Aqueous Reactions & Solutions Stoichiometry

CHAPTER 4 :AQUEOUS REACTIONS AND SOLUTION STOICHIOMETRY

QBA Miguel A. Castro Ramírez

http://www.colamsc.com/index.php?e4183c4ebf0f1cd56bbe0f966169ac1d=8103d50a48ab8a2255d0b2d9929f9dd6

GENERAL PROPERTIES OF AQUEOUS SOLUTIONS

Solutions

• Solutions are defined as homogeneous mixtures of two or more pure substances.

• The solvent is present in greatest abundance.

• All other substances are solutes.

• Aqueous Solution: A solution which water is the dissolving medium.


Electrolytic Properties

• Electrolyte : A substance whose aqueous solutions contain ions

: Substances that dissolves in water to give an electrically conducting solution.

•Nonelectrolyte : A substance that does not form ions in solution

: Substances that dissolves in water to give an a nonconducting or very poorly

condusting solution.


Ionic Compounds in Water

• Water is a very effective solvent of ionic compounds

• Electrically neutral molecule, but one end of the molecule (O atom) rich in electron and has a partial negative charge (δ-)

• The other end (H atoms), has a partial positive charge (δ+)


• Cation (+) are attractive to (δ-) and anions (-) are attractive to (δ+)• As an ionic compound dissolves, the ions become surrounded by H2O molecules.

• The ions are said to be solvated. • The solvation process helps stabilize ions in solution and prevents cations and anions from recombining.• The ions and their shells (water molecules) are free to move, the ions become dispersed uniformly throughout the solution.


Molecular Compounds in Water

• Dissolves in water, solution usually consist of intact molecules dispersed throughout the solution.• So, molecular compounds are non-electrolytes.• Example: Table sugar (sucrose) and methanol.• A few molecular substances have aqueous solutions that contain ions.• Example: HCl – Ionizes and dissociates into H+(aq) and Cl-(aq) ions.

Strong and Weak Electrolytes


• Two categories: Weak and strong- Differ in the extend to which they conduct electricity.

• Strong electrolytes: Dissociates completely when dissolved in water.• All soluble ionic compounds are electrolytes Eg: NaCl• Strong acid, strong base and soluble ionic compounds.

• Weak electrolytes: Only dissociates partially when dissolved in water.

• E.g: CH3COOH: Most of the solute is present as CH3COOH molecules. Only small fraction of the acid is present as H+(aq) and CH3COO-(aq) ions.


• For strong electrolyte: Single arrow represent the ionization of strong electrolytes.

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

• The absence of a reverse arrow indicates that the H+ and Cl- ions have no tendency to recombine in water.

PRECIPITATION REACTIONS

• Precipitation Reaction: Reaction that result in the formation of an insoluble product.• Precipitate: An insoluble solid formed by a reaction in solution. Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)

• To predict whether certain combinations of ions form insoluble compounds, we must consider some guidelines concerning the solubilities of common ionic compounds.

Solubility Guidelines for Ionic Compounds


• Solubility: The amount of substance that can be dissolved in a given quantity of solvent at the given temperature.

• Insoluble: The attraction between the oppositely charged ions in the solid is too great for the water molecules to separate the ions to any significant extend- substance remains undissolved.

• No rules based on physical properties to predict the solubility of compound.


• Based on the experimental observation only.

• All common ionic compounds of alkali metals ions and the NH4

+ are soluble in water.


• To predict the forming of precipitation:1) Note the ion present in the reactant2) Consider the possible combinations of the cations and anions3) Determine whether the any of the product is insoluble

Example:Mg(NO3)2 react with NaOH. Will the precipitate form?

• Existing ions: Mg2+, NO3-, Na+ and OH-

2) Possible reaction: Mg2+ and OH- ; Na+ and NO3

-

3) Products: Mg(OH)2 and NaNO3 - Mg(OH)2 is insoluble.

- NaNO3 is soluble 4) Balanced equation:

Mg(NO3)2 (aq) + 2NaOH(aq) Mg(OH)2 (s) + NaNO3 (s)


Exchanged (Metathesis) Reactions

• Metathesis comes from a Greek word that means “to transpose.”•

• To complete and balance a metathesis reactions:

1. Use the chemical formulas of the reactants to determine the present ions

2. Write the chemical formulas of the products by combining the cation from one reactant with anion from another reactant.

3. Balance the equation.

AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)


The molecular equation

Shows all of the reactants and products as intact, undissociated compounds.

AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)

The total ionic equation

Shows all of the soluble ionic substances dissociated into ions.

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

Writing Equations for Aqueous Ionic Reactions


Writing Equations for Aqueous Ionic Reactions

The net ionic equation

• Eliminates the spectator ions and shows the actual chemical change taking place.

• Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions

Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) →

AgCl (s) + K+(aq) + NO3-(aq)

Ag+(aq) + Cl-(aq) → AgCl (s)

1. Write a balanced molecular equation.2. Dissociate all strong electrolytes.3. Cross out anything that remains

unchanged from the left side to the right side of the equation.

4. Write the net ionic equation with the species that remain.


Writing Net Ionic Equations


EXAMPLE

Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction does occur, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions.(a) potassium fluoride(aq) + strontium

nitrate(aq) (b) ammonium perchlorate(aq) + sodium bromide(aq)

Using Molecular Depictions to Understand a Precipitation Reaction


Consider the molecular views of the reactants for a precipitation reaction.

(a) A: KCl, Na2SO4, MgBr2, or Ag2SO4?

(b) B: NH4NO3, MgSO4, Ba(NO3)2, or CaF2?

(c) Name precipitate and speactator ions from reaction. Write balanced molecular, total ionic, and net ionic equations.

ACID-BASE REACTIONS

• Acids: Substance that ionize in aqueous solutions to form H+ ions, thereby increasing the concentration of H+ (aq) ions.• Acids : Called as proton donors. •Molecules from different acids can ionize to form different numbers of H+ ions.

There are only seven strong acids:• Hydrochloric (HCl)• Hydrobromic (HBr)• Hydroiodic (HI)

• Nitric (HNO3)

• Sulfuric (H2SO4)

• Chloric (HClO3)

• Perchloric (HClO4)

ACID-BASE REACTIONS

ACID-BASE REACTIONS

• Base: Substance that accept or react with H+ ions.• Produce OH- ions when dissolve in water• The strong bases are the soluble metal salts of hydroxide ion

Example: Ca(OH)2

• Compounds that do not contain OH- also can be base.

Example: NH3 is common base. When added to water, it accepts and H+ ion and produces OH- ion.

NH3 (aq) + H2O (l) NH4+ (aq) + OH-

(aq)

- NH3 is a weak electrolyte. Because only small fraction of NH3 forms NH4

+ and OH- ions.

ACID-BASE REACTIONS

ACID BASE

Arrhenius

Substances that increase the concentration of H+ when dissolved in water

Substances that increase the concentration of OH− when dissolved in water

Brønsted and Lowry

Proton donorsProton acceptors

Neutralization Reactions and Salts

ACID-BASE REACTIONS

• Properties of acidic solutions are different from the basic one.• Neutralization reaction occurs when a solution of acid and base are mixed.• The products of the reaction have none of the characteristic properties of either the acidic or basic solution.• Neutralization process will produce salt and water as products.

• CH3COOH (aq) + NaOH (aq) →CH3COONa (aq) + H2O (l)

Neutralization Reactions and Salts

ACID-BASE REACTIONS

• When a strong acid reacts with a strong base, the net ionic equation is:HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) →

Na+ (aq) + Cl- (aq) + H2O (l)

H+ (aq) + OH- (aq) → H2O (l)

Acid-Base Reactions with Gas Formation

ACID-BASE REACTIONS

• Some metathesis reactions do not give the product expected.

• When a carbonate or bicarbonate reacts with an acid, first gives the carbonic acid (H2CO3) and salt.

HCl(aq) + NaHCO3(aq) NaCl(aq) + H2CO3

• Carbonic acid (H2CO3) then decomposed and the products become salt, carbon dioxide, and water

H2CO3(aq) H2O(l) + CO2(g)

• The overall reaction is summarized:

Molecular eq: NaHCO3 (aq) + HBr (aq) →NaBr (aq) + CO2 (g) + H2O (l)

Total ionic eq: Na+(aq) + HCO3-(aq) + H+(aq) + Br-

(aq) Na+(aq) + Br-(aq) + CO2

(g) + H2O (l)

Net ionic eq: HCO3-(aq) + H+(aq) + CO2 (g) +

H2O (l)

ACID-BASE REACTIONS

OXIDATION-REDUCTION (REDOX) REACTIONS

• Which electrons are transferred between reactants.

• Corrosion: The conversion of a metal into a metal compound by a reaction between the metal and its enviroment.

• When a metal corrodes, it loses electrons and form cations.

• Example: Calcium vigorously attacked by acids to form calcium ion:

Ca(s) + 2H+ (aq) Ca2+ (aq) + H2 (g)

• Oxidation: Loss of electron by a substance• The term oxidation is used because the

1st reactions of this sort to be studied thoroughly were reactions with oxygen.

• Many metals react directly with O2 in air to form metal oxide.

• Example: 2Ca(s) + O2(g) 2CaO(s)

• Ca losses e- to produce CaO which is ionic compound.

• Ca is oxidized (lose e-) while O2 gain e- and transformed to O2- ion. O2 is said to be reduced.

• Reduction: The gain of electron by a substance.

• The oxidation and reduction process in in-situ process.



Oxidation Numbers

• To determine if an oxidation-reduction reaction has occurred, we assign an oxidation number (also called oxidation states) to each element in a neutral compound or charged entity.

• The oxidation numbers of certain atoms change in an oxidation-reduction reaction.

• Oxidation occurs when the oxidation number increase, while reduction occurs when the oxidation number decreases.

• We use the following rules for assigning the oxidation numbers:

• Elements in their elemental form have an oxidation number of 0.• The oxidation number of a monatomic ion is the same as its charge. (e.g; K+ has an oxidation number of +1) – We write the sign +/- 1st before write the number.• Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.

• Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.• Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.


• Fluorine always has an oxidation number of −1.

• The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

• The sum of the oxidation numbers in a neutral compound is 0.

• The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.


EXAMPLE:

Determine the oxidation number (O.N.) of each element in these compounds:


(a) zinc chloride

(b) sulfur trioxide

(c) nitric acid

EXERCISE:

Which reactions are redox reactions:


(a) CaO(s) + CO2(g) CaCO3(s)

(b) 4 KNO3(s) 2 K2O(s) + 2 N2(g) + 5 O2(g)(c) NaHSO4(aq) + NaOH(aq) Na2SO4(aq) + H2O(l)


Oxidation of Metals by Acids and Salts

• The reaction of an acid or metal salt conforms to the general following pattern:

A + BX AX + B Zn(s) + 2 HBr(aq) ZnBr2(aq) + H2(g)

• These reactions are called displacement reactions because the ionic solution is displaced or replaced through oxidation of an element.

• Many metals undergo displacement reactions with an acid to produce salt and H2 gas.


• In displacement reactions, ions oxidize an element.• The ions, then, are reduced.• To show that the redox reaction have occurred, the oxidation number is shown below:

Mg (s) + 2HCl(aq) MgCl2(aq) +H2(g)

0 +1-1 +2 -1 0


• Metal can also be oxidized by aqueous solutions of various salt.• Example:

Molecular Eq: Fe(s) + Ni(NO3)2(aq) Fe(NO3)2(aq) + Ni(s)

Net ionic Eq: Fe(s) + Ni2+ (aq) Fe2+ (aq) + Ni (s)

• The oxidation of Fe to form Fe2+ in this reaction is accompanied by the reduction of Ni2+ to Ni.• Whenever a substance is oxidized, some other substance must be reduced.


The Activity Series

• A list of metals arranged in order of decreasing ease of oxidation is called: Activity series • TOP: Most easily oxidized (1A) : Active metals• BOTTOM: Stable (8B n 1B) : Noble metal• Can be used to predict the outcome of reactions between metals and either metal salts or acids. • Any metal on the list can be oxidized the ions of elements below it.


• For example: Copper is above silver in the series. Thus, copper metal will be oxidized by silver ions:

Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)


• Only those metals above hydrogen in the activity series are able to react with acids to form H2. • Example: Ni(s) + 2HCl(aq) NiCl2 (aq) + H2 (g)

• Because elements below hydrogen in the activity series are not oxidized by H+, Cu doesn’t react with HCl(aq). But does react with HNO3.

• This reaction however, is not a simple oxidation of Cu by the H+ ions of the acid. Instead, the metal is oxidized to Cu2+ by the nitrate ion of the acid, accompanied by the formation of brown NO2(g):Cu(s) + 4 HNO3(aq) Cu(NO3)2 (aq) + 2H2O (l) +

2 NO2(g)

CONCENTRATIONS OF SOLUTIONS

• Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different.• Molarity is one way to measure the concentration of a solution

• A molar of solutions (1 M) contains 1 mol of solutes in every 1 L solution.

moles of solute

volume of solution in litersMolarity (M) =

Molarity

What is the difference between 0.5 mol of H2SO4 and 0.5 Molar of H2SO4?

A •Weigh the solid needed.•Transfer the solid to a volumetric flask that contains about half the final volume of solvent.

B Dissolve the solid thoroughly by swirling.

C Add solvent until the solution reaches its final volume.


Laboratory preparation of molar solutions

EXAMPLE:

What is the molarity of an aqueous solution that contains 0.715 mol of glycine (H2NCH2COOH) in 495 mL?


EXAMPLE:

How many grams of solute are in 1.75 L of 0.460 M sodium monohydrogen phosphate buffer solution?

EXERCISE:

An experiment calls for the addition to a reaction vessel of 0.184 g of sodium hydroxide. How many mililiters of 0.150 M NaOH should be added.


Expressing the Concentration of an Electrolyte


• When ionic compound dissolves, the relative concentrations of the ions introduced into the solution depend on the chemical formula of the compound.

• Example: 1 M solution of NaCl is 1 M in Na+ ions and 1 M in Cl- ions.

: 1 M of Na2SO4 is 2 M of H+ ions and 1 M in SO4

2- ions.

• The conc of an electrolyte solution can be specified either in term of the compound used to make the solution (1 M of Na2SO4) or in terms of the ions that the solution contains (2M Na+ and 1M SO4

2-)

Dilution

• Stock solutions: Solutions that are routinely used in the lab are often purchased or prepared in concentrated form.

• Dilution: The process of obtained lower concentration from high concentration solutions by adding water.

• One can also dilute a more concentrated solution by– Using a pipet to deliver a volume of the

solution to a new volumetric flask, and– Adding solvent to the line on the neck

of the new flask.



• The molarity of the new solution can be determined from the equation

Moles solute before dilution = moles solute after dilution

Mc × Vc = Md × Vd,

where Mc and Md are the molarity of the concentrated and dilute solutions, respectively, and Vc and Vd are the volumes of the two solutions.

• The concentration can be in L or mL as long as the unit is the same in both side.

(1.00 M)(Vconc) = (0.100 M)(250 mL)

Vconc = 25 mL (1.00 M)(Vconc) = (0.100 M)(0.25 mL)

Vconc = 0.025 mL

SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS

• Problem solving procedure : Outline of the procedure used to solve stoichiometry problems that involve measured (lab) units of mass, solution conc (Molarity) or volume.


EXAMPLE:

How many grams of NaOH are needed to neutralize 20.0 mL of 0.150 M H2SO4 solution?


Titration

• To determine the concentration of particular solute in a solution : Titration.

• Titration: Acid-base, precipitation or redox reaction.

• Equivalence point: The point which stoichiometry equivalent quantities are brought together.

• Indicators: - Used to determine the end point of the reaction.

- Color changes of the indicator showing that reaction had occur.

- Example: Phenolphthalein: Acidic- Colorless, Basic- pink.


EXAMPLE:

• What volume of 0.128 M HCl is needed to neutralize 2.87 g of Mg(OH)2?

EXERCISE

A solution of 100 mL of 0.200 M KOH is mixed with a solution of 200.00 mL of 0.150 M NiSO4.

e) Write the balanced chemical equation for the reaction.f) What precipitate forms?g) What is the limiting reactant?h) How many grams of the precipitate form?


I am only one;But still I am one;I cannot do everything;But still I can do something.I will not refuse to do something I can do.

-Helen Keller-

Aqueous Reactions & Solutions Stoichiometry

Documents

Transcript of Aqueous Reactions & Solutions Stoichiometry