Atomic Structure Chapter 1 PDF

COLLEGE OF ENGINEERING Department of Mechanical and Industrial Engineering ATOMIC STRUCTURE Smallest unit of matter  Lesson Plan Introduction to Atom, elements, compounds and mixtures. Various Theorie...

COLLEGE OF ENGINEERING Department of Mechanical and Industrial Engineering ATOMIC STRUCTURE Smallest unit of matter  Lesson Plan Introduction to Atom, elements, compounds and mixtures. Various Theories of Atomic structure Bohr Model [Numerical on transitions], Somerfield Model Schrodinger Model Dual nature of electron (De Broglie) [Evidences, examples] Heisenberg Uncertainty Principle [Evidences, examples] Concept of atomic orbitals Quantum Numbers Pauli’s Exclusion Principle Hund’s Rule Numerical (Energy levels, electronic configuration)  Let’s Refresh Atom: An atom is the smallest particle of an element that can exist and still have properties of element. Atoms consist of principally three fundamental particles (building blocks of atoms): electrons, protons and neutrons. Ions: are charged particles carrying either a negative or a positive charge. Ions are called as Cation (+) or Anion (-) Molecule: A molecule is formed when the atoms of the same or different elements combine. Identify the names of fundamental particles that makes protons, neutrons & electrons.  Rutherford model In 1911 Ernest Rutherford, carried out an experiment to test Thomson’s Plum Pudding Rutherford Experiment: involved directing 𝛼 model. He expected the 𝛼 particles to travel through the particles at a thin sheet of metal foil. foil with, at the most, very minor deflections in their paths. 3/7/2021 5 Gold Foil Experiment 6 Observation Conclusion ⮚ Most of the space inside the atom is ⮚ Most of the alpha particles passed empty undeflected through the foil ⮚ As only few positively charged particles ⮚ A small fraction was deflected by were deflected, this shows that positive charge is not spread uniformly inside the small angles atom ⮚ Very few alpha particles (1 in ⮚ Volume occupied by nucleus is negligibly small as compared to the size 20,000) bounced back by getting of the atom. deflected by 180ᵒ  Rutherford model (Limitations) Examine the failure of Rutherford model. Rutherford was unable to explain the stability of atom. Did not explain about the arrangement of electrons around the nucleus and energies of these electrons  Bohr’s model Positive charge at the nucleus. Who was Neils Bohr? In Electrons revolve around the nucleus which year he founded this model? in a fix orbits. These orbits/energy levels are quantized. (?) Electrons don’t lose energy in these orbits. Certain amount of energy (How much ?) is required for electron to move to higher energy level. After 10-5 sec it loses same energy and come back to previous state.  Describe Limitations The Bohr Model considers electrons to have both a known radius and orbit, which is impossible as the electrons are continuously moving. The Bohr Model is very limited in terms of size of the atom. The Bohr Model does not account for the fact that accelerating electrons do not emit electromagnetic radiation. This model could not explain the ability of atoms to form molecules by chemical bonds.  Remember According to the Bohr model of the which orbit has the minimum energy? 1.K 2.L 3.M 4.N  Calculate What is the change in energy when the electron relaxes from n=3 to n=2? Calculate the energy required by an electron to jump from level 1 to level 4?  Somerfield model He suggested that each shell is made of subshells. Like 2nd shell is made up of 1s and 1p subshell. Evaluate how Sommerfield model was different from Bohr model?  Schrodinger Model In 1926 Erwin Schrödinger, an Austrian physicist, developed electron cloud model. Schrödinger combined the mathematical equations for the behavior of waves (with the de Broglie equation) to generate a mathematical model for the distribution of electrons in an atom. This mathematical equations describe the possibility of finding an electron in a certain position.  Schrodinger Model… This model can be portrayed as a nucleus surrounded by an electron cloud. Where the cloud is most dense, the probability of finding the electron is greatest, and conversely, the electron is less likely to be in a less dense area of the cloud. Describe Schrodinger Electron cloud Model  Evaluate 1. How Bohr model is different from electron cloud model? (2 points) 2. How negative cloud density is related with probability of finding of electrons? 16  De Broglie In 1924 de Broglie (In 1929 he won Nobel award for his PhD thesis) reasoned that if light wave can behave like a stream of particles (photons) then perhaps electrons can possess wave properties. He said an electron bound to the nucleus behave like a stationary wave (a wave which do not travel). de Broglie reason lead to the conclusion that wave particle duality exists for particles. And he presented the formula: λ = h/mv [h Plank’s constant, m is particle mass, v is particle velocity]  De Broglie….. Significance of de-Broglie equation: Its true for all particles. But NOT significant for big particles or we can say in our daily life, it is not observable. Its significant for subatomic particles like electrons, protons etc. Verification of de-Broglie equation. i. Particle nature (Photoelectric effect) ii. Wave nature (Davison Germer experiment) Video on Wave nature of particle https://www.youtube.com/watch?v=MTuyEn-ngIQ  Solved Example If an electron is moving with a velocity of 2*10 6 m/s, find the wavelength of the electron. Given Data: Velocity of the electron, v =2×106 ms-1 Mass of electron, m =9.1×10-31 Kg Planck’s Constant, h = 6.62607015×10 −34 Js The de-Broglie wavelength is given by λ = h/mv = 6.62607015×10−34 /(2×106)(9.1×10-31 ) λ = 0.364×109m 19  Now You Calculate If Ahmed kicked a football of 0.3 kg with a speed of 300 m/s, then calculate the wavelength associated with the football. Also comment on the answer. What is given data here? 20  Heisenberg Uncertainty In our daily we can precisely determine and predict the instant position as well as velocity of a body. But in case of small particles like electrons it’s position and velocity at a given instant can not be determined with absolute accuracy. It is called Heisenberg uncertainty principle. It’s a direct consequence of the dual nature of matter and radiation. Mathematically Δx * Δp ≥ h/4π Here, Δx is uncertainty in position, Δv, is uncertainty in velocity, m is mass, h is planks constant.  Understand, Evaluate, Analyze Learn, why it is impossible to observe Heisenberg principle in daily life? How Bohr’s model fail in Heisenberg principle? Why electron can not exist inside nucleus? Collect argument, & evidences  Solved Example Uncertainty of the position of a particle with mass = 3.0×10−30 kg is 2.0×10−10 m. If the uncertainty in momentum is 60% of the momentum, what is the speed of this particle? Given Data??? Solution: The momentum of the particle is, p=mv, Here, v is the speed of the particle. As per Uncertainty principle: Δx⋅Δp≥h4π Since, Δp=60%  0.6p  p=0.6mv; Δx⋅(0.6mv)=h/4π 2.0×10−10×0.6×3×10−30×v =6.626×10−34 J.s/4π v≈1.5×105m/s 23  Now you calculate Position of a chloride ion on a material can be determined to a maximum error of 1μm. If the mass of the chloride ion is 5.86 × 10-26Kg, what will be the error in its velocity measurement? Answer 9*10-4 m/s More problems at Heisenberg Uncertainty Principle - Detailed Explanation, Fo rmula and Derivation, (byjus.com) Problem 2 is given here for you to solve. 24  Concept of Atomic Orbital ORBIT ORBITAL It is a region of space around the It is well-defined circular path nucleus where the probability of followed by electron around nucleus. finding an electron is maximum. It represents two dimensional It represents three dimensional motion motion of electron around nucleus. of electron around nucleus. The maximum no. of electrons in an The maximum no. of electrons in an orbit is 2n2. orbital is 2. Orbit is circular in shape. Orbitals have different shapes. Shapes of Orbitals 26  Quantum Numbers Quantum Numbers: Exact location and energy state of an electron. [Full profile of an electron] n = Principal quantum number. Tell about main energy level. Always integer l = Subsidiary quantum number. Tell about sub energy level within main energy level. Values range from 0 to (n-1) ml = Magnetic quantum number. Tell about orientation of orbital in space. Values range from +l to – l including zero. ms = Spin quantum number. Tell about spin state of an electron. [+1/2, -1/2]  Examples How write “l” and “ml” values? ▪ For n=4, ‘l’ values will be 0,1,2 & 3 ▪ If l=0 🡪 s orbital 🡪 ml = 0 ▪ If l=1 🡪 p orbital 🡪 ml = -1,0,+1 ▪ If l=2 🡪 d orbital 🡪 ml = -2,-1,0,+1,+2 ▪ If l=3 🡪 f orbital 🡪 ml = -3,-2,- 1,0,+1,+2,+3 ▪ Write all quantum numbers for 3p 1 electron. Test your understanding  Test your understanding Write all quantum number for following electrons: 2s1 (n=2, l=0, ml = 0, ms = +1/2) 2p3 (n=2, l = 1, ml= -1, 0, +1; ms = +1/2, +1/2, +1/2 ) 3d6 4d2  Aufbau Principle The orbitals of lower energy are filled in first with the electrons and only then the orbitals of high energy are filled. (n+l) rule: Electron will prefer the orbital where value of n+l is smaller. If n+l values are same then electron will prefer to enter in orbital with smaller n value. For example electrons will enter in 3p and then in 4s. After that 3d.  Examples Why is electronic configuration of Cr and Cu are special?  Test your understanding A. Write full electronic configuration for following elements: Z= 12 and Z= 36 B. Another way to write configuration in short in Noble gas notations. Write for Z=12, and Z= 20  Pauli’s exclusion Principle This principle was formulated by Austrian physicist Wolfgang Pauli in 1925 for electrons. It is impossible for two electrons of a poly-electron atom to have the same values of the four quantum numbers. For example, if two electrons reside in the same orbital, and if their n, ℓ, and mℓ values are the same, then their ms must be different, and thus the electrons must have opposite half-integer spins of 1/2 and −1/2.  Hund’s Rule of maximum multiplicity Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.