History of Atomic Models PDF

# History of Atomic Models ## Timeline: 400 BC * **Scientist:** Democritus * **(Greek Philosopher)** Democritus was a Greek philosopher who was the first person to use the term *atom* (*atomos*: meaning indivisible). He thought that if you take a piece of matter and divide it and continue to divide...

# History of Atomic Models ## Timeline: 400 BC * **Scientist:** Democritus * **(Greek Philosopher)** Democritus was a Greek philosopher who was the first person to use the term *atom* (*atomos*: meaning indivisible). He thought that if you take a piece of matter and divide it and continue to divide it you will eventually come to a point where you could not divide it any more. This fundamental or basic unit was what Democritus called an atom. ## Timeline: 350 B.C. * **Scientist:** Aristotle * Aristotle modified an earlier theory that matter was made of four "elements": earth, fire, water, air. * Aristotle was wrong. However, his theory persisted for 2000 years. ## Timeline: 1800's * **Scientist:** John Dalton * John Dalton was the first to adapt Democritus' theory into the first modern atomic model. ## Dalton's Atomic Model 1. Atoms were solid spheres. 2. All substances are made of atoms; atoms are small particles that cannot be divided or destroyed. 3. All atoms of an element are identical in size, mass, shape and other physical properties. 4. Atoms of different elements have different in size, mass, shape weights and different chemical properties. 5. Atoms combine to form compounds ## Timeline: 1890's * **Scientist:** J.J. Thomson * J.J. Thomson was a physicist who is credited for discovering the electron. He used his research on cathode ray tube technology in this discovery. A cathode ray tube is made of glass containing two thin pieces of metal, called electrodes, sealed in it. The electrical discharge through the gases could be observed only at very low pressures and at very high voltages. The pressure of different gases could be adjusted by evacuation. When sufficiently high voltage is applied across the electrodes, current starts flowing through a stream of particles moving in the tube from the negative electrode (cathode) to the positive electrode (anode). These were called cathode rays or cathode ray particles. <start_of_image> Diagrams of cathode rays with labels: * **cathode (-)** * **anode (+)** * **high voltage** * **fluorescent coating** * **to vacuum pump** * The flow of current from cathode to anode was further checked by making a hole in the anode and coating the tube behind anode with phosphorescent material zinc sulphide. When these rays, after passing through anode, strike the zinc sulphide coating, a bright spot on the coating is developed. * The results of these experiments are summarized below. * The cathode rays start from cathode and move towards the anode. * These rays themselves are not visible but their behavior can be observed with the help of certain kind of materials (fluorescent or phosphorescent) which glow when hit by them. * In the absence of electrical or magnetic field, these rays travel in straight lines * In the presence of electrical or magnetic field, the behavior of cathode rays are similar to that expected from negatively charged particles, suggesting that the cathode rays consist of negatively charged particles, called electrons. * The characteristics of cathode rays (electrons) do not depend on the material of electrodes and the nature of the gas present in the cathode ray tube. Thus, we can conclude that electrons are basic constituents of all the atoms. ## Thomson's Atomic Model * Thomson proved that an atom can be divided into smaller parts. * He proposed that an atom possesses a spherical shape (radius approximately 10^{-10} m) in which the electrons are embedded into it in such a manner as to give the most stable electrostatic arrangement. * Many different names were given to this model, for example, plum pudding, raisin pudding or watermelon. ## Rutherford's Nuclear Model of Atom * Rutherford and his students bombarded very thin gold foil with α-particles. A stream of high energy α-particles from a radioactive source was directed at a thin foil (thickness ~ 100 nm) of gold metal. The thin gold foil had a circular fluorescent zinc sulphide screen around it. * Whenever an α-particle struck the screen (Without putting a gold slide), a tiny flash of light was produced at that (A) point. * After putting the foil of gold it was found that: * most of the α-particles passed through the gold foil undeflected * Small fraction of the α-particles was deflected by small angles. * Very few α-particles (~ 1 in 20,000) bounced back, that is, were deflected by nearly 180°. Diagram of Rutherford's scattering experiment: * **Gold foil** * **Source of alpha particles** * **Lead plate** * **Photographic plate** * On the basis of the observations, Rutherford drew the following conclusions regarding the structure of the atom: * Most of the space in the atom is empty as most of the α-particles passed through the foil undeflected. * A few positively charged α-particles were reflected. The reflection must be due to a collision with a large body that can not be radiated or penetrate. * A few positively charged α-particles were deflected. The deflection must be due to a large repulsion force showing the atom as Thomson had proved. The positive charge has to be concentrated in a very small volume that repelled and deflected the positively charged α-particles. This very small portion of the atom was called nucleus. ## Rutherford's Nuclear Model of Atom * **The Atom:** Although it has very small size, it has a complicated structure that resembles the solar system in which electrons revolve around the central nucleus in orbits as planets revolve around the sun. * **The Nucleus:** Is much smaller than the atom. Located in the centre of the atom with a positive charge. There is a big space between the nucleus and orbits of electrons, so most of the atom is a space. Most mass of the atom is concentrated in the nucleus and the mass of electrons is very small and can be neglected. * **Electrons:** * Have negligible mass compared to that of the nucleus. * The number of electrons with negative charge are equals to the number of protons with positive charge so the atom is electrically neutral. * Electrons revolve around the nucleus in a fixed orbit. * The electrons are affected by two forces equal in strength but in opposite directions, which are: * Attraction Force of the nucleus to electrons. * Centrifugal force due to velocity of electron around the nucleus. ## Objection to the Rutherford Model: * According to Maxwell's theory: if a charged electron particle moves around another particle charged with an opposite charge, the electron loses part of its energy coming out in the form of radiation, thereby reducing the orbit gradually. And that is mean the electron finally fall in nucluse (and that never done). ## Atomic spectra and its explanation (Bohr theory) * The study of atomic spectra is considered the key which solved the puzzle of atomic structure. That was the work of... * On heating atoms of a pure element in gaseous or vapour state to high temperature or exposing them to low pressure inside electric discharge tube, they emit radiation known as line spectrum. On examining this radiant light by a spectroscope, we observe a group of small number of restricted coloured lines separated by dark areas so it’s called line spectrum. It was found experimentally that the spectral lines are essential characteristic for each element i.e. there are no two elements that have the same spectral lines. * **Spectrum:** The visible part of the light that analyzed using an optical analyzer (S.S) * **Atomic spectra are of two types:** * **Linear emission spectrum:** Are separate colored lines present in dark areas radiated by the excited element's * **Linear absorption spectrum:** Black lines in the continuous spectrum ## Bohr's Atom * Bohr’s model was based on his observations on the spectrum produced by heating the hydrogen atom. When heating this atom emits radiation, then we expose it to a triple prism. The triple prism makes the white light deviate, producing all the colors of the visible spectrum. Each color corresponds to a specific amount of energy; however, when light emitted from a hydrogen atom is passed through a triple prism, only certain colors of light are seen. That’s why Bohr says that electrons have specific quantities of energy in an atom, and by the colors emitted by a hydrogen atom. ## Bohr's Model * Studying the line spectra of hydrogen atoms, Bohr was able to reach his atomic model ## Bohr Postulates * Bohr adopted some of Rutherford’s postulates about atomic structure; * A positively charged nucleus exists in the center of the atom. * Atom is electrically neutral as number of protons equals to number of electron’s. * Electrons orbit the nucleus only in a definite allowed energy levels and they cannot be found at intermediate distances. * Electrons revolve around the nucleus in orbits due to centrifugal and attraction forces. * Electrons orbit the nucleus in a rapid movement without gaining or losing energy. * Each electron in the atom has a definite amount of energy depending on the distance between its energy level and the nucleus; the energy of any level increases as its radius increases. * The maximum number of energy levels in atoms in their ground state (unexcited) is only seven (K, L, M, N, O, P, Q). Each level has energy expressed by a whole number called principle Quantum. Number from (1 to 7). * When atom is excited by heating (Quantum) or by electric discharge the electron will transfer to a higher Energy. level agrees with the absorbed quantum. The excited electron in the higher Energy. level is then unstable, so it returns to its original level losing the same quantum of energy, which it gained during excitation in the form of radiation have spectrum line. ## REMARK: * Quantum is defined as the amount of energy gained or lost when an electron jumps from one Energy level to another. * The difference in energy between levels (Quantum) is not equal i.e. the difference in this energy decreases further from the nucleus. * The electron does not move from its level to another unless the energy absorbed or emitted is equal to the difference in energy between 2 levels i.e. one quantum. ## The success of the Bohr model * It explained hydrogen atom spectrum. * He introduced the idea of quantum number to detect energy of electrons in energy levels. * He proved that electrons during rotation around the nucleus in ground state do not radiate energy, so they will not fall back to the nucleus. * A reconciliation between Rutherford and Maxwell ## The Limitation of the Bohr model * Bohr failed to explain the spectrum of any other element even that of Helium except hydrogen. * He considered the electron as negative charged particle only and did not consider that it also has wave properties. * He postulated that it is possible to determine precisely both speed and location of an electron at the same time. This is experimentally impossible. * He described the electron when moving in a circular planer orbit. Later it was confirmed that hydrogen atom has 3... ## Sommerfeld atomic model * An improved model of the Bohr atomic model was proposed by Summerfield in 1916. Summerfield agreed to all the hypotheses in Bohr’s theory except for the circular orbits of electron motion. According to Bohr’s theory, when an electron falls from a higher energy to a lower energy level, it radiates one spectral line, but using spectral devices with high analytical power, shows that each spectral line consists of at least two spectral lines separated by very little distance. ## So Sommerfield made a modification of Bohr’s theory based on: * Each of the energy levels produced by Bohr’s model includes a number of sub-levels equal to the level number so that one orbit is circular and the rest is oval. * To determine the energy of the oval orbit, sommerfield entered a new quantum number known as the number of osmotic quantities (K), which is an integer that expresses the energy of the electron in the oval orbit (sub-level number). The oval in which the electron moves. * The speed of an electron’s movement changes according to its position in the orbit so that it increases as it approaches nucleus. * Using n and K, Summerfield was able to determine the value between the axis length of the orbital using the simple equation(K=n-1) * b/a=(K+1)/n * a,b= the axis of the oval orbital Diagram of the oval orbital where n = 3 and K = 1, 2, 3: ## 1-The wave nature of the electron * All previously experimental considered the electron just a negatively charged particle but de Brawley assumed that the electron has a dual nature. * De Brawley considered that Every moving body (such as electron or the nucleus of an atom or whole molecule) is associated with (accompanied by) a wave motion (or matter waves) which has some properties of light waves. * It is a material particle. * It has wave properties. Tables comparing electromagnetic and matter waves: | Electromagnetic Wave | Matter wave | |---|---| | They are separating from the moving body | They are not separating from the moving body | | Their speed is equal to the speed of light | Their speed is not equal to the speed of light | ## 2- The principle of uncertainty * According to Heisenberg’s uncertainty rule, it is impossible to know the velocity and the location of the electron around the nucleus at the same time. * But, it is correct is to say that the electron may be somewhere around the nucleus as much as possible. That is called the electron clouds. ## 3- Quantum Numbers (Schrödinger equation) * In 1926, the Austrian scientist, Schrodinger applying the ideas of Planck, Einstein, De Broglie, and Heisenberg established the wave mechanical theory of the atom and managed to derive a wave equation that could describe the electron wave motion in the atom. On solving Schrodinger’s equation:. * It is possible to determine the allowed energy levels and to define the region of space around the nucleus where it is most probable to find the electron in each energy level. As a result of Schrodinger's work, our concept of the electronic motion around the nucleus has changed. Instead of speaking about the stable circular "orbits" of particular radii, and the areas between these orbits as being completely forbidden for electrons. The concept of electron cloud used to express the region of space around the nucleus where the possibility of finding the election in all distances, and directions. Inside the electron cloud there are areas that have a great possibility of finding an electron in it called orbital. Diagram of the electron cloud around the nucleus: ## REMARK: * **Electron cloud** is defined as the region of space around the nucleus where the possibility of finding the election in all distances, and directions. * **Orbital** is defined as: The area of space inside the electron cloud where there is a great probability for finding electrons. The mathematical solution of the Schrodinger equation introduced four numbers which are called quantum numbers. ## Quantum Numbers: * They define the energy, shape, number and direction of orbitals. ## 1- Principle Q.no (n). * The principal quantum number (n) describes the distance of the electron from the nucleus. * The order of principle energy levels their number in the heaviest known atom in the ground state is seven. * The number of electrons required to fill a given energy level = two times the square of the level no (2n²). Table of energy levels with the number of electrons that each level can hold: | Energy Level | Quantum number (n) | Number of Electrons | |---|---|---| | 1st E.L | K | 2 | | 2nd E.L | L | 8 | | 3rd E.L | M | 18 | | 4th E.L | N | 32 | ## 2- Subsidiary Q.no (L). * Used to detect the number of sub levels in each energy level. * The energy sub levels take the symbols s, p, d, f. This is shown by the scientist Somerfield. When he used a spectroscope which has a high resolving power, he found that the single line (which represents electron transition between two different energy levels) is indeed a number of fine spectral lines which represents electron transition between very near energy levels (sublevels). * The number of sublevels in each energy level = order of principle energy level (n). * The energy of sub levels of same Energy level is not equal. * Values of subsidiary quantaum (L) =(n-1). Table showing the sub levels in each energy level: | Energy Level | Quantum number (n) | Sub levels | |---|---|---| | 1st E.L | K | 1s | | 2nd E.L | L | 2s, 2p | | 3rd E.L | M | 3s, 3p, 3d | | 4th E.L | N | 4s, 4p, 4d, 4f | ## 3-Magnetic Q number (m) * Used to detect no of orbitals in each energy sub level and their direction in space, which equal to (m=2l +1) where (l) is the value of the number of Subsidiary quantum, and the value of the number of magnetic quantum range between (-1) and (+1). Diagram of an orbital with labels: * **z** * **y** * **s** * **x** * **Sublevel (S)** (m = 2 x 0 +1 = 1) So the level (s) has one direction in the space and as such it has a spherical shape around the nucleus. * **Sublevel (P)** has 3 orbitals (m = 2 x 1+1 = 3). Therefore, the level (p) has three directions in the space (Px, Py, Pz) is perpendicular to the other two. Also P consists of two dumb or two bell shaped in contact with each other and each dumb can contain an electron and these two dumb are meeting head to head at a point where the electron is difficult to exist (zero electron density). Diagram of the p orbitals: * **PX** * **PY** * **PZ** * **Sublevel (d)** has 5 orbitals (m = 2 x 2+1 = 5) Therefore, the level (d) has five directions. * **Sublevel (f)** has 7 orbitals (m = 2 x 3+1 = 7) Therefore, the level (f) has seven directions. ## 4-Spain Q number (m) * Any orbital contain two electrons each electron spain around its axis during orbits around nucleus . Althought the electrons in the same orbitals carry the same negative charge we might expect them to repel. Yet due to the spain of electron around its axis a magnetic field will be arised so one electron spins around its axis clockwise while the other electron spins anti clockwise in order to from 2 opposite magnetic fields to decrease the force of repulsion between them which keep the atom stable. * **Spain Q number** is used to detects the direction in which the electron spins around its axis during its rotation around the nucleus. Diagram showing the clockwise and anticlockwise spins of electrons around the nucleus: * **+1/2** * **-1/2** ## Principles of distributing electrons * There are two important rules which must be considered in distributing electrons in the atom. These rules are: ## 1-Building-up principle * It states that electrons must fill the lower energy sub-levels with lower (n+L) first and then the higher-energy sub-levels.and if we have two orbital with same (n+L), the electron prefers to fill the orbital with lower(n). ## 2-Hund's Rule * No electron pairing takes place in a given sublevel until each orbital contains one electron. Diagram showing how the electrons fill up the orbitals in: * Nitrogen * Oxygen The diagram shows: * **2px** * **2py** * **2pz** * **n = 1** * **n = 2** * **Nitrogen** * **Oxygen** ## 1s <2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p