1.Atomic structure.pdf

Chemistry : Atomic Structure Course Number: 0531-201 What is Chemistry? Chemistry, the science that deals with the properties, composition, and structure of substances (defined as elements and compounds), the transformations they undergo, and the energy that is released or absorbed during...

Chemistry : Atomic Structure Course Number: 0531-201 What is Chemistry? Chemistry, the science that deals with the properties, composition, and structure of substances (defined as elements and compounds), the transformations they undergo, and the energy that is released or absorbed during these processes. Why should we read chemistry?  Atomic Structure According to John Dalton, an atom is the smallest unit of matter which takes part in a chemical reaction. Fundamental particles are three types 1. Stationary fundamental particle i. Electron (e-). J.J Thomson discovered electrons in 1897 Mass of electron = 9.107 x 10-28 g or =9.107x10-31 kg , Charge of electron = 1.602x10-19 C ii. Proton (p or, H+) E.Goldstein discovered proton in 1836 Mass of proton =1.672x10-24 g =1.672x10-27 kg Charge of proton=1.602 x10-19 C iii. Neutron (n) James Chadwick discovered neutron in 1932 Mass of neutron =1.675x10-24 =1.675x10-27 kg Charge of neutron =carry no charge i.e. neutral  Atomic Structure Fundamental particles are three types 2. Temporary particles i. meson, neutrino,antineutrino, positron and so on 3. Composite particle i. heavy particle e.g. deuteron particle and α-particle Q1-What is the basic structure of atom?  Atomic Structure Daltons Modern Theory In 1803 British scientist John Dalton gave a theory. He was known as the father of Modern Chemistry. 1. The first part of his theory,all matter is made of atoms, which are indivisible. 2. The second part of the theory says all atoms of a given element are identical in mass and properties.  Atomic Structure What is Rutherford Atomic Model ?. Rutherford's atomic model became known as the nuclear model. 1. The nucleus of an atom is composed of proton & neutron 2. The nucleus is surrounded by negatively charged particles called electrons 3. The electrons revolve around the nucleus in a fixed circular path at very high speed. These fixed circular paths were termed as “orbits.”  Atomic Structure What is the Limitation of Rutherford Atomic Model ? Rutherford proposed that the electrons revolve around the nucleus in fixed paths called orbits. According to Maxwell, accelerated charged particles emit electromagnetic radiations and hence an electron revolving around the nucleus should emit electromagnetic radiation. This radiation would carry energy from the motion of the electron which would come at the cost of shrinking of orbits. Ultimately the electrons would collapse in the nucleus. Calculations have shown that as per the Rutherford model, an electron would collapse into the nucleus in less than 10-8 seconds. So the Rutherford model was not in accordance with Maxwell’s theory and could not explain the stability of an atom This model did not say anything about the arrangement of electrons in an atom Atomic Structure Limitation- Atomic Structure What is Bohrs’ Theory? Postulates are- 1. According to Bohr‟s theory , electrons revolve in definite circular orbits around the nucleus. While revolving in these discrete orbits the electrons do not radiate energy and this is the reason why electrons do not fall into the nucleus. Such orbits are called stationary orbits. And the orbits or shells are represented by the letters K, L, M, N… or the numbers, n=1, 2, 3, 4, ….. So on. 2. An electron can move from a higher energy level (E higher) to a lower energy level (E lower) by emitting appropriate energy. Similarly, an electron can also move from a lower energy level (E lower) to a higher energy level (E higher) by absorbing appropriate energy. The energy absorbed or emitted is equal to the difference between the energies of the two energy levels, i.e., ∆E = E higher - E lower = hν Where h = Planck‟s constant and ν = frequency of electro-magnetic radiation. 3. The angular momentum of the electron is an integral multiple of 𝑛ℎ /2𝜋. i,e,Angular momentum, mvr = 𝑛ℎ/2𝜋 where, m = mass of electron, v= velocity of electron, r= radius of orbit, h = Planck‟s constant =6.626 x 10-34 Js and n= 1,2,3…… 4. It explains the stability of an atom & also the line spectrum of hydrogen atom. Here, release or absorbed energy, ∆E = E higher - E lower = hν  Bohr’s atomic model: Limitation- ▪ Bohr's model of an atom failed to explain the 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. Nature Of Electron An electron is a subatomic particle, with a negative electric charge, found in the outer regions of atoms. It plays a fundamental role in the behavior of matter and electricity. Discovered in the late 19th century, electrons are an essential component of the atomic structure, orbiting the positively charged nucleus. However, wave-particle duality is a profound concept in the realm of quantum mechanics. It challenges the classical notion that particles are purely discrete entities or waves are purely continuous phenomena. Instead, it suggests that particles, like electrons, can exhibit both wave-like and particle-like characteristics, depending on how they are observed or measured. Nature Of Electron What is the De Broglie wavelength, and why is it important? The De Broglie wavelength is a fundamental concept, representing the wave-like nature of particles, including electrons. It’s inversely proportional to momentum and central to understanding wave-particle duality. Heisenberg’s Uncertainty Principle In everyday life, calculating the speed and position of a moving object is relatively straightforward. We can measure a car traveling at 60 miles per hour or a tortoise crawling at 0.5 miles per hour and simultaneously pinpoint where the objects are located. But in the quantum world of particles, making these calculations is not possible due to a fundamental mathematical relationship called the uncertainty principle. Heisenberg Uncertainty Principle states that it is impossible to determine the exact position and the momentum of a sub-atomic particle simultaneously with great accuracy. The principle is named after German physicist Werner Heisenberg, who proposed the uncertainty principle in the year 1927. Heisenberg’s uncertainty principle limits the accuracy of simultaneous measurement of position and momentum. The more precise our location measurement is, the less precise our momentum measurement will be, and vice versa. According to Heisenberg’s uncertainty principle, it is impossible to estimate both the position and velocity accurately for a particle that displays both particle and wave characteristics Heisenberg Uncertainty Formula This uncertainty relationship says that, if ∆x is small I.e. if the position of a particle is measured more accurately then ∆p would be large I.e. momentum is measured less accurately. On the other hand if ∆p is small i.e. if momentum is measured more accurately, then, ∆x would be large I.e. position of a particle is measured less accurately.  Schrödinger Equation In 1926, Austrian physicist Erwin Rudolph Joseph Alexander Schrödinger presented a mathematical concept of wave-particle dualism. The Schrodinger wave equation is a mathematical expression that describes the wave function of a quantum system in the form of the partial differential equation. It describes the probability of finding a particle at a particular position and with particular momentum.. Ref: https://sciencequery.com/what-is-schrodinger-wave-equation/#Q_A Quantum Numbers Definition: Quantum numbers are a set of values that describe the state of an electron including its distance from the nucleus, the orientation and type of orbital where its likely to be found, and its spin. Type: Quantum numbers are four types. They are given below: i. Principal quantum number ii. Subsidiary quantum number iii. Magnetic quantum number iv. Spin quantum number i. Principal Quantum Number: It describes the size of the orbital. It is expressed in term of n. Here, n=1,2,3………. Quantum Numbers ii. Subsidiary quantum number: If describes the shape of the orbital. It is expressed in term of ℓ. Here, ℓ = from 0 to (n-1) If ℓ = 0,1,2,3 then the orbitals can be expressed as s, p, d, f respectively. Quantum Numbers iii. Magnetic Quantum Number: This quantum number has been proposed to account for the splitting up of spectral lines (Zeeman Effect). It describes the orientation in space of a particular orbital. It is expressed in term of m. Here m = 0, ± ℓ SOLVED PROBLEM. List all possible values of l and m for n = 2. SOLUTION. Here, the principal quantum number n = 2. The azimuthal quantum number can have only two values. These are 0 and 1. When l = 0 m=0 And l = 1 m = + 1, 0, – 1 What designation are given to the orbitals having (a) n=2 l=1 (b) n=1 l=0 (c) n=3 l=2 (d) n=4 l=3 SOLUTION (a) when n = 2 and l = 1 the orbital is 2p (b) when n = 1 and p = 0 the orbital is 1s (c) when n = 3 and l = 2 the orbital is 3d (d) when n = 4 and l = 3 the orbital is 4f Quantum Numbers Quantum Numbers Quantum Numbers iv. Spin Quantum Number: It describes the spin of the electron within the orbital. It is expressed in term of S= ± 1/2  Electronic Configuration of Atom The electron configuration of an element describes how electrons are distributed in its atomic orbitals. Electron configurations of atoms follow a standard notation in which all electron- containing atomic subshells (with the number of electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is 1s22s22p63s1. Notation 1. The electron configuration of an atom is written with the help of subshell labels. 2. These labels contain the shell number (given by the principal quantum number), the subshell name (given by the azimuthal quantum number) and the total number of electrons in the subshell in superscript. 3. For example, if two electrons are filled in the „s‟ subshell of the first shell, the resulting notation is „1s2‟. 4. With the help of these subshell labels, the electron configuration of magnesium (atomic number 12) can be written as 1s2 2s2 2p6 3s2 Ref: https://byjus.com/chemistry/electron-configuration/#Writing_Electron_Configurations Electronic Configuration of Atom 1. Aufbau Principle 2. Pauli Exclusion Principle 3. Hund’s Rule Aufbau Principle Aufbau principle states that an electron occupies an orbital in the order of lowest to highest energy orbital. „Aufbau‟ is a german word that means construction or build up. This is the reason, it is also called a building-up principle or construction principle. Basically, this principle explains how electrons are distributed among energy levels in the ground states of atoms. The electrons are filled in an orbital following Pauli‟s exclusion principle as well as Hund‟s rule. According to the Aufbau principle, “electrons are filled in an orbital in the increasing order of orbital energy level“. In other words, electrons are filled in the orbital in such a way that the orbital with the lowest energy gets filled first, then to the next higher energy level, and so on i.e. in the pattern of their increasing energy. The correct sequence of filling of electrons in an orbital depends on two factors. Orbital having the lowest value of (n+l) is filled first In the case where two orbitals have the same (n+l) value, the orbital with the lowest „n‟ value will be filled first by electrons. Aufbau Principle diagram The order in which electrons are filled in atomic orbitals as per the Aufbau principle is illustrated below- It explains the manner in which the electrons are filled in the orbital. It is used to determine the electronic configuration of an atom or ion. It is important to note that there exist many exceptions to the Aufbau principle such as chromium and copper. These exceptions can sometimes be explained by the stability provided by half-filled or completely filled subshells. The increasing order of energy of various orbitals is as follows : 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s...... Electronic Configuration of Atom 2. Pauli Exclusion Principle ▪ The Pauli exclusion principle states that a maximum of two electrons, each having opposite spins, can fit in an orbital. ▪ This principle can also be stated as “no two electrons in the same atom have the same values for all four quantum numbers”. ▪ Therefore, if the principal, azimuthal, and magnetic numbers are the same for two electrons, they must have opposite spins. 3. Hund’s Rule This rule describes the order in which electrons are filled in all the orbitals belonging to a subshell. It states that every orbital in a given subshell is singly occupied by electrons before a second electron is filled in an orbital. In order to maximize the total spin, the electrons in the orbitals that only contain one electron all have the same spin (or the same values of the spin quantum number). An illustration detailing the manner in which electrons are filled in compliance with Hund‟s rule of maximum multiplicity is provided above. Electron Configuration of Hydrogen The atomic number of hydrogen is 1. Therefore, a hydrogen atom contains 1 electron, which will be placed in the s subshell of the first shell/orbit. The electron configuration of hydrogen is 1s1, as illustrated below. Electron Configuration of Oxygen The atomic number of oxygen is 8, implying that an oxygen atom holds 8 electrons. Its electrons are filled in the following order: K shell – 2 electrons L shell – 6 electrons Therefore, the electron configuration of oxygen is 1s2 2s2 2p4, as shown in the illustration provided below. Chlorine Electronic Configuration Chlorine has an atomic number of 17. Therefore, its 17 electrons are distributed in the following manner: K shell – 2 electrons L shell – 8 electrons M shell – 7 electrons The electron configuration of chlorine is illustrated below. It can be written as 1s22s22p63s23p5 or as [Ne]3s23p5 Grossary- Q1 _What is meant by the electronic configuration of an element? The electronic configuration of an element is a symbolic notation of the manner in which the electrons of its atoms are distributed over different atomic orbitals. While writing electron configurations, a standardized notation is followed in which the energy level and the type of orbital are written first, followed by the number of electrons present in the orbital written in superscript. For example, the electronic configuration of carbon (atomic number: 6) is 1s22s22p2. Q2 _What are the three rules that must be followed while writing the electronic configuration of elements? The three rules that dictate the manner in which electrons are filled in atomic orbitals are: The Aufbau principle: electrons must completely fill the atomic orbitals of a given energy level before occupying an orbital associated with a higher energy level. Electrons occupy orbitals in the increasing order of orbital energy level. Pauli’s exclusion principle: states that no two electrons can have equal values for all four quantum numbers. Consequently, each subshell of an orbital can accommodate a maximum of 2 electrons and both these electrons MUST have opposite spins. Hund‟s rule of maximum multiplicity: All the subshells in an orbital must be singly occupied before any subshell is doubly occupied. Furthermore, the spin of all the electrons in the singly occupied subshells must be the same (in order to maximize the overall spin). Q3_Why are electronic configurations important? Electron configurations provide insight into the chemical behaviour of elements by helping determine the valence electrons of an atom. It also helps classify elements into different blocks (such as the s-block elements, the p- block elements, the d-block elements, and the f-block elements). This makes it easier to collectively study the properties of the elements. Q4_List the electron configurations of all the noble gases. The electronic configurations of the noble gases are listed below. Helium (He) – 1s2 Neon (Ne) – [He]2s22p6 Argon (Ar) – [Ne]3s23p6 Krypton (Kr) – [Ar]3d104s24p6 Xenon (Xe) – [Kr]4d105s25p6 Radon (Rn) – [Xe]4f145d106s26p6 Q5_What is the electronic configuration of copper? The electronic configuration of copper is [Ar]3d104s1. This configuration disobeys the aufbau principle due to the relatively small energy gap between the 3d and the 4s orbitals. The completely filled d-orbital offers more stability than the partially filled configuration