BALANCE ALL REACTIONS

4)
(a) Solutions of tin(II) chloride and iron(III) chloride are mixed.
(b) Solutions of cobalt(II) nitrate and sodium hydroxide are mixed.
(c) Ethene gas is burned air.
(d) Equal volumes of equimolar solutions of phosphoric acid and potassium hydroxide are mixed.
(e) Solid calcium sulfite is heated in a vacuum.
(f) Excess hydrochloric acid is added to a solution of diamminesilver(I) nitrate.
(g) Solid sodium oxide is added to distilled water.
(h) A strip of zinc is added to a solution of 6.0-molar hydrobromic acid.

4. a) A small piece of calcium metal is added to hot distilled water.
b) Butanol is burned in air.
c) Excess concentrated ammonia solution is added to a solution of nickel(II) sulfate.
d) A solution of copper(II) chloride is added to a solution of sodium sulfide.
e) A solution of tin(II) nitrate is added to a solution of silver nitrate.
f) Excess hydrobromic acid solution is added to a solution of potassium hydrogen carbonate.
g) Powdered strontium oxide is added to distilled water.
h) Carbon monoxide gas is passed over hot iron(III) oxide.

Unit 3
Thermochemistry
Basic Terms

·         System – the part of the universe that we are studying.

·         Surroundings – the rest of the universe.

·         Universe – includes both system and surroundings

·         Univ. = Sys + Surr

 
Basic Concepts

·         Open System – exchanges energy and matter with its surroundings.

–Cup of coffee no lid.

·         Closed system – exchanges only energy with its surroundings.

–Sun tea

·         Isolated system – exchanges no energy or matter.

–Closed thermos (for a short period)

 
Basic Concepts

·         Law of Conservation of Energy

–Energy can be converted from one form to another but can neither be destroyed or created.

·         Heat – the transfer of energy due to differences in temperature.(q)

·         Important note – the sign of all calculations is extremely important, it reflects energy flow.

 
Basic Concepts

·         In thermodynamics we always calculate a change from final to initial conditions.

Ex. ΔT = Tf – Ti

·         Exothermic – energy flows out of the system to the surroundings.

·         Endothermic – energy flows into the system from the surroundings.

 
Energy (E)

·         Potential energy – capacity to do work based on position or composition.

·         Kinetic Energy – capacity of moving objects to do work.

·         Work (w) – force x distance or

                                                            W = -PΔV

Energy (E)
 
Enthalpy (H)
 
Enthalpy

                        ΔH = ΔHproducts – ΔHreactants

 
Calorimetry

·         Heat Capacity – the heat required to raise the temp of a system 1oC is given by

                                                C = q/ΔT

Specific Heat – the heat capacity of 1 gram of a substance

                        q = mass x specific heat x ΔT

 
Hess’s Law of Heat Summation

·         The heat of a reaction is constant, whether the reaction is carried out in 1 step or a series of steps.

·         Rules for Hess’s Law problems

–If the reaction is reversed, reverse the sign of ΔH.

–If the reaction is multiplied, multiply ΔH by that same factor.

 
Hess’s Law

·         Ex. Calculate the enthalpy change for the oxidation of ethanol to acetic acid.

            C2H5OH(l) + O2(g) à CH3COOH(l) + H2O(l)

Given the following:

C2H5OH(l) + O2(g) à CO2(g) + H2O(l) ΔH = -1370 kJ/mol

CH3COOH(l) + O2(g) à CO2(g) + H2O(l) ΔH = -874kJ/mol

 
Hess’s Law

·         Make the two equations equal your main equation.

            C2H5OH(l) + O2(g) à CH3COOH(l) + H2O(l)

C2H5OH(l) + O2(g) à CO2(g) + H2O(l) ΔH = -1370 kJ/mol

CO2(g) + H2O(l) à CH3COOH(l) + O2(g) DH = 874 kJ/mol

Add the values for ΔH = -1370 + 874 = -496kJ/mol
 
Standard Enthalpies of Formation (ΔHof)

·         ΔHof is defined as the change in enthalpy that accompanies the formation of one mol of a compound from its elements with all substances in their standard states.

·         A degree sign always means standard 25oC and 1 atm. If a solution it is 1M.

 
Standard Enthalpies of Formation

·         ΔHof = ∑npΔHof (products ) –       ∑nr ΔHof (reactants )

 Elements in their standard states have enthalpies equal to zero.

 
Standard Enthalpies of Formation

·         EX. Calculate ΔHof for the following reaction.

·         CH4(g) + O2(g) à CO2(g) + H2O(g)

·         Balanced CH4(g) + 2O2(g) à CO2(g) + 2H2O(g)

·         From appendix 4 in your book pg. A21

            ΔHof values

            CH4 = -75kJ/mol        O2 = 0kJ/mol

            CO2 = -393.5kJ/mol    H2O = -242kJ/mol

 
Standard Enthalpies of Formation

·         Use equation and plug in

·         ΔHof = ∑ np ΔHof (products ) –  ∑nr ΔHof (reactants )

ΔHof = [1(-393.5) + 2(-242)] – [1(-75) + 2(0)]

ΔHof = -802.5kJ
Read and be able to answer questions on 6.5 and 6.6
 


1) A 1.2156-gram sample of a mixture of CaCO3 and Na2SO4 was analyzed by dissolving the sample and completely precipitating the Ca2+ as CaC2O4. The CaC2O4 was dissolved in sulfuric acid and the resulting H2C2O4 was titrated with a standard KMnO4 solution.

(a) On a page of your answer booklet, write the balanced equation for the titration reaction, shown balanced below.

2MnO4¯ + 5H2C2O4 + 6H+ ----> 2Mn2+ + 10CO2 + 8H2O

Indicate which substance is the oxidizing agent and which substance is the reducing agent.
(b) The titration of the H2C2O4 obtained required 35.62 milliliters of 0.1092-molar MnO4¯ solution. Calculate the number of moles of H2C2O4 that reacted with the MnO4¯.
(c) Calculate the number of moles of CaCO3 in the original sample.
(d) Calculate the percentage by weight of CaCO3 in the original sample.

 
 2 Write the formulas to show the reactants and the products for any FIVE of the laboratory situations described below. Answers to more than five choices will not be graded. In all cases, a reaction occurs. Assume that solutions are aqueous unless otherwise indicated. Represent substances in solution as ions if the substances are extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. You BALANCEthe equations
a.)    A solution of sodium iodide is added to a solution of lead (II) acetate.
 b.)     Pure solid phosphorus (white form) is burned in air.
 c.)     Solid cesium oxide is added to water.
 d.)     Excess concentrated hydrochloric acid is added to a 1.0 M solution of cobalt (II) chloride.
 e.)    Solid sodium hydrogen carbonate (sodium bicarbonate) is strongly heated.
 f.)      An excess of hydrochloric acid is added to solid zinc sulfide.
 g.)     Acidified solutions of potassium permanganate and iron (II) nitrated are mixed together.
 h.) A solution of potassium hydroxide is added to solid ammonium chloride.

Write Reactions  YOU MUST BALANCE ALL REACTIONS

 4.             a)             Sulfur dioxide gas is bubbled into distilled water.

b)                   A drop of potassium thiocyanate solution is added to a solution of iron(III) nitrate.

c)                   A piece of copper wire is placed in a solution of silver nitrate.

d)                   Solutions of potassium hydroxide and propanoic acid are mixed.

e)                   A solution of iron(II) chloride is added to an acidified solution of sodium dichromate.

f)                    Chlorine gas is bubbled through a solution of potassium bromide.

g)                   Solutions of strontium nitrate and sodium sulfate are mixed.

 Powdered magnesium carbonate is hearted strongly.

 
4)

 (a) Solutions of tin(II) chloride and iron(III) chloride are mixed.

(b) Solutions of cobalt(II) nitrate and sodium hydroxide are mixed.

(c) Ethene gas is burned air.

(d) Equal volumes of equimolar solutions of phosphoric acid and potassium hydroxide are mixed.

(e) Solid calcium sulfite is heated in a vacuum.

(f) Excess hydrochloric acid is added to a solution of diamminesilver(I) nitrate.

(g) Solid sodium oxide is added to distilled water.

(h) A strip of zinc is added to a solution of 6.0-molar hydrobromic acid.

 


4) Give the formulas to show the reactants and the products for FIVE of the following chemical reactions. Each of the reactions occurs in aqueous solution unless otherwise indicated. Represent substances in solution as ions if the substance is extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. In all cases a reaction occurs. You need not balance.

Example: A strip of magnesium is added to a solution of silver nitrate.

 

Mg + Ag+ ---> Mg2+ + Ag

(a) Excess sodium cyanide is added to a solution of silver nitrate.

(b) Solutions of manganese(II) sulfate and ammonium sulfide are mixed.

(c) Phosphorous(V) oxide powder is sprinkled over distilled water.

(d) Solid ammonium carbonate is heated.

(e) Carbon dioxide gas is bubbled through a concentrated solution of potassium hydroxide.

(f) A concentrated solution of hydrochloric acid is added to solid potassium permanganate.

(g) A small piece of sodium metal is added to distilled water.

(h) A solution of potassium dichromate is added to an acidified solution of iron(II) chloride.


Acitivity (Elecrtromotive) Series of Metals & Nonmetals
 (Listed by decreasing chemical activity)
Metals
Nonmetals
Li
F
K
Cl
Ba
Br
Ca
I
Na
 
Mg
 
Al
 
Zn
 
Cr
 
Fe
 
Cd
 
Co
 
Ni
 
Sn
 
Pb
 
H2
 
Cu
 
Hg
 
Ag
 
Au
 
                                                        
 
Solubility Rules
 

1. All salts of group IA and the ammonium ion are soluble.

2. All acetates and nitrates are soluble.

3. All binary compounds of group VIIA are soluble except silver,mercury (I), and lead.

4. All sulfates are soluble except those of Ba, Sr, Pb, Ca, and Mercury (I).

5. All carbonates, hydroxides, oxides, sulfides, and phosphates are insoluble except those of group IA and the ammonium ion.

           
General classes of decomposition reactions:
 
1. When some acids are heated, they

    decompose to form water and an acidic

    oxide.
 

      H2CO3(aq) ---> CO2(g) + H2O(l)

 
2. When some metallic hydroxides are

    heated, they decompose to form a

    metallic oxide and water.

 
    Ca(OH)2(s) ---> CaO(s)   + H2O(g)
 
3. When some metallic carbonates are

    heated, they decompose to form a

    metallic oxide and carbon dioxide.

   Li2CO3(s) -----> Li2O(s)   + CO2(g)

 
 
4. When some metallic chlorides are

     heated, they decompose to form

     metallic chlorides and oxygen.

 

 2KClO3(s) -----> 2KCl(s) + 3O2(g)

 

5. Most metallic oxides are stable, but a

    few decompose when heated.

 

      2HgO(s) ---> 2Hg(s) + O2(g)

 
6. Some compounds cannot be

    decomposed by heat, but can be

    decomposed into their elements by

    electricity (electrolysis).
 

     2NaCl(l) ----2Na(s) + Cl2(g)
 

General classes of synthesis reactions.

 
1. Two or more elements combine to
     form a compound.
 

                   Fe(s) + S(l) ---> FeS(s)

 

2. An acid anhydride, nonmetallic

    oxide, combines with water to give an

    acid.
  

       SO2(g) + H2O(l) ----> H2SO4(aq)

 
 
 

3. A basic anhydride, metallic oxide,

    combines with water to form a base.

        

     Na2O(cr) + H2O(l) ---> 2NaOH(aq)

 

4. A basic oxide combines with a

    nonmetallic oxide to form a salt (ionic

    compound).
      

         CO2(g) + Na2O(s) ---> Na2CO3(s)

General types of single displacement

 

1. An active metal will repalce the metallic ion in a compound of a less active

    metal.

Fe(s) + Cu(NO3)2(aq)---> Fe(NO3)2(aq) + Cu(s)

 
2. Some active metals (Groups IA & IIA)

    will react with water to give a metallic

     hydroxide and hydrogen gas.

 

   Ca(s) + 2H2O(l) ---> Ca(OH)2(aq) + H2(g)

 
3. Active metals will replace hydrogen in

    acids to give a salt and hydrogen gas.

 

      Zn(cr) + 2HCl(aq) --> ZnCl2(aq) + H2(g)

 
4. An active nonmetal will replace a less

    active nonmetal ( Group VIIA - the

    Halogens).
 

    Cl2(g) + 2NaBr(aq) ---> 2NaCl(aq) + Br2(g)
 

General procedure for solving stoichiometry problems:

 
1. Wrtie the balanced equation for the reaction.
 
2. Convert given material to moles.
 

3. Determine the mole ratio from the coefficients of the balanced equation and    

    convert from moles of given material to moles of required material.

 
4. Convert to desired units.