CHEM 120: Introduction to Inorganic Chemistry

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CHEM 120 Introduction to Inorganic Chemistry

• Instructor Upali Siriwardane (Ph.D., Ohio State
  University)
• CTH 311, Tele 257-4941, e-mail
  upali_at_chem.latech.edu
• Office hours 1000 to 1200 Tu Th 800-900
  and 1100-1200 M,W, F

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Chapters Covered and Test dates

• Tests will be given in regular class periods 
  from  930-1045 a.m. on the following days
• September 22,     2004 (Test 1) Chapters 1 2
• October 8,         2004(Test 2)  Chapters  3,
  4
• October 22,         2004 (Test 3) Chapter  5 6
  
• November 12,     2004 (Test 4) Chapter  7 8
• November 15 Brief survey of chpater 9-10
• November 17,      2004 MAKE-UP Comprehensive
  test (Covers all chapters 1-8)
• Grading
• ( Test 1 Test 2 Test3 Test4 Test5)
  x.70 Homework quiz average x 0.30 Final
  Average
•                               5

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Chapter 8. Chemical and Physical Change Energy,
Rate, and Equilibrium

• Thermodynamics
• 1. Endothermic and exothermic based on heat flow
  between a system and its surroundings.
• 2. Enthalpy (DH), entropy (DS) , and free energy
  (DG)
• 3. Experiments to get thermochemical information
  and fuel values.
• Reaction Rates
• 4. Reaction rate and the role of kinetics in
  chemical and physical change.
• 5. Activation energy and the activated complex
  and effects on reaction rate.
• 6. Predict the affect of, concentration,
  temperature, and catalysis on the rate of a
  chemical reaction.
• 7. Write rate equations for elementary processes.
  
• Chemical Equilibrium
• 8. Equilibrium chemical reactions.
• 9. Equilibrium-constant expressions and
  equilibrium constants.
• 10.LeChatelier's principle for predicting
  equilibrium position.

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Thermodynamics

• Thermodynamics is study of energy, heat and work.
  
• As chemists we are interested in heat changes in
  chemical and physical changes.
• Aim to predict whether a change (both physical
  and chemical ) will occur spontaneously when left
  to itself--do we have to do anything other than
  mix the reactants together to make it occur?

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Chemical reaction and energy

• In a chemical reaction heat is released or
  absorbed by our system to or from the
  surroundings. In a chemical reaction convert
  energy in bonds into heat energy (and vice
  versa).
• We can measure the energy changes in these
  processes.

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Universe

• System
• Surroundings
• Universe

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First law of thermodynamics

• Law of Conservation of Energy

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Thermochemistry

• Heat changes during chemical reactions
• Thermochemical equation. eg.
• H2 (g) O2 (g) ---gt 2H2O(l) DH - 256 kJ
• DH is called the enthalpy of reaction.
• if DH is reaction is called endothermic
• if DH is - reaction is called exothermic

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Exothermic processes

• In an exothermic process

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Exothermic rxn

• Reactants
• 
  DE
• products

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Endothermic Reaction

• In an endothermic reaction

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Endothermic process

• 
  products
• 
  DE
• reactants

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Enthalpy, H

• Define enthalpy (H)
• Normally talk about a change in enthalpy (DH,
  DHo)

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Enthalpy, H

• If DH is negative

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Exothermic Reaction

• Reactants
• 
  DH
• products

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Enthalpy, H

• If DH is positive

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Endothermic process

• 
  products
• 
  DH
• reactants

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Are these exo- or endothermic?

• When solid NaOH is dissolved in water the soln
  gets hotter.
• S(s) O2(g) g SO2(g) DHo -71kcal
• N2(g) 2O2(g 16.2kcal g 2NO2(g)

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Spontaneous Processes

• A spontaneous reactions is one that occurs
  without us having to do anything to it (once it
  has started). (No external energy input)
• Name some spontaneous processes.
• Note that one direction is spontaneous, the
  reverse is not.

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• We want to come up with a way of predicting
  whether something will be spontaneous.
• It has been observed that many exothermic
  processes are spontaneous.
• Question Does exothermicity guarantee that
  something is spontaneous?

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• Look at H2O(s) g H2O(l) DHorxn1.44kcal

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Entropy, S

• Entropy (S, So) is a measure of the disorder, or
  randomness of a system.
• The greater

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Entropy info
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• For a rxn D Srxn sum of entropy of all the
  products minus the sum of the entropy of all the
  reactants.

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Get entropy increases

• When go from a system of

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Is there an entropy increase when

• A log burns in a fireplace
• Water vapor condenses on a cold surface
• A solid metal melts
• Water boils

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Does a positive entropy change insure a
spontaneous process

• Look at H2O(l) g H2O(s)
• DSlt0 but

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Second law of thermodynamics

• The entropy of the universe increases in a
  spontaneous process and is equal to zero in a
  system at equilibrium.
• Not always easy to calculate the entropy change
  of the universe.
• Define a new quantity, G, Gibbs free energy. G
  refers to the system we are studying.

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• DG D H - T D S
• This equation combines the exothermicity and
  positive entropy criteria.
• Criterion for spontaneity
• If DG lt 0
• If DG gt 0
• If DG 0

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How does DG D H - T D Swork?

• Look at H2O(l) g H2O(g) water boiling
• DH 10.6kcal and DS 0.0284kcal
• DG DH -T DS
• at 50oC DG
• at 100oC DG
• at 120oC DG

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DG D H - T D S

• An exothermic rxn that has a positive entropy
  change is
• An endothermic rxn that has a negative entropy
  change is

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DG D H - T D S

• An exothermic rxn that has a negative entropy
  change is spontaneous
• An endothermic rxn that has a positive entropy
  change is spontaneous

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Calorimetry how to measure heat changes in
reactions

• Measure heat change (temp inc or dec) in a
  quantity of water or solution that is in contact
  with the reaction of interest and is isolated
  from the surroundings.

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Constant Volume (Bomb) Calorimeter

• Used for combustion reactions, etc.
• CxHy (xy/2)O2 g xCO2 yH2O
• Measure ?t of H2O bath, calculate ?Hsurr,
  determine ?Hrxn

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To calculate the amt of heat (Q) absorbed or
released

• Q m x SH x ?T
• Q heat change ?T temperature change
  Tfinal-Tinitial
• Add heat, temp inc remove heat temp dec
• Specific heat (SH) the amount of heat (cal)
  required to raise the temperature of one g of a
  substance by 1C

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• Units of SH cal/(g oC) SH of H2O 1.00
  cal/(g oC) of Al 0.21 cal/(g oC)
• Calc the amt of heat liberated (in cal and kcal)
  from 366 g of aluminum when it cools from 77.0oC
  to 12.0oC.
• 10kJ of heat is supplied to 1000g of H2O and to
  1000g of Al. Calc the inc in temp for both.

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• To raise the temp of a mass of water from 25oC to
  50oC requires 7.5 kcal. What is the mass of the
  water?
• A sample of Al weighs 67 g. If 854 cal of heat
  are required to raise the temp of this sample
  from 25oC to 85oC, calculate the specific heat of
  aluminum.

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Fuel value

• Fuel value is amt of energy per gram of food.
• One nutritional Calorie (C) 1 kcal 1000 cal
  1 cal 4.184 J 1 kcal1 Cal 4.184 kJ
• 8.7 A 1.00 g sample of a candy bar (which
  contains a lot of sugar was burned in a bomb
  calorimeter. A 3.0oC temp increase was observed
  for 1.00 x 103 g of water. The entire candy bar
  weighed 2.5 ounces. Calc the fuel value (in
  nutritional Calories) of the sample and the total
  caloric content of the candy bar.

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• 8.8 If the fuel value of 1.00g of a certain
  carbohydrate is 3.00 nutritional Calories, how
  many grams of water must be present in the
  calorimeter to record a 5.00oC change in temp?

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Kinetics (Reaction Rates)

• Thermodynamics tells us whether a reaction should
  occur spontaneously, but does not tell us how
  fast the reaction will occur. Kinetics tells us
• For example, thermodynamics says that diamond
  will spontaneously change into graphite.
• Kinetics tells us

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• Look at H2(g) I2(g) g 2HI(g)
• Lets envision how the rreaction might occur on a
  molecular basis.
• To react
• 1. H2 must collide with I2
• 2. There must be enough energy to break a H-H and
  a I-I bond to initiate reaction
• 3. The molecules must collide in the correct
  geometry.

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• A molecule that is moving has kinetic energy
  faster the motion, the greater the KE. When
  molecules collide some of the KE is changed into
  vibrational energy of the bonds . Sometimes there
  is enough energy gained through collision to
  break a bond and initiate reaction. If the energy
  is not enough to break the bond, the molecules
  bounce of one another with no reaction occurring.
• So there is some minimum collision energy below
  which no reaction occurs.

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Activation energy, Ea
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Factors that affect reaction rate

• I. Structure of reacting species A.
  oppositely charged species often react
  faster than neutral species. B. bond
  strength can influence rxn rate (Ea)
  C. size and shape of molecule can be
  important

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• II. Concentration of reactants if increase the
  conc of the reactants

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• III. Increase the temperature

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• IV. Physical state of reactants reactions in
  solution (liquids) are often very fast. In the
  solid state molecules have limited motion, in
  the gas phase have large distances between
  molecules and not many collisions so these
  reactions may be slower. In the liquid phase the
  molecules (ions) are able to move and are close
  to each other.

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• V. Add a catalyst. A catalyst speeds up the rate
  of reaction by
• . The catalyst generally works by giving a
  different pathway (of lower energy) for the
  reaction to occur. A catalyst is not used up
  (consumed) in a reaction. The catalyst appears to
  be unchanged at the end.
• Biological catalysts enzymes

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Mathematical representation of reaction rate

• Object is to develop a mathematical relationship
  btn rate and concs of various species know as
  rate equation (law)
• for A B C g products we say the
• or

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• For a rate law where rate k
  AxByCz
• say

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Some sample rate laws

• H2(g) I2(g) g 2HI(g) rate kH2I2
• H2(g) I2(g) g 2HI(g) high P, Au catalyst rate
  k
• 2N2O5 g 4NO2 O2 rate kN2O5

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• 2NO(g) O2(g) g 2NO2(g) rate kNO2O2
• CHCl3(g) Cl2(g) g CCl4(g) HCl(g) rate
  kCHCl3Cl21/2

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• A reaction is found to be second order in A and
  third order in B. Write the rate equation.
• The rate equation and the rate constant have to
  be determined experimentally.

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• Write the general form of the rate equation for
• CH4(g) O2(g) g CO2(g) 2H2O(g)
• 2NO2(g) g 2NO(g) O2(g)