Category: Chemistry
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Rydberg Constant: Definition, Equation, and Examples
Continue ReadingRydberg Constant Definition
The Rydberg formula is a mathematical formula that is used to measure the wavelength of light resulting from an electron jumping between the energy states of an atom.
When an electron moves from one atomic orbital to another orbital, the energy of electron changes. When the electron moves from a high-energy orbital to a lower energy orbit, then a photon of light is produced. When the electron jumps from a low energy level to a higher energy level, then a photon of light is absorbed by the atom.
What is Rydberg Constant?
Rydberg constant (or represented by the symbol R∞ or RΗ ), where R∞ is for heavier atoms and RH for hydrogen atoms, is a fundamental constant of atomic physics that was proposed by the Swedish physicist named Johannes Rydberg, who defined the wavelengths or frequencies of light in various sequences of associated spectral lines, most remarkably those emitted by hydrogen atoms in the Balmer series. The value of this constant is usually based on the fact that the light emitted by the nucleus of the atom is exceptionally enormous as compared with a single orbiting electron (hence the infinity symbol ∞ is used).
The constant can be expressed as follows;
R∞ = mee4/8ε20h3c
Here;
R∞ represents Rydberg constant
Me represents the rest mass of the electron
e represents the elementary charge
ℇ0 represents permittivity of free space
h represents Planck constant
c represents the speed of light in a vacuum
The value of the Rydberg constant R∞ is equal to 10,973,731.56816 per metre. When it is used in this form in the mathematical explanation of series of spectral lines, then the result we get will be equal to the number of waves per unit length or the wavenumbers. Thus, multiplication with the speed of light results in the frequencies of the spectral lines.
Value of Rydberg Constant
The accepted values of the Rydberg constant, R∞, are as follows;
Rydberg Constant represented in nm is 10 973 731.568 548(83) m-1.
Rydberg Constant represented in Joules is 2.179 871 90(17).10-18 J.
Rydberg Constant represented in Electron Volt is 13.605662285137 eV.
Rydberg Constant represented in Tons of TNT is 5.2100191204589E-28 tTNT.
Rydberg Constant represented in ergs is 2.179872E-11 ergs.
Rydberg Constant and Hydrogen Spectrum
The hydrogen atoms of the molecule dissociates when an electric discharge is passed through a gaseous molecule of hydrogen.Thus, it results in the release or emission of electromagnetic radiation produced by the energetically excited hydrogen atoms. The hydrogen emission spectrum generally contains radiation of discrete frequencies.
This series of the hydrogen emission spectrum is termed as the Balmer series. This is the only series of electromagnetic spectrum that lies in the visible region. The numeric value, 109,677 cm-1, is termed as the Rydberg constant for hydrogen. The Balmer series is fundamentally the part of hydrogen emission spectrum accountable for the excitation electron that moves from higher energy shell to the second shell. Likewise, other transitions also have their own series names.
Some of them are listed below,
• The transition of electron from higher shell to first shell is termed as Lyman series
• The transition of electron from other higher shell to second shell is termed Balmer series
• The transition of electron from other higher shell to third shell is termed Paschen series
• The transition of electron from other higher shell to 4th shell is termed Bracket series
• The transition of electron from other higher shell to 5th shell is termed Pfund serie
Rydberg Constant Equation
Atomic hydrogen demonstrates emission spectrum. This spectrum enfolds several spectral series. Once the electrons in the gas are excited, then they make transitions between different energy levels. These spectral lines are thus the consequence of such electron transitions between different energy levels exhibited by Neils Bohr. The wavelengths of these spectral series is calculated by Rydberg formula.
Johannes Rydberg tried to find a mathematical relationship between spectral lines. He ultimately discovered that there is an integer relationship between the wavenumbers of consecutive lines.
His conclusions were linked with Bohr’s model of the atom to create this formula which is;
1/λ = RZ2(1/n12 – 1/n22)
Where,
λ represents the wavelength of the photon (wavenumber = 1/wavelength)
R represents Rydberg’s constant (1.0973731568539(55) x 107 m-1)
Z represents atomic number of the given atom
n1 and n2 are the integers for orbits or shell where n2 > n1.
It was later found that n2 and n1 were associated with the principal quantum number or also termed as energy quantum number. This formula hence, works quite well for transitions between energy levels of a hydrogen atom where only one electron is involved. For atoms with multiple electrons or heavier atoms, this method starts to break down and often give inappropriate results. The reason for the inaccuracy is because the amount of screening for inner electrons or outer electron transitions often varies. The equation is too basic to compensate for these differences.
For most questions, we will deal with hydrogen so the formula which can be used is as follows:
1/λ = RH(1/n12 – 1/n22)
where;
RH represents Rydberg’s constant
The value of Z for hydrogen is 1
Rydberg Constant
The Rydberg constant holds high position in atomic physics as it is linked with basic fundamental atomic constants, that are e, h, c, and me.
Examples:
1. Determine the wave length of light emitted by the electron in Hydrogen atom when electron jumps from n=4 energy level to n=2 energy level.
Using Rydberg equation;
1/ λ = R (1/22-1/n2)
1/ λ =1.0974×107(1/4-1/16)
= 2057625m-1
Thus,
λ = 1/2057625 = 4.86×10-7m =486 nm.
2. Determine the wavelength of the electromagnetic radiation that is emitted from an electron when it moves from n = 3 to n = 1.
Using the Rydberg equation;
1/λ = R (1/n12 – 1/n22)
Now puting the values of n1 and n2
n1 =1, n2 is equal to 3
R for Rydberg constant = 1.0974 x107m-1
Thus,
1/λ = (1.0974 x 107) (1/12 – 1/32)
The wavelength (λ) = 1.025 x 10-7 m
Rydberg Constant Citations
- New value for the Rydberg constant from the hydrogen Balmer- beta transition. Phys Rev Lett . 1987 Mar 30;58(13):1293-1295.
- Rydberg constant and fundamental atomic physics. Phys Rev A Gen Phys . 1989 Mar 15;39(6):2888-2898.
- The Rydberg constant and proton size from atomic hydrogen. Science . 2017 Oct 6;358(6359):79-85.
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Wavenumber: Definition, Properties, and Examples
Continue ReadingWavenumber
Wavenumber, also called wave number, a unit of frequency, often used in atomic, molecular, and nuclear spectroscopy, equal to the true frequency divided by the speed of the wave and thus equal to the number of waves in a unit distance.
Physicists and pharmacists often use two different types of wavenumber;
The Spatial Wavenumber or also called as spatial frequency is referred to as the number of wavelengths per unit distance. Angular wavenumber is generally used in physics and geophysics. Fundamentally, the equations for both angular and spatial wavenumber are the same except the fact that the angular wavenumber uses 2π in its numerator as this is the number of radians in a complete circle (that is 180 x 2 = 360°).
The Angular Wavenumber is also termed as circular wavenumber gives us the number of radians (a measure of angle) per unit distance. Spatial wavenumber is generally used in chemistry.
Waves can define sound, light, or the wavefunction of given particles, but each wave has a wavenumber.
Theoretically, the wavenumber is also termed as propagation number or angular wavenumber is referred to as the number of the complete cycle of a wave over its wavelength. It is denoted as a scalar quantity and is represented by the symbol k and the mathematical depiction is as follows:
k=1/λ
Where,
• k represents the wavenumber
• λ represents the wavelength
Wavenumber Formula
Using the equation mentioned above to calculate the spatial wavenumber (ν)
ν = 1 / 𝜆
= f / v
Where,
𝜆 represents wavelength
f represents frequency
v represents the speed of the wave.
k = 2π / 𝜆
where
𝜆 = v/f
Thus
k = 2πf / v
This equation is used to calculate angular wavenumber (k).
Wavenumber Formula in Spectroscopy
Wavenumber is a term that is used in spectroscopy to describe a frequency that has been separated by the speed of light in a vacuum. The formula of Wave number in spectroscopy and chemistry fields is given as follows;
¯v¯ = 1/ λ = ω/ 2πc = v/ c
Where,
¯v¯represents the spectroscopy wavenumber.
Λ represents the wavelength often called as spectroscopic wavenumber
Thus,
ω= 2 πv is the angular frequency
Spectroscopic wavenumber can also be converted into energy per photon by using Planck’s relation as given below;
E = hcv¯
Where
E represents the energy per photon
h represents the reduced Planck’s constant = 6.62607004 × 10-34 m2 kg / s
c represents the speed of light
v¯ represents the Spectroscopic Wavenumber
Spectroscopic wavenumber can also be converted to the wavelength of light as mentioned below;
λ (1n/v¯)
Where,
λ represents the wavelength
n represents the refractive index of the medium
v¯represents the Spectroscopic Wavenumber
The SI unit of measurement of Spectroscopic Wavenumber is often expressed as the reciprocal of meter that is written as m-1.
• The CGS unit of measurement of Spectroscopic Wavenumber is often expressed as the reciprocal of a centimeter that is written as cm-1.
Wavenumber Formula for Wave Equations
The formula for wave number in theoretical physics is given by
k = 2π/λ = ω/vp
Where;
k represents the angular wavenumber
λ represents the wavelength
ω = 2πv represents the angular frequency
When an electromagnetic wave propagate at the speed of light or c in a vacuum, then the wave equation k is represented as follows;
k =E/hc
Where,
k represents the angular wave number
E represents the energy of the wave
h represents the Planck’s constant which is equal to 6.62607004 × 10-34 m2 kg / s
c represents the speed of light
Applications of Wavenumber
• A wavenumber helps to calculate the spatial frequency.
• Apart from spatial frequency, wavenumber is also used to explain other quantities for example optics and wave scatterings in physics.
• Wavenumbers and wave vectors are often used to explain in X-ray diffraction and neutron diffraction, electron diffraction, and also in elementary particles in physics.
• Group velocity can also be explained with the help of a wavenumber.
Wavenumber Examples
Example 1: Calculate the Angular Wavenumber if the W wvelength of the Light wave is given as 500 Nanometers.
The formula for angular wavenumber is as follows;
k = 2π/λ
Where,
λ represents the wavelength of the light wave and is given as 500 nanometers which is further equal to 500 × 10-9 m.
[we know that 1nm =10−9m]
Now substituting the values in the formula to get the angular wavenumber as follows:
k = 2π/500×10−9
Thus,
k = 12.56 x 106 m-1
Example 2: Calculate the wavelength, frequency, and wavenumber of a light wave whose time period is given as 5.0×10−10 s.
The frequency or represented by symbol v of the light wave is given by;
1/ period = 1 / 5.0×10−10 seconds = 2×109 hertz or Hz
The wavelength of the light wave is given by;
c/v = 3×108 m / 2×109/s = 15×10−2 m. ( c is the speed of light)
The wavenumber of the light wave is given by;
νˉ= 1/λ= 1/ 15.0×10−2 =6.6 /m
Example 3: Calculate the frequency and wavenumber of radiation with wavelength 380 nm.
Given that; wavelength or
λ = 380nm = 380×10−9m [ we know that 1nm=10−9m]
Speed of light or c = 3×108 m/sec
Thus the Frequency (v) is given as,
v = c/ λ = 3×108 ms-1 / 380×10−9m
= 7.89 × 1014 hertz or Hz.
Wavenumber Citations
- Linear-in-wavenumber actively-mode-locked wavelength-swept laser. Opt Lett . 2020 Oct 1;45(19):5327-5330.
- Frequency-wavenumber domain analysis of guided wavefields. Ultrasonics . 2011 May;51(4):452-66.
- Wavenumber selection method to determine the concentration of cocaine and adulterants in cocaine samples. J Pharm Biomed Anal . 2018 Apr 15;152:120-127.
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Absorption Spectra: Definition, Properties, and Examples
Continue ReadingAbsorption Spectra Definition
Atomic spectra is referred to as the study of atoms (and their atomic ions) through their interaction with electromagnetic radiation. When a beam of light travels from one medium to another, it either bends in the direction of the normal or away from the normal. The speed of light is thus dependent on the nature of the medium through which it passes.
Absorption Spectra Characteristics
• It was observed that when a ray of white light falls on a prism it usually experience refraction twice. First, when it travels from the rarer medium (that is air) to a denser medium (that is glass) and secondly when it passes from the denser medium (that is glass) to a rarer medium (that is air).
• Lastly, a band of colours is observed, commonly termed as a spectrum, which is formed out of a ray of white light. On observing this spectrum more thoroughly, it was seen that the colour having a smaller wavelength deviates the most and vice versa.
• Therefore, a spectrum of colours is seen which ranges from the colour red to violet, where the red colour having the longest wavelength deviates the least. This kind of spectrum is often termed a continuous spectrum.
• The spectrum of the electromagnetic radiation released or absorbed by an electron during its transitions between different energy levels in an atom is termed as atomic spectra.
Types of Atomic Spectra
There are three types of atomic spectra as given below;
Emission Spectra
Absorption Spectra
Continuous Spectra
Spectra and Spectroscopy
Spectrum is broadly used in the field of optics and in many other fields. Spectrum displays a varied range of wavelengths having different frequency radiations. A rainbow is referred to as a spectrum that contains different wavelengths of light. The spectrum of light formed from the rainbow is generally referred to as VIBGYOR.
A spectroscope or Spectrograph is the device that is used to separate the radiations of different wavelengths. A spectrometer is a scientific instrument that helps to separate and measure spectral components of this physical phenomenon. Spectroscopy is the branch of science that generally deals with the study of the spectrum.
Classification of Spectra
Spectra is categorized into two types as mentioned below:
• Emission spectra
• Absorption spectra
i. Emission Spectra
The emission spectrum is commonly formed by the radiation emitted or produced by an electron in the excited molecules or atoms and this is termed as the emission spectrum. When an atom or molecule absorbs energy, then the electrons are excited to a higher energy level and when the electron falls back to its lower energy level, light is emitted or produced, which generally has the energy equivalent, to the difference between higher and the lower states energy.
Due to the availability of numerous states of energy, an electron thus can undergo many transitions, each transition gives rise to a unique wavelength that encompasses the emission spectrum. The emission spectrum is thus formed by the frequencies obtained from these emitted light.
Based on the source, the emission spectrum is further categorized into;
a) Continuous Spectrum: When the spectrum has no breaks or openings between their wavelength range then this type of spectrum is termed as a continuous spectrum. For instance; A rainbow.
b) Line Spectrum: When the spectrum has a discrete or distinct line that is atoms emit light only at specific wavelengths with dark spaces between them, then this type of spectrum is termed as a line spectrum. For instance; Hydrogen line spectrum.
Absorption Spectra
This kind of spectrum is created by the frequencies of light that is transmitted with dark bands when energy is absorbed by the electrons generally in the ground state to reach higher energy level states. This is the kind of spectrum that is produced when atoms absorb energy.
When light from any source is passed through the chemical solution, then a pattern comprising of dark lines is observed. This pattern is further analysed with the help of a spectroscope.
The dark line pattern is seen precisely in the same place where coloured lines in the emission spectrum were observed. The spectrum hence attained is termed as the absorption spectrum.
Emission spectra, unlike absorption spectrum, emit all the colours in an electromagnetic spectrum, whereas few colours in the absorption spectrum may be absent due to the redirection of absorbed photons.
Absorption Spectroscopy
Absorption spectroscopy is referred to as a spectroscopic technique that is used for measuring the absorption of radiation when it interacts with the sample.
Absorption spectroscopy is linked to the absorption spectrum because the sample thus used interacts with photons produced from the radiating field.
Applications Absorption Spectroscopy
a) Chemical Analysis: The numerical nature of absorption spectroscopy makes it a perfect choice for chemical analysis. The absorption spectrum of the different compounds can be differentiated from one another using absorption spectroscopy. This is only possible because of the specificity nature of the absorption spectrum.
b) An application of absorption spectroscopy is an infrared gas analyser that is often used for detecting the pollutants present in the air or atmosphere. It also helps to distinguish between pollutants from nitrogen, oxygen etc.
c) Remote Sensing: Absorption spectroscopy is also used for analytical purposes such as for measuring the presence of dangerous elements.
Emission Spectra vs Absorption Spectra
The key difference between emission and absorption spectra is that an emission spectrum consists of different coloured lines, however, an absorption spectrum consists of dark-coloured lines in their spectrum. Other differences between absorption and emission spectrum are mentioned below;
Emission Spectra Absorption Spectra The emission spectrum is formed when atoms release energy Absorption Spectra is formed when atoms absorb energy Generally comprises of coloured lines in the spectrum Generally comprises of dark lines or gaps in its spectrum The type of photons emitted from the emission spectrum is used in estimating the different kind of elements from which a substance is made as each element emits different amount of energy and generally has a unique emission level. The wavelengths of light absorbed is used in calculating the number of substances that is present in the sample Absorption Spectra Citations
- Basic aspects of absorption and fluorescence spectroscopy and resonance energy transfer methods. Methods Cell Biol . 2008;84:213-42.
- UV-VIS absorption spectroscopy: Lambert-Beer reloaded. Spectrochim Acta A Mol Biomol Spectrosc . 2017 Feb 15;173:965-968.
- Absorption spectroscopy. Methods Enzymol . 2011;500:59-75.
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Balmer Series: Definition, Equation, and Examples
Continue ReadingBalmer Series
Balmer spectral series was initially noticed by Johann Balmer during the years 1885, Therefore the series is named after him that is the balmer series. Balmer series is exhibited when an electron shift takes place from higher energy states (that is ni =3,4,5,6,7,…) to lower energy states that are nf = 2 energy states.
The wavelength of the Balmer series generally falls in the visible part of the electromagnetic spectrum( having wavelength 400nm to 740nm) or is abbreviated as EM spectrum. In astronomy, the existence of Hydrogen is noticed by using the H-Alpha line of the Balmer series.
Balmer Series Formula
In the year 1885, based on experimental observations on the hydrogen atom, Balmer advised the formula for linking the wavenumber of the spectral lines emitted to the energy shells involved in an electron shift. This formula is given as:
λ = C(m2/m2-n2)
Here, λ represents the observed wavelength
C represents a constant (364.50682 nm)
n represents the lower energy level = 2, and m represents the higher energy level, which commonly has a value greater than 3.
The above observation was further refined by Johannes Rydberg, where R represents the Rydberg constant.
1/λ = R ((1/nf2) – (1/ni2))
According to the Balmer series, nf is always equal to 2. This equation was hence combined with the Bohr model to analyze the energy that is needed to shift an electron between its initial energy level and final energy levels.
ΔE = Rhc ((1/nf2) – (1/ni2))
Paschen Series (nf = 3): The series was initially noticed during the years 1908, by a German physicist named Friedrich Paschen. The Paschen series is exhibited when electron transition takes place from higher energy states (that is ni =4,5,6,7,8,…) to lower energy states that is nf =3 energy state. The wavelength of the Paschen series generally falls in the Infrared region of the electromagnetic spectrum.
Brackett Series (nf = 4): The series was initially noticed during the years 1922, by a famous American physicist Friedrich Sumner Brackett. Brackett series is exhibited when electron shift takes place from higher energy states (that is ni =5,6,7,8,9…) to lower energy state that is nf =4 energy state. The wavelength of the Brackett series generally falls in the Infrared region of the electromagnetic spectrum.
Pfund Series (nf = 5): The series was initially noticed during the years 1924, by a famous scientist named August Harman Pfund. Pfund series is exhibited when an electron shift takes place from higher energy states(that is ni = 6,7,8,9,10,…) to nf=5 energy state. The wavelength of the Pfund series usually falls in the Infrared region of the electromagnetic spectrum.
The Bohr Model
In 1913, Niels Bohr a famous scientist and chemist proposed a model for the hydrogen atom having the atomic number 1. He stated that the electrons present in an atom revolve around the nucleus in discrete paths called orbit or an energy shell. The electron while in its stationary or its rest state cannot produce energy, but can only absorb energy when it moves from one orbit to another orbit.
The quantum number, abbreviated as n is used to designate the different energy states. The lowest energy state is termed as the ground state, in which n is always equal to one. The excited states are further equal to 2, 3, 4, and so on. When the electron present at the ground state absorbs energy which is equivalent to the difference between the ground state of the electron and the second state the electron by absorbing a photon. The electron thus turns out to be more excited and displays transitions from the ground state to the n= 2 excited states.
According to Bohr, the potential energy (P.E) of an electron present in the nth level is measured by using the following equation mentioned below:
En = -(Rhc/n2)
where En represents the potential energy,
R represents the Rydberg constant which is equal to 1.0974 × 107 m-1
h represents Planck’s constant which is equal to 6.62607004 × 10-34 m2·kg/s),
c represents the speed of light (~ 3 × 108 m/s).
The electrons can also spontaneously return to the ground state or any other lower excited state. When this happens, then some amount of energy is produced can be depicted in the form of the emitted photon. The energy of the photon is thus always equal to the energy difference between the higher and lower energy states. Subsequently, different types of atoms have different energy levels and the light emitted from each transition varies for every atom.
Hydrogen Spectrum
As it is now known that electrons in an atom or a molecule absorb energy and become excited, and then they transfer themselves from a lower energy level to a higher energy level, and radiation or energy is released when they come back to their original ground states. This great phenomenon is also the same for the emission spectrum through hydrogen atom as well , and therefore it is termed as the hydrogen emission spectrum.
The Hydrogen Atom
The hydrogen atom is referred to as the simplest atomic system that is found in nature, therefore it produces the simplest of these series. When the beam of light or radiation is allowed to enter the device via a slit, then each distinct component of the light or radiation can be depicted in the form of images of the source. These images are pictured when resolved under the spectroscope. The images received will be in the form of parallel lines that are organized next to each other with consistent spacing. The lines seen will be apart in the higher wavelength side and then they come closer progressively when shifted from higher to lower wavelength side. The shortest wavelength will thus hold the least spaced spectral lines.
Rydberg Formula
The wavelengths of the spectral series are commonly calculated by using the Rydberg formula.
Scientifically, it can be expressed as-
1/λ = R ((1/n2f – (1/n2f))
Where,
• 𝜆 represents the wavelength
• R represents the Rydberg constant has the value 1.09737✕107 m-1
• Z represents the atomic number of an atom
• ni represents the lower energy level
• nf represents the higher energy level
Note: This equation mentioned above is valid only for Hydrogen and Hydrogen like elements.
Lyman series (nf =1)
This spectral series was projected during the years 1906-1914, by the famous scientist Theodore Lyman. According to Bohr’s model, the Lyman series is exhibited when electron shift takes place from higher energy states (that is ni = 2,3,4,5,6,…) to nf =1 energy state. The wavelength of the Lyman series generally falls in the Ultraviolet band.
Balmer Series Citations
- Measurements of aluminum and hydrogen microplasma. Appl Opt . 2007 Jul 1;46(19):4026-31.
- Measurement of the deuterium Balmer series line emission on EAST. Rev Sci Instrum . 2016 Nov;87(11):11D616.
- Balmer series Hbeta measurements in a laser-induced hydrogen plasma. Appl Opt . 2003 Oct 20;42(30):5992-6000.
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Electromagnetic Radiation: Definition, Properties, and Examples
Continue ReadingElectromagnetic Radiation
Electromagnetic radiation is a spectrum that consists of radio waves, microwaves, infrared waves, visible light, ultraviolet radiation, X-rays, and gamma rays.
Electromagnetic radiation has a particle as well as a wave nature therefore this makes it interesting to study its nature in quantum theory.
Electromagnetic Radiation Properties
• The oscillating charged particles creates an oscillating electric and magnetic fields which are perpendicular or right angles to each other and both are perpendicular to the direction of propagation of the wave as well.
• Electromagnetic waves generally do not require a medium which means that they can travel in a vacuum.
• There are numerous types of electromagnetic radiation, that differ from one another in terms of their wavelength or frequency.
• Electromagnetic radiation is commonly categorized based on several properties such as frequency, wavelength, amplitude etc.
Electromagnetic Radiation Formula
Frequency is referred to as the number of waves that passes through a given point in one second. A general equation that is related to the speed of light, frequency, and wavelength of electromagnetic radiation is as follows:
c = ν 𝝀
Where,
• c represents the speed of light
• ν represents the frequency of the electromagnetic wave
• 𝝀 or lambda represents the wavelength of the electromagnetic wave.
Apart from the above-mentioned parameters frequency and wavelength, some other factors are also used to classify electromagnetic radiation. One of these factors is the wavenumber. Scientifically, the wavenumber is equal to the reciprocal of the wavelength. It is represented in the SI unit as m.
ν = 1/Wavelength
Dual Behaviour of Electromagnetic Radiation
Electromagnetic Radiation was assumed to have a wave nature only thus with the help of wave nature we can clearly explain a phenomenon like interference and diffraction. But Wave nature of Electromagnetic Radiation was unable to describe few things such as Black body radiation & the photoelectric effect.
In 1900, Planck stated the quantum theory and was successful in explaining blackbody radiation. According to this theory, atoms or molecules release or absorb energy only in discrete amounts termed as quantum. Quantum is referred to as the smallest amount of energy that is absorbed or released in the form of electromagnetic radiation.
Further, Einstein explained the Photoelectric effect by using Planck Quantum theory. He proposed that when a photon falls on the surface of a metal, then the complete photon’s energy is transferred to the electron.
Now based on the above observations of both Planck Quantum theory & Einstein Theory of Photoelectric effect, it was found that Electromagnetic Radiation behaves like particles or photons as well. Now the particle nature was not much reliable with the known wave nature of light. Consequently, this caused striking confusion among the scientists. The only solution to this problem was to accept the dual nature of Electromagnetic Radiation.
Thus, electromagnetic radiation posses dual nature;
Wave Nature
Particle Nature
Particle Nature of Electromagnetic Radiation
i. Photoelectric Effect
The photoelectric effect is referred to the emission of electrons when electromagnetic radiation, such as light, hits a substance. Electrons emitted in this effect are termed as photoelectrons. Though, this phenomenon of photoelectric effect can be explained only by the particle nature of light, in which light can be pictured as a stream of particles of electromagnetic energy. These particles of light are termed as photons.
Photons are explained below;
• Photons are elementary particles. It is referred to as a quantum of light.
• The energy of a photon is , E = hf Where h represents Planck’s constant F represents wave frequency E represents photon energy
• A photon generally remains unaffected by electric and magnetic fields. Photon is electrically neutral in nature.
• A photon is massless that is it has zero mass.
• Photons, unlike atoms can be formed or destroyed when radiation is produced or absorbed.
ii. Black Body Radiation
When the black body is heated, it becomes red-hot. In simple words, it releases red coloured light. When the temperature is increased further, then the colour of the radiation emitted changes as follows, first from red to yellow then to white and lastly to purple as the temperature increases. This states that the wavelength of radiation produced by the black body decreases with a rise in temperature.
Wave Nature of Electromagnetic Radiation
Wave theory of radiation was unable to explain the phenomena of the photoelectric effect and also the black body radiation.
Major points of electromagnetic wave theory include;
The energy that is emitted from a source is in the form of radiation and is also termed as radiant energy. These radiations comprise of electric and magnetic fields which oscillate perpendicular (or at right angles ) to each other and also is perpendicular to the direction of propagation of radiation.
These radiations or electromagnetic radiations (or electromagnetic waves) travel with the velocity of light and also possesses wave character.
Characteristics of a wave:
• Wavelength: The wavelength of a wave is referred to as the distance between two consecutive crest or trough It is represented by a symbol lambda ( 𝝀 ) and is generally expressed in cm or m.
• Frequency: The frequency of a wave is referred to as a number of waves that passes through a point in one second It is represented by a symbol v (nu) and is generally expressed in its SI unit that is hertz or abbreviated as Hz.
• Velocity: The velocity of a wave is referred to as the linear distance which is travelled by a wave in one second It is represented by a symbol v and is commonly expressed in centimetre per second or metre per second (cm/sec or m/sec).
• Amplitude: The amplitude of a wave is referred to as the height of the crest as well as the depth of trough true It is generally represented by symbol a and is expressed in the metre , centimetre or the units of length.
• Wavenumber: Wavenumber is referred to as the number of waves that is present in 1 cm length.
Electromagnetic Radiation Citations
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Atomic Model: Definition, Properties, Types, and Examples
Continue ReadingAtomic Model
Atomic structure is defined as the structure of an atom containing a nucleus present in the center in which the protons or positively charged particles and neutrons (neutral) are present. The negatively charged particles are termed as electrons and they revolve around the nucleus.
In the 1880s, the first scientific theory of atomic structure was explained by John Dalton. A variety of different models have been evolved over the past decades to understand the functions of an atom. As a result, there are five basic atomic models which helped us to describe and comprehend the structure of the atom. Each of these models mentioned below had its own advantages and drawbacks.
The five atomic models that shaped the modern atomic theory are:
• John Dalton’s atomic model
• J.J. Thomson’s atomic model
• Ernest Rutherford’s atomic model
• Niels Bohr’s atomic model
• Quantum Numbers/model
I. Dalton Atomic Model
The English chemist and scientist named John Dalton stated that all matter is made up of atoms, which are undividable. He also proposed that all the atoms present in an element are the same, but the atoms of different elements generally differ in their size and mass. The following are the postulates of Dalton’s theory;
• All matter is made up of particles called atoms.
• Atoms are indivisible particles.
• Specific elements generally have only one type of atom present in them.
• Each atom has a constant mass respectively that differs from element to element.
• Atoms can neither be formed nor can be destroyed but can be transformed from one form to another.
Drawbacks of Dalton Atomic Model:
• The theory was not able to describe the existence of isotopes.
• Dalton’s atomic theory does not explain the existence of subatomic particles. Dalton’s atomic theory projected that the atoms were indivisible. However, the discovery of subatomic particles (for example, protons, electrons, and neutrons) disproved this postulate.
• Dalton’s atomic theory failed to explain isobars( two different elements having the same mass number. For Instance: 40Ar and 40Ca)
II. Thomson Atomic Model
The English chemist and scientist Sir Joseph John Thomson described the structure of the atom in the early 1900s.
He was awarded the Nobel prize later for the finding of “electrons”. His work is chiefly based on an experiment titled a cathode ray experiment. The working of this experiment is as follows:
Cathode Ray Experiment: It has a tube made of glass which further has two openings, one opening is for the vacuum pump and the other one is for the inlet through which a gas is pumped.
A high voltage electric current is passed through a discharge tube that contains gas at a very low pressure, a green glow is thus seen at the other end of the discharge tube. This green glow or fluorescence observed is the result of the rays which are released from the cathode towards the anode. These rays are termed as cathode rays.
Conclusions: Based on this observation from his cathode ray experiment, Thomson defined the atomic structure as a positively charged sphere that contains negatively charged particles called electrons.
It is usually stated as the “plum pudding model” because it can be pictured as a plum pudding dish where the pudding represents the positively charged atom and the plum pieces in it represent the electrons.
Drawbacks of Thomson’s Atomic Model: The theory did not mention anything about the nucleus present in the center of an atom.
III. Rutherford Atomic Model
Rutherford, a famous scientist revised the structure of an atom with the discovery of another subatomic particle named as a Nucleus. His atomic model is built on the Alpha ray scattering experiment.
Alpha Ray Scattering Experiment Structure:
• Rutherford took a gold foil as he wanted a fragile layer.
• In this experiment, fast-moving alpha particles were bombarded on a thin gold foil.
• Alpha particles are referred to the helium ions with a +2 charge and thus have a significant amount of energy.
• Rutherford predicted that the alpha particles would pass through the gold foil but some of the particles deflected and striked the fluorescent screen.
Conclusions:
• Since most of the rays passed straight through the gold foil, Rutherford thus observed that most of the space inside the atom is vacant or empty.
• Few rays which got reflected is because of the repulsion with some other positive charge present inside the atom.
• 1/1000th of rays got forcefully deflected because of the presence of a very strong positive charge confined in the center of the atom. He named this strong positive charge as “nucleus”.
• He stated that most of the charge and the mass of the atom is present in the center (Nucleus).
Rutherford’s Structure of Atom Based on the above comments and assumptions, Rutherford projected his own atomic structure which is;
• The nucleus is present at the center of an atom, generally where most of the charge and mass of the atom is concentrated.
• Electrons revolve around the nucleus (which is present in the center) in circular paths called orbits, just like the planets revolve around the sun.
Limitations of Rutherford Atomic Model:
• If electrons present in an atom revolve around the nucleus, then they have to spend energy, as a result, a lot of energy will be spent by the electrons, and ultimately, electrons will lose all their energy and will fall into the nucleus thus Rutherford was unable to explain the stability of atoms.
IV. Bohr Atomic Model
Neils Bohr model is the most widely used atomic model which helped to define the atomic structure of an element that is built on Planck’s theory of quantization. Bohr’s Postulates:
• The electrons present inside the atoms are positioned in discrete orbits termed as “stationary orbits”.
• The energy levels of these shells can be denoted as quantum numbers.
• Electrons can travel to higher levels by absorbing energy and can also move to lower energy levels by emitting or releasing their energy.
• When an electron stays in its rest or stationary form, then there will be no absorption or release of energy.
Drawbacks of Bohr’s Atomic Model:
• Bohr’s atomic structure is applicable only for single-electron species. For instance; H, He+, etc.
• Bohr’s theory was unable to explain both Stark and Zeeman’s effects.
V. Quantum Numbers
• Principal Quantum number (n): It represents the orbital number or the shell number of the electron.
• Azimuthal Quantum numbers (l): It represents the orbital (sub-orbit) of the electron.
• Magnetic Quantum number: It represents the number of energy states present in each orbit.
• Spin Quantum number(s): It represents the direction of spin, that is when S = -½ the spin of an electron is Anticlockwise and ½ then the spin of an electron is Clockwise.
Electronic Configuration of an Atom:
The electrons are usually filled in the s, p, d, f orbits as per the following rule; 1.
Aufbau’s principle: The filling of electrons must take place by following the ascending order of the energy orbitals that is;
• Initially, the lower energy orbital should be filled and then the higher energy levels.
• Ascending order of energy orbitals is as follows 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on. 2.
Pauli’s exclusion principle: this principle states that no two electrons can have all the four quantum numbers mentioned above to be the same or similar.
NOTE: If two electrons are positioned in the same energy state then they should be positioned with opposite spines.
Atomic Model Citations
- Nagaoka’s atomic model and hyperfine interactions. Proc Jpn Acad Ser B Phys Biol Sci . 2016;92(4):121-34.
- Dalton’s disputed nitric oxide experiments and the origins of his atomic theory. Chemphyschem . 2008 Jan 11;9(1):106-10.
- Atomic Theory and Multiple Combining Proportions: The Search for Whole Number Ratios. Ann Sci . 2015 Apr;72(2):153-69.
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Millikan Oil Drop Experiment: Definition, and Examples
Continue ReadingMillikan Oil Drop Experiment Definition
In this experiment, Millikan permitted the charged tiny droplets of oil to pass through a hole into an electric field. Then, by changing the strength of the electric field, the charge over an oil droplet was calculated, which always results as an integral value of ‘e.’
What is Millikan Oil Drop Experiment?
Millikan oil drop experiment is accomplished by Millikan and Harvey Fletcher in 1909 to measure the charge of an electron. This experiment proved to be helpful in the physics community.
Robert Andrews Millikan was a famous American physicist and was awarded the Nobel Prize for Physics in 1923 for his work on the elementary electronic charge and the photoelectric effect.
In 1909, Millikan performed a series of experiments to find the electric charge which is carried by an electron. He initially began his experiment by measuring the path of charged water droplets in an electric field. The results projected that the charge present on the droplets is a simple multiple of the basic electric charge, but the test’s result was not precise enough to be considerable.
To obtain more specific results, in 1910 he performed his famous oil-drop experiment in which he substituted water (which tends to evaporate quickly) with oil, and this experiment is further explained below;
Millikan Oil Drop Experiment Apparatus
The apparatus for the experiment was created by Millikan and Fletcher. It is consists of two metal plates that are held at some distance by an insulated rod. Four holes were made in the plate, out of which three were only allowed to pass the light through them and the fourth one is used to view through the microscope.
Ordinary oil was not used for this experiment as it tends to evaporate by the heat of the light and therefore an error could be caused in Millikan Oil Drop Experiment. So, the oil having low vapor pressure was used, the same that is used in a vacuum apparatus.
• A specific type of oil as mentioned above is sprayed into the chamber, where the drops attain electrical charge.
• The droplets are then allowed to enter the space present between the plates and, as they were charged, they could be effortlessly controlled by altering the voltage across the plates.
• Mainly, the oil drops were allowed to fall between the plates having no electric field. They then rapidly reached terminal velocity due to the presence of friction of the air in the chamber.
• The field was then turned on and it was huge enough, thus some of the drops started to rise. This is because of the presence of upwards electric force, (FE) on them which is greater than the downwards gravitational force, g.
• Millikan’s experiment was actually meant to have the drops to fall at a constant rate. At this constant rate, the gravitational force present on the drop and the force of the electric field or upwards electric force on the drop is equal.
• Millikan then repeated this same experiment for over 150 oil drops out of which he selected 58 oil drops results and then with the help of these observations he determined the highest common factor.
Millikan Oil Drop Experiment Calculation
As mentioned above,
Fup = Fdown
Fup = Q. E
Fdown = m.g
Where, Q represents an electron’s charge
E represents the electric field
m represents the droplet’s mass, and
g represents gravity.
Q⋅E = mg
Therefore,
Q = mg /E
It can be said that an electron charge is measured by Millikan. Millikan stated that all drops had charges that were equal to 1.6x 10-19 C multiples.
Millikan Oil Drop Experiment Conclusion
The charge present on an oil droplet is always equal to an integral value of e (1.6 x 10-19). Hence, the assumption of Millikan’s Oil Drop Experiment displays that the charge is quantized, that is the charge present on any particle is always be an integral multiple of e.
Millikan’s oil drop experiment was a vibrant demonstration of the quantization of charge.
The experiment has since been commonly conducted by many physics undergraduates, though it is quite expensive and to get the precise result is quite difficult.
Millikan Oil Drop Experiment Importance
Millikan’s experiment is quite crucial to study because it establishes the charge over an electron.
Millikan used a simple device in which he adjusted the actions of electric, gravitational, and air drag forces.
With the help of the apparatus, he was successful in estimating the charge on an electron that is equal to 1.60 × 10-19 C.
Reason for Using Oil Drops
Oil drops are used in Millikan oil-drop experiment because oil drops usually retain their mass over some time when it is exposed to higher temperatures. Likewise, he used an atomizer for ultra-fine droplets. Therefore, he preferred oils over water because water changes its state or form at much higher temperatures.
Furthermore, it is extensively known that oil tends to retain the exact volume, omass, and weight. This property of oil enabled him to have a precise measurement of the charge. Other liquids present in nature may separate, disintegrate or evaporate.
Millikan Oil Drop Experiment Citations
- Millikan recharged. Nature Physics volume 8, page106 (2012)
- Millikan “oil drop” stabilized by growth. Science . 1979 Jan 26;203(4378):353-4.
- Performing the Millikan experiment at the molecular scale: Determination of atomic Millikan-Thomson charges by computationally measuring atomic forces. J Chem Phys . 2017 Oct 28;147(16):161726.
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Electromagnetic Waves: Definition, Properties, and Examples
Continue ReadingElectromagnetic Waves Definition
Electromagnetic radiation involves electromagnetic waves, which are coordinated oscillations of both electric and magnetic fields. It can further be said that electromagnetic waves are the composition of oscillating electric and magnetic fields.
Electromagnetic radiation or electromagnetic waves are produced due to periodic change of electric or magnetic field. In a vacuum or void, electromagnetic waves generally travel with the speed of light which is represented as c.
1. The position of an electromagnetic waves present within the electromagnetic spectrum can be categorized by its frequency of oscillation or wavelength.
2. Some sources of electromagnet radiation include; the cosmos (for instance – the sun and stars), radioactive elements, and manufactured devices.
3. EM displays a dual wave and particle nature.
What is Electromagnetic Waves?
Electromagnetic radiations are produced when an atomic particle, such as an electron, is accelerated and is moved by an electric field. This movement produces oscillating electric and magnetic fields, which are at right angles or 90 degrees to each other and usually travel in a bundle of light energy termed as photons.
A wavelength is defined as the distance between two consecutive troughs (peak) of a wave. This distance is commonly measured in meters (m).
Frequency is referred to as the number of waves formed in a given period of time. It is generally measured in hertz (Hz).
Mathematical Representation of Electromagnetic Wave
In the electromagnetic wave, E is referred to as the electric field vector and B is defined as the magnetic field vector.
Therefore, the direction of propagation of the electromagnetic wave is thus given by vector cross product of the electric field and magnetic field. It is shown below: E→×B→
Graphical Representation of Electromagnetic Waves
Electromagnetic waves are commonly shown by a sinusoidal graph. It comprises of time-varying electric and magnetic fields which are perpendicular or at right angles to each other and are also perpendicular (right angles) to the direction of propagation of waves. The uppermost point of the wave is named as crest whereas the lowest point is termed as a trough. The waves travel at a fixed velocity of 3 x 108 m.s-1 in vacuum.
The Electromagnetic Spectrum
EM radiation spans a widespread range of wavelengths and frequencies. This range is termed as the electromagnetic spectrum. The EM spectrum is usually divided into seven regions, based on decreasing wavelength and increasing energy and frequency. The common terms are: radio waves, microwaves, infrared (IR), ultraviolet (UV), X-rays, gamma rays and visible light. Normally, lower-energy radiation such as radio waves, is represented in frequency and microwaves, infrared, visible and UV light are generally expressed as wavelength and higher-energy radiation, like X-rays and gamma rays are stated in energy per photon.
Radio Waves
Radio waves has the lowest range of the EM spectrum, with frequencies equal to 30 gigahertz (GHz), and wavelengths of 10 millimetres or 0.4 inches.
Uses: Radio waves are used mainly for communications including voice, data, and entertainment media.
Microwaves
Microwaves lie in between the EM spectrum of radio and infrared waves. They have frequencies from about 3 GHz up to about 30 trillion hertz, or 30 terahertz (THz), and wavelengths of 10 mm (0.4 inches) to 100 micrometres (μm) or 0.004 inches.
Uses: Microwaves are commonly used for high-bandwidth communications, radar and is also used as a heat source for microwave ovens and industrial applications.
Infrared
Infrared is in the range of the EM spectrum that lies between microwaves and visible light. IR express frequencies from 30 THz up to about 400 THz and wavelengths of about 100 μm (0.004 inches) to 740 nanometres (nm), or 0.00003 inches.
Uses: IR light is generally invisible to naked eyes, but if the intensity is enough, then it can be felt as heat.
Visible Light
Visible light lies in the middle of the infrared and ultraviolet waves of EM spectrum. It displays the frequencies of around 400 THz to 800 THz and wavelengths of nearly 740 nm (0.00003 inches) to 380 nm (.000015 inches).
Uses: visible light can be referred to as the wavelengths which are visible to naked eyes.
Ultraviolet
Ultraviolet light lies in the range of the EM spectrum between visible light and X-rays. Its frequencies are around 8 × 1014 to 3 × 1016 Hz and wavelengths of about 380 nm (.000015 inches) to about 10 nm (0.0000004 inches).
Uses: UV light is referred to as a constituent of sunlight; though, it is not visible to the naked eye. These radiations further have several medical and industrial applications.
• These rays are germicidal in nature, kills bacteria, viruses and moulds present in the air, water and on surfaces.
• It is also used to spot phony bank notes, as these fakes notes turn fluorescent in colour under UV light whereas real notes don’t turn fluorescent under the UV light.
X-rays
X-rays are widely classified into two types, soft X-rays, and hard X-rays. Soft X-rays have the range that lie between UV and gamma rays. Soft X-rays have frequencies of about 3 × 1016 to 1018 Hz and wavelengths of about 10 nm (4 × 10−7 inches) to 100 picometers (pm), or 4 × 10−8 inches. Hard X-rays further reside in the same region of the EM spectrum as gamma rays. The only difference between them is; X-rays are produced by accelerating electrons, whereas gamma rays are formed by nuclei of atom.
Uses: The most vital use of X-rays is that it is used to detect bone fracture.
Gamma-rays
Gamma-rays generally have frequencies greater than 1018 Hz and wavelengths less than 100 pm (4 × 10−9 inches).
Uses: Gamma radiation damage the living cells and tissues, which is beneficial for killing of cancer cells in small doses. But it is tremendously hazardous to humans in uncontrolled amount.
Electromagnetic Waves Citations
- Benefits and hazards of electromagnetic waves, telecommunication, physical and biomedical: a review. Eur Rev Med Pharmacol Sci . 2019 Apr;23(7):3121-3128.
- Electromagnetic radiation. Br Med Bull . 2003;68:157-65.
- Effect of electromagnetic waves on human reproduction. Ann Agric Environ Med . 2017 Mar 31;24(1):13-18.
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Electron Charge: Definition, Properties, and Examples
Continue ReadingElectron Charge
An electron is referred to as a subatomic (smaller than an atom) particle that carries a single unit of negative electricity, denoted as – e. Moreover, the charge of the electron is always equal to the magnitude of the elementary charge (e) which is equal to 1.6 x 10-19 C but bearing a negative sign. Thus,
Charge on the electron (- e) = – 1.6021 x 10-19 Coulomb
Coulomb: In the International System of Units or abbreviated as SI units, one column is defined as the amount of electric charge carried by a current of 1 ampere flowing for 1 second.
Symbol: C
The charge on a single electron (e) is referred to as the unit electrical charge. It is stated as a negative polarity. The charge on an electron is equal, though, opposite to the positive charge on a proton (positively charged particles).
Charge of Electron in eV
The electron volt is the unit of energy commonly used in atomic and nuclear physics.
1 Electron volt (eV) = 1.6 x 10-12 erg
Here the value of 1 erg = 10-7 Joule
So, the charge of the electron in the electron volt = 1.6021 x 10-19 Joule
Mass of Electron
The mass of an electron is equal to 0.000548 amu or atomic mass unit when they orbit the nucleus and also have a charge of -1. The invariant mass of the electron is written as, m (-e) and is equal to 9.1 x 10-31 kg Here, the invariant mass is referred to as the mass of a stationary electron. The mass of an electron in amu or atomic mass units is given by, 5.489×10-4 amu Mass of an electron in eV is given by 0.511 MeV.
Electrical Charges
The steady flow of electrons is termed as current. Current is generally what flows through electrical wires and electronics items, from light bulbs to televisions in our house.
The charge q of a body is given by, q = ne
Where n= number of electrons transferred
e is equal to the charge on each electron
The SI unit of charge is written as coulomb or C.
The CGS (centimeter-gram-second) unit of charge is equal to 1 e.s.u or electrostatic unit of charge.
Properties of Electric Charges
Electric charge, like mass and volume, is referred to as the fundamental physical property of matter.
SI unit of electric charge is equal to Coulomb (C), which represents 6.242×1018e, where e is equal to the charge of a proton. Electric charges can either be positive or negative; a singular proton has a charge equal to 1.602×10−19 C, whereas an electron has a charge equal to -1.602×10−19 C.
The electric charge is generally measured in Coulombs (C). Charges can either be positive or negative.
Electric charge (unlike mass) is independent of speed.
Atomic Electrical Charges
Atoms are referred to as the fundamental building blocks of matter and contain three types of particles mentioned below:
Protons
Neutrons
Electrons
Of these three subatomic particle types mentioned above, two particles mainly protons and electrons carry a net electric charge, whereas neutrons are neutral and carry no charge.
Both protons and electrons have a charge that is generally quantized. That is, the magnitude of their respective charges are equal to each other that is 1. This standard value is thus equal to about 1.6×10-19 Coulombs.
Protons
Protons are present with the neutrons in the center of an atom or nucleus. Protons generally have a charge of +1 and a mass of 1 atomic mass unit or amu, which is nearly equal to 1.66×10-24 grams. The number of protons present in an atom states the identity of that particular element ( for instance, an atom carrying 1 proton is hydrogen, and an atom having two protons is helium).
Electrons
Electrons is a negatively charged subatomic particles commonly found in the periphery of the atom and have a charge equal to -1. They are much smaller as compared to protons; their mass is 1183611836 atomic mass units or amu. Electrons present in an atom move around in the space outside the nucleus like a cloud. The negatively charged electronic cloud specifies the area of the space where electrons are likely to be found in an atom.
Properties of Electrons
Electrons can either be free (that is not attached to any atom) or may be found bound to an atom’s nucleus. The charge on a single electron is generally considered as the unit electrical charge. The mass of an electron at rest is represented as me and is approximately 9.11 x 10-31 kilogram (kg).
• Electrons are accountable for the negatively charged component present in an atom.
• Electrons are commonly attracted to positively charged particles termed as protons.
Neutrons
Neutrons are referred to as a subatomic particles that are one of the chief constituents present in the atomic nuclei. They are generally represented by the symbol n or no. Neutrons do not have any net electric charge linked with them. However, they do have a mass that is slightly greater in magnitude as compared to a proton. Neutrons and protons together are referred to as nucleons, because they usually behave in a similar manner inside the atomic nuclei of atoms. The atomic mass of a neutron can be approximated to one atomic mass unit.
Ions
In the ground state, an atom will always have an equal number of protons and electrons, and therefore will have a net charge equal to 0. Though the electrons can be shifted from one atom to another, it is possible for atoms to develop a charge. Atoms present in such a state are named as ions.
Nucleus
The nucleus of an atom consists of positively charged particles termed as protons and neutral particles named as neutrons.
Electron Charge Citations
- Measuring the Electron’s Charge and the Fine-Structure Constant by Counting Electrons on a Capacitor. J Res Natl Inst Stand Technol . Mar-Apr 1992;97(2):299-304.
- Experimental charge density from electron microscopic maps. Protein Sci . 2017 Aug;26(8):1619-1626.
- Chemical bonding in view of electron charge density and kinetic energy density descriptors. J Comput Chem . 2009 May;30(7):1093-102.
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Oxidation: Definition, Classification, and Examples
Continue ReadingOxidation Definition
Oxidation can be defined as the gaining of an oxygen atom and its reverse is losing an oxygen atom is called as Reduction. Thus, the process in which electrons are lost in a reaction, is called as Oxidation. The reaction in which electrons are gained is called as Reduction.
Oxidation Example
For example, in this reaction, CuO+Mg →MgO+Cu , electrons are gained by copper and lost by magnesium.
Another example could be that of hydrofluoric acid where fluorine is reduced and hydrogen is oxidized, H2+F2→2HF. Thus, to make hydrofluoric gas, hydrogen and fluorine are required.
i). Using oxygen to carry out oxidation: Oxidation using oxygen has been previously done and it is known that the first oxidizing agent at that time was oxygen. However, oxygen is added when there is loss of electron. An example of oxidation is 2Fe+O2 → Fe2O3 In this case oxidation of iron leads to the formation of iron oxide, also known as rust.
ii). Using hydrogen to carry out oxidation: This is the opposite of the oxidation of oxygen. In oxidation there is gain in oxygen and loss of hydrogen, whereas in reduction, there is loss of oxygen and gain in hydrogen atoms. An example is CH3CH2OH → CH3CHO. There is loss of hydrogen as ethanol gets oxidized.
iii). Oil rig chemistry: Oxidation and reduction can be understood with oil rig. “Oil” means oxidation is loss and “rig” means reduction of electrons. It is concerned with electrons only not hydrogen or oxygen. For example, if hydrogen has lost a electron, it has been oxidized (H+). If chlorine gains an electron, it is called as reduced (Cl–). This is called as the oil rig concept.
Oxidation Process
Oxidation and reduction can take place on their own and does not happen hand in hand, but are considered as half- half reactions which completes to form a complete reaction. These reactions tend to possess a charge and the charge needs to be balanced in a reaction as electrons are lost or added in a reaction. Thu, forming the overall net charge.
Oxidizing and Reducing Agents
In a redox reaction, electrons are passed to the oxidant by the one getting reduced. Thus, in such cases, lost electrons will be oxidized by reducing agent and electrons gained will be reduced by the oxidizing agent.
Oxidizers: those substances are called oxidizers which loses electrons or donates electron to another compound is called as Oxidizing agent. Thus, they themselves get reducing by losing an electron. They can be referred to as electron acceptor. Example are O2, F2 of elements such as H2O2 and MnO–.
Reducers: They reduce other compounds by gaining electrons, thus can be termed as reducers or reducing agents. They are referred to as electron donor. Example are iron, magnesium, sodium and others. Reducing agents oxidize themselves by donating electrons to other compounds.
Oxidation Reduction Reaction (Redox Reaction)
A reaction is said to be redox if oxidation and reduction both of them take place simultaneously in a reaction. Example are metal corrosion, photosynthesis, respiration, browning of fruits, burning and others are the example of redox reaction.
Classification of Redox Reaction
Three types of redox reactions takes place and they are:
i). Oxygen Atom Transfer: In this reaction, carbon reacts with mercury oxide to form carbon dioxide and mercury. The charge on mercury oxide was +2 whereas now that charge has moved to carbon dioxide. Thus, carbon is oxidized.
C + 2HgO → CO2 + Hg
ii). Hydrogen Atom Transfer: In this case, hydrogen atoms are removed from hydrazine and transferred to oxygen molecule. Thus, oxygen gets reduced and hydrazine loses hydrogen atoms and thus gets oxidized.
N2H4+O2 →N2+2H2O
iii). Electron Transfer: Zinc and copper react with each other in a solution to form zinc ion and copper metal. Thus, copper gets reduced while zinc gets oxidized. The +2 net charge on copper is transferred to zinc ion.
Stoichiometric Basis
There is no information available on what is the mechanism behind these redox reactions and thus it is said that stoichiometry gives the molecules their characteristics. Therefore, there are stoichiometric redox reaction which contain hydrogen, oxygen and electron transfer.
Oxidation State Change
A few theories have been made on oxidation and reduction. It states that each atom possess a nucleus which has positive charge and electrons which are negatively charged surrounding the nucleus. Thus, maintaining the balance. Thus, when a molecule gains or donates an electron, it obtains an oxidation number showing the electron number which can bind to the other molecules. When a molecules undergoes oxidation or reduction its oxidation state is determined. For example, Fe3+ is the state in which oxygen is present. Thus, +3 is the oxidation state of Fe.
Examples of Oxidation
The only oxidizer which can oxidize all the metals as well as non-metals is oxygen, which forms oxides.
S + O2 → SO2
4Li + 02 → 2Li2O
These are some of the example of oxygen forming oxides.
Oxidation in Chemistry
Metal displacement is a brilliant example of oxidation in chemistry. Metal displacement occurs when one metal occupies the position of other metal. For example,
Zn + CuSO4 → ZnSO4 + Cu
In this reaction, zinc reacts with copper sulphate to form zinc sulphate. It is a redox reaction, where copper get positioned at the place of zinc and copper metal are freed.
The ionic equation can be written as:
Zn + Cu2+ → Zn2+ + Cu
The two half reactions are:
In this reaction copper gets reduced, Cu2+ + 2e- → Cu
In this reaction, zinc gets oxidized, Zn →Zn2+ +2e-
Another example is denitrification, where nitrate is reduced to nitrogen. The reaction is:
2NO3– +10e- +12H+ → N2 + 6H2O
Hydrocarbon oxidation in presence of oxygen will form water and then lead to the formation of alcohol, carboxylic acid, ketone and later form peroxide.
Oxidation in Biology
Redox reaction is important for various biological processes such as photosynthesis and anaerobic respiration. Example of anaerobic oxygen is when glucose is oxidized to carbon dioxide and oxygen gets reduced.
C6H12O6 + 6O2 → 6CO2 + 6H2
For cell respiration the reaction is;
6CO2 + 6H2O + LIGHT ENERGY → C6H12O6 + 6O2
During photosynthesis, to form sugar, carbon dioxide gets reduced to form oxygen. The opposite of photosynthesis results in the formation of carbon dioxide and water. For reduction of NAD+ to NADH, carbon which is reduced can be used.
Oxidation in Geology
The role of redox reaction in geology is for mineral deposition, production and for mobilization of minerals. Whenever the rocks are about to undergo oxidation, their color changes to red. Rock will look white in color or green when, some liquid oozes out of it, which could be uranium. The deposition of such compounds on rocks and marbles are redox reaction examples.
Oxidation Citations
- Bio-inspired Nonheme Iron Oxidation Catalysis: Involvement of Oxoiron(V) Oxidants in Cleaving Strong C-H Bonds. Angew Chem Int Ed Engl . 2020 May 4;59(19):7332-7349.
- Hydroxylamine driven advanced oxidation processes for water treatment: A review. Chemosphere . 2021 Jan;262:128390.
- Recent Advances in Copper Catalyzed Alcohol Oxidation in Homogeneous Medium. Molecules . 2020 Feb 9;25(3):748.
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