Reference Electrode

A reference electrode is an electrode which has a stable and well-known electrode potential. The high stability of the electrode potential is usually reached by employing a redox system with constant (buffered or saturated) concentrations of each participants of the redox reaction.
There are many ways reference electrodes are used. The simplest is when the reference electrode is used as a half cell to build an electrochemical cell. This allows the potential of the other half cell to be determined. An accurate and practical method to measure an electrode's potential in isolation (absolute electrode potential) has yet to be developed.

Aqueous reference electrodes

Common reference electrodes and potential with respect to the standard hydrogen electrode:

• Standard hydrogen electrode (SHE) (E=0.000 V) activity of H⁺=1
• Normal hydrogen electrode (NHE) (E ≈ 0.000 V)concentration H⁺=1
• Reversible hydrogen electrode (RHE) (E=0.000 V - 0.0591*pH)
• Saturated calomel electrode (SCE) (E=+0.241 V saturated)
• Copper-copper(II) sulfate electrode(CSE) (E=+0.314 V)
• Silver chloride electrode (E=+0.197 V saturated)
• pH-electrode (in case of pH buffered solutions, see buffer solution)
• Palladium-hydrogen electrode
• Dynamic hydrogen electrode (DHE)

• 

  Standard Hydrogen Electrode
  
• 

  Cu-Cu(II) reference electrode
  
• 

  Ag-AgCl reference electrode
  

Nonaqueous reference electrodes

While it is convenient to compare between solvents to qualitatively compare systems it is not quantitatively meaningful. Much as pK_a are related between solvents, but not the same, so is the case with E°. While the SHE might seem to be a reasonable reference for nonaqueous work as it turns out the platinum is rapidly poisoned by many solvents including acetonitrile causing uncontrolled drifts in potential. Both the SCE and saturated Ag/AgCl are aqueous electrodes based around saturated aqueous solution. While for short periods it may be possible to use such aqueous electrodes as references with nonaqueous solutions the long-term results are not trustworthy. Using aqueous electrodes introduces undefined, variable, and unmeasurable junction potentials to the cell in the form of a liquid-liquid junction as well as different ionic composition between the reference compartment and the rest of the cell. The best argument against using aqueous reference electrodes with nonaqueous systems, as mentioned earlier, is that potentials measured in different solvents are not directly comparable.
The potential for the Fc0/+ couple is sensitive to solvent.

	Fc0/+ couple, NBu4PF6 at 298 °C
solvent	E
MeCN	0.40
CH2Cl2	0.46
THF	0.56
DMF	0.45
acetone	0.48

A Quasi-Reference Electrode (QRE) avoids the issues mentioned above. A QRE with ferrocene or similar internal standard (cobaltocene) referenced back to ferrocene is ideal for nonaqueous work. Since the early 1960s ferrocene has been gaining acceptance as the standard reference for nonaqueous work for a number of reasons, and in 1984, IUPAC recommended ferrocene (II/III) as a standard redox couple. The preparation of the QRE electrode is simple, allowing for a fresh reference to be prepared with each set of experiments. Since QREs are made fresh, there is also no concern with improper storage or maintenance of the electrode. QREs are also more affordable than other reference electrodes.
To make a quasi-reference electrode (QRE):

1. Insert a piece of silver wire into concentrated HCl then allow the wire to dry on a KimWipe. This forms an insoluble layer of AgCl on the surface of the electrode and gives you a Ag/AgCl wire. Repeat dipping every few months or if the QRE starts to drift.
2. Obtain a Vycor glass frit (4 mm diameter) and glass tubing of similar diameter. Attach Vycor glass frit to the glass tubing with heat shrink Teflon tubing.
3. Rinse then fill the clean glass tube with supporting electrolyte solution and insert Ag/AgCl wire.
4. The ferrocene (II/III) couple should lie around 400 mV versus this Ag/AgCl QRE in an acetonitrile solution. This potential will varying up to 200 mV with the specific undefined conditions. Thus adding an internal standard such as ferrocene at some point during the experiment is always necessary.

Pseudo-reference electrodes

A pseudo-reference electrode is a term that is not well defined and borders on having multiple meanings since pseudo and quasi are often used interchangeably. They are a class of electrodes named pseudo-reference electrodes because they do not maintain a constant potential but vary predictably with conditions. If the conditions are known, the potential can be calculated and the electrode can be used as a reference. Most electrodes work over a limited range of conditions, such as pH or temperature, outside of this range the electrodes behavior becomes unpredictable. The advantage of a pseudo-reference electrode is that the resulting variation is factored into the system allowing researchers to accurately study systems over a wide range of conditions.
Yttria-stabilized zirconia (YSZ) membrane electrodes were developed with a variety of redox couples, e.g., Ni/NiO. Their potential depends on pH. When the pH value is known, these electrodes can be employed as a reference with notable applications at elevated temperatures.

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Auxochrome

An auxochrome (Gr. Auxanein:to increase, chroma:colour) is a group of atoms attached to a chromophore which modifies the ability of that chromophore to absorb light. They themselves fail to produce the colour; but when present along with the chromophores in an organic compound intensifies the colour of the chromogen. Examples include the hydroxyl group (-OH), the amino group (-NH₂), and an aldehyde group (-CHO).
An auxochrome is a functional group of atoms with nonbonded electrons which, when attached to a chromophore, alters both the wavelength and intensity of absorption. If these groups are in direct conjugation with the pi-system of the chromophore, they may increase the wavelength at which the light is absorbed and as a result intensify the absorption. A feature of these auxochromes is the presence of at least one lone pair of electrons which can be viewed as extending the conjugated system by resonance.

Effects on chromophore

It increases the color of any organic compound. For example, benzene does not display color as it does not have a chromophore; but nitrobenzene is pale yellow color because of the presence of a nitro group (-NO₂) which acts as a chromophore. But Para-hydroxynitrobenzene exhibits a deep yellow color, in which -OH group acts as an auxochrome. Here the auxochrome (-OH) is conjugated with the chromophore -NO₂. Similar behavior is seen in azobenzene which has a red color, but para-hydroxy azobenzene is dark red in color.
The presence of an auxochrome in the chromogen molecule is essential to make a dye. However, if an auxochrome is present in the meta position to the chromophore, it does not affect the color.
An auxochrome is known as a compound that produces a bathochromic shift, also known as red shift because it increases the wavelength of absorption, therefore moving closer to infrared light. Woodward-Fieser rules estimate the shift in wavelength of maximum absorption for several auxochromes attached to a conjugated system in an organic molecule.
An auxochrome helps a dye to bind to the object that is to be colored. Electrolytic dissociation of the auxochrome group helps in binding and it is due to this reason a basic substance takes an acidic dye.

Explanation for the colour modification

A molecule exhibit colour because it absorbs colours only of certain frequencies and reflects or transmits others. They are capable of absorbing and emitting light of various frequencies. Light waves with frequency very close to their natural frequency are absorbed readily. This phenomenon, known as resonance, means that the molecule can absorb radiation of a particular frequency which is same as the frequency of electron movement within the molecule. Chromophore is the part of the molecule where the energy difference between two different molecular orbitals falls within the range of the visible spectrum and hence absorbs some particular colours from visible light. Hence the molecule appears coloured. When auxochromes are attached to the molecule, the natural frequency of the latter gets changed and thus the colour gets modified. Different auxochromes produce different effects in the chromophore which inturn causes absorption of light from other parts of the spectrum. Normally, auxochromes which intensifies the colour are chosen.

Classification

There are mainly two types of auxochromes:

• Acidic: -COOH, -OH, -SO₃H
• Basic: -NHR, -NR₂, -NH₂

Chromophore

A chromophore is the part of a molecule responsible for its colour. The colour arises when a molecule absorbs certain wavelengths of visible light and transmits or reflects others. The chromophore is a region in the molecule where the energy difference between two different molecular orbitals falls within the range of the visible spectrum. Visible light that hits the chromophore can thus be absorbed by exciting an electron from its ground state into an excited state.
In biological molecules that serve to capture or detect light energy, the chromophore is the moiety that causes a conformational change of the molecule when hit by light.

Chemical structure of beta-carotene. The eleven conjugated double bonds that form the chromophore of the molecule are highlighted in red.

Conjugated pi-bond system chromophores

Conjugated chromophore that straightens in response to a photon γ (light), of the correct wavelength: 11-cis-retinal becomes all-trans-retinal

In the conjugated chromophores, the electrons jump between energy levels that are extended pi orbitals, created by a series of alternating single and double bonds, often in aromatic systems. Common examples include retinal (used in the eye to detect light), various food colorings, fabric dyes (azo compounds), pH indicators, lycopene, β-carotene, and anthocyanins. Various factors in a chromophore's structure go into determining at what wavelength region in a spectrum the chromophore will absorb. Lengthening or extending a conjugated system with more unsaturated (multiple) bonds in a molecule will tend to shift absorption to longer wavelengths. Woodward-Fieser rules can be used to approximate ultraviolet-visible maximum absorption wavelength in organic compounds with conjugated pi-bond systems.
Some of these are metal complex chromophores, which contain a metal in a coordination complex with ligands. Examples are chlorophyll, which is used by plants for photosynthesis and hemoglobin, the oxygen transporter in the blood of vertebrate animals. In these two examples, a metal is complexed at the center of a tetrapyrrole macrocycle ring: the metal being iron in the heme group (iron in a porphyrin ring) of hemoglobin, or magnesium complexed in a chlorin-type ring in the case of chlorophyll. The highly conjugated pi-bonding system of the macrocycle ring absorbs visible light. The nature of the central metal can also influence the absorption spectrum of the metal-macrocycle complex or properties such as excited state lifetime. The tetrapyrrole moiety in organic compounds which is not macrocyclic but still has a conjugated pi-bond system still acts as a chromophore. Examples of such compounds include bilirubin and urobilin, which exhibit a yellow color.

Auxochrome

An auxochrome is a functional group of atoms attached to the chromophore which modifies the ability of the chromophore to absorb light, altering the wavelength or intensity of the absorption.

Halochromism in chromophores

Halochromism occurs when a substance changes color as the pH changes. This is a property of pH indicators, whose molecular structure changes upon certain changes in the surrounding pH. This change in structure affects a chromophore in the pH indicator molecule. For example, phenolphthalein is a pH indicator whose structure changes as pH changes as shown in the following table:

Structure
pH	0−8.2	8.2−12.0
Conditions	acidic or near-neutral	basic
Color name	colorless	pink to fuchsia
Color

In a pH range of about 0-8, the molecule has three aromatic rings all bonded to a tetrahedral sp³ hybridized carbon atom in the middle which does not make the π-bonding in the aromatic rings conjugate. Because of their limited extent, the aromatic rings only absorb light in the ultraviolet region, and so the compound appears colorless in the 0-8 pH range. However as the pH increases beyond 8.2, that central carbon becomes part of a double bond becoming sp² hybridized and leaving a p orbital to overlap with the π-bonding in the rings. This makes the three rings conjugate together to form an extended chromophore absorbing longer wavelength visible light to show a fuchsia color.At pH ranges outside 0-12, other molecular structure changes result in other color changes.

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Periodic Table

Visit this blog and at the end of the page you will find a periodic table.
Just click on any element in the table and find exact atomic number, melting and freezing point and also the weight of that element. I am sure you will find it useful.
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Nernst Equation

The Nernst Equation enables one to determine electromotive forces (emf) of many processes, for instance the resting potential of cell membranes. We can then deduce the biological standard potentials which are important in studying biological process such as action potential during a spike of a neuron in response to a stimulus.

Introduction

The Nernst Equation is derived from the emf and the Gibbs energy under non-standard conditions.

Eo=Eoreduction−Eooxidation 

When Eo is positive, the reaction is spontaneous. When Eo is negative, the reaction is not spontaneous. Since the change in Gibbs free energy, ΔG, is also related to spontaneity of a reaction, therefore, ΔG and E are related. Specifically,

ΔG=−nFE 

where, n is # of electrons transferred in the reaction, F is the Faraday constant (96500 C/mol) and E is potential difference. Under standard conditions, this equation is then

ΔGo=−nFEo.

Since,

ΔG=ΔGo+RTlnQ(1)

Substituting ΔG=−nFE and ΔGo=−nFEo into equation (1), we have:

−nFE=−nFEo +RTlnQ

Divide both sides of the equation above by −nF, we have

E=Eo −RTnFlnQ(2)

Equation (2) can be rewritten in the form of log base 10:

E=Eo −2.303RTnFlogQ(3)

At standard temperature T = 298K, the 2.303RTF equals .0592 V, so equation (3) turns into:

E = Eo −.0592VnlogQ

The equation above indicates that the electrical potential of a cell depends upon the reaction quotient Q of the reaction. As the redox reaction proceeds, reactants are consumed, thus concentration of reactants decreases. Conversely, the products concentration increases due to the increased in products formation. As this happens, cell potential gradually decreases until the reaction is at equilibrium, at which ΔG =0.

Example 1

For the Zn-Cu redox reaction: Zn(s)+Cu2+(aq)→Zn2+(aq)+Cu(s)Eocell=+1.10V SOLUTION Initially, [Cu2+] = [Zn2+] = 1.0 M at standard T = 298K As the reaction proceeds, [Cu2+] decreases as [Zn2+] increases. Lets say after one minute, [Cu2+] = 0.05 M while [Zn2+] = 5.0 M. According to Nernst, cell potential after 1 minute is: E = Eo −0.0592VnlogQ E = 1.10V −0.0592V2log5.0M.05 M E=1.04V As you can see, the initial cell potential is E =1.10V, after 1 minute, the potential drops to 1.04 V. As the reaction continues to progress, more Cu2+ will be consumed and more Zn2+ will be generated. As a result, the cell potential continues to decrease and when the cell potential drops down to 0, the concentration of reactants and products stops changing. This is when the reaction is at equilibrium.

At equilibrium, the reaction quotient Q=Keq. Also, at equilibrium, ΔG =0 and ΔG=−nFE, so E =0.   

Therefore, substituting Q=Keq and E =0 into the Nernst equation, we have:

0=Eo−RTnFlnKeq

At standard conditions, the equation above simplifies into:

0=Eo −0.0592nlogKeq

This equation can be rearranged into: 

logKeq=nEo0.0592

The equation above indicates that the equilibrium constant K_eq is proportional to the standard potential of the reaction. Specifically, when:

• K>1,Eo>0, reaction favors products formation.
• K<1,Eo<0, reaction favors reactants formation.

This result fits Le Châtlier's Principle,which states that when a system at equilibrium experiences a change, the system will minimize that change by shifting the equilibrium in the opposite direction.

Biological Application of the Nerst Equation

The emf can be determined measuring the current in a galvanic cell. This method is important when determining the standard potential of oxidizing agents. But our interest is the biological standard potentials such as neurons firing.

Action Potential of a Neuron Cell

Perhaps the most fascinating system in our body is the nervous system. The neuron is the basic operating unit of the nervous system and its mechanism is still under a lot of research. Electrolytes Sodium and Potassium are the most prevalent electrolytes in our body. Intracellular concentrations of Potassium is higher inside the cell in comparison to extracellular. For Sodium it's the opposite more extracellular and less intracellular. This difference in concentration is the mechanics of how a neuron and other cells have resting potentials. ATPase pumps and selective membrane channels allow for the concentration gradient difference. With the Nernst Equation we can deduce the membrane potential of a neuron for our discussion. The concentration of potassium inside is 150 mM, and 15 mM outside. Plugging these values in the equation above out comes the neuron potential which can be depolarize in response to ...

  

References

1. Atkins, Peter and de Paula, Julio. Physical Chemistry for the Life Sciences. New York: W.H. Freeman and Company. p. 214-222.