Chapter 4 Types of Chemical Reactions and Solution Stoichiometry 2 PDF

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CHEMISTRY 120/121

Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

Overview: 

Selective Precipitation



Acid-Base Reactions



Oxidation-Reduction Reaction



Balancing Redox Reactions

Types of Chemical Reactions and Solution Stoichiometry Ch 4.7-4.12 Selective Precipitation We can use the fact that salts have different solubility to separate mixtures of ions. e.g. Suppose we have an aqueous solution containing the cations Ag+, Ba2+, and Fe3+ and the anion NO3-. We want to separate the cations by precipitating them one at a time, a process called ____________________ ____________________. How can this goal be accomplished? *How much solid NaCl must be added to 1.50 L of a 0.100 M AgNO3 solution to precipitate all the Ag+ in the form of AgCl?

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Acid-Base Reactions Acid-Base reactions are also called ____________________ reactions because the acid and base neutralize each other’s properties.

The net ionic equation for an acid-base reaction is

(as long as the salt that forms is soluble in water) Recall: Acids ionize in water to form H+ ions. More precisely, the ____________________ from the acid molecule is donated to a water molecule to form hydronium ion, ____________________. Note: Most chemists use H+ and H3O+ interchangeably. Bases dissociate in water to form OH- ions. Bases, like NH3, that do not contain OH- ions, produce OH- by pulling H+ off water molecules.

In the reaction of an acid with a base, the H+ from the acid combines with the OH- from the base to make ____________________. The cation from the base combines with the anion from the acid to make the ____________________.

Types of acid/base reactions: HCl/NaOH 

Strong base + strong acid (____________________ reaction): HCl(aq) + NaOH(aq)  NaCl(aq) + H2O H+ + Cl- + Na+ + OH-  H2O + Na+ + ClH+ + OH-  H2O i.e. products: salt + water



Strong base + weak acid: KOH/CH3COOH HC2H3O2 (CH3COOH) + K+ + HO-  K+ + H2O + C2H3O2HO- + HC2H3O2  H2O + C2H3O210 J. Zhou Summer 2018

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Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

Strong acid + weak base: HCl/NH3 H+ + Cl- + NH3  NH4+ + ClH+ + NH3  NH4+ (reacts completely)

Example: Predict the products of the following reactions and write balance molecular, complete ionic and net ionic equations for each reaction.

a)

HCl (aq) + Ba(OH)2 (aq) →

b)

H2SO4 (aq) + Ca(OH)2 (aq) →

Acid-Base Titrations When a substance being analyzed contains an acid, the amount of acid present is usually determined by titration with a standard solution containing hydroxide ions.

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Example: 25.00 mL of a 0.350 M NaOH solution is needed to neutralize 87.50 mL of an HCl solution. Calculate the concentration of the HCl solution?

Example: 1.3009 g of KHP (KHC8H4O4) is dissolved in water and titrated with an unknown concentration of NaOH. 41.20 mL of NaOH is added to reach the endpoint. What is the concentration of the NaOH solution? Note that KHP is a monoprotic acid.

Redox Reactions Many chemical reactions involve the transfer of electrons between atoms or ions. The flow of electrons is associated with electricity. Reactions in which electrons are transferred from one atom to another are called ____________________ - ____________________ (____________________) reactions. Atoms that lose electrons are being ____________________; atoms that gain electrons are being ____________________. ____________________ is the process that occurs when the oxidation number of an element ____________________ and an element ____________________ electrons. 

A compound gains oxygen



A compound loses hydrogen



A half-reaction has electrons as products

e.g.

2Na (s) + Cl2 (g) → 2NaCl (s)

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We can assign oxidation number of an atom in a molecule or ion in order of priority: 1) Free elements have an oxidation state = 0. Na = 0 and Cl2 = 0 in 2 Na(s) + Cl2(g) 2) Monatomic ions have an oxidation state equal to their charge. Na = +1 and Cl = −1 in NaCl 3a) The sum of the oxidation states of all the atoms in a compound is 0. Na = +1 and Cl = −1 in NaCl, (+1) + (−1) = 0 3b) The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion. N = +5 and O = −2 in NO3–, (+5) + 3(-2) = −1 4a) Group 1 metals have an oxidation state of +1 in all their compounds. Na = +1 in NaCl 4b) Group 2 metals have an oxidation state of +2 in all their compounds. Mg = +2 in MgCl2 5) In their compounds, nonmetals have oxidation states according to the table below. Nonmetal

Oxidation State

Example

F

−1

CF4

H

+1

CH4

O

−2

CO2

Group 17

−1

CCl4

Group 16

−2

CS2

Group 15

−3

NH3

Example: Determine the oxidation states of all the atoms in a propanoate ion, C3H5O2-

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Recall: Oxidation occurs when an atom’s oxidation state increases during a reaction. Reduction occurs when an atom’s oxidation state decreases during a reaction. e.g.

3Cl2

+ I-

+ 3H2O

→

6Cl- + IO3- + 6H+

We generally split the redox reaction into two separate ____________________- reactions.

Oxidation and reduction must occur simultaneously. If an atom loses electrons, another atom must take them. The reactant that reduces an element in another reactant is called the ____________________ agent. The reducing agent contains the element that is oxidized. The reactant that oxidizes an element in another reactant is called the ____________________ agent, The oxidizing agent contains the element that is reduced.

Example: Assign oxidation states, determine the elements oxidized and reduced, and determine the oxidizing agent and reducing agent in the following reactions. a) Sn4+ + Ca → Sn2+

b) F2

+ S

+ Ca2+

→ SF 4

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Balancing Redox Reactions Some redox reactions can be balanced simply be inspection, but many are hard to balance using that method. 

Many are written as net ionic equations



Many have elements in multiple compounds.

The main principle is that electrons are transferred—so if we can find a method that will keep track of the electrons; it will allow us to balance the equation.

Balancing Redox Reactions by the Half-Reaction Method: 

In this method, the reaction is broken down into two-half reactions, one for oxidation and another for reduction.



Each half-reaction includes electrons. o Electrons go on the product side of the oxidation half-reaction—loss of electrons o Electrons go on the reactant side of the reduction half-reaction—gain of electrons



Each half-reaction is balanced for its atoms.



Then the two half-reactions are adjusted so that the electrons lost and gained will be equal when combined.

1) Assign oxidation states a. Determine element oxidized and element reduced 2) Write oxidation and reduction half-reactions, including electrons a. Oxidation electrons on right; reduction electrons on left 3) Balance half-reactions by mass a. First balance elements other than H and O b. Add H2O where O is needed c. Add H+ where H is needed 4) Balance half-reactions by charge a. Balance charge by adjusting electrons 5) Balance charge by adjusting electrons 6) Add half-reactions 7) Check! (If reaction was done in base, neutralize H+ with OH-)

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Example: Balance the following equations in acidic solution. a) Fe2+

b) I-

+ MnO4- →

+ Cr2O72- →

Fe3+

+ Mn2+

Cr3+ + I2

Example: Balance the following redox reaction occurring in basic solution. CrI3 (s)

+ Cl2 (g) → CrO42- + IO4- + Cl-

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