The nucleus a chemist s view
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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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The nucleus a chemist s view

The Nucleus: A Chemist’s View


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


Types of radiation

Types of Radiation


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


A decay series

A radioactive nucleus reaches a stable state by a series of steps

A Decay Series


Half life concept

Half-life Concept


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)


Sample half lives

Sample Half-Lives


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

Fission


Fission processes

Fission Processes

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


A fission reactor

A Fission Reactor


Fusion

Fusion


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


The nucleus a chemist s view

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


The nucleus a chemist s view

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


The nucleus a chemist s view

e-

  • Electricity travels in a complete circuit

H+

MnO4-

Fe+2


The nucleus a chemist s view

  • Instead of a salt bridge

Porous Disk

H+

MnO4-

Fe+2


The nucleus a chemist s view

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


The nucleus a chemist s view

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.


The nucleus a chemist s view

  • For the last reaction

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

Fe+2

Cu2+


The nucleus a chemist s view

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


The nucleus a chemist s view

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.


The nucleus a chemist s view

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.


The nucleus a chemist s view

1.10

e-

e-

Zn

Cu

1.0 M Cu+2

1.0 M Zn+2

Cathode

Anode


The nucleus a chemist s view

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


The nucleus a chemist s view

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.


The nucleus a chemist s view

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)


The nucleus a chemist s view

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.


The nucleus a chemist s view

  • 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.


The nucleus a chemist s view

  • 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


The nucleus a chemist s view

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.)


The nucleus a chemist s view

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.)


The nucleus a chemist s view

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.


The nucleus a chemist s view

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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