Structure Of Atom

Chemistry Structure Of Atom

Structure Of Atom:

  • Atoms are composed of three type of particles: protons, neutrons, and electron.
  • Protons and neutrons are responsible for most of the atomic mass .The mass of an electron is very small (9.108 X 10-28 grams).
  • Both the protons and neutrons reside in the nucleus.
  • Protons have a positive (+) charge, neutrons have no charge –they are neutral.
  • Electrons reside in orbitals around the nucleus. They have a negative charge (-).
  • It is the number of protons that determines the atomic number, e.g., H = 1. The number of protons in an element is constant (e.g., H=1, Ur=92) but neutron number may vary, so mass number (protons + neutrons) may vary.

Different Models of the atom:

Dalton’s model of the atom

  • John Dalton proposed that all matter is composed of very small things which he called atoms.
  • This was not a completely new concept as the ancient Greeks (notably Democritus) had proposed that all matter is composed of small, indivisible (cannot be divided) objects.
  • When Dalton proposed his model electrons and the nucleus were unknown.

Thomson’s model of the atom

  • After the electron was discovered by J.J. Thomson in 1897, people realised that atoms were made up of even smaller particles than they had previously thought.
  • However, the atomic nucleus had not been discovered yet and so the “plum pudding model” was put forward in 1904.
  • In this model, the atom is made up of negative electrons that float in a “soup” of positive charge, much like plums in a pudding or raisins in a fruit cake.
  • In 1906, Thomson was awarded the Nobel Prize for his work in this field. However, even with the Plum Pudding Model, there was still no understanding of how these electrons in the atom were arranged.
  • The discovery of radiation was the next step along the path to building an accurate picture of atomic structure.
  • In the early twentieth century, Marie and Pierre Curie, discovered that some elements (the radioactive elements) emit particles, which are able to pass through matter in a similar way to X–rays.
  • It was Ernest Rutherford who, in 1911, used this discovery to revise the model of the atom.

Rutherford’s model of the atom

  • Rutherford carried out some experiments which led to a change in ideas around the atom.
  • His new model described the atom as a tiny, dense, positively charged core called a nucleus surrounded by lighter, negatively charged electrons.
  • Another way of thinking about this model was that the atom was seen to be like a mini solar system where the electrons orbit the nucleus like planets orbiting around the sun.
  • This model is sometimes known as the planetary model of the atom.

Bohr”s model

        
  • In atomic physics, the  Bohr model or Bohr diagram, introduced by Niels Bohr in 1913, depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus—similar to the structure of the Solar System, but with attraction provided by electrostatic forces rather than gravity.
  • Rutherford’s model introduced nuclear model of atom, in which he explained that a nucleus (positively charged) is surrounded by negatively charged electrons.
  • Bohr modified this atomic structure model by explaining that electrons move in fixed orbital’s (shells) and not anywhere in between and he also explained that each orbit (shell) has a fixed energy level.
  • Rutherford basically explained nucleus of an atom and Bohr modified that model into electrons and their energy levels.
  • Bohr’s model consists of a small nucleus (positively charged) surrounded by negative electrons moving around the nucleus in orbits. Bohr found that an electron located away from the nucleus has more energy, and electrons close to the nucleus have less energy.

Postulates of Bohr’s Model of an Atom:

  • In an atom, electrons (negatively charged) revolve around the positively charged nucleus in a definite circular path called as orbits or shells.
  • Each orbit or shell has a fixed energy and these circular orbits are known as orbital shells.
  • The energy levels are represented by an integer (n=1, 2, 3…) known as quantum number. This range of quantum number starts from nucleus side with n=1 having the lowest energy level. The orbits n=1, 2, 3, 4… are assigned as K, L, M, N…. shells and when an electron attains the lowest energy level it is said to be in the ground state.
  • The electrons in an atom move from lower energy level to higher energy level by gaining the required energy and an electron moves from higher energy level to lower energy level by losing energy.

Limitations of Bohr’s Model of an Atom:

  • Bohr’s model of an atom failed to explain Zeeman Effect (effect of magnetic field on the spectra of atoms).
  • It also failed to explain the Stark effect (effect of electric field on the spectra of atoms).
  • It violates the Heisenberg Uncertainty Principle.
  • It could not explain the spectra obtained from larger atoms.

James Chadwick model

  • Rutherford predicted (in 1920) that another kind of particle must be present in the nucleus along with the proton.
  • He predicted this because if there were only positively charged protons in the nucleus, then it should break into bits because of the repulsive forces between the like-charged protons.
  • To make sure that the atom stays electrically neutral, this particle would have to be neutral itself.
  • In 1932 James Chadwick discovered the neutron and measured its mass.

Isotopes:

  • Any of two or more forms of a chemical element, having the same number of protons in the nucleus, or the same atomic number, but having different numbers of neutrons in the nucleus, or different atomic weights.
  • There are 275 isotopes of the 81 stable elements, in addition to over 800 radioactive isotopes, and every element has known isotopic forms. Isotopes of a single element possess almost identical properties.
  • The isotope of an element is defined by the nucleon number, which is the sum of the number of protons and the number of neutrons in the atomic nucleus.
  • The nucleon number is customarily written as a superscript preceding the chemical symbol for the element. For example, 16O represents oxygen-16, which has 8 protons and 8 neutrons, while 12 C represents carbon-12, with 6 protons and 6 neutrons.
  • These are the most common naturally occurring isotopes of oxygen and carbon, respectively.
  • Some carbon-14 is found in nature. An atom of carbon-14 contains 6 protons and 8 neutrons and is denoted 14 C. Over time, 14 C decays into 12 C.

Isotopes of Hydrogen

  • The three are all isotopes of hydrogen.
  • They have the same atomic number, or number of protons,  but different atomic masses.
  • The number of neutrons can be calculated by calculating the difference between the atomic mass and atomic number.
  •  For the isotopes of hydrogen, they have varying number of neutrons.
  • For protium, the number of neutrons is zero; for Deuterium, the number of neutrons is one; and for Tritium, the number of neutrons is two.

Isotopes of Carbon

  • There are three isotopes of carbon: carbon-12, carbon-13 and carbon-14. The numbers that are after the carbon refer to the atomic mass.
  • The most common and abundant isotope of carbon is carbon-12.
  • The least abundant form of carbon is carbon-14, with an abundance of less than 0.0001%. If we calculate the number of neutrons for each carbon isotope, we can see that they differ from each other.
  • For carbon-12, we have 6 neutrons; for carbon-13, we have 7 neutrons; and for carbon-14, we have 8 neutrons.

Types Of Isotopes

Stable Isotopes

  • Stable isotopes have a stable proton-neutron combination and do not display any sign of decay.
  • This stability comes from the amount of neutrons present in an atom.
  • If an atom has too many or too few neutrons, it is unstable and tends to disintegrate. Since stable isotopes do not decay, they do not produce radiation or its associated health risks.

Uses of Stable Isotopes

  • Scientists performing environmental and ecological experiments use stable isotopes of oxygen, hydrogen, sulfur, nitrogen and carbon.
  • For example, in geochemistry, scientists study the chemical composition of geological materials such as minerals and rocks.
  • Stable isotopes are dependable tools for determining many facts about geological materials, such as their age and where they came from.

Radioactive Isotopes

  • Radioactive isotopes have an unstable combination of protons and neutrons.
  • These isotopes decay, emitting radiation that includes alpha, beta and gamma rays.
  • Scientists classify radioactive isotopes according to their creation process: long-lived, cosmogenic, anthropogenic and radiogenic.
  • Long-lived radioactive isotopes emerged during the creation of the solar system,
  • Cosmogenic radioactive isotopes occur as a reaction of the atmosphere to cosmic rays emitted by stars.
  • Anthropogenic isotopes come from human-made nuclear activities, such as weapons testing and nuclear fuel production.
  • Radiogenic isotopes are the end result of radioactive decay.

Uses of Radioactive Isotopes

  • Radioactive isotopes find uses in agriculture, food industry, pest control, archaeology and medicine.
  • Radiocarbon dating, which measures the age of carbon-bearing items, uses a radioactive isotope known as carbon-14.
  • In medicine, gamma rays emitted by radioactive elements are used to detect tumors inside the human body.
  • Food irradiation — the process of exposing food to a controlled level of gamma rays — kills many types of bacteria, making food safer to eat.

Isobars:

  • one of two or more atoms having equal atomic weights but different atomic numbers.
  • Their chemical property is different because there is the difference in the number of electrons.
  • It has same atomic mass but different atomic no. Because an additional number of neutrons compensates the difference in the number of nucleons.
  • The example of two Isotopes and Isobars is iron and nickel. Both have the same mass number which is 58 whereas the atomic number of iron is 26, and the atomic number of nickel is 28.

Examples of Isobars:

(i)   32 Ce76, 34Se76

(ii) 26Fe58, 27 Ni58 

(iii) 18Ar40 19K40 20Ca40

(iv)11Na24 12Mg24

(v) 27Co64 28Ni64

Isotones:

  • Isotone, any of two or more species of atoms or nuclei that have the same number of neutrons.
  • Thus, chlorine-37 and potassium-39 are isotones, because the nucleus of this species of chlorine consists of 17 protons and 20 neutrons, whereas the nucleus of this species of potassium contains 19 protons and 20 neutrons.
  • S-36, Cl-37, Ar-38, K-39, Ca-40. These nuclei contain 20 neutrons each, but a different number of protons: sulphur 16, chlorine 17, argon 18, potassium 19 and calcium 20 protons.

Examples:

Isotones of Carbon:

  • Isotones are atoms which have same number of neutrons but different atomic number and mass number. Since they have different number of protons and electrons, they do not show same chemical and physical properties.
  • For example: Isotones of carbon. Oxygen, Nitrogen and Carbon are isotones of each other.
 O  N  C
 Neutrons  8  8   8
 Protons  8  7  6
 Mass Number  16  15  14
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