Structure Of An Atom Class 11 Notes Chemistry Chapter 2

structure of an atom class 11 notes

This article is on the Structure Of An Atom Class 11 Notes of Chemistry. The notes on the structure of an atom of class 11 chemistry have been prepared with great care for the students so that they can give a quick glance of the chapter. The Notes on the Structure of an atom of Class 11 has been divided into two articles. This article is on Dalton’s Atomic Theory, Discovery of electrons, protons and neutrons, Atomic models, electromagnetic radiations and Atomic spectrum. The second article (Part 2) includes Bohr’s model, Heisenberg’s Uncertainty principle, Quantum mechanical model, Quantum numbers, Pauli’s principle, Aufbau principle and Hund’s Rule.

Structure of An Atom (Atomic Structure) Part 1

The word atom in Greek means indivisible. The concept of atomic theory of matter was put on firm footing by John Dalton with the atomic theory that was developed by him during the years 1803–1808.

Dalton’s Atomic Theory

Dalton’s atomic theory (1808) is based on the following two laws: the law of conservation of mass and law of definite proportions. He also proposed the law of multiple proportions, as a logical consequence of his theory.

(a) Every element is composed of extremely small indestructible particles called atoms (‘atom’ in Greek means indivisible).

(b) Atoms of any one element are all similar but they differ from atoms of another element.

(c) Atoms of each element are fundamental particles, have a characteristic mass but do not have any structure.

(d) Atoms of various elements take part in a chemical reaction to form compound (which is called molecule).

(e) In any compound, the relative number and kinds of atoms are constant.


Cathode Rays – Discovery of Electron

In the mid-1800s, scientists (William Crookes 1879, Julius Plucker 1889) started to study the discharge of electricity through partially evacuated tubes. Gases are normally poor conductor of electricity and they do not conduct electricity under normal pressure even with an applied potential of 1000 volt. However when the pressure was reduced to 10–2 mm at a potential of 1000 volt, Crookes and Plucker noticed that:

(i) From the cathode surface, a glow surrounding the cathode began and the space left between the glow and cathode was named Crookes dark space. Under this condition, electric current starts to flow from one electrode to another.

(ii) When the pressure is sufficiently low, the glow fills the whole of the tube.

Subsequently, Thomson (1897) carried out the discharge through a vacuum tube which was filled with a gas at very low pressure (10-2 to 10-3 mm); he noticed the emission of invisible rays which produced fluorescence on the glass and influenced photographic plate. These rays were called cathode rays.

cathode ray discharge tube

Important characteristics of cathode rays are as follows:

(a) Cathode rays are emitted from the surface of the cathode; their direction is not affected by the position of the anode.

(b) Cathode rays produce fluorescence when they fall on certain substances like ZnS. The colour of fluorescence varies with the chemical nature of the substance.

(c) Cathode rays are deflected, both, by electric and magnetic fields.

(d) Cathode rays produce X-rays when they strike a metal target of a high atomic number such as tungsten, which is highly penetrating. In 1897 J. J. Thomson determined the charge/mass value of the electron by studying the deflection of cathode rays in electric and magnetic fields. The value of e/m has been found to be – 1.7588 × 108 C/g.

structure of an atom class 11 notes

Millikan’s Oil-Drop Experiment

In 1909, the first precise measurement of the charge on an electron was made by Robert A. Millikan using his oil drop experiment. The charge on the electron was calculated to be – 1.6022 × 10-19 C. An electron has the smallest charge known; so it was, designated as a unit negative charge.

Mass of the electron: The mass of the electron is calculated from the values of e/m.

m = e /(e/m) = – 1.6022 × 10-19/ – 1.7588 × 108 = 9.1096 × 10-28 g or 9.1096 × 10-31 kg.

This is known as the rest mass of the electron; that is, the mass of the electron when it is moving with low speed. For calculating the mass of a moving electron, the following formula is used.

Mass of moving electron = rest mass of electron/ √1 – (υ/c)2, where υ is the velocity of the electron and c is the velocity of light. When υ is equal to c, mass of the moving electron is infinity and when the velocity of the electron becomes greater than c, the mass of the electron becomes imaginary.

Note: An electron can thus be defined as a subatomic particle which carries charge -1.60 ×10-19 C, that is, one unit negative charge and has mass 9.1 × 10-19 g, that is, 1/1837 the mass of the hydrogen atom (0.000549 amu).

Discovery of Protons – Positive Rays

Goldstein (1886) repeated the discharge tube experiment but he used perforated cathode and noticed the emission of positive rays or canal rays. These rays do not originate from the anode, and so it is wrong to call them anode rays. The specific charge (e/m) of canal rays particles varied with the nature of gas and was found to be maximum if H2 was used. The positive rays particles were thus, called positively charged gaseous atoms left after the removal of an electron or ionized gaseous atoms. However, if H2 gas is used in discharges, the positive rays particles are named as protons (usually represented as P). Thus, a subatomic particle, that is a fundamental constituent of all matter, is called a proton; it has a mass 1.673 × 10-27 kg and charge + 1.603 × 10-19 C.

Discovery of Neutron

Chadwick bombarded beryllium with a stream of α-particles. He observed that the produced penetrations were not affected by electric and magnetic fields. These radiations consisted of neutral particles, which were called neutrons.

94 Be(Beryllium) + 42He (α-particle) → 126C (Carbon) + 10n (Neutron)

The mass of the neutron was determined to be 1.675 × 10-24 g, that is, nearly equal to the mass of the proton. Thus a neutron is a subatomic particle having a mass of 1.675 × 10-24 g, approximately 1 amu or nearly equal to the mass of a proton or a hydrogen atom and carrying no electrical charge. The e/m of a neutron is zero.


Thomson’s Model

After the discovery of protons and electrons, Thomson in 1898 proposed a watermelon model; the atom is considered as a sphere of positive charge with the electrons distributed within the sphere of radius of 10-10 m so as to give the most stable electrostatic agreement. In this model, the atom is visualized as a pudding or cake of positive charge with raisins (electrons) embedded in it. A major point of this model is that the mass of atom is considered to be spread uniformly over the atom.

Thomson’s Model

Rutherford’s Model

Rutherford’s Experiment

Rutherford conducted a series of experiments using α-particles. A beam of α-particles was directed against a thin foil of gold, platinum, silver, or copper. The foil was surrounded by a circular fluorescent zinc screen. Whenever an α-particle struck the screen, it produced a flash of light.

Rutherford’s Model

The following observations were made:

(a) Most of the α-particles went straight without suffering any deflection.

(b) Some of them were deflected through small angles.

(c) A very small number (about 1 in 20,000) did not pass through the foil at all but suffered large deflections or even rebound.

Following conclusions were drawn from these observations:

(a) Many of the particles went straight through the metal foil undeflected, indicating that there must be very large empty space within the atom.

(b) Some of the α-particles were deflected from their original paths through moderate angles, indicating that the whole of the positive charge is concentrated in a space called nucleus. It is proposed to be present at the center of the atom.

(c) A very small number of the α-particles suffered strong deflections or even rebound on their path indicating that the nucleus is rigid and α-particles recoil due to direct collision with the positively charged heavy mass.

Rutherford proposed a model of the atom known as the nuclear atomic model. As per this model,

(a) An atom consists of a positively charged heavy nucleus where all the protons and neutrons are present. The number of positive charge on the nucleus is different for different atoms.

(b) The volume of the nucleus is very small and is only a very small fraction of the total volume of the atoms.

Thus, the diameter of an atom is 100,000 times the diameter of the nucleus.
The radius of a nucleus is proportional to the cube root of the mass number.
The radius of the nucleus = 1.33 × 10-13 × A1/3 cm where A is the mass number.

(c) There is an empty space around the nucleus called extra nuclear part, where electrons are present. The number of electrons in an atom is always equal to the number of protons present in the nucleus. The volume of the atom is about 1015 times the volume of the nucleus.

(d) The electrons revolve around the nucleus in closed orbits with high speeds. Centrifugal force is acting on the revolving electrons and is being counterbalanced by the force of attraction between electrons and the nucleus.

Atomic Number and Mass Number

Atomic number of an element = Total number of protons present in the nucleus

= Total number of electrons present in the neutral atom
Mass number of an element = No. of protons + No. of neutrons

Isotopes, Isobars, Isotones

Such atoms of the same element having the same atomic number but different mass numbers are known as isotopes. Such atoms of different elements having different atomic numbers, but have the same mass number, e.g. 4018Ar, 4019K, 4020Ca are known as isobars. Such atoms of different elements containing the same number of neutrons, e.g. 146C, 157N, 168O, are called isotones.

Drawbacks of Rutherford’s Model

(a) According to wave theory, when a charged particle moves under the influence of an attractive force, it loses energy continuously in the form of electromagnetic radiations. Thus, the electron which moves in an attractive field (created by protons present in the nucleus) will emit radiations. As a result, the electron will lose energy at every turn and move closer and closer to the nucleus following a spiral path and finally fall into the nucleus, thereby making the atom unstable. But the atom is quite stable meaning the electrons do not fall into the nucleus, thus this model does not explain the stability of the atom.

(b) If the electrons lose energy continuously, the observed spectra should be continuous. But the observed spectra consists of well-defined lines of definite frequency. Hence, in an atom, the loss of energy by the electrons is not continuous.


Electromagnetic Radiations (EMR) are energy radiations which do not need any medium for propagation, e.g. visible, ultraviolet, X-rays, etc. Following are the important characteristics of EMR:

(a) All electromagnetic radiations or waves travel with the velocity of light.

(b) These consist of electric and magnetic fields that oscillate in directions perpendicular to each other and perpendicular to the direction in which the wave is travelling.

Electromagnetic Radiations

Electromagnetic Spectrum

The arrangement of various types of electromagnetic radiations in the order of their increasing or decreasing wavelengths or frequencies is known as the electromagnetic spectrum.

Characteristics of Waves

A wave is always characterized by the following six characteristics:

(a) Wavelength is the distance between two nearest crests or nearest troughs and is denoted by λ (lambda) and is measured in units of centimetre (cm), angstrom (Å), micrometre, (µm) or nanometre (nm).

(b) Frequency is defined as the number of waves passing through a point in one second and is denoted by the symbol (nu) and is measured in terms of cycles (or waves) per second (cps) or Hertz (Hz).
λν = Distance travelled in one second = Velocity = c or ν = c λ

(c) Velocity is defined as the distance covered in one second by the wave and is denoted by the letter ‘c’. All electromagnetic waves travel with the same velocity, that is, 3 × 1010 cm/sec. λν = 3 × 1010. Thus, a wave of higher frequency has a shorter wavelength while a wave of lower frequency has a longer wavelength.

(d) Wavenumber is the number of wavelength per centimeter. It is the reciprocal of wavelength and is denoted by the symbol ν (nu bar). ν =1/λ. It is expressed in cm-1 or m-1.

(e) Amplitude is defined as the height of the crest or depth of the trough of a wave and is denoted by the letter ‘A’. It determines the intensity of the radiation.

(f) Time period is the time taken by one wave to complete a cycle or vibration and is denoted by T. T = 1/ν. Unit: Second per cycle.

electromagnetic spectrum


In 1905, Einstein pointed out that light can be considered to consist of a stream of particles, called photons. The energy of each photon of light depends on the frequency of the light, that is, E = hv. According to Einstein, energy is also related as E = mc2 where m is the mass of a photon. Thus, he pointed out that light has a wave as well as particle characteristics (dual nature). Though the wave theory successfully explains many properties of electromagnetic radiations such as reflection, refraction, diffraction, interference, polarization etc. it fails to explain some phenomena like blackbody radiation, photoelectric effect etc.

structure of an atom class 11 notes

A new theory which is known as the quantum theory of radiation was presented by Max Planck in 1901, to explain the blackbody radiation and photoelectric effect. According to this theory, a hot body emits radiant energy not continuously but discontinuously in the form of small packets of energy called quantum (quanta in plural). The energy associated with each quantum of given radiation is proportional to the frequency of the emitted radiation. E ∝ ν or E = hν where h is a constant known as Plank’s constant. Its numerical value is 6.624 × 10-27 erg/sec. The energy emitted or absorbed by a body can be either one quantum or any whole number multiple of hν, that is, 2hν, 3hν, 4hν,…, nhν quanta of energy.

Photoelectric Effect

The emission of electrons from a metal surface, when exposed to light radiations of appropriate wavelength, is called the photoelectric effect. The emitted electrons are called photoelectrons. Work function or threshold energy may be defined as the minimum amount of energy required to eject electrons from a metal surface. According to Einstein, the maximum kinetic energy of the ejected electron = Absorbed energy – Work function ½ mv2max = hν – hν0 = hc[1/λ – 1/λ0], where ν0 and λ0 are threshold frequency and threshold wavelength respectively.

Photoelectric Effect

Stopping potential: Stopping potential is the minimum potential at which the photoelectric current becomes zero. If ν0 is the stopping potential, then ν0 = h(ν – ν0).

Laws of Photoelectric Effect

(a) Rate of emission of photoelectrons from a metal surface is directly proportional to the intensity of incident light.

(b) The maximum kinetic energy of photoelectrons is directly proportional to the frequency of incident radiation; also, it is independent of the intensity of light used.

(c) There is no time lag between the incidence of light and emission of photoelectrons.

(d) For the emission of photoelectrons, the frequency of incident light must be equal or greater than the threshold frequency.


The impression produced on a screen when radiations of particular wavelengths are passed through a prism or diffraction grating is known as spectrum. It is broadly of two types: Emission spectra and Absorption spectra. Differences between the emission spectrum and the absorption spectrum.

Emission Spectra

It is obtained from the substances which emit light on excitation, that is, either by heating the substances on a film or by passing electric discharge through a thin filament of high melting point metal. Emission spectra are of two types:

(a) Continuous spectra: When white light is allowed to pass through a prism, the light gets resolved into several colours. This spectrum is a rainbow of colours, meaning violet merges into blue, blue into green and so on. This is a continuous spectrum.

(b) Discontinuous spectra: When gases or vapours of a chemical substance are heated in an electric arc or in a Bunsen flame, light is emitted. When a ray of this light is passed through a prism, a line spectrum is produced.

This spectrum consists of a limited number of lines, each of which corresponds to a different wavelength of light. Each element has a unique line spectrum. The spectrum of hydrogen is an example of the line emission spectrum or atomic emission spectrum. If an electric discharge is passed through hydrogen gas at low pressure, a bluish light is emitted. If a ray of this light is passed through a prism, a discontinuous line spectrum of many isolated sharp lines is obtained. The wavelengths of the different lines show that these lines are in the visible, ultraviolet and infrared regions. The lines observed in the hydrogen spectrum can be classified into six series.

Absorption Spectra

When the radiations from a continuous source like a hot body (sunlight) containing the quanta of all wavelengths pass through a sample of hydrogen gas, then the wavelength missing in the emergent light gives dark lines on the bright background. This type of spectrum that contains a lesser number of wavelengths in the emergent light than in incident light is called an absorption spectrum.

Absorption Spectra

Ritz Mathematical Formula

Ritz presented a mathematical formula to find the wavelengths of various hydrogen lines.

ū = 1/λ = v/c = (1/n12 – 1/n22) where, R is universal constant, known as Rydberg constant. Its value is 0.9678 cm-1, n1 and n2 are integers (such that n2 > n1). For a given spectral series, n1 remains constant while n2 varies from line to line in the same series. The value of n1 is 1, 2, 3, 4 and 5 for the Lyman, Balmer, Paschen, Brackett and Pfund series respectively. n2 is greater than n1 by at least 1.

Note: (i) Atoms give line spectra while molecules give band spectra.
(ii) Balmer, Paschen, Brackett, Pfund series are found in the emission spectrum.

This article has tried to highlight the basic concepts of the structure of an atom in the form of notes for class 11 students in order to understand the basic concepts of the chapter. The notes on the structure of an atom have not only been prepared for class 11 but also for the different competitive exams such as iit jee, neet, etc.

Check Part 2 of the chapter here: Structure of Atom (Atomic Structure) Part 2

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