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TRANSCRIPT
Chapter 6
Solutions
Denniston Topping Caret
6th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
6.1 Properties of Solutions
• Solution - homogeneous mixture• Solute - the substance in the mixture present
in lesser quantity• Solvent - the substance present in the largest
quantity• Aqueous solution - solution where the
solvent is water• Solutions can be liquids as well as solids and
gases
Examples of Solutions
• Air - oxygen and several trace gases are dissolved in the gaseous solvent, nitrogen
• Alloys - brass and other homogeneous metal mixtures in the solid state
• Focus on liquid solutions as many important chemical reactions take place in liquid solutions
6.1
Prop
ertie
s of S
olut
ions
• Clear, transparent, no visible particles• May have color• Electrolytes are formed from solutes that are
soluble ionic compounds• Nonelectrolytes do not dissociate
• Volumes of solute and solvent are not additive– 1 L ethanol + 1 L water does not give 2 L of solution
6.1
Prop
ertie
s of S
olut
ions
)(-Cl )(Na )NaCl( OH2 aqaqs +→ +
General Properties of Liquid Solutions
6.1
Prop
ertie
s of S
olut
ions Solutions and Colloids
• Colloidal suspension - contains solute particles which are not uniformly distributed– Due to larger size of particles (1nm - 200 nm)– Appears identical to solution from the
naked eye– Smaller than 1 nm, have solution– Larger than 1 nm, have a precipitate
6.1
Prop
ertie
s of S
olut
ions Tyndall Effect
• Tyndall effect - the ability of a colloidal suspension to scatter light– See a haze when shining light through the mixture
– Solutions: light passes right through without scattering
Colloidal suspension Solution-light as haze, scatters light -no haze
6.1
Prop
ertie
s of S
olut
ions Degree of Solubility
• Solubility - how much of a particular solute can dissolve in a certain solvent at a specified temperature
• Factors which affect solubility:1 Polarity of solute and solvent
• The more different they are, the lower the solubility2 Temperature
• Increase in temperature usually increases solubility3 Pressure
• Usually has no effect• If solubility is of gas in liquid, directly proportional
to applied pressure
6.1
Prop
ertie
s of S
olut
ions Saturation
• Saturated solution - a solution that contains all the solute that can be dissolved at a particular temperature
• Supersaturated solution - contains more solute than can be dissolved at the current temperature
• How is this done? • Heat solvent, saturate it with solute then cool slowly• Sometimes the excess will precipitate out• If it doesn’t precipitate, the solution will be
supersaturated
6.1
Prop
ertie
s of S
olut
ions Solubility and Equilibrium
• If excess solute is added to a solvent, some dissolves
• At first, rate of dissolution is large
• Later, reverse reaction – precipitation – occurs more quickly
• When equilibrium is reached the rates of dissolution and precipitation are equal, there is some dissolved and some undissolved solute
• A saturated solution is an example of a dynamic equilibrium
6.1
Prop
ertie
s of S
olut
ions Solubility of Gases: Henry’s Law
• Henry’s law – the number of moles of a gas dissolved in a liquid at a given temperature is proportional to the partial pressure of the gas above the liquid
• Gas solubility in a liquid is directly proportional to the pressure of the gas in the atmosphere in contact with the liquid
• Gases are most soluble at low temperatures• Solubility decreases significantly at higher
temperatures– Carbonated beverages – CO2 solubility less when warm– Respiration – facilitates O2 and CO2 exchange in lungs
6.2 Concentration Based on Mass
• Concentration - amount of solute dissolved in a given amount of solution
• Concentration of a solution has an effect on – Physical properties
• Melting and boiling points– Chemical properties
• Solution reactivity
• Amount of solute = mass of solute in grams• Amount of solution = volume in milliliters
• Express concentration as a percentage by multiplying ratio by 100% = weight/volume percent or % (W/V)
%100solution of smilliliter
solute of gramsVW% ×=
solution ofamount solute ofamount ion concentrat =
Weight/Volume Percent6.
2 C
once
ntra
tion
Bas
ed o
n M
ass
6.2
Con
cent
ratio
n B
ased
on
Mas
sCalculating Weight/Volume
PercentCalculate the percent composition or % (W/V) of 2.00 x 102 mL containing 20.0 g sodium chloride
20.0 g NaCl, mass of solute
2.00 x 102 mL, total volume of solution
% (W/V) = 20.0g NaCl / 2.00 x 102 mL x 100%
= 10.0% (W/V) sodium chloride
6.2
Con
cent
ratio
n B
ased
on
Mas
sCalculate Weight of Solute from
Weight/Volume PercentCalculate the number of grams of glucose in 7.50 x 102 mL of a 15.0% solution
15.0% (W/V) = Xg glucose/7.50 x 102 mL x 100%
Xg glucose x 100% = (15.0% W/V)(7.50 x 102 mL) Xg glucose = 113 g glucose
%100solution of smillilitersolute of grams
VW% ×=
%100solutions grams
solute gramsWW% ×=
6.2
Con
cent
ratio
n B
ased
on
Mas
sWeight/Weight Percent
• Weight/weight percent is most useful for solutions of 2 solids whose masses are easily obtained
• Calculate % (W/W) of platinum in gold ring with 14.00 g Au and 4.500 g Pt[4.500 g Pt / (4.500 g Pt + 14.00 g Au)] x 100%= 4.500 g / 18.50 g x 100% = 24.32% Pt
6.2
Con
cent
ratio
n B
ased
on
Mas
sParts Per Thousand (ppt) and
Parts Per Million (ppm)
• As percentage is the number of parts of solute in 100 parts of solution, ppt and ppm change the calculation only by orders of magnitude– ppt = g solute / g solution x 103 ppt– ppm = g solute / g solution x 106 ppt– ppt and ppm are most often used for very
dilute solutions
6.3 Concentration of Solutions: Moles and Equivalents
• Chemical equations represent the relative number of moles of reactants producing products
• Many chemical reactions occur in solution where it is most useful to represent concentrations on a molar basis
• The most common mole-based concentration unit is molarity
• Molarity– Symbolized M– Defined as the number of moles of solute per
liter of solution
Molarity6.
3 M
oles
and
Equ
ival
ents
solution Lsolute moles=M
• Calculate the molarity of 2.0 L of solution containing 5.0 mol NaOH
• Use the equation
• Substitute into the equation:MNaOH = 5.0 mol solute
2.0 L solution= 2.5 M
Calculating Molarity from Moles6.
3 M
oles
and
Equ
ival
ents
solution Lsolute moles=M
• If 5.00 g glucose are dissolved in 1.00 x 102 mL of solution, calculate molarity, M, of the glucose solution
• Convert from g glucose to moles glucose– Molar mass of glucose = 1.80 x 102 g/mol5.00 g x 1 mol / 1.80 x 102 g = 2.78 x 10-2 mol glucose– Convert volume from mL to L1.00 x 102 mL x 1 L / 103 mL = 1.00 x 10-1 L
• Substitute into the equation:
Mglucose = 2.78 x 10-2 mol glucose1.00 x 10-1 L solution
= 2.78 x 10-1 M
Calculating Molarity From Mass6.
3 M
oles
and
Equ
ival
ents
solution Lsolute moles=M
solution Lsolute moles=M6.
3 M
oles
and
Equ
ival
ents Dilution
Dilution is required to prepare a less concentrated solution from a more concentrated one
– M1 = molarity of solution before dilution– M2 = molarity of solution after dilution– V1 = volume of solution before dilution– V2 = volume of solution after dilution
moles solute = (M)(L solution)
• In a dilution will the number of moles of solute change? – No, only fewer per unit
volume
• So,
• Knowing any three terms permits calculation of the fourth
M1V1 = M2V2
6.3
Mol
es a
nd E
quiv
alen
ts Dilution
• Calculate the molarity of a solution made by diluting 0.050 L of 0.10 M HCl solution to a volume of 1.0 L– M1 = 0.10 M molarity of solution before dilution– M2 = X M molarity of solution after dilution– V1 = 0.050 L volume of solution before dilution– V2 = 1.0 L volume of solution after dilution
• Use dilution expression
• X M = (0.10 M) (0.050 L) / (1.0 L)
0.0050 M HCl OR 5.0 x 10-3 M HCl
M1V1 = M2V2
6.3
Mol
es a
nd E
quiv
alen
ts Calculating Molarity After Dilution
6.3
Mol
es a
nd E
quiv
alen
ts Representation of Concentration of Ions in Solution
Two common ways of expressing concentration of ions in solution:
1. Moles per liter (molarity)• Molarity emphasizes the number of
individual ions
2. Equivalents per liter (eq/L)• Emphasis on charge
ionon charges ofnumber (g)ion of massmolar ionan of equivalent One =6.
3 M
oles
and
Equ
ival
ents Comparison of Molarity and
Equivalents1 M Na3PO4
• What would the concentration of PO43- ions be?
• 1 M
• Equivalent is defined by the charge
• One Equivalent of an ion is the number of grams of the ion corresponding to Avogadro’s number of electrical charges
6.3
Mol
es a
nd E
quiv
alen
ts Molarity vs. Equivalents – 1 M Na3PO4
• 1 mol Na+ = 1 equivalent Na+
• 1 mol PO43- = 3 equivalents PO4
3-
• Equivalents of Na+?– 3 mol Na+ = 3 equivalents of Na+
• Equivalents of PO43-?
– 1 mol PO43- = 3 equivalents of PO4
3-
6.3
Mol
es a
nd E
quiv
alen
ts Calculating Ion Concentration• Calculate eq/L of phosphate ion, PO4
3- in a solution with 5.0 x 10-3 M phosphate
• Need to use two conversion factors:– mol PO4
3- mol charge
– mol charge eq PO43
5.0 x 10-3 mol PO43- x 3 mol charge x 1 eq
1 L 1 mol PO43- 1mol charge
• 1.5 x 10-2 eq PO43- /L
6.4 Concentration-Dependent Solution Properties
• Colligative properties - properties of solutions that depend on the concentrationof the solute particles, rather than the identity of the solute
• Four colligative properties of solutions1. vapor pressure lowering2. boiling point elevation3. freezing point depression4. osmotic pressure
Vapor Pressure of a LiquidConsider Raoult’s law in molecular
terms• Vapor pressure of a solution
results from escape of solvent molecules from liquid to gas phase
• Partial pressure of gas phase solvent molecules increases until equilibrium vapor pressure is reached
• Presence of solute molecules hinders escape of solvent molecules, lowering equilibrium vapor pressure6.
4 C
once
ntra
tion-
Dep
ende
nt
Solu
tion
Prop
ertie
s
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Vapor Pressure Lowering• Raoult’s law - when a nonvolatile solute is added
to a solvent, vapor pressure of the solvent decreases in proportion to the concentration of the solute
• Solute molecules (red below) serve as a barrier to the escape of solvent molecules resulting in a decrease in the vapor pressure
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Freezing Point Depression and Boiling Point Elevation
• Freezing point depression may be explained considering the equilibrium between solid and liquid states– Solute molecules interfere with the rate at which
liquid water molecules associate to form the solid state
• Boiling point elevation can be explained considering the definition as the temperature at which vapor pressure of the liquid equals the atmospheric pressure– If a solute is present, then the increase in boiling
temperature is necessary to raise the vapor pressure to atmospheric temperature
• Freezing point depression (DTf) - is proportional to the number of solute particles – Solute particles, not just solute
• How does an electrolyte behave?– Dissociate into ions
• An equal concentration of NaCl will affect the freezing point twice as much as glucose (a nonelectrolyte)
• Each solvent has a unique freezing point depression constant or proportionality factor
DTf=kf m6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Freezing Point Depression
• Boiling point elevation (DTb) - is proportional to the number of solute particles
• An electrolyte will affect boiling point to a greater degree than a nonelectrolyte of the same concentration
• Each solvent has a unique boiling point elevation constant
DTb=kb m
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Boiling Point Elevation
Molality• Solute concentration is expressed in
mole-based units– Number of particles is critical, not the mass
of solute
• Molality (m) = moles of solute per kg of solvent– The denominator is in kg solvent, not in kg
solution
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
solvent kgsolute moles Molality =
Osmotic Pressure• Some types of membranes appear impervious
to matter, but actually have a network of small holes called pores
• These pores may be large enough to permit small solvent molecules to move from one side of the membrane to the other
• Solute molecules cannot cross the membrane as they are too large
• Semipermeable membrane - allows solvent but not solute to diffuse from one side to another
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Osmotic Pressure
• Osmosis - the movement of solvent from a dilute solution to a more concentratedsolution through a semipermeable membrane
• Requires pressure to stop this flow
• Osmotic pressure () - the amount of pressure required to stop the flow across a semipermeable membrane
• Osmolarity - the molarity of particles in solution– Osmol, used for osmotic pressure
calculation
=MRT
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Osmotic Pressure
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Calculating OsmolarityCalculate the osmolarity of 5.0 x 10-3 M Na3PO4
Na3PO4 is an ionic compound forming electrolytes
1 mol Na3PO4 yields 4 product ions[5.0 x 10-3 mol Na3PO4 / L] x 4 mol particles
1 mol Na3PO4
= 2.0 x 10-2 mol particles / L2.0 x 10-2 mol particles / L = 2.0 x 10–2 osmol
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Calculating Osmotic PressureCalculate the osmotic pressure of a 5.0 x 10-2 Msolution of NaCl at 25°C (298 K)
Definition of osmotic pressure, =MRTM is osmolarity =[5.0 x 10-2 mol NaCl / L] x 2 mol particles
1 mol NaCl = 1.0 x 10-1 mol particles / L
Substitute into osmotic pressure equation: = 1.0 x 10-1 mol particles x 0.0821 L x atm x 298K
Liter K x mol = 2.4 atm
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Tonicity and the Cell• Living cells contain aqueous solution and these cells are
also surrounded by aqueous solution• Cell function requires maintenance of the same osmotic
pressure inside and outside the cell• Solute concentration of fluid surrounding cells higher
than inside results in a hypertonic solution causing water to flow into the surroundings, causing collapse = crenation
• Solute concentration of fluid surrounding cells too low, results in a hypotonic solution causing water to flow into the cell, causing rupture = hemolysis
• Isotonic solutions have identical osmotic pressures and no osmotic pressure difference across the cell membrane
IsotonicHemolysisCrenation
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
Tonicity and the Cell
Pickling Cucumber in Hypertonic Brine Due to Osmosis
6.4
Con
cent
ratio
n-D
epen
dent
So
lutio
n Pr
oper
ties
6.5 Water as a Solvent• Water is often referred to as the “universal
solvent”• Excellent solvent for polar molecules • Most abundant liquid on earth• 60% of the human body is water
– transports ions, nutrients, and waste into and out of cells
– solvent for biochemical reactions in cells and digestive tract
– reactant or product in some biochemical processes
6.6 Electrolytes in Body FluidsCATIONS IN THE BLOOD and CELLS• Na+ and K+ two most important cations
Na+
K+
Blood Cells
135 meq/L 3.5 – 5.0 meq/L
10 meq/L 125 meq/L
• Active transport - the transporting of Na+ and K+ ions across the cell membrane
• Cellular energy must be expended to make concentration of ions different on each side of the cell membrane
• This is accomplished via large protein molecules embedded in cell membranes
6.6
Elec
troly
tes i
n B
ody
Flui
ds
• Danger to the body occurs when Na+
and K+ both in blood and in cells becomes too high or low
- Na+ too low: - Decrease of urine output
- Dry mouth
- Flushed skin
- Fever
- Na+ too high:- Confusion, stupor, or coma6.
6 El
ectro
lyte
s in
Bod
y Fl
uids
- K+ too high:- Death by heart failure
- K+ too low:- Death by heart failure
ANIONS IN THE BLOOD
• Cl-
- acid/base balance- maintenance of osmotic pressure- oxygen transport by hemoglobin6.
6 El
ectro
lyte
s in
Bod
y Fl
uids
• HCO3-
- Form in which most waste CO2 is carried out of the body
PROTEINS IN THE BLOOD• Blood clotting factors
• Antibodies
• Albumins (carriers of nonpolar substances which cannot dissolve in water)
• Proteins are transported as a colloidal suspension
• The blood also transports nutrients and waste products6.
6 El
ectro
lyte
s in
Bod
y Fl
uids
Chapter 7
Energy, Rate, and Equilibrium
Denniston Topping Caret
6th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
7.1 Thermodynamics• Thermodynamics – the study of energy,
work, and heat– applied to chemical change
• Calculate the quantity of heat obtained from combustions of one gallon of fuel oil
– applied to physical change• Determine the energy released by boiling water
• The laws of thermodynamics help us to understand why some chemical reactions occur and others do not
Basic Concepts – from Kinetic Molecular Theory– Molecules and atoms in a reaction mixture are
in constant, random motion
– Molecules and atoms frequently collide with each other
– Only some collisions, those with sufficient energy, will break bonds in molecules
– When reactant bonds are broken, new bonds may be formed and products result
The Chemical Reaction and Energy7.
1 Th
erm
odyn
amic
s
7.1
Ther
mod
ynam
ics
Change in Energy and Surroundings
• Absolute value for energy stored in a chemical system cannot be measured
• Can measure the change in energy during these chemical changes
• System – contains the process under study
• Surroundings – the rest of the universe
7.1
Ther
mod
ynam
ics
Changes in the System• Energy can be lost from the system to the
surroundings
• Energy may be gained by the system at the expense of the surroundings– This energy change is usually in the form of heat
– This change can be measured
• First Law of Thermodynamics – energy of the universe is constant
• This law is also called the Law of Conservation of Energy
• Where does the reaction energy come from that is released and where does the energy go when it is absorbed?7.1
Ther
mod
ynam
ics
Law of Conservation of Energy
Changes in Chemical Energy
• Consider the reaction converting AB and CD to AD and CB
• Each chemical bond is stored chemical energy
• If a reaction will occur – Bonds must break– Breaking bonds requires energy7.
1 Th
erm
odyn
amic
s
A-B + C-D A-D + C-B
If the energy required to break the bonds is less than the energy released when the bonds are formed, there is a net release of energy
– This is called an Exothermic reaction– Energy is a product in this reaction
These bonds must be broken in the
reaction, requiringenergy
These bonds are formed, releasing
energy
7.1
Ther
mod
ynam
ics
A-B + C-D A-D + C-B
Exothermic Reactions
• If the energy required to break the bonds is larger than the energy released when the bonds are formed, there will need to be an external supply of energy– This is called an Endothermic reaction
These bonds must be broken in the
reaction, requiringenergy
These bonds are formed, releasing
energy
7.1
Ther
mod
ynam
ics
A-B + C-D A-D + C-B
Endothermic Reactions
Endothermic ReactionDecomposition
22 kcal + 2NH3(g) N2(g) + 3H2(g)
7.1
Ther
mod
ynam
ics
Exothermic ReactionCombustion
CH4(g) + 2O2(g)CO2(g) + 2H2O(g) + 211 kcal
7.1
Ther
mod
ynam
ics
Enthalpy• Enthalpy – represents heat energy
• Change in Enthalpy (DHo) – energy difference between the products and reactants of a chemical reaction
• Energy released, exothermic reaction, enthalpy change is negative – In the combustion of CH4, DHo = -211 kcal
• Energy absorbed, endothermic, enthalpy change is positive.– In the decomposition of NH3, DHo=+22 kcal
7.1
Ther
mod
ynam
ics
Spontaneous and Nonspontaneous Reactions
• Spontaneous reaction - occurs without any external energy input
• Most, but not all, exothermic reactions are spontaneous
• Thermodynamics is used to help predict if a reaction will occur
• Another factor is needed, Entropy
DSo is positive
DSo is negative
7.1
Ther
mod
ynam
ics
Spontaneous and Nonspontaneous Reactions
7.1
Ther
mod
ynam
ics
Spontaneous and Nonspontaneous Reactions
Are the following processes exothermic or endothermic?
– Fuel oil is burned in a furnace
– C6H12O6(s) 2C2H5OH(l) + 2CO2(g)
DH=-16 kcal– N2O5(g) + H2O(l) 2HNO3(l) + 18.3 kcal
7.1
Ther
mod
ynam
ics
Entropy
• The second law of thermodynamics – the universe spontaneously tends toward increasing disorder or randomness
• Entropy (So) – a measure of the randomness of a chemical system• High entropy – highly disordered system, the
absence of a regular, repeating pattern
• Low entropy – well organized system such as a crystalline structure
• No such thing as negative entropy
DSo of a reaction = So(products) -So(reactants)
7.1
Ther
mod
ynam
ics
Entropy of Reactions
• A positive DSo means an increase in disorder for the reaction
• A negative DSo means a decrease in disorder for the reaction
All of these processes have a positive DSo
7.1
Ther
mod
ynam
ics
Processes Having Positive Entropy
Phase change
Melting
Vaporization
Dissolution
7.1
Ther
mod
ynam
ics
Entropy and Reaction Spontaneity
• If exothermic and positive DSo…
SPONTANEOUS
• If endothermic and negative DSo…
NONSPONTANEOUS
• For any other situations, it depends on the relative size of DHo and DSo
Greatest Entropy
• Which substance has the greatest entropy?
– He(g) or Na(s)
– H2O(l) or H2O(g)
7.1
Ther
mod
ynam
ics
• Free energy (DGo) – represents the combined contribution of the enthalpy and entropy values for a chemical reaction
• Free energy predicts spontaneity of chemical reactions– Negative DGo…Always Spontaneous
– Positive DGo…Never Spontaneous
DGo = DHo - TDSo
T in Kelvin
7.1
Ther
mod
ynam
ics
Free Energy
Free Energy and Reaction Spontaneity
• Need to know both DH and DS to predict the sign of DG, making a statement on reaction spontaneity
• Temperature also may determine direction of spontaneity DH +, DS - : DG always +, regardless of TDH -, DS + : DG always -, regardless of TDH +, DS + : DG sign depends on TDH -, DS - : DG sign depends on T7.
1 Th
erm
odyn
amic
s
7.2 Experimental Determination of Energy Change in Reactions
• Calorimetry – the measurement of heat energy changes in a chemical reaction
• Calorimeter – device which measures heat changes in calories
• The change in temperature is used to measure the loss or gain of heat
• Change in temperature of a solution, caused by a chemical reaction, can be used to calculate the gain or loss of heat energy for the reaction– Exothermic reaction – heat released is absorbed
– Endothermic reaction – reactants absorb heat from the solution
• Specific heat (SH) – the number of calories of heat needed to raise the temperature of 1 g of the substance 1 oC
sss SHT ×∆×= mQ
Heat Energy in Reactions7.
2 D
eter
min
atio
n of
Ene
rgy
Cha
nge
in R
eact
ions
• Specific heat of the solution along with the total number of g solution and the temperature change permits calculation of heat released or absorbed during the reaction
• S.H. for water is 1.0 cal/goC
• To determine heat released or absorbed, need:– specific heat– total number of grams of solution– temperature change (increase or decrease)
Heat Energy in Reactions7.
2 D
eter
min
atio
n of
Ene
rgy
Cha
nge
in R
eact
ions
• Q is the product – ms is the mass of solution in the calorimeterDTs is the change in temperature of the solution from
initial to final state– SHs is the specific heat of the solution
• Calculate with this equation
– Units are: calories = gram x ºC x calories/gram - ºC
Calculation of Heat Energy in Reactions
7.2
Det
erm
inat
ion
of E
nerg
y C
hang
e in
Rea
ctio
ns
sss SHT ×∆×= mQ
7.2
Det
erm
inat
ion
of E
nerg
y C
hang
e in
Rea
ctio
nsCalculating Energy Involved in
Calorimeter ReactionsIf 0.10 mol HCl is mixed with 0.10 mol KOH in a “coffee cup” calorimeter, the temperature of 1.50 x 102 g of the solution increases from 25.0oC to 29.4oC. If the specific heat of the solution is 1.00 cal/goC, calculate the quantity of energy evolved in the reactionDTs = 29.4oC - 25.0oC = 4.4oC
Q = ms x DTs x SHs
= 1.50 x 102 g solution x 4.4oC x 1.00 cal/goC
= 6.6 x 102 cal
7.2
Det
erm
inat
ion
of E
nerg
y C
hang
e in
Rea
ctio
nsCalculating Energy Involved in
Calorimeter ReactionsIs the reaction endothermic or exothermic
– 0.66 kcal of heat energy was released to the surroundings, the solution
– The reaction is exothermic
What would be the energy evolved for each mole of HCl reacted?– 0.10 mol HCl used in the original reaction – [6.6 x 102 cal / 0.10 mol HCl] x 10 = 6.6 kcal
7.2
Det
erm
inat
ion
of E
nerg
y C
hang
e in
Rea
ctio
nsBomb Calorimeter and
Measurement of Calories in FoodsNutritional Calorie
(large “C” Calorie) = – 1 kilocalorie (1kcal)
– 1000 calories• the fuel value of
food• Bomb Calorimeter
is used to measure nutritional Calories
7.2
Det
erm
inat
ion
of E
nerg
y C
hang
e in
Rea
ctio
nsCalculating the Fuel Value of
Foods1 g of glucose was burned in a bomb calorimeter. 1.25 x 103 g H2O was warmed from 24.5oC to 31.5oC. Calculate the fuel value of the glucose (in Kcal/g).DTs = 31.5oC - 24.5oC = 6.1oC
Surroundings of calorimeter is water with specific heat capacity = 1.00 cal/g H2O oCFuel Value =
= 1.25 x 103 g H2O x 6.1oC x 1.00 cal/g H2O oC= 7.6 x 103 cal
7.6 x 103 cal x 1 Calorie / 103 cal = 7.6 nutritional Calories
sss SHT ×∆×= mQ
7.3 Kinetics
• Thermodynamics determines if a reaction will occur spontaneously but tells us nothing about the amount of time the reaction will take
• Kinetics – the study of the rate (or speed) of chemical reactions– Also supplies an indication of the mechanism –
step-by-step description of how reactants become products
Kinetic Information• Kinetic information represents changes over
time, seen here:– disappearance of reactant, A– appearance of product, B
7.3
Kin
etic
s
7.3
Kin
etic
sAlternative Presentation of
Kinetic DataRather than the graph shown before, this figure demonstrates the change from purple reactant to green product over time from the molecular perspective
7.3
Kin
etic
sKinetic Data Assessed by Color
Change•Change in color over time can be used to monitor the progress of a chemical reaction•The rate of color change can aid in calculating the rate of the chemical reaction
7.3
Kin
etic
sThe Chemical Reaction
• C-H and O=O bonds must be broken • C=O and O-H bonds must be formed
• Energy is required to break the bonds– This energy comes from the collision of the molecules
– If sufficient energy available, bonds break and atoms recombine in a lower energy arrangement
– Effective collision is one that produces product molecules
– leads to a chemical reaction
CH4(g) + 2O2(g) CO2(g) +2H2O(g) + 211 kcal
7.3
Kin
etic
sActivation Energy and the
Activated Complex• Activation energy – the minimum amount of
energy required to initiate a chemical reaction
• Picture a chemical reaction in terms of changes in potential energy occurring during the reaction
– Activated complex – an extremely unstable, short-lived intermediate complex
– Formation of this activated complex requires energy (Ea) to overcome the energy barrier to start the reaction
• Reaction proceeds from reactants to products via the activated complex
• Activated complex – can’t be isolated from the reaction mixture
• Activation energy (Ea) is the difference between the energy of the reactants and that of the activated complex
7.3
Kin
etic
sActivation Energy and the
Activated Complex
•To be an Exothermic reaction requires a net release of energy (DHo)
7.3
Kin
etic
sActivation Energy in the Endothermic Reaction
• This figure diagrams an endothermic reaction
• Reaction takes place slowly due to the large activation energy required
• The energy of the products is greater than that of the reactants
7.3
Kin
etic
sFactors That Affect Reaction
Rate
1. Structure of the reacting species
2. Concentration of reactants
3. Temperature of reactants
4. Physical state of reactants
5. Presence of a catalyst
7.3
Kin
etic
sStructure of Reacting Species
• Oppositely charged species react more rapidly• Dissociated ions in solution whose bonds are already
broken have a very low activation energy• Ions with the same charge do not react• Bond strength plays a role
• Covalent molecules bonds must be broken with the activation energy before new bonds can be formed
• Magnitude of the activation energy is related to bond strength
• Size and shape influence the rate• Large molecules may obstruct the reactive part of the
molecule• Only molecular collisions with correct orientation lead to
product formation
7.3
Kin
etic
sThe Concentration of Reactants
• Rate is related to the concentration of one or more of the reacting substances
• Rate will generally increase as concentration increases– Higher concentration means more reactant
molecules per unit volume
– More reactant molecules means more collisions per unit time
7.3
Kin
etic
sThe Temperature of Reactants
• Rate increases as the temperature increases– Increased temperature relates directly to
increased average kinetic energy
– Greater kinetic energy increases the speed of particles
– Faster particles increases likelihood of collision
• Higher Kinetic Energy means a higher percentage of these collisions will result in product formation
7.3
Kin
etic
sThe Physical State of Reactants
Reactions occur when reactants can collide frequently with sufficient energy to react•Solid state: atoms, ions, compounds are close together but restricted in motion
•Gaseous state: particles are free to move but often are far apart causing collisions to be relatively infrequent
•Liquid state: particles are free to move and are in close proximity
•Reactions to be fastest in the liquid state and slowest in the solid state
• Liquid > Gas> Solid
7.3
Kin
etic
sThe Presence of a Catalyst
• Catalyst – a substance that increases the reaction rate– Undergoes no net change– Does not alter the final product of the reaction– Interacts with the reactants to create an alternative
pathway for product production
7.3
Kin
etic
sUse of a Solid Phase Catalyst
Haber Process is a synthesis of ammonia facilitated by a solid phase catalyst
– Diatomic gases bind to the surface– Bonds are weakened– Dissociation of diatomic gases and reformation of NH3– Newly formed NH3 leaves the solid surface with
catalyst unchanged
N2+3H2 2NH3
• Consider a decomposition reaction with the following balanced chemical equation:
• When heated N2O5 decomposes to 2 products – NO2 and O2
• When holding all factors constant, exceptconcentration, rate of reaction is proportional to the concentration
)(O)(NO4)(ON2 2252 ggg +→∆
7.3
Kin
etic
sMathematical Representation of
Reaction Rate
• Reaction rate is proportional to reactant concentration –– Concentration of N2O5 is denoted as [N2O5]– Replace proportionality symbol with = and
proportionality constant k– k is called the rate constant
)(O)(NO4)(ON2 2252 ggg +→∆
]O[Nrate 52∝
]Ok[N rate 52=
7.3
Kin
etic
sMathematical Representation of
Reaction Rate
• For a reaction Aproducts we write the equation:rate = k[A]n
• This is called the rate equation (or rate law)
• The exponent n is the order of the reaction– If n=1, first order– If n=2, second order
– n must be determined experimentally
– This exponent is not the same as the coefficient of the reactant in the balanced equation
7.3
Kin
etic
sRate Equation
• For the equation A + B productsthe rate equation is:
– rate = k[A]n[B]m
• What would be the general form of the rate equation for the reaction:
CH4+2O2CO2+2H2O
– Rate = k[CH4]n[O2]m
• Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product
7.3
Kin
etic
sRate Equation
• Write the form of the rate equation for the oxidation of ethanol (C2H5OH)
• The reaction has been experimentally determined to be first order in ethanol and third order in oxygen– Rate expression involves only the reactants– Concentrations: [C2H5OH][O2]– Raise each to exponent corresponding to its order
rate = k [C2H5OH][O2]3
– Remember that 1 as an exponent is understood and NOT written
– Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product
7.3
Kin
etic
sWriting Rate Equations
7.4 Equilibrium
Rate and Reversibility of Reactions• Equilibrium reactions – chemical reactions
that do not go to completion– Completion – all reactants have been converted
to products– Equilibrium reactions are also called Incomplete
reactions– Seen with both physical and chemical processes
• After no further observable change, measurable quantities of reactants andproducts remain
7.4
Equi
libriu
mPhysical Equilibrium
• Physical equilibria are reversible reactions– Dissolved oxygen in lake water– Stalactite and stalagmite formation– Sugar dissolved in water
• Reversible reaction – a process that can occur in both directions– Use the double arrow symbol
• Dynamic equilibrium – the rate of the forward process in a reversible reaction is exactly balanced by the rate of the reverse process
1. If add 2-3 g of sugar into 100 mL water• All will dissolve with stirring in a short time• No residual solid sugar, sugar dissolved completely
Sugar (s) Sugar (aq)
2. If add 100 g of sugar in 100 mL of water
• Not all of it will dissolve even with much stirring• Over time, you observe no further change in the
amount of dissolved sugar• Appears nothing is happening – Incorrect!
7.4
Equi
libriu
mSugar in Water
2. If add 100 g of sugar in 100 mL of water• Appears nothing is happening – Incorrect!
– Individual sugar molecules are constantly going into and out of solution
– Both happen at the same rate– Over time the amount of sugar dissolved in the
measured volume of water does not change
• An equilibrium situation has been established
• Some molecules dissolve and others return to the solid state – the rate of each process is equal
Sugar(s) Sugar(aq)
7.4
Equi
libriu
mSugar in Water
sugar(s) sugar(aq)
7.4
Equi
libriu
mDynamic Equilibrium
• The double arrow serves as an indicator of– a reversible process– an equilibrium process– the dynamic nature of the process
• Continuous change is taking place without observable change in the amount of sugar in either the solid or the dissolved form
)][sugar()][sugar(Keq s
aqkk
r
f ==
7.4
Equi
libriu
m
• ratef = forward rate rater = reverse rate
• at equilibrium: ratef = rater
• ratef = kf[sugar(s)]
• rater = kr[sugar(aq)]
• kf[sugar(s)]=kr[sugar(aq)]• Equilibrium constant (Keq) – ratio of the two rate
constants
Equilibrium Constant
7.4
Equi
libriu
mChemical Equilibrium
The Reaction of N2 and H2
N2(g) + 3H2(g) 2NH3(g)
• Mix components at elevated temperature
• Some molecules will collide with sufficient energy to break N-N and H-H bonds
• Rearrangement of the atoms will produce the product NH3
7.4
Equi
libriu
mChemical Equilibrium
• Initially the forward reaction is rapid– Reactant concentrations are
high– Product concentration
negligible• Forward reaction rate
decreases with time– Concentrations of reactants
are decreasing– Product concentration
increasing
N2(g) + 3H2(g) 2NH3(g)
Equilibrium occurs when the rate of reactant depletion is equal to the rate of product depletion Rates of forward and reverse reactions are Equal
Chemical Equilibrium
Basic equation divides into 2 parts:
• forward rxn: N2(g) + 3H2(g) 2NH3(g)
• reverse rxn: 2NH3(g)N2(g) + 3H2(g)
• ratef = kf[N2]n[H2]m
• rater = kr[NH3]p
• ratef=rater
N2(g) + 3H2(g) 2NH3(g)
m2
n2
p3
eq ]H[][N][NHK ==
r
f
kk7.
4 Eq
uilib
rium
m2
n2
p3
eq ]H[][N][NHK ==
r
f
kk
• The exponents in the rate expression are numerically equal to the coefficients
• Keq is a constant at constant temperature
322
23
eq ]H][[N][NHK =
7.4
Equi
libriu
mChemical Equilibrium
ba
dc
B][[A]D][C][Keq =
aA + bB cC + dD
7.4
Equi
libriu
mThe Generalized Equilibrium-
Constant Expression for a Chemical Reaction
• A and B are reactants• C and D are products• a, b, c, and d are the coefficients of the
balanced equation
7.4
Equi
libriu
mWriting Equilibrium-Constant
Expressions• Equilibrium constant expressions can only be
written after a correct, balanced chemical equation
• Each chemical reaction has a unique equilibrium constant value at a specified temperature
• The brackets represent molar concentration
• All equilibrium constants are shown as unitless
• Only the concentration of gases and substances in solution are shown• Concentration for pure liquids and solids are not shown
7.4
Equi
libriu
mWriting an Equilibrium-Constant
ExpressionWrite an equilibrium-constant expression for the reversible reaction:
H2(g) + F2(g) 2HF(g)•No solids or liquids are present
– All reactants and products appear in the expression– Numerator term is the product term [HF]2
– Denominator term is the reactants [H2] and [F2] – Each term contains an exponent identical to the
corresponding coefficient in the balanced equationKeq = [HF]2
[H2][F2]
7.4
Equi
libriu
mWriting an Equilibrium-Constant
ExpressionWrite an equilibrium-constant expression for the reversible reaction:
MnO2(s) + 4H+(aq) + 2Cl-(aq) Mn2+(aq) + Cl2(g) + H2O(l)
• MnO2 is a solid
• H2O(l) is a product, but negligible compared to solvent water– Numerator term is the product terms [Mn2+] and [Cl2]– Denominator term is the reactants [H+]4 and [Cl-]2
– Each term contains an exponent identical to the corresponding coefficient in the balanced equation
Keq = [Mn2+] [Cl2][H+]4 [Cl-]2
7.4
Equi
libriu
mInterpreting Equilibrium Constants
• Reversible arrow in chemical equation indicates equilibrium exists
• The numerical value of the equilibrium constant tells us the extent to which reactants have converted to products
1. Keq greater than 1 x 102
• Large value of Keq indicates numerator (product term) >>> denominator (reactant term)
• At equilibrium mostly product present
7.4
Equi
libriu
mInterpreting Equilibrium Constants
2. Keq less than 1 x 10-2
• Small value of Keq indicates numerator (product term) <<< denominator (reactant term)
• At equilibrium mostly reactant present
3. Keq between 1 x 10-2 and 1 x 102
• Equilibrium mixture contains significant concentration of both reactants and products
2NO2(g) N2O4(g)
7.4
Equi
libriu
mCalculating Equilibrium Constants
• A reversible reaction is allowed to proceed until the system reaches equilibrium
• Amount of reactants and products no longer changes
• Analyze reaction mixture to determine molar concentrations of each product and reactant
HI placed in a sealed container and comes to equilibrium; equilibrium reaction is:
2HI(g) H2(g) + I2(g)
•Equilibrium concentrations:– [HI] = 0.54 M– [H2] = 1.72 M– [I2] = 1.72 M
•Substitute concentrations:7.4
Equi
libriu
mCalculating an Equilibrium
Constant
Keq = [H2] [I2][HI]2
Keq= [1.72] [1.72] = 2.96[0.54]2
0.29= 10.1 or 1.0 x 101
2 significant fig res
7.4
Equi
libriu
mLeChateleir’s Principle
• LeChateleir’s Principle – if a stress is placed on a system at equilibrium, the system will respond by altering the equilibrium composition in such a way as to minimize the stress
• If reactants and products are present in a fixed volume and more NH3 is added into the container, the system will be stressed– Stressed = the equilibrium will be disturbed
N2(g) + 3H2(g) 2NH3(g)
7.4
Equi
libriu
mLeChateleir’s Principle
• Adding NH3 to the system causes stress– To relieve stress, remove as much of added
material as possible by converting it to reactants
• Adding N2 or H2 to the system causes stress also– To relieve stress, remove as much of added
material as possible by converting it to product
N2(g) + 3H2(g) 2NH3(g)Equilibrium shifted
Product introduced:Reactant introduced:
N2(g) + 3H2(g) 2NH3(g)
7.4
Equi
libriu
mEffect of Concentration
• Adding or removing either reactants or products at a fixed volume is saying that the concentration is changed
• Removing material decreases concentration
• System will react to this stress to return concentrations to the appropriate ratio
A B C
A: Reaction at equilibriumB: Shift to reactant with more red colorC: Shift to product with loss of red color
• Exothermic reactions: treat heat as a productN2(g) + 3H2(g) 2NH3(g) + 22 kcal
7.4
Equi
libriu
mEffect of Heat
• Addition of heat is treated as increasing the amount of product
• More product shifts equilibrium to the left
– Increases amount of reactants
– Decreases amount of product Heat favors the blue species
while cold favors the pink
• Endothermic Reaction – treat heat as a reactant
39 kcal + 2N2(g) + O2(g) 2NH3(g)
• This reaction shift will shift to the right if heat is added by increasing the temperature
7.4
Equi
libriu
mEffect of Heat
7.4
Equi
libriu
mEffect of Pressure
• Pressure affects the equilibrium only if one or more substances in the reaction are gases
• Relative number of gas moles on reactant and product side must differ
• When pressure goes up…shift to side with less moles of gas
• When pressure goes downs…shifts to side with more moles of gas
N2(g) + 3H2(g) 2NH3(g)
• If increase pressure, which way will the equilibrium shift?– Increased pressure favors decreased volume with more
product (2 moles) formed and less reactant (4 moles)
2HI(g) H2(g) + I2(g)• If increase the pressure in this reaction, which way
will the equilibrium shift?– No shift in equilibrium as both reactant and product
have 2 moles of gas
7.4
Equi
libriu
mEffect of Pressure
7.4
Equi
libriu
mEffect of a Catalyst
• A catalyst has no effect on the equilibrium composition
• It increases the rate of both the forward and reverse reaction to the same extent
• While equilibrium composition and concentration do not change, equilibrium is reached in a shorter time
Chapter 8
Acids and Bases and Oxidation-Reduction
Denniston Topping Caret
6th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
8.1 Acids and Bases
• Acids: Taste sour, dissolve some metals, cause plant dye to change color
• Bases: Taste bitter, are slippery, are corrosive
• Two theories that help us to understand the chemistry of acids and bases1. Arrhenius Theory2. Brønsted-Lowry Theory
8.1
Aci
ds a
nd B
ases
• Acid - a substance, when dissolved in water, dissociates to produce hydrogen ions– Hydrogen ion: H+ also called “protons”
HCl is an acid:HCl(aq) H+(aq) + Cl-(aq)
Arrhenius Theory of Acids and Bases
8.1
Aci
ds a
nd B
ases
Arrhenius Theory of Acids and Bases
• Base - a substance, when dissolved in water, dissociates to produce hydroxide ions
NaOH is a baseNaOH(aq) Na+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Arrhenius Theory of Acids and Bases
• Where does NH3 fit? • When it dissolves in water it has basic
properties but it does not have OH- ions in it
• The next acid-base theory gives us a broader view of acids and bases
8.1
Aci
ds a
nd B
ases
Brønsted-Lowry Theory of Acids and Bases
• Acid - proton donor• Base - proton acceptor
– Notice that acid and base are not defined using water
– When writing the reactions, both accepting and donation are evident
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
What donated the proton? HClIs it an acid or base? Acid
What accepted the proton? H2OIs it an acid or base? Base
Brønsted-Lowry Theory of Acids and Bases
base
8.1
Aci
ds a
nd B
ases
acid
.
base acidNH3(aq) + H2O(l) NH4
+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Brønsted-Lowry Theory of Acids and Bases
Now let us look at NH3 and see why it is a base
Did NH3 donate or accept a proton? Accept
Is it an acid or base? Base
What is water in this reaction? Acid
Acid-Base Properties of Water
• Water possesses both acid and base properties– Amphiprotic – a substance possessing both acid
and base properties– Water is the most commonly used solvent for
both acids and bases– Solute-solvent interactions between water and
both acids and bases promote solubility and dissociation
8.1
Aci
ds a
nd B
ases
8.1
Aci
ds a
nd B
ases
Acid and Base Strength
• Acid and base strength – degree of dissociation– Not a measure of concentration– Strong acids and bases – reaction with water is
virtually 100% (Strong electrolytes)
8.1
Aci
ds a
nd B
ases
Strong Acids and Bases
• Strong Acids:– HCl, HBr, HI Hydrochloric Acid, etc.– HNO3 Nitric Acid– H2SO4 Sulfuric Acid– HClO4 Perchloric Acid
• Strong Bases:– NaOH, KOH, Ba(OH)2
– All metal hydroxides
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)
8.1
Aci
ds a
nd B
ases
Weak Acids
• Weak acids and bases – only a small percent dissociates (Weak electrolytes)
• Weak acid examples:– Acetic acid:
– Carbonic Acid:
• Weak base examples:– Ammonia:
– Pyridine:
– Aniline:C6H5NH2(aq) + H2O(l) C6H5NH3
+(aq) + OH-(aq)
C5H5NH2(aq) + H2O(l) C5H5NH3+(aq) + OH-(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Weak Bases
• The acid base reaction can be written in the general form:
• Notice the reversible arrows• The products are also an acid and base
called the conjugate acid and base
acid baseHA + B A- + HB+
8.1
Aci
ds a
nd B
ases
Conjugate Acids and Bases
acid base
• Conjugate Acid – what the base becomes after it accepts a proton.
• Conjugate Base – what the acid becomes after it donates its proton
• Conjugate Acid-Base Pair – the acid and base on the opposite sides of the equation
base acid
HA + B A- + HB+
8.1
Aci
ds a
nd B
ases
HA + B A- + HB+
8.1
Aci
ds a
nd B
ases
Acid-Base Dissociation
• The reversible arrow isn’t always written– Some acids or bases essentially dissociate 100%– One way arrow is used
• HCl + H2O Cl- + H3O+
– All of the HCl is converted to Cl-
– HCl is called a strong acid – an acid that dissociates 100%
• Weak acid - one which does not dissociate 100%
8.1
Aci
ds a
nd B
ases
Conjugate Acid-Base Pairs
• Which acid is stronger:
HF or HCN? HF
• Which base is stronger:
CN- or H2O? CN -
8.1
Aci
ds a
nd B
ases
Acid-Base PracticeWrite the chemical reaction for the following acids
or bases in water. Identify the conjugate acid base pairs.
1. HF (a weak acid)
2. H2S (a weak acid)
3. HNO3 (a strong acid)
4. CH3NH2 (a weak base)
Note: The degree of dissociation also defines weak and strong bases
• Pure water is virtually 100% molecular• Very small number of molecules dissociate
– Dissociation of acids and bases is often called ionization
• Called autoionization• Very weak electrolyte
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
The Dissociation of Water
• H3O+ is called the hydronium ion• In pure water at room temperature:
– [H3O+] = 1 x 10-7 M– [OH-] = 1 x 10-7 M
• What is the equilibrium expression for:
Remember, liquids are not included in equilibrium expressions
]OH][O[HK -3eq
+=H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
8.1
Aci
ds a
nd B
ases
Hydronium Ion
• This constant is called the ion product for water and has the symbol Kw
• Since [H3O+] = [OH-] = 1.0 x 10-7 M, what is the value for Kw?– 1.0 x 10-14
– It is unitless
]OH][O[HK -3w
+=
8.1
Aci
ds a
nd B
ases
Ion Product of Water
8.2 pH: A Measurement Scale for Acids and Bases
• pH scale – a scale that indicates the acidity or basicity of a solution– Ranges from 0 (very acidic) to 14 (very basic)
• The pH scale is rather similar to the temperature scale assigning relative values of hot and cold
• The pH of a solution is defined as:pH = -log[H3O+]
• Use these observations to develop a concept of pH– if know one concentration, can calculate the
other– if add an acid, [H3O+] and [OH-] – if add a base, [OH-] and [H3O+] – [H3O+] = [OH-] when equal amounts of acid
and base are present• In each of these cases 1 x 10-14 = [H3O+][OH-]8.
2 pH
: A M
easu
rem
ent
Scal
e fo
r Aci
ds a
nd B
ases A Definition of pH
• pH of a solution can be:– Calculated if the concentration of either is
known• [H3O+] • [OH-]
– Approximated using indicator / pH paper that develops a color related to the solution pH
– Measured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH8.
2 pH
: A M
easu
rem
ent
Scal
e fo
r Aci
ds a
nd B
ases Measuring pH
• How do we calculate the pH of a solution when either the hydronium or hydroxide ion concentration is known?
• How do we calculate the hydronium or hydroxide ion concentration when the pH is known?
• Use two facts:
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
es Calculating pH
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
esCalculating pH from Acid
MolarityWhat is the pH of a 1.0 x 10-4 M HCl solution?
– HCl is a strong acid and dissociates in water– If 1 mol HCl is placed in 1 L of aqueous
solution it produces 1 mol [H3O+]– 1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M
= -log [H3O+]= -log [1.0x10-4]= -[-4.00] = 4.00
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
es Calculating [H3O+] from pHWhat is the [H3O+] of a solution with pH = 6.00?
• 4.00 = -log [H3O+] • Multiply both sides of equation by –1
• -4.00 = log [H3O+]• Take the antilog of both sides
• Antilog –4.00 = [H3O+]• Antilog is the exponent of 10
• 1.0 x 10-4 M = [H3O+]
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
es Calculating the pH of a BaseWhat is the pH of a 1.0 x 10-3 M KOH solution?• KOH is a strong base (as are any metal hydroxides)• 1 mol KOH dissolved and dissociated in aqueous
solution produces 1 mol OH-
• 1.0 x 10-3 M KOH solution has [OH-] = 1.0 x 10-3 M
• Solve equation for [H3O+] = 1 x 10-14 / [OH-]• [H3O+] = 1 x 10-14 / 1.0 x 10-3 = 1 x 10-11
• pH = -log [1 x 10-11]= 11.00
1 x 10-14 = [H3O+][OH-]
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
esCalculating pH from Acid
MolarityWhat is the pH of a 2.5 x 10-4 M HNO3 solution?
• We know that as a strong acid HNO3dissociates to produce 2.5 x 10-4 M [H3O+]
• pH = -log [2.5 x 10-4]
• = 3.60
pH = -log[H3O+]
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
es Calculating [OH-] from pH
What is the [OH-] of a solution with pH = 4.95?• First find [H3O+]• 4.95 = -log [H3O+]• [H3O+] = 10-4.95
• [H3O+] = 1.12 x 10-5
• Now solve for [OH-]• [OH-] = 1 x 10-14 / 1.12 x 10-5
= 1.0 x 10-9
pH = -log[H3O+]
1 x 10-14 = [H3O+][OH-]
The pH Scale8.
2 pH
: A M
easu
rem
ent
Scal
e fo
r Aci
ds a
nd B
ases
1.0 x 100 0.001.0 x 10-1 1.001.0 x 10-2 2.001.0 x 10-3 3.001.0 x 10-4 4.001.0 x 10-5 5.001.0 x 10-6 6.001.0 x 10-7 7.00
For a strong acidHCl molarity pH
Mor
e Aci
dic
1.0 x 100 14.001.0 x 10-1 13.001.0 x 10-2 12.001.0 x 10-3 11.001.0 x 10-4 10.001.0 x 10-5 9.001.0 x 10-6 8.001.0 x 10-7 7.00
For a strong baseNaOH molarity pH
Mor
e ba
sic
Each 10 fold change in concentration changes the pH by one unit
8.2
pH: A
Mea
sure
men
t Sc
ale
for A
cids
and
Bas
es
8.3 Reactions Between Acids and Bases
• Neutralization reaction – the reaction of an acid with a base to produce a salt and water
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)Acid Base Salt Water
• Break apart into ions:H+ + Cl- + Na+ + OH-Na+ + Cl- + H2O
• Net ionic equation– Show only the changed components– Omit any ions appearing the same on both sides of
equation = Spectator Ions
H+ + OH- H2O
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases • The net ionic neutralization reaction is more
accurately written:H3O+(aq) + OH-(aq) 2H2O(l)
• This equation applies to any strong acid / strong base neutralization reaction
• An analytical technique to determine the concentration of an acid or base is titration
• Titration involves the addition of measured amount of a standard solution to neutralize the second, unknown solution
• Standard solution – solution of known concentration
Net Ionic Neutralization Reaction
Buret – long glass tube calibrated in mL which contains the standard solution
Flask contains a solution of unknown concentration plus indicator
Indicator – a substance which changes color as pH changes
Standard solution is slowly added until the color changes
The equivalence point is when the moles of H3O+
and OH- are equal
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Acid – Base Titration
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Relationship Between pH and Color in Acid-Base Indicators
Determine the Concentration of a Solution of Hydrochloric Acid• Place a known volume of acid whose concentration
is not known into a flask• Add an indicator, experience guides selection, here
phenol red is good• Known concentration of NaOH is placed in a buret• Drip NaOH into the flask until the indicator
changes color
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Determine the Concentration of a Solution of Hydrochloric Acid
• Indicator changes color – equivalence point is reached – Mol OH- = Mol H3O+ present in the unknown acid
• Volume dispensed from buret is determined• Calculate acid concentration:
– Volume of Hydrochloric Acid: 25.00 mL– Volume of NaOH added: 35.00 mL– Concentration of NaOH: 0.1000 M– Balanced reaction shows that HCl and NaOH react 1:1 8.
3 R
eact
ions
Bet
wee
n A
cids
and
Bas
es
Determine the Concentration of a Solution of Hydrochloric Acid
• 35.00 mL NaOH x 1L NaOH x 0.1000 mol NaOH103 mL NaOH 1L NaOH= 3.500 x 10-3 mol NaOH
• 3.500 x 10-3 mol NaOH x 1 mol HCl1 mol NaOH
= 3.500 x 10-3 mol HCl– this amount of HCl is contained in 25.00 mL
• 3.500 x 10-3 mol HCl x 103 mL HCl25.00 mL HCl 1 L HCl
• = 1.400 x 10-1 mol HCl / L HCl = 0.1400 M HCl8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Determine the Concentration of a Solution of Hydrochloric Acid
• Alternative strategy to solve the acid concentration(Macid)(Vacid) = (Mbase)(Vbase)
• (Macid) = (Mbase)(Vbase)(Vacid)
• (Macid) = (0.1000 M) (35.00 mL)(25.00 mL)
• = 0.1400 M HCl
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Calculating the Concentration of Sodium Hydroxide
Calculate [NaOH] if 25.00 mL of this solution were required to neutralized 45.00 mL of 0.3000 M HCl•Alternative strategy to solve the acid concentration
(Macid)(Vacid) = (Mbase)(Vbase)
•(Mbase) = (Macid)(Vacid)(Vbase)
•(Mbase) = (0.3000 M) (45.00 mL)(25.00 mL)
= 0.5400 M NaOH
• The previous examples have the acid and base at a 1:1 combining ratio– Not all acid-base pairs do this
• Polyprotic substance – donates or accepts more than one proton per formula unit– Hydrochloric acid is monoprotic, producing one H+
ion for each unit of HCl
– Sulfuric acid is diprotic, each unit of H2SO4produces 2 H+ ions
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2 H2O(l)8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Polyprotic Substances
Step 1.
H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq)
Step 2.
HSO4-(aq) + H2O(l) SO4
2-(aq) + H3O+(aq)
• In Step 1 H2SO4 behaves as a strong acid –dissociating completely
• In Step 2 HSO4-( behaves as a weak acid –
reversibly dissociating, note the double arrow8.3
Rea
ctio
ns B
etw
een
Aci
ds a
nd B
ases
Dissociation of Polyprotic Substances
8.4 Acid-Base Buffers
• Buffer solution - solution which resists large changes in pH when either acids or bases are added
• These solutions are frequently prepared in laboratories to maintain optimum conditions for chemical reactions
• Buffers are also used routinely in commercial products to maintain optimum conditions for product behavior
8.4
Aci
d-B
ase
Buf
fers • Buffers act to establish an equilibrium between a
conjugate acid – base pair• Buffers consist of either
– a weak acid and its salt (conjugate base)– a weak base and its salt (conjugate acid)
– Acetic acid (CH3COOH) with sodium acetate (CH3COONa)
• An equilibrium is established in solution between the acid and the salt anion
• A buffer is LeChatelier’s Principle in action
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
The Buffer Process
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4
Aci
d-B
ase
Buf
fers
Addition of Base (OH-) to a Buffer Solution
• Adding a basic substance to a buffer causes changes– The OH- will react with the H3O+ producing water– Acid in the buffer system dissociates to replace
the H3O+ consumed by the added base– Net result is to maintain the pH close to the initial
level
• The loss of H3O+ (the stress) is compensated by the dissociation of the acid to produce more H3O+
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)8.4
Aci
d-B
ase
Buf
fers
Addition of Acid (H3O+) to a Buffer Solution
• Adding an acidic substance to a buffer causes changes– The H3O+ from the acid will increase the overall
H3O+
– Conjugate base in the buffer system reacts with the H3O+ to form more acid
– Net result is to maintain the H3O+ concentration and the pH close to the initial level
• The gain of H3O+ (the stress) is compensated by the reaction of the conjugate base to produce more acid
8.4
Aci
d-B
ase
Buf
fers
Buffer Capacity• Buffer Capacity – a measure of the
ability of a solution to resist large changes in pH when a strong acid or strong base is added
• Also described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing pH
• Buffering process is an equilibrium reaction described by an equilibrium-constant expression
– In acids, this constant is Ka
• If you want to know the pH of the buffer, solve for [H3O+], then calculate pH
COOH]CH[]COOCH][OH[K
3
-33
a
+
=
8.4
Aci
d-B
ase
Buf
fers
Preparation of a Buffer Solution
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
8.4
Aci
d-B
ase
Buf
fers
Calculating the pH of a Buffer Solution
Calculate the pH of a buffer solution in which – Both the acetic acid (acid) and sodium acetate (salt)
concentrations are 2.0 x 10-2 M– Sodium acetate, the salt, is also the conjugate base
– Ka = 1.75 x 10-5
– [H3O+] = [acid]Ka[conjugate base]
= {(2.0 x 10-2 M) x (1.75 x 10-5)} / 2.0 x 10-2 M= 1.75 x 10-5
– pH = -log 1.75 x 10-5 = 4.76
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
COOH]CH[]COOCH][OH[K
3
-33
a
+
=
8.4
Aci
d-B
ase
Buf
fers
Henderson-Hasselbalch Equation• Solution of equilibrium-constant expression
and pH can be combined into one operation• Henderson-Hasselbalch Equation is this
combined expression• Using these two equations:
– pKa = -log Ka just as pH = -log[H3O+]
– pKa = pH – log ( [CH3COO-] / [CH3COOH] )
– Henderson-Hasselbalch –
– pH = pKa + log( [CH3COO-] / [CH3COOH] )
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)COOH]CH[
]COOCH][OH[K3
-33
a
+
=
8.4
Aci
d-B
ase
Buf
fers
Henderson-Hasselbalch Equation
• pH = pKa + log( [CH3COO-] / [CH3COOH] ) can be rewritten pH = pKa + log ( [conjugate base] / [weak acid])
8.5 Oxidation-Reduction Processes
• Oxidation-reduction processes are responsible for many types of chemical change
• Oxidation - defined by one of the following – loss of electrons – loss of hydrogen atoms– gain of oxygen atoms
• Example: NaNa+ + e-
– Oxidation half reaction
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
s• Reduction - defined by one of the
following:– gain of electrons– gain of hydrogen– loss of oxygen
• Example: Cl + e- Cl-
– Reduction half reaction
• Cannot have oxidation without reduction.
Oxidation-Reduction Processes
Na + Cl Na+ + Cl-
Oxidizing Agent• Is reduced• Gains electrons• Causes oxidation
Reducing Agent• Is oxidized• Loses electrons• Causes reduction
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sOxidation and Reduction as Complementary Processes
Na Na+ + e-
Cl + e- Cl-
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sApplications of Oxidation and
Reduction• Corrosion - the deterioration of metals
caused by an oxidation-reduction process– Example: rust (oxidation of iron)
4Fe(s) + 3O2(g) 2Fe2O3(s)
• Combustion of Fossil Fuels– Example: natural gas furnaces
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
s• Bleaching
• Most bleaching agents are oxidizing agents
• The oxidation of the stains produces compounds that do not have color– Example: Chlorine bleach - sodium
hypochlorite (NaOCl)
Applications of Oxidation and Reduction
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sBiological Processes
• Respiration– Electron-transport chain of aerobic
respiration uses reversible oxidation and reduction of iron atoms in cytochrome c
• Metabolism– Break down of molecules into smaller pieces
by enyzmes
• Is Zn oxidized or reduced?• Oxidized
• Copper is reduced8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sVoltaic Cells
• Voltaic cell – electrochemical cell that converts stored chemical energy into electrical energy
• Let’s consider the following reaction:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sVoltaic Cells
• If the two reactants are placed in the same flask they cannot produce electrical current
• A voltaic cell separates the two half reactions
• This makes the electrons flow through a wire to allow the oxidation and reduction to occur
Zn Zn2+ + 2e-
Oxidationanode – electrode
where oxidation occurs
Cu2+ + 2e- CuReduction
cathode – electrode where reduction occurs8.
5O
xida
tion-
Red
uctio
n Pr
oces
ses
Voltaic Cell Generating Electrical Current
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sVoltaic Cell Generating
Electrical Current
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sSilver Battery
• Batteries use the concept of the voltaic cell
• Modern batteries are– Smaller
– Safer
– More dependable
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sElectrolysis
• Electrolysis reactions – uses electrical energy to cause nonspontaneous oxidation-reduction reactions to occur
• These reactions are the reverse of a voltaic cell– Rechargeable battery
• When powering a device behaves as voltaic cell• With time the chemical reaction nears completion• Battery appears to “run down”• Cell reaction is reversible when battery attached
to charger
8.5
Oxi
datio
n-R
educ
tion
Proc
esse
sComparison of Voltaic and
Electrolytic Cells
Chapter 9
The Nucleus, Radioactivity and Nuclear Medicine
Denniston Topping Caret
6th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
9.1 Natural Radioactivity• Radioactivity – process by which atoms
emit energetic particles or rays• Radiation – the particles or rays emitted
– comes from the nucleus• Nuclear symbols – what we use to designate
the nucleus– Atomic symbol– Atomic number– Mass number
9.1
Nat
ural
Rad
ioac
tivity
B115
atomic symbol
atomic number number of protons
mass numbernumber of
protons and neutrons
Nuclear Symbols
B115
9.1
Nat
ural
Rad
ioac
tivity Writing Nuclear Symbols
• This defines an isotope of boron
• In nuclear chemistry often called a nuclide
• This is not the only isotope of boron– boron-10 also exists– How many protons and neutrons does
boron-10 have?• 5 protons, 5 neutrons
Three Isotopes of Carbon• Each nucleus contains the same number of protons• Only the number of neutrons is different• With different numbers of neutrons the mass of
each isotope is different
9.1
Nat
ural
Rad
ioac
tivity
9.1
Nat
ural
Rad
ioac
tivity Unstable Isotopes
• Some isotopes are stable
• The unstable isotopes are the ones that produce radioactivity
• To write nuclear equations we need to be able to write the symbols for the isotopes and the following:– alpha particle
– beta particles
– gamma rays
α α He He 42
42
2 42
+9.1
Nat
ural
Rad
ioac
tivity Alpha Particles
• Alpha particle (a) – 2 protons, 2 neutrons
• Same as He nucleus (He2+)
• Slow moving, and stopped by small barriers
• Symbolized in the following ways:
β β e 01-
01−
9.1
Nat
ural
Rad
ioac
tivity Beta Particles and Positrons
• Beta particles (b) – fast-moving electron
• Emitted from the nucleus as a neutron is converted to a proton
• Higher speed particles, more penetrating than alpha particles
• The symbol is…
g
9.1
Nat
ural
Rad
ioac
tivity Gamma Rays
• Gamma Rays (g) – pure energy (electromagnetic radiation)
• Highly energetic
• The most penetrating form of radiation
• Symbol is simply…
9.1
Nat
ural
Rad
ioac
tivity Properties of Alpha, Beta, and
Gamma Radiation
• Ionizing radiation – produces a trail of ions throughout the material that it penetrates
• The penetrating power of the radiation determines the ionizing damage that can be caused
• Alpha particle < beta particle < gamma rays
9.2 Writing a Balanced Nuclear Equation
• Nuclear equation – used to represent nuclear change
• In a nuclear equation, you do not balance the elements, instead...– the total mass on each side of the reaction
arrow must be identical– the sum of the atomic numbers on each side of
the reaction arrow must be identical
9.2
Writ
ing
a B
alan
ced
Nuc
lear
Equ
atio
ns He Th U 42
23490
23892 +→
238 = 234 + 4
92 = 90 + 2
mass number
atomic number
Alpha Decay
• Upon decomposition, nitrogen-16 produces oxygen-16 and a beta particle
• In beta decay, one neutron in nitrogen-16 is converted to a proton and the electron, the beta particle is released
e ON 01-
168
167 +→
9.2
Writ
ing
a B
alan
ced
Nuc
lear
Equ
atio
nsBeta Decay
Positron Emission• A positron has same mass as an electron, or beta
particle, BUT opposite charge (+)• Unlike beta decay, the product nuclide has the
same mass number as the parent BUT the atomic number has decreased by one
01115
116
01
115
116
B C
e B C
+→
+→
9.2
Writ
ing
a B
alan
ced
Nuc
lear
Equ
atio
ns
• Gamma radiation occurs to increase the stability of an isotope – The energetically unstable isotope is called a
metastable isotope
• The atomic mass and number do not change
• Usually gamma rays are emitted along with alpha or beta particles
Tc Tc 9943
m9943 +→9.
2 W
ritin
g a
Bal
ance
d N
ucle
ar E
quat
ions
Gamma Production
• To predict the product, simply remember that the mass number and atomic number are conserved
• What is the identity of X?
e01-
23992 XU +→
93239Np9.
2 W
ritin
g a
Bal
ance
d N
ucle
ar E
quat
ions
Predicting Products of Nuclear Decay
9.3 Properties of RadioisotopesNuclear Structure and Stability• Binding Energy – the energy that holds the
protons, neutrons, and other particles together in the nucleus
• Binding energy is very large• When isotopes decay (forming more stable
isotopes,) binding energy is released
9.3
Prop
ertie
s of
Rad
iois
otop
es Important factors for stable isotopes– Ratio of neutrons to protons– Nuclei with large number of protons (84 or more)
tend to be unstable– The “magic numbers” of 2, 8, 20, 50, 82, or 126 help
determine stability – these numbers of protons or neutrons are stable
– Even numbers of protons or neutrons are generally more stable than those with odd numbers
– All isotopes (except 1H) with more protons than neutrons are unstable
Stable Radioisotopes
9.3
Prop
ertie
s of
Rad
iois
otop
esHalf-Life
• Half-life (t1/2) – the time required for one-half of a given quantity of a substance to undergo change
• Each radioactive isotope has its own half-life – Ranges from a fraction of a second to a
billion years
– The shorter the half-life, the more unstable the isotope
Half-Lives of Selected Radioisotopes
9.3
Prop
ertie
s of
Rad
iois
otop
es
9.3
Prop
ertie
s of
Rad
iois
otop
esDecay Curve for the Medically Useful Radioisotope Tc-99m
9.3
Prop
ertie
s of
Rad
iois
otop
esPredicting the Extent of
Radioactive DecayA patient receives 10.0 ng of a radioisotope with a half-life of 12 hours. How much will remain in the body after 2.0 days, assuming radioactive decay is the only path for removal of the isotope from the body.• Calculate n, the number of half-lives elapsed using the half-life as the conversion factorn = 2.0 days x 1 half-life / 0.5 days = 4 half lives
• Calculate the amount remaining10.0 ng 5.0 ng 2.5 ng 1.3 ng 0.63 ng
1st half-life 2nd half-life 3rd half-life 4th half-life
• 0.63 ng remain after 4 half-lives
Radiocarbon Dating
• Radiocarbon dating – the estimation of the age of objects through measurement of isotopic ratios of carbon– Ratio of carbon-14 and carbon-12
• Basis for dating:– Carbon-14 (a radioactive isotope) is
constantly being produced by neutrons from the sun
H C n N 11
146
10
147 +→+
9.5
Rad
ioca
rbon
Dat
ing
• Living systems are continually taking in carbon– The ratio of carbon-14 to carbon-12 stays
constant during its lifetime
• Once the living system dies, it quits taking in the carbon-14– The amount of carbon-14 decreases
according to the reaction:
e N C 01-
147
146 +→
• The half-life of carbon-14 is 5730 years– This information is used to calculate the age
Radiocarbon Dating
9.4 Nuclear PowerEnergy Production
E = mc2
• Equation by Albert Einstein shows the connection between energy (E) and mass (m)
• c is the speed of light • The equation shows that a very large amount of
kinetic energy can be formed from a small amount of matter– Release this kinetic energy to convert liquid water into
steam– The steam drives an electrical generator producing
electricity
9.4
Nuc
lear
Pow
er • Fission (splitting) occurs when a heavy nuclear particle is split into smaller nuclei by a smaller nuclear particle
•Accompanied by a large amount of energy
•Is self-perpetuating
•Can be used to generate steam
energy n 3 Ba Kr U U n 10
14156
9236
23692
23592
10 +++→→+
Nuclear Fission
9.4
Nuc
lear
Pow
erFission of Uranium-235
• Chain reaction – the reaction sustains itself by producing more neutrons
9.4
Nuc
lear
Pow
erRepresentation of the “Energy Zones” of a Nuclear Reactor
• A nuclear power plant uses a fissionable material as fuel– Energy released by the fission heats water– produces steam– drives a generator or turbine– converts heat to electrical energy
• Fusion (to join together) – combination of two small nuclei to form a larger nucleus
• Large amounts of energy is released
• Best example is the sun
• An Example:
• No commercially successful plant exists in U.S.
energy n He H H 10
42
31
21 ++→+
9.4
Nuc
lear
Pow
erNuclear Fusion
9.4
Nuc
lear
Pow
erBreeder Reactors
• Breeder reactor – fission reactor that manufactures its own fuel
• Uranium-238 (non-fissionable) is converted to plutonium-239 (fissionable)
• Plutonium-239 undergoes fission to produce energy
9.5 Medical Applications of Radioactivity
• Modern medical care uses the following:– Radiation in the treatment of cancer– Nuclear medicine - the use of
radioisotopes in the diagnosis of medical conditions
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
• Based on the fact that high-energy gamma rays cause damage to biological molecules
• Tumor cells are more susceptible than normal cells
• Example: cobalt-60• Gamma radiation can cure cancer but
can also cause cancer
Cancer Therapy Using Radiation
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
Nuclear Medicine• The use of isotopes in diagnosis
• Tracers – small amounts of radioactive substances used as probes to study internal organs
• Nuclear imaging – medical techniques involving tracers
• Example:– Iodine concentrates in the thyroid gland.– Using radioactive 131I and 125I will allow the study
of how the thyroid gland is taking in iodine
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
Tracer Studies• Isotopes with short half-lives are preferred for
tracer studies. Why?– They give a more concentrated burst
– They are removed more quickly from the body
• Examples of imaging procedures:– Bone disease and injury using technetium-99m
– Cardiovascular disease using thallium-201
– Pulmonary disease using xenon-133
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
Making Isotopes for Medical Applications
• Artificial radioactivity – a normally stable, nonradioactive nucleus is made radioactive
• Made in two ways:
• In core of a nuclear reactor
• In particle accelerators – small nuclear particles are accelerated to speeds approaching the speed of light and slammed into another nucleus
• Tracer in the liver
• Used in the diagnosis of Hodgkin’s disease
Au n Au 19879
10
19779 →+
Ga p Zn 6731
11
6630 →+
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
Examples of Artificial Radioactivity
9.5
Med
ical
App
licat
ions
of
Rad
ioac
tivity
Preparation of Technetium-99m
• Some isotopes used in nuclear medicine have such a short half-life that they need to be generated on site
• 99mTc has a half-life of only 6 hours
e Tc Mo 01-
99m43
9942 +→
9.6 Biological Effects of Radiation
Radiation Exposure and SafetyThe Magnitude of the Half-Life• Isotopes with short half-lives have one major
disadvantage and one major advantage– Disadvantage: larger amount of radioactivity per
unit time– Advantage: if accident occurs, reaches
background radiation levels more rapidly
9.6
Bio
logi
cal E
ffec
ts o
f R
adia
tion
Shielding• Alpha and beta particles need low level
of shielding: lab coat and gloves• Lead, concrete or both required for
gamma rays
Distance from the Radioactive Source• Doubling the distance from the source
decreases the intensity by a factor of 4
Radiation Exposure and Safety
9.6
Bio
logi
cal E
ffect
s of
Rad
iatio
nRadiation Exposure and Safety
Time of Exposure• Effects are cumulative
Types of Radiation Emitted• Alpha and beta emitters are generally less
hazardous then gamma emitters
Waste Disposal• disposal sites are considered temporary
9.7 Measurement of Radiation
Nuclear Imaging• Isotope is administered• Isotope begins to concentrate in the organ• Photographs (nuclear images) are taken at
periodic intervals• Emission of radioactive isotope creates the
image
• Computers and television are coupled
• Gives a continuous and instantaneous record of the voyage of the isotope throughout the body– Gives increased
sensitivity– CT scanner is an
example
Computer Imaging9.
7 M
easu
rem
ent o
f R
adia
tion
9.7
Mea
sure
men
t of
Rad
iatio
nThe Geiger Counter
• Detects ionizing radiation
• Has largely been replaced by more sophisticated devices
9.7
Mea
sure
men
t of
Rad
iatio
nFilm Badges
• A piece of photographic film that is sensitive to energies corresponding to radioactive emissions
• The darker the film, when developed, the longer the worker has been exposed
9.7
Mea
sure
men
t of
Rad
iatio
nUnits of Radiation Measurement
The Curie• The amount of radioactive material
that produces 3.7 x 1010 atomic disintegrations per second
• Independent of the nature of the radiation
The Roentgen• The amount of radiation needed to
produce 2 x 109 ion pairs when passing through one cm3 of air at 0oC
• Used for very high energy ionizing radiation only
9.7
Mea
sure
men
t of
Rad
iatio
nUnits of Radiation Measurement
9.7
Mea
sure
men
t of
Rad
iatio
n Rad – Radiation absorbed dosage
• The dosage of radiation able to transfer 2.4 x 10-3 cal of energy to one kg of matter
• This takes into account the nature of the absorbing material
Units of Radiation Measurement
9.7
Mea
sure
men
t of
Rad
iatio
nThe Rem
• Roentgen Equivalent for Man
• Obtained by multiplication of the rad by a factor called the relative biological effect (RBE)• RBE = 10 for alpha particles• RBE = 1 for beta particles
• Lethal dose (LD50) - the acute dosage of radiation that would be fatal for 50% of the exposed population– LD50 = 500 rems
Units of Radiation Measurement
Chapter 10
An Introduction to Organic Chemistry:The Saturated Hydrocarbons
Denniston Topping Caret
6th Edition
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
10.1 The Chemistry of CarbonWhy are there so many organic compounds?1. Carbon forms stable, covalent bonds with other
carbon atoms• Consider three allotropic forms of elemental carbon
– Graphite in planar layers– Diamond is a three-dimensional network– Buckminsterfullerene is 60 C in a roughly spherical
shape
Why are there so many organic compounds?
2. Carbon atoms form stable bonds with other elements, such as:
– Oxygen– Nitrogen– Sulfur– Halogen
• Presence of these other elements confers many new physical and chemical properties on an organic compound
Why are there so many organic compounds?
3. Carbon atoms form double or triple bonds with:
– Other carbon atoms (double & triple)
– Oxygen (double only)– Nitrogen (double & triple)
• These combinations act to produce a variety of organic molecules with very different properties
CH2 O
CH2 NH
CH2 CH2
CH CH
CH N
double bonds
triple bonds
Why are there so many organic compounds?
4. Carbon atoms can be arranged with these other atoms; is nearly limitless
– Branched chains– Ring structures– Linear chains
• Two organic compounds may even have the same number and kinds of atoms but completely different structures and thus, different properties
– These are called isomers
CH3CH2CH2CH3 CH2 CH2
CH2
CH2
CH2
Isomers
• Many carbon compounds exist in the form of isomers• Isomers are compounds with the same molecular formula but different structures• An isomer example: both are C4H10but have different structures
– Butane– Methylpropane
Isomers
All have the same molecular formula: C4H8
Important Differences Between Organic and Inorganic Compounds
• Bond type–Organics have covalent bonds
• Electron sharing–Inorganics usually have ionic bonds
• Electron transfer• Structure
–Organics• Molecules• Nonelectrolytes
–Inorganics• Three-dimensional crystal structures• Often water-soluble, dissociating into ions -
electrolytes
Important Differences Between Organic and Inorganic Compounds
• Melting Point & Boiling Point–Organics have covalent bonds
• Intermolecular forces broken fairly easily–Inorganics usually have ionic bonds
• Ionic bonds require more energy to break• Water Solubility
–Organics• Nonpolar, water insoluble
–Inorganics• Water-soluble, readily dissociate
Comparison of Major Properties of Organic and Inorganic Compounds
Families of Organic Compounds
• Hydrocarbons contain only carbon and hydrogen
• They are nonpolar molecules– Not soluble in water – Are soluble in typical nonpolar organic
solvents • Toluene• Pentane
Families of Organic Compounds
• Hydrocarbons are constructed of chains or rings of carbon atoms with sufficient hydrogen atoms to fulfill carbon’s need for four bonds
• Substituted hydrocarbon is one in which one or more hydrogen atoms is replaced by another atom or group of atoms
Division of the Family of Hydrocarbons
Hydrocarbon Saturation• Alkanes are compounds that contain only
carbon-carbon and carbon-hydrogen single bonds– A saturated hydrocarbon has no double or
triple bonds• Alkenes and alkynes are unsaturated
because they contain at least one carbon to carbon double or triple bond
Cyclic Structure of Hydrocarbons
• Some hydrocarbons are cyclic– Form a closed ring– Aromatic hydrocarbons contain a benzene
ring or related structure
Common Functional Groups
10.2 Alkanes
• The general formula for a chain alkane is CnH2n+2– In this formula n = the number of carbon atoms
in the molecule• Alkanes are saturated hydrocarbons
– Contain only carbon and hydrogen– Bonds are carbon-hydrogen and carbon-carbon
single bonds
Formulas Used in Organic Chemistry
• Molecular formula - lists kind and number of each type of atom in a molecule, no bonding pattern
• Structural formula - shows each atom and bond in a molecule
• Condensed formula - shows all the atoms in a molecule in sequential order indicating which atoms are bonded to which
• Line formula - assume a carbon atom at any location where lines intersect– Assume a carbon at the end of any line – Each carbon in the structure is bonded to the correct
number of hydrogen atoms
The Tetrahedral Carbon Atom
(a) Lewis dot structure(b)The tetrahedral shape around the carbon atom(c) The tetrahedral carbon drawn with dashes and
wedges(d)The stick drawing of the tetrahedral carbon
atom(e) Ball and stick model of methane
Drawing Methane and Ethane
C
H
HH
CH
HH
109.5 o
Staggered form of ethane
H
HH
H
in plane
in front of plane
behind plane
Comparison of Ethane and Butane Structures
Names and Formulas of the First Ten Straight-Chain Alkanes
Comparison of Physical Properties of Five Isomers of Hexane
Compare the basic linear structure of hexane– All other isomers have one or more carbon atoms
branching from the main chain– Branched-chain forms of the molecule have a much
smaller surface area• Intermolecular forces are weaker• Boiling and melting points are lower than straight chains
Physical Properties of Organic Molecules
1. Nonpolar2. Not water soluble3. Soluble in nonpolar organic solvents4. Low melting points5. Low boiling points6. Generally less dense (lighter) than water7. As length (molecular weight) increases, melting
and boiling points increase as does the density
Properties of Alkanes
-250
-200
-150
-100
-50
0
50
100
150
200
0 1 2 3 4 5 6 7 8 9 10
Number of Carbons in Chain
Melting PointBoiling Point
Properties of Alkanes
• Most of the alkanes are hydrophobic: water hating
• Straight chain alkanes comprise a homologous series: compounds of the same functional class that differ by a –CH2- group
• Nonpolar alkanes are:– Insoluble in water (a highly polar solvent)– Less dense than water and float on it
Alkyl Groups
• An alkyl group is an alkane with one hydrogen atom removed
• It is named by replacing the -ane of the alkane name with -yl
• Methane becomes a methyl group
or CH3H CH
HH
H CH
H
Alkyl Groups
• All six hydrogens on ethane are equivalent• Removing one H generates the ethyl group• All 3 structures shown at right are the same
CH
CHH
HH
H
CH3CH2CH3CH2
C2H5
Names and Formulas of the First Five Alkyl Groups
Alkyl Group Classification
• Alkyl groups are classified according to the number of carbons attached to the carbon atom that joins the alkyl group to a molecule
• All continuous chain alkyl groups are 1º• Isopropyl and sec-butyl are 2º groups
Iso- Alkyl Groups• Propane: removal of a hydrogen generates
two different propyl groups depending on whether an end or center H is removed
n-propyl isopropyl
CH3CH CH3CH3CH2CH2
CH3 CH2 CH3
Sec- Alkyl Groups• n-butane gives two butyl groups depending
on whether an end (1º) or interior (2º) H is removed
sec-butyln-butyl
CH3 CH2 CH2 CH3
CH3 CH CH2 CH3CH3 CH2 CH2 CH2
Structures and Names of Some Branched-Chain Alkyl Groups
CH3 CH CH3
CH3
CH3 CH CH2
CH3
CH3 C CH3
CH3
More Alkyl Group Classification
• Isobutane gives two butyl groups depending on whether a 1o or 3o H is removed
isobutyl t-butyl
1o C 3o C
Nomenclature
• The IUPAC (International Union of Pure and Applied Chemistry) is responsible for chemical names
• Before learning the IUPAC rules for naming alkanes, the names and structures of eight alkyl groups must be learned
• These alkyl groups are historical names accepted by the IUPAC and integrated into modern nomenclature
Carbon Chain Length and Prefixes
IUPAC Names for Alkanes
1. The base or parent name for an alkane is determined by the longest chain of carbon atoms in the formula
– The longest chain may bend and twist, it is seldom horizontal
– Any carbon groups not part of the base chain are called branches or substituents
– These carbon groups are also called alkyl groups
IUPAC Names for Alkanes
• Rule 1 applied – Find the longest chain in each molecule
• A=7 B=8
CH3
CH2
CH2CH2CH CH2CH3
CH3
CH3CHCH2
CH3
CH2CHCH2CH2
CH2
CH3
CH3
A
B
IUPAC Names for Alkanes
2. Number the carbon atoms in the chain starting from the end with the first branch
– If both branches are equally from the ends, continue until a point of difference occurs
IUPAC Names for AlkanesNumber the carbon atoms correctly• Left: first branch is on carbon 3• Right: first branch is on carbon 3 (From
top) not carbon 4 (if number from right)
CH3
CH2
CH2CH2CH CH2CH3
CH3
CH3CHCH2
CH3
CH2CHCH2CH2
CH2
CH3
CH3
1
2
3 4 5
6 7 8
this branch would be on C-4if you started at correct C-8
123
45
6
7
IUPAC Names for Alkanes
3. Write each of the branches/substituents in alphabetical order before the base/stem name (longest chain)
– Halogens usually come first– Indicate the position of the branch on the main
chain by prefixing its name with the carbon number to which it is attached
– Separate numbers and letters with a hyphen– Separate two or more numbers with commas
IUPAC Names for Alkanes
Name : 4-ethyl-2-methylhexane
CH3CH2CH CH2CH CH3
CH3CH2
CH3
IUPAC Names for Alkanes• Hyphenated and number prefixes are
not considered when alphabetizing groups– Name the compound below
– 5-sec-butyl-4-isopropylnonane
CH CHCH3
CH3 CH2 CH2 CH3
CHCHCH3
CH2 CH3CH2 CH2 CH2 CH3
IUPAC Names for Alkanes
• When a branch/substituent occurs more than once– Prefix the name with
• di• tri• tetra
– Then list the number of the carbon branch for that substituent to the name with a separate number for each occurrence
• Separate numbers with commas
• e.g., 3,4-dimethyl or 4,4,6-triethyl
IUPAC Names for Alkanes
5-ethyl-2,3-dimethylheptaneethyl>dimethyl
CH3CHCH3 CH CH2CH
CH2CH3
CH3
CH2CH3
Name
Practice: IUPAC Name
6-ethyl-6-isobutyl-3,3-dimethyldecane
CH3CCH3
CH2
CH3
CH2CH2C CH2CH3
CH2
CH CH3CH3
CH2CH2CH2CH3
Name1
2
3 4 5 6
7 8 9 10
Structural Isomers• Constitutional/Structural Isomers differ in how
atoms are connected– Two isomers of butane have different physical
properties– The carbon atoms are connected in different
patterns
ButaneBp –0.4 oCMp –139 oC
IsobutaneBp –12 oC
Mp –145 oC
CH3 CH2 CH2 CH3 CH3 CH CH3
CH3
10.3 Cycloalkanes
• Cycloalkanes have two less hydrogens than the corresponding chain alkane– Hexane=C6H14; cyclohexane=C6H12
• To name cycloalkanes, prefix cyclo- to the name of the corresponding alkane– Place substituents in alphabetical order before
the base name as for alkanes– For multiple substituents, use the lowest
possible set of numbers; a single substituent requires no number
Cycloalkane Structures
Cyclopropane
Cyclobutane
Cyclohexane
Type of Formula: Structural Condensed Line
Naming a Substituted Cycloalkane
Name the two cycloalkanes shown below• Parent chain 6 carbon ring 5 carbon ring
cyclohexane cyclopentane• Substituent 1 chlorine atom a methyl group
chloro methyl• Name Chlorocyclohexane Methylcyclopentane
cis-trans Isomers in Cycloalkanes• Atoms of an alkane can rotate freely around the
carbon-carbon single bond having an unlimited number of arrangements
• Rotation around the bonds in a cyclic structure is limited by the fact that all carbons in the ring are interlocked– Formation of cis-trans isomers, geometric isomers, is a
consequence of the lack of free rotation• Stereoisomers are molecules that have the same
structural formulas and bonding patterns, but different arrangements of atoms in space– cis-trans isomers of cycloalkanes are stereoisomers
whose substituents differ in spatial arrangement
cis-trans Isomers in Cycloalkanes• Two groups may be on the same side (cis) of the imagined
plane of the cycloring or they may be on the opposite side (trans)
• Geometric isomers do not readily interconvert, only by breaking carbon-carbon bonds can they interconvert
10.4 Conformations of Alkanes• Conformations differ only in rotation about carbon-
carbon single bonds • Two conformations of ethane and butane are shown
– The first (staggered form) is more stable because it allows hydrogens to be farther apart and thus, the atoms are less crowded
Two Conformations of Cyclohexane
Chair form (more stable) Boat formA
EE
A
AE
E A
E
E
A
A
H HH H
H H
HH H
H
HH
E=equitorial A=axial
10.5 Reactions of Alkanes
• Alkanes, cycloalkanes, and other hydrocarbons can be: – Oxidized (by burning) in the presence of excess
molecular oxygen, in a process called combustion
– Reacted with a halogen (usually chlorine or bromine) in a halogenation reaction
Alkane Reactions
The majority of the reaction of alkanes are combustion reactions
– Complete CH4 + 2O2 CO2 + 2H2O Complete combustion produces
– Carbon dioxide and water
– Incomplete 2CH4 + 3O2 2CO + 4H2O• Incomplete combustion produces
– Carbon monoxide and water– Carbon monoxide is a poison that binds
irreversibly to red blood cells
HalogenationHalogenation is a type of substitution reaction, a reaction that results in a replacement of one group for another
– Products of this reaction are: • Alkyl halide or haloalkane• Hydrogen halide
– This reaction is important in converting unreactive alkanes into many starting materials for other products
– Halogenation of alkanes ONLY occurs in the presence of heat and/or light (UV)
HH +Br2
heat orlight +HBr
BrH
Petroleum Processing
Fraction Boiling Pt Range ºC
Carbon size
Typical uses
Gas -164-30 C1-C4 Heating, cooking
Gasoline 30-200 C5-C12 Motor fuel
Kerosene 175-275 C12-C16 Fuel for stoves, diesel and jet engines
Heating oil Up to 375 C15-C18 Furnace oil
Lubricating oil
350 and up C16-C20 Lubrication, mineral oil
Greases Semisolid C18-up Lubrication, petroleum jelly
Paraffin (wax)
Melts at 52-57 C20-up Candles, toiletries
Pitch / tar Residue in boiler High Roofing, asphalt