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The Nucleus : A Chemist’s View. Nuclear Symbols. Mass number, A (p + + n o ). Element symbol. Atomic number, Z (number of p + ). Balancing Nuclear Equations.  A reactants =  A products. 235 + 1 = 142 + 91 + 3(1). 92 + 0 = 56 + 36 + 3(0).

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nuclear symbols
Nuclear Symbols

Mass number, A

(p+ + no)

Element symbol

Atomic number, Z

(number of p+)

balancing nuclear equations
Balancing Nuclear Equations

Areactants = Aproducts

235 + 1 = 142 + 91 + 3(1)

92 + 0 = 56 + 36 + 3(0)

Zreactants = Zproducts

balancing nuclear equations 2
Balancing Nuclear Equations #2

222

226 = 4 + ____

222

Rn

86

86

88 = 2 + ___

Atomic number 86 is radon, Rn

balancing nuclear equations 3
Balancing Nuclear Equations #3

95

235 + 1 = 139 + 2(1) + ____

95

Y

39

39

92 + 0 = 53 + 2(0) + ____

Atomic number 39 is yttrium, Y

alpha decay
Alpha Decay

Alpha production (a):

an alpha particle is a

helium nucleus

Alpha decay is limited to heavy, radioactive

nuclei

alpha radiation
Alpha Radiation

Limited to VERY large nucleii.

beta decay
Beta Decay

Beta production (b):

A beta particle is an

electron ejected from

the nucleus

Beta emission converts a neutron to a proton

beta radiation
Beta Radiation

Converts a neutron into a proton.

gamma ray production
Gamma Ray Production

Gamma ray production (g):

Gamma rays are high energy photons produced in association with other forms of decay.

Gamma rays are massless and do not, by themselves, change the nucleus

deflection of decay particles
Deflection of Decay Particles

attract

Opposite charges_________ each other.

repel

Like charges_________ each other.

positron production
Positron Production

Positron emission:

Positrons are the anti-particle of the electron

Positron emission converts a proton to a neutron

electron capture
Electron Capture

Electron capture: (inner-orbital electron is captured by the nucleus)

Electron capture converts a proton to a neutron

nuclear stability
NuclearStability

Decay will occur in such a way as to return a nucleus to the band (line) of stability.

The most stable nuclide is Iron-56

If Z > 83, the nuclide is radioactive

decay kinetics
Decay Kinetics

Decay occurs by first order kinetics (the rate of decay is proportional to the number of nuclides present)

N0 = number of nuclides

present initially

N = number of nuclides

remaining at time t

k = rate constant

t = elapsed time

calculating half life
Calculating Half-life

t1/2 = Half-life (units dependent on rate constant, k)

nuclear fission and fusion
Nuclear Fission and Fusion

Fusion:Combining two light nuclei to form a heavier, more stable nucleus.

Fission: Splitting a heavy nucleus into two nuclei with smaller mass numbers.

energy and mass
Energy and Mass

Nuclear changes occur with small but measurable losses of mass. The lost mass is called the mass defect, and is converted to energy according to Einstein’s equation:

DE = Dmc2

Dm = mass defect

DE = change in energy

c = speed of light

Because c2 is so large, even small amounts of mass are converted to enormous amount of energy.

fission processes
Fission Processes

A self-sustaining fission process is called a chain reaction.

review
Review
  • Oxidation reduction reactions involve a transfer of electrons.
  • OIL- RIG
  • Oxidation Involves Loss
  • Reduction Involves Gain
  • LEO-GER
  • Lose Electrons Oxidation
  • Gain Electrons Reduction
slide28
Solid lead(II) sulfide reacts with oxygen in the air at high temperatures to form lead(II) oxide and sulfur dioxide. Which substance is a reductant (reducing agent) and which is an oxidant (oxidizing agent)?
  • PbS, reductant; O2, oxidant 
  • PbS, reductant; SO2, oxidant 
  • Pb2+, reductant; S2- oxidant 
  • PbS, reductant; no oxidant 
  • PbS, oxidant; SO2, reductant
applications
Applications
  • Moving electrons is electric current.
  • 8H++MnO4-+ 5Fe+2 +5e-® Mn+2 + 5Fe+3 +4H2O
  • Helps to break the reactions into half reactions.
  • 8H++MnO4-+5e-® Mn+2 +4H2O
  • 5(Fe+2® Fe+3 + e- )
  • In the same mixture it happens without doing useful work, but if separate
slide30

e-

e-

e-

e-

e-

  • Connected this way the reaction starts
  • Stops immediately because charge builds up.

H+

MnO4-

Fe+2

galvanic cell
Galvanic Cell

Salt Bridge allows current to flow

H+

MnO4-

Fe+2

slide32

e-

  • Electricity travels in a complete circuit

H+

MnO4-

Fe+2

slide33

Instead of a salt bridge

Porous Disk

H+

MnO4-

Fe+2

slide34

e-

e-

e-

e-

Anode

Cathode

e-

e-

Reducing Agent

Oxidizing Agent

cell potential
Cell Potential
  • Oxidizing agent pulls the electron.
  • Reducing agent pushes the electron.
  • The push or pull (“driving force”) is called the cell potential Ecell
  • Also called the electromotive force (emf)
  • Unit is the volt(V)
  • = 1 joule of work/coulomb of charge
  • Measured with a voltmeter
slide36

0.76

H2 in

Cathode

Anode

H+ Cl-

Zn+2 SO4-2

1 M ZnSO4

1 M HCl

standard hydrogen electrode
Standard Hydrogen Electrode
  • This is the reference all other oxidations are compared to
  • Eº = 0
  • º indicates standard states of 25ºC, 1 atm, 1 M solutions.

H2 in

H+ Cl-

1 M HCl

cell potential1
Cell Potential
  • Zn(s) + Cu+2 (aq)® Zn+2(aq) + Cu(s)
  • The total cell potential is the sum of the potential at each electrode.
  • Eºcell = EºZn® Zn+2 + EºCu+2® Cu
  • We can look up reduction potentials in a table.
  • One of the reactions must be reversed, so change it sign.
cell potential2
Cell Potential
  • Determine the cell potential for a galvanic cell based on the redox reaction.
  • Cu(s) + Fe+3(aq)® Cu+2(aq) + Fe+2(aq)
  • Fe+3(aq)+ e-® Fe+2(aq) Eº = 0.77 V
  • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V
  • Cu(s) ® Cu+2(aq)+2e-Eº = -0.34 V
  • 2Fe+3(aq)+ 2e-® 2Fe+2(aq) Eº = 0.77 V
reduction potential
Reduction potential
  • More negative Eº
    • more easily electron is added
    • More easily reduced
    • Better oxidizing agent
  • More positive Eº
    • more easily electron is lost
    • More easily oxidized
    • Better reducing agent
line notation
Line Notation
  • solid½Aqueous½½Aqueous½solid
  • Anode on the left½½Cathode on the right
  • Single line different phases.
  • Double line porous disk or salt bridge.
  • If all the substances on one side are aqueous, a platinum electrode is indicated.
slide42

For the last reaction

  • Cu(s)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)½Pt(s)

Fe+2

Cu2+

slide43
In a galvanic cell, the electrode that acts as a source of electrons to the solution is called the __________; the chemical change that occurs at this electrode is called________.  

a.  cathode, oxidation  

b.  anode, reduction  

c.  anode, oxidation  

d.  cathode, reduction

slide44
Under standard conditions, which of the following is the net reaction that occurs in the cell?

Cd|Cd2+ || Cu2+|Cu  

a.  Cu2+ + Cd → Cu + Cd2+

b.  Cu + Cd → Cu2+ + Cd2+

c.  Cu2+ + Cd2+ → Cu + Cd  

d.  Cu + Cd 2+ → Cd + Cu2+

galvanic cell1
Galvanic Cell
  • The reaction always runs spontaneously in the direction that produced a positive cell potential.
  • Four things for a complete description.
  • Cell Potential
  • Direction of flow
  • Designation of anode and cathode
  • Nature of all the components- electrodes and ions
practice
Practice
  • Completely describe the galvanic cell based on the following half-reactions under standard conditions.
  • MnO4- + 8 H+ +5e-® Mn+2 + 4H2O Eº=1.51 V
  • Fe+3 +3e-® Fe(s) Eº=0.036V
potential work and d g
Potential, Work and DG
  • emf = potential (V) = work (J) / Charge(C)
  • E = work done by system / charge
  • E = -w/q
  • Charge is measured in coulombs.
  • -w = q E
  • Faraday = 96,485 C/mol e-
  • q = nF = moles of e- x charge/mole e-
  • w = -qE = -nFE= DG
potential work and d g1
Potential, Work and DG
  • DGº = -nFEº
  • if Eº > 0, then DGº < 0 spontaneous
  • if Eº< 0, then DGº > 0 nonspontaneous
  • In fact, reverse is spontaneous.
  • Calculate DGº for the following reaction:
  • Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq)
  • Fe+2(aq)+ e-® Fe(s) Eº = 0.44 V
  • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V
cell potential and concentration
Cell Potential and Concentration
  • Qualitatively - Can predict direction of change in E from LeChâtelier.
  • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s)
  • Predict if Ecell will be greater or less than Eºcell if [Al+3] = 1.5 M and [Mn+2] = 1.0 M
  • if [Al+3] = 1.0 M and [Mn+2] = 1.5M
  • if [Al+3] = 1.5 M and [Mn+2] = 1.5 M
the nernst equation
The Nernst Equation
  • DG = DGº +RTln(Q)
  • -nFE = -nFEº + RTln(Q)
  • E = Eº - RTln(Q) nF
  • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) Eº = 0.48 V
  • Always have to figure out n by balancing.
  • If concentration can gives voltage, then from voltage we can tell concentration.
the nernst equation1
The Nernst Equation
  • As reactions proceed concentrations of products increase and reactants decrease.
  • Reach equilibrium where Q = K and Ecell = 0
  • 0 =Eº- RTln(K) nF
  • Eº= RTln(K) nF
  • nF Eº= ln(K) RT
batteries are galvanic cells
Batteries are Galvanic Cells
  • Car batteries are lead storage batteries.
  • Pb +PbO2 +H2SO4®PbSO4(s) +H2O
batteries are galvanic cells1
Batteries are Galvanic Cells
  • Dry Cell Zn + NH4+ +MnO2 ® Zn+2 + NH3 + H2O + Mn2O3
batteries are galvanic cells2
Batteries are Galvanic Cells
  • Alkaline Zn +MnO2 ® ZnO+ Mn2O3 (in base)
batteries are galvanic cells3
Batteries are Galvanic Cells
  • NiCad
  • NiO2 + Cd + 2H2O ® Cd(OH)2 +Ni(OH)2
corrosion
Corrosion
  • Rusting - spontaneous oxidation.
  • Most structural metals have reduction potentials that are less positive than O2 .
  • Fe ® Fe+2+2e-Eº= 0.44 V
  • O2 + 2H2O + 4e- ® 4OH- Eº= 0.40 V
  • Fe+2 + O2 + H2O ® Fe2O3 + H+
  • Reactions happens in two places.
slide57

Fe2+

Rust

e-

Iron Dissolves- Fe ® Fe+2

O2 + 2H2O +4e-® 4OH-

Salt speeds up process by increasing conductivity

Water

Fe2+ + O2 + 2H2O ® Fe2O3 + 8 H+

preventing corrosion
Preventing Corrosion
  • Coating to keep out air and water.
  • Galvanizing - Putting on a zinc coat
  • Has a lower reduction potential, so it is more easily oxidized.
  • Alloying with metals that form oxide coats.
  • Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.
electrolysis
Electrolysis
  • Running a galvanic cell backwards.
  • Put a voltage bigger than the potential and reverse the direction of the redox reaction.
  • Used for electroplating.
slide60

1.10

e-

e-

Zn

Cu

1.0 M Cu+2

1.0 M Zn+2

Cathode

Anode

slide61

A battery >1.10V

e-

e-

Zn

Cu

1.0 M Cu+2

1.0 M Zn+2

Cathode

Anode

calculating plating
Calculating plating
  • Have to count charge.
  • Measure current I (in amperes)
  • 1 amp = 1 coulomb of charge per second
  • q = I x t
  • q/nF = moles of metal
  • Mass of plated metal
  • How long must 5.00 amp current be applied to produce 15.5 g of Ag from Ag+
calculating plating1
Calculating plating
  • Current x time = charge
  • Charge ∕Faraday = mole of e-
  • Mol of e- to mole of element or compound
  • Mole to grams of compound

Or the reverse if you want time to plate

slide64
Calculate the mass of copper which can be deposited by the passage of 12.0 A for 25.0 min through a solution of copper(II) sulfate.
slide65
How long would it take to plate 5.00 g Fe from an aqueous solution of Fe(NO3)3 at a current of 2.00 A?
other uses
Other uses
  • Electrolysis of water.
  • Separating mixtures of ions.
  • More positive reduction potential means the reaction proceeds forward.
  • We want the reverse.
  • Most negative reduction potential is easiest to plate out of solution.
redox
Redox

Know the table

2. Recognized by change in oxidation state.

3. “Added acid”

4. Use the reduction potential table on the front cover.

5. Redox can replace. (single replacement)

slide68
6. Combination Oxidizing agent of one element will react with the reducing agent of the same element to produce the free element.

I- + IO3- + H+ ® I2 + H2O

7. Decomposition.

a) peroxides to oxides

b) Chlorates to chlorides

c) Electrolysis into elements.

d) carbonates to oxides

examples
Examples
  • A piece of solid bismuth is heated strongly in oxygen.
  • A strip or copper metal is added to a concentrated solution of sulfuric acid.
  • Dilute hydrochloric acid is added to a solution of potassium carbonate.
slide70
Hydrogen peroxide solution is added to a solution of iron (II) sulfate.
  • Propanol is burned completely in air.
  • A piece of lithium metal is dropped into a container of nitrogen gas.
  • Chlorine gas is bubbled into a solution of potassium iodide.
examples1
Examples
  • A stream of chlorine gas is passed through a solution of cold, dilute sodium hydroxide.
  • Asolution of tin ( II ) chloride is added to an acidified solution of potassium permanganate
  • A solution of potassium iodide is added to an acidified solution of potassium dichromate.
slide72
Magnesium metal is burned in nitrogen gas.
  • Lead foil is immersed in silver nitrate solution.
  • Magnesium turnings are added to a solution of iron (III) chloride.
  • Pellets of lead are dropped into hot sulfuric acid
  • Powdered Iron is added to a solution of iron(III) sulfate.
a way to remember
A way to remember
  • An Ox – anode is where oxidation occurs
  • Red Cat – Reduction occurs at cathode
  • Galvanic cell- spontaneous- anode is negative
  • Electrolytic cell- voltage applied to make anode positive
slide74
A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.
  • (a) Draw a diagram of this cell.
  • (b) Describe what is happening at the cathode (Include any equations that may be useful.)
slide75
A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.
  • (c) Describe what is happening at the anode. (Include any equations that may be useful.)
slide76
A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.
  • (d) Write the balanced overall cell equation.
  • (e) Write the standard cell notation.
slide77
A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.

(f) The student adds 4 M ammonia to the copper sulfate solution, producing the complex ion Cu(NH3)+ (aq). The student remeasures the cell potential and discovers the voltage to be 0.88 volt. What is the Cu2+ (aq) concentration in the cell after the ammonia has been added?

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