxidants react with undesirable free radicals. In industry: to produce iron, steel, other metals, and consumer goods.

Redox Titrations In a redox titration, one reactant (often the titrant) is an oxidizing agent and the other reactant is a reducing agent. Permanganate ion, usually from KMnO4, is one of the most commonly used oxidizing agents in the chemical laboratory and makes an excellent titrant. The math is very similar to that used in acid-base titrations.

Calculation Example 1 Example 4.12: A g sample of an iron ore is dissolved in acid, and the iron is converted entirely to Fe+2 (aq). To titrate the resulting solution, L of M KMnO4 (aq) is required. What is the mass percent of iron in the ore? 5Fe+2 (aq) + MnO4-1 (aq) + 8 H+1 (aq) → 5Fe+3 (aq) + Mn+2 (aq) + 4H2O (l)

Calculation Example 2 Example 4.12A: Suppose the titration in Example 4.12 was carried out with M K2Cr2O7 (aq) rather than KMnO4 (aq). What volume of K2Cr2O7 (aq) would be required? 6Fe+2 (aq) + Cr2O7-2 (aq) + 14H+1 (aq) → 6Fe+3 (aq) + 2Cr+3 (aq) + 7H2O (l)

Stop! Complete Worksheet #5 by next class period and finish your titration lab report!

Today… Turn in: Get out WS#5 to Check
Put Lab Report in basket (rubric on top) Our Plan: Questions on WS#5 Worksheet Race Crash Course Video Begin Report for Investigation 8 Homework (Write in Planner): Lab Report Due Friday

Quick Review Label the oxidation numbers for each species and identify which substance is oxidized and which is reduced in the following examples. Cu (s) + AgNO3 (aq) → Ag (s) + Cu(NO3)2 (aq) SeO3-2 (aq) + I-1 (aq) + H+1 (aq) → Se (s) + I2 (g) + H2O (l)

Crash Course Water & Solutions
Good introduction to redox lab:

Today… Turn in: Nothing Our Plan: Investigation 8 Test Review
Homework (Write in Planner): Test Next Class Breakfast Club Friday at 6:30 am!

Today… Turn in: Lab Report – rubric on top Our Plan:
Questions on Test Review Unit 2 Test Homework (Write in Planner): Watch the Alkanes Notes online and complete the booklet to p. 7

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