Chapter 4Reactions in Aqueous Solutions

Many chemical and almost all biological reactions occur in the aqueous medium

Substances (solutes) that dissolve in water (solvent) can be divided into two categories: Electrolytes Non-Electrolytes

Three Major Types of Reactions Precipitation Reaction – the product an

insoluble substance separates from the solution

Acid/Base Reactions – A proton transfer from an acid to a base

Oxidation/Reduction (Redox) “the bane of the AP Test” – Electrons are transferred from a reducing agent to an oxidizing agent

Solution Stoichiometry Quantitative studies with known

concentrations (Molarity) of solutions Gravimetric Analysis Titrations

General Properties of a Solution Solution – a homogenous mixture of two or

more substances Solution may be gaseous (air), solid (alloy) or

Liquid (salt water)

In this chapter we will deal only with aqueous solutions Most common Solvent - Water

Electrolytes versus Nonelectrolytes Electrolytes – Ionic compounds that

completely or partially dissociate in solution with the ability to pass electric current in solution Acids/Bases will ionize in solution, therefore

electricity can be conducted Non-Electrolytes – Molecular compounds

that do not dissociate in solution, therefore no electric current can be pass

Ionic Compounds in Solution Water is a great solvent for ionic

compounds because it is polar, the positive end attracts the Negative Ion and vice versa

Acids and Bases as Electrolytes Some acid/bases competely dissociate in solution

HCl HNO3 H2SO4 Ba(OH)2 NaOH

While others only partially dissociate CH3COOH HF HNO2 HN3

Writing Partial Dissociation Equations

Partial dissociation equations are written with a double arrow, indicating a reversible reaction

Write partial dissociation for CH3COOH

Precipitation Reactions A double replacement reaction (metathesis)

in which a product is insoluble

Solubility RulesIn water at 25 Degrees

All common compounds of Group I and ammonium ions are soluble.

All nitrates, acetates, and chlorates are soluble. All binary compounds of the halogens (other than F) with

metals are soluble, except those of Ag, Hg(I), and Pb. All sulfates are soluble, except those of barium, strontium,

calcium, lead, silver, and mercury (I). The latter three are slightly soluble.

Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.

Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.

Soluble or Insoluble at 25 Degrees Celsius in Water

PbSO4

BaCO3

Li3PO4

FeS Ca(OH)2

Co(NO3)3

Net Ionic Equations Write the correctly balanced equation and

decide on state of each product Write free state of all ions and insoluble

product Cancel out spectator ions – anyone not part

of the reaction Check charges and balancing in net ionic

Practice Net IonicPredict, Balance and write net ionic

1) Lead Nitrate and Potassium Iodide

2) Barium Chloride and Sodium Sulfate

3) Potassium Phosphate and Calcium Nitrate

4) Aluminum Nitrate and Sodium Hydroxide

Acid – Base Reactions Acids react with metal such as Zn, Mg and

Fe to produce hydrogen gas Acids react with carbonates and

bicarbonates to produce carbon dioxide gas, water and the salt

Bronsted Acid and Bases Bronsted Acid is a proton donor Bronsted Base is a proton acceptor

HCl (aq) H+ (aq) + Cl-(aq)

In water the H+ attracts to the water molecule producing the hydronium ion

Monoprotic Acids Each unit of acid yields one hydrogen ion

upon ionization

Diprotic Acids Each unit of the acid gives up two hydrogen

ions in two separate steps (they strip)

Triprotic Acids Yield three hydrogen ions in three separate

steps (they strip)

 

Bronsted Acid is a proton donorBronsted Base is a proton acceptor

Classify each of the following as an Bronsted acid or Bronsted base, explain your reasoning based on the definition HBr

SO-2

4

HI

HCO-

3

NO2

Neutralization ReactionAcid and Base will form Salt and

Water

Write the net ionic for the following Hydrochloric acid and Sodium Hydroxide Sulfuric acid and Aluminum Hydroxide

Acid – Base Reactions Leading to Formation of a Gas

Certain Salts – Carbonates, bicarbonates, sulfites and sulfides react with acids to form gaseous products

Oxidation Numbers Oxidation Reaction – refers to half reaction that

involves loss of electrons Reduction reaction – refers to a half reaction that

involves the gain of electrons The extent of oxidation in a redox reaction must

be equal to the extent of reduction; that is the number of electrons lost by a reducing agent must be equal to the number of electrons gained by an oxidizing agent

The half-reactions of a redox reaction or oxidation-reduction reaction

Oxidation Number The number of charges the atom would

have in a molecule if electrons are transfer completely

The convention is that the cation is written first in a formula, followed by the anion.

For example, in NaH, the H is H-; in HCl, the H is H+. The oxidation number of a free element is always 0. The atoms in He and N2, for example, have oxidation numbers of 0. The oxidation number of a monatomic ion equals the charge of the ion. For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3. The usual oxidation number of hydrogen is +1. The oxidation number of hydrogen is -1 in compounds containing elements that are

less electronegative than hydrogen, as in CaH2. The oxidation number of oxygen in compounds is usually -2. Exceptions include OF2, since F is more electronegative than O, and BaO2, due to

the structure of the peroxide ion, which is [O-O]2-. The oxidation number of a Group IA element in a compound is +1. The oxidation number of a Group IIA element in a compound is +2. The oxidation number of a Group VIIA element in a compound is -1, except

when that element is combined with one having a higher electronegativity. The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl. The sum of the oxidation numbers of all of the atoms in a neutral compound is

0. The sum of the oxidation numbers in a polyatomic ion is equal to the charge of

the ion. For example, the sum of the oxidation numbers for SO42- is -2.

Assign oxidation numbers to all the elements in the following compounds

Na2O

HNO2

Cr2O7-2

PF3

MnO4-

Arrange the following species in order of increasing oxidation number of the sulfur atoms

H2S

SO2

SO3

S8

H2SO4

S-2

HS-

Concentration Molarity = moles of solute/liters of solution

Example: How many grams of potassium dichromate are required to prepare a 125ml solution whose concentration is 1.83M

Concentration In a biochemical assay a chemist needs a

to add 4.07g of glucose to a reaction mixture. Calculate the volume in milliliters the volume of a 3.16M glucose she should use

Dilution of Solutions The procedure of making a less

concentrated solution from a high concentration solution

MiVi = MfVf

Dilution Problem Describe hou you would prepare 2.50 * 102

ml of a 2.25M H2SO4 solution, starting with a 7.41 M stock solution of H2SO4

Dilution Problem #2 How would you prepare a 200ml of a .866M

KOH solution, starting with 5.07M stock solution

Acid – Base Titrations In a titration a solution of an accurately

known concentration, called the standard is added gradually to another solution of unknown until reaction is neutralized (equivalence point)

Indicators are used to color the reaction when it is complete

Titration Problem In a titration experiment, a student finds

that 25.46ml of a NaOH solution is needed to neutralize .6092g of KHP. What is the concentration of the NaOH solution?

Titration Problem #2 How many milliliters of a .836M NaOH

solution is needed to neutralized 25ml of a .335M of H2SO4?

Solution Stoichiometry

When sodium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 150ml 3M of silver nitrate?

1. When Magnesium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 4.5M in 250ml of silver nitrate? What is the name of the other product of the reaction?