CHAPTER 02 > Structure of Atom 27
CHAPTER > 02
Structure of Atom
KEY NOTES
Å The first atomic theory proposed by Dalton in 1808,
regarded the atom as the ultimate indivisible particle of Atomic Models
matter. Thomson’s Model of Atom
Å A large number of subatomic particles have been discovered Å Thomson in 1898 assumed that an atom is a sphere of positive
but only electron, proton and neutron are of great charged uniformly distributed with the electrons scattered as
importance among them and hence are called fundamental points throughout the sphere. This was also known as plum
particles. pudding, raisin pudding or watermelon model. An important
feature of this model is that the mass of atom is assumed to be
Discovery of Subatomic Particles uniformly distributed over the atom.
Å Cathode rays or electrons were discovered by J.J. Thomson Å Henri Becqueral observed that there are certain elements
in 1897, by utilising Faraday’s study of electrical discharge which emit radiation on their own. These elements were
in partially evacuated tube known as cathode rays tubes. named as radioactive elements and the phenomenon is called
Å The specific charge is the ratio of charge to mass of an radioactivity.
electron, i.e. e /me ratio. Å Rutherford’s nuclear model of atom Rutherford bombarded
Å By carrying out accurate measurements, Thomson was able very thin gold foil with α-particle.
to determine the value of e /me as 1758820
. × 1011 C kg − 1 . Å Thomson’s model of atom was proved wrong by Rutherford’s
Å Millikan devised a method known as oil drop experiment alpha-particles scattering experiment carried in 1909.
to determine the charge on the electron to be − 1.6 × 10− 19 C. Å The main features of this model are :
Å The mass of electron was determined by combining — Most of the particles passed the foil undeflected, which
Millikan’s and Thomson’s value of e /me ratio which comes indicated that most of the space in atom is empty.
out to be 91094
. × 10− 31 kg. — Some of them were deflected, but only at small angles. This
shows that there is something positively charged at the
Å Canal rays (or anode rays or positive rays) were
centre.
discovered by Goldstein. These rays consist of positively
charged particles called protons.
— Few particles were deflected at large angles. It means that in
the atom, mass and positive charge is centrally located in
Å Unlike cathode rays, the e /me value of canal rays depends extremely small region called nucleus. The nucleus is
upon the nature of gas taken in the tube. surrounded by electrons that move around the nucleus with
Å Neutrons are neutral particles and discovered by a very high speed in circular paths called orbits.
Chadwick. These are the heaviest particles of the atom. — Electrons and the nucleus are held together by electrostatic
Å The discovery of sub-atomic particles led to the proposal of force of attraction.
various atomic models to explain the structure of atom.
,Atomic and Mass Number Particle Nature of Electromagnetic
Å Atomic number ( Z) = number of protons in the Radiation : Planck’s Quantum Theory
nucleus of an atom = number of electrons in a neutral
Å The ideal body, which emits and absorbs radiations of all
atom
frequencies is called a black body and the radiation emitted by
Å The protons and neutrons present in the nucleus are such a body is called black body radiation.
collectively known as nucleons. The total number of
Å According to Planck’s quantum theory, the radiant energy
nucleons is termed as mass number of the atom.
which is emitted or absorbed in the atom of small discrete
Mass number ( A) = number of protons ( Z) + number packets of energy known as quantum and in case of light, the
of neutrons (n). quantum of energy is called photons.
c
Isobars and Isotopes E = hν or E = h
λ
Å Isobars are the atoms with same mass number, but
where, h = Planck’s constant = 6.63 × 10− 34 Js
different atomic number for, e.g. 14 14
6 C and 7 N.
E = Energy of photon or quantum
Å The species with same atomic number but different
Å If n is the number of quanta of a particular frequency and ET be
mass number are called isotopes, e.g. 6 C12 and 6 C14 .
the total energy, then
Å Hydrogen has three isotopes protium (11 H, only one ET = nhν
proton), deuterium (12 D, one proton and one neutron)
Photoelectric Effect
and tritium (13 T, one proton and two neutrons). Å The phenomenon of ejection of electrons from a metal surface
Drawbacks of Rutherford’s Model when a light of certain frequency strikes on its surface is called
photoelectric effect.
Rutherford’s model was failed to account for the
stability of the atom. Also, it did not explain about the Å For each metal, there is a characteristic minimum frequency,
electronic structure of atoms. known as threshold frequency ( ν 0) below, which photoelectric
effect is not observed.
Development Leading to the Bohr’s Å When a photon of sufficient energy strikes an electron in the
Model of Atom atom of the metal, it transfers its energy instantaneously to the
Å Nature of electromagnetic radiation and experimental electron during the collision and the electron is ejected.
results regarding atomic spectra play an important Å The number of electrons ejected is proportional to the intensity
role in the development of Bohr’s model. or brightness of light.
Å Light, X-rays and γ-rays are the examples of radiant Å Following the conservation of energy principle, the kinetic
energy. energy of ejected electron is given by the equation,
Maxwell in 1856 showed that radiant energy has wave 1
Å
hν = h ν 0 + m e ν 2
properties and called them electromagnetic waves or 2
electromagnetic radiations. Å Dual nature of electromagnetic radiation Light possesses both
Å There are many types of electromagnetic radiations particle and wave like properties. Whenever radiation interacts
which differ from one another in wavelength or with matter, it displays particles, like properties in contrast to
frequency. These constitute electromagnetic the wave like properties.
spectrum.
Atomic Spectra
Å The small portion in the electromagnetic spectrum
Å The pictorial representation of arrangement of various types of
around 1015 Hz is called visible light.
EMR in their increasing order of wavelength (or decreasing
Å All these radiations travel with the speed of light and order of frequency) is known as spectrum.
do not require any medium for their propagation or Å The spectrum of white light ranges from violet at 7.50 × 1014 Hz
transmission.
to red at 4 × 1014 Hz, such a spectrum is called continuous
Å The frequency ( ν), wavelength ( λ ) and velocity of light
spectrum.
(c) are related by the equation,
Å The spectrum of radiation emitted by a substance that has
c = νλ
absorbed energy in increasing order of wavelengths or
Å Wave number is defined as the number of decreasing frequencies is called as an emission spectrum.
wavelengths per unit length. Its commonly used unit Atoms, molecules or ions that have absorbed radiation are said
is cm − 1 . to be ‘excited’.
KEY NOTES
, Å An absorption spectrum is like the photographic negative — The frequency of radiation absorbed or emitted when
of an emission spectrum. transition occurs between two stationary states that
Å The study of emission or absorption spectra is referred to differ in energy by ∆E is given by :
as spectroscopy. ∆ E E2 − E1
v= =
Å The emission spectra of gas phase do not show a h h
continuous spread of wavelength from red to violet, rather where, E1 and E2 are the energies of the lower and higher
they emit light only at specific wavelengths with dark allowed energy states respectively. This expression is
spaces between them. Such spectra are called line spectra commonly known as Bohr’s frequency rule.
or atomic spectra. — The angular momentum of an electron is quantised. In a
Å Line emission spectra are of great interest in the study of given stationary state, it can be expressed as,
electronic structure and also used in chemical analysis to h
mevr = n ⋅ , n = 1, 2, 3 …
identify unknown atoms. 2π
Å When an electric discharge is passed through gaseous — The stationary states for electron are numbered,
hydrogen, the H2 molecules dissociate and the n = 1, 2, 3... . These integral numbers are known as
energetically excited hydrogen atoms produced emit principle quantum number.
electromagnetic radiation of discrete frequencies. — The radii of the stationary states are expressed as :
Å The line spectra of hydrogen lies in three regions of
electromagnetic spectrum, i.e. infrared, visible and rn = n2 a0
UV-region. where, a0 = 52.9 pm.
Å The set of lines in the visible region is known as Balmer Thus, the radius of the first stationary state called the
series. Bohr orbit is 52.9 pm.
Å The Swedish spectroscopist, Johannes Rydberg noted that — Energy of an electron in nth state is given as,
all series of lines in the hydrogen spectrum could be 1
described by the following formula. En = − R H 2
n
1 1
ν = 109677 2 − 2 cm− 1 n = 1, 2, 3 …
n1 n2
where, R H is called Rydberg constant and its value is
where, n1 = 1, 2 … equal to 2.18 × 10− 18 J.
n2 = n1 + 1, n1 + 2 … Å Bohr’s theory can also be applied to the ions containing
Å The value 109677 cm − 1 is called the Rydberg constant. only one electron, for example He+ , Li 2 + , Be3 + and so on.
Å The first five series of lines that correspond to Å The energy of the stationary states associated with these
n1 = 1, 2, 3 , 4, 5 are known Lyman, Balmer, Paschen, hydrogen like species is given by,
Bracket and Pfund series respectively. Z2
En = − 2.18 × 10− 18 2 J
Series n1 n2 Spectral Regions n
Lyman 1 2, 3, 4K Ultraviolet and radii by the expression,
n2
Balmer 2 3, 4, 5K Visible rn = 52.9 pm
Z
Paschen 3 4, 5, 6K Infrared
Line Spectrum of Hydrogen
Brackett 4 5, 6, 7K Infrared Å The energy gap between the two orbits (higher and lower
Pfund 5 6, 7K Infrared orbits) is given by equation,
1 1
∆E = 2.18 × 10− 18 J 2 − 2
Bohr’s Model for Hydrogen Atom ni n f
Å The main postulates for Bohr’s model are :
Å The frequency associated with the absorption and emission
— The electron in the hydrogen atom can move around the of the photon can be evaluated by using equation,
nucleus in a circular path of fixed radius and energy.
∆E R H 1 1
These paths are called orbits, stationary states or allowed ν= = 2
− 2
energy states. h h ni n f
— Energy is emitted when an electron jumps from higher Å In case of absorption spectrum, n f > ni and in case of
energy level to lower energy level. emission spectrum ni > n f .
KEY NOTES
CHAPTER > 02
Structure of Atom
KEY NOTES
Å The first atomic theory proposed by Dalton in 1808,
regarded the atom as the ultimate indivisible particle of Atomic Models
matter. Thomson’s Model of Atom
Å A large number of subatomic particles have been discovered Å Thomson in 1898 assumed that an atom is a sphere of positive
but only electron, proton and neutron are of great charged uniformly distributed with the electrons scattered as
importance among them and hence are called fundamental points throughout the sphere. This was also known as plum
particles. pudding, raisin pudding or watermelon model. An important
feature of this model is that the mass of atom is assumed to be
Discovery of Subatomic Particles uniformly distributed over the atom.
Å Cathode rays or electrons were discovered by J.J. Thomson Å Henri Becqueral observed that there are certain elements
in 1897, by utilising Faraday’s study of electrical discharge which emit radiation on their own. These elements were
in partially evacuated tube known as cathode rays tubes. named as radioactive elements and the phenomenon is called
Å The specific charge is the ratio of charge to mass of an radioactivity.
electron, i.e. e /me ratio. Å Rutherford’s nuclear model of atom Rutherford bombarded
Å By carrying out accurate measurements, Thomson was able very thin gold foil with α-particle.
to determine the value of e /me as 1758820
. × 1011 C kg − 1 . Å Thomson’s model of atom was proved wrong by Rutherford’s
Å Millikan devised a method known as oil drop experiment alpha-particles scattering experiment carried in 1909.
to determine the charge on the electron to be − 1.6 × 10− 19 C. Å The main features of this model are :
Å The mass of electron was determined by combining — Most of the particles passed the foil undeflected, which
Millikan’s and Thomson’s value of e /me ratio which comes indicated that most of the space in atom is empty.
out to be 91094
. × 10− 31 kg. — Some of them were deflected, but only at small angles. This
shows that there is something positively charged at the
Å Canal rays (or anode rays or positive rays) were
centre.
discovered by Goldstein. These rays consist of positively
charged particles called protons.
— Few particles were deflected at large angles. It means that in
the atom, mass and positive charge is centrally located in
Å Unlike cathode rays, the e /me value of canal rays depends extremely small region called nucleus. The nucleus is
upon the nature of gas taken in the tube. surrounded by electrons that move around the nucleus with
Å Neutrons are neutral particles and discovered by a very high speed in circular paths called orbits.
Chadwick. These are the heaviest particles of the atom. — Electrons and the nucleus are held together by electrostatic
Å The discovery of sub-atomic particles led to the proposal of force of attraction.
various atomic models to explain the structure of atom.
,Atomic and Mass Number Particle Nature of Electromagnetic
Å Atomic number ( Z) = number of protons in the Radiation : Planck’s Quantum Theory
nucleus of an atom = number of electrons in a neutral
Å The ideal body, which emits and absorbs radiations of all
atom
frequencies is called a black body and the radiation emitted by
Å The protons and neutrons present in the nucleus are such a body is called black body radiation.
collectively known as nucleons. The total number of
Å According to Planck’s quantum theory, the radiant energy
nucleons is termed as mass number of the atom.
which is emitted or absorbed in the atom of small discrete
Mass number ( A) = number of protons ( Z) + number packets of energy known as quantum and in case of light, the
of neutrons (n). quantum of energy is called photons.
c
Isobars and Isotopes E = hν or E = h
λ
Å Isobars are the atoms with same mass number, but
where, h = Planck’s constant = 6.63 × 10− 34 Js
different atomic number for, e.g. 14 14
6 C and 7 N.
E = Energy of photon or quantum
Å The species with same atomic number but different
Å If n is the number of quanta of a particular frequency and ET be
mass number are called isotopes, e.g. 6 C12 and 6 C14 .
the total energy, then
Å Hydrogen has three isotopes protium (11 H, only one ET = nhν
proton), deuterium (12 D, one proton and one neutron)
Photoelectric Effect
and tritium (13 T, one proton and two neutrons). Å The phenomenon of ejection of electrons from a metal surface
Drawbacks of Rutherford’s Model when a light of certain frequency strikes on its surface is called
photoelectric effect.
Rutherford’s model was failed to account for the
stability of the atom. Also, it did not explain about the Å For each metal, there is a characteristic minimum frequency,
electronic structure of atoms. known as threshold frequency ( ν 0) below, which photoelectric
effect is not observed.
Development Leading to the Bohr’s Å When a photon of sufficient energy strikes an electron in the
Model of Atom atom of the metal, it transfers its energy instantaneously to the
Å Nature of electromagnetic radiation and experimental electron during the collision and the electron is ejected.
results regarding atomic spectra play an important Å The number of electrons ejected is proportional to the intensity
role in the development of Bohr’s model. or brightness of light.
Å Light, X-rays and γ-rays are the examples of radiant Å Following the conservation of energy principle, the kinetic
energy. energy of ejected electron is given by the equation,
Maxwell in 1856 showed that radiant energy has wave 1
Å
hν = h ν 0 + m e ν 2
properties and called them electromagnetic waves or 2
electromagnetic radiations. Å Dual nature of electromagnetic radiation Light possesses both
Å There are many types of electromagnetic radiations particle and wave like properties. Whenever radiation interacts
which differ from one another in wavelength or with matter, it displays particles, like properties in contrast to
frequency. These constitute electromagnetic the wave like properties.
spectrum.
Atomic Spectra
Å The small portion in the electromagnetic spectrum
Å The pictorial representation of arrangement of various types of
around 1015 Hz is called visible light.
EMR in their increasing order of wavelength (or decreasing
Å All these radiations travel with the speed of light and order of frequency) is known as spectrum.
do not require any medium for their propagation or Å The spectrum of white light ranges from violet at 7.50 × 1014 Hz
transmission.
to red at 4 × 1014 Hz, such a spectrum is called continuous
Å The frequency ( ν), wavelength ( λ ) and velocity of light
spectrum.
(c) are related by the equation,
Å The spectrum of radiation emitted by a substance that has
c = νλ
absorbed energy in increasing order of wavelengths or
Å Wave number is defined as the number of decreasing frequencies is called as an emission spectrum.
wavelengths per unit length. Its commonly used unit Atoms, molecules or ions that have absorbed radiation are said
is cm − 1 . to be ‘excited’.
KEY NOTES
, Å An absorption spectrum is like the photographic negative — The frequency of radiation absorbed or emitted when
of an emission spectrum. transition occurs between two stationary states that
Å The study of emission or absorption spectra is referred to differ in energy by ∆E is given by :
as spectroscopy. ∆ E E2 − E1
v= =
Å The emission spectra of gas phase do not show a h h
continuous spread of wavelength from red to violet, rather where, E1 and E2 are the energies of the lower and higher
they emit light only at specific wavelengths with dark allowed energy states respectively. This expression is
spaces between them. Such spectra are called line spectra commonly known as Bohr’s frequency rule.
or atomic spectra. — The angular momentum of an electron is quantised. In a
Å Line emission spectra are of great interest in the study of given stationary state, it can be expressed as,
electronic structure and also used in chemical analysis to h
mevr = n ⋅ , n = 1, 2, 3 …
identify unknown atoms. 2π
Å When an electric discharge is passed through gaseous — The stationary states for electron are numbered,
hydrogen, the H2 molecules dissociate and the n = 1, 2, 3... . These integral numbers are known as
energetically excited hydrogen atoms produced emit principle quantum number.
electromagnetic radiation of discrete frequencies. — The radii of the stationary states are expressed as :
Å The line spectra of hydrogen lies in three regions of
electromagnetic spectrum, i.e. infrared, visible and rn = n2 a0
UV-region. where, a0 = 52.9 pm.
Å The set of lines in the visible region is known as Balmer Thus, the radius of the first stationary state called the
series. Bohr orbit is 52.9 pm.
Å The Swedish spectroscopist, Johannes Rydberg noted that — Energy of an electron in nth state is given as,
all series of lines in the hydrogen spectrum could be 1
described by the following formula. En = − R H 2
n
1 1
ν = 109677 2 − 2 cm− 1 n = 1, 2, 3 …
n1 n2
where, R H is called Rydberg constant and its value is
where, n1 = 1, 2 … equal to 2.18 × 10− 18 J.
n2 = n1 + 1, n1 + 2 … Å Bohr’s theory can also be applied to the ions containing
Å The value 109677 cm − 1 is called the Rydberg constant. only one electron, for example He+ , Li 2 + , Be3 + and so on.
Å The first five series of lines that correspond to Å The energy of the stationary states associated with these
n1 = 1, 2, 3 , 4, 5 are known Lyman, Balmer, Paschen, hydrogen like species is given by,
Bracket and Pfund series respectively. Z2
En = − 2.18 × 10− 18 2 J
Series n1 n2 Spectral Regions n
Lyman 1 2, 3, 4K Ultraviolet and radii by the expression,
n2
Balmer 2 3, 4, 5K Visible rn = 52.9 pm
Z
Paschen 3 4, 5, 6K Infrared
Line Spectrum of Hydrogen
Brackett 4 5, 6, 7K Infrared Å The energy gap between the two orbits (higher and lower
Pfund 5 6, 7K Infrared orbits) is given by equation,
1 1
∆E = 2.18 × 10− 18 J 2 − 2
Bohr’s Model for Hydrogen Atom ni n f
Å The main postulates for Bohr’s model are :
Å The frequency associated with the absorption and emission
— The electron in the hydrogen atom can move around the of the photon can be evaluated by using equation,
nucleus in a circular path of fixed radius and energy.
∆E R H 1 1
These paths are called orbits, stationary states or allowed ν= = 2
− 2
energy states. h h ni n f
— Energy is emitted when an electron jumps from higher Å In case of absorption spectrum, n f > ni and in case of
energy level to lower energy level. emission spectrum ni > n f .
KEY NOTES
