2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp CHAPTER: 2.0 Atomic Structure MODE CHECK LIST 2.1 Bohr’s atomic model (a) Describe Bohr’s atomic model. Lecture (b) Explain the existence of energy levels in an atom. Lecture (c) Calculate the energy of an electrons using : En= -RH( 1/n2 ) , RH=2.18 x 10-18J Lecture (d) Describe the formation of line spectrum of hydrogen atom. Lecture (e) Illustrate the formation of Lyman, Balmer, Paschen, Brackett and Pfund series. Lecture (f) Calculate the energy change of an electron during transition : ∆E=( 1/ni 2 - 1/nf 2 ) , RH=2.18 x 10-18J Tutorial (g) Calculate the photon of energy emitted by an electron that produces a particular wavelength during transition. ∆E=hv, v=c/λ Tutorial (h) Perform calculations involving the Rydberg equation. 1/λ=RH( 1/n1 2 - 1/n2 2 ),RH=1.097 x 107m-1 and n1<n2 Tutorial (i) Calculate the ionisation energy of hydrogen atom from Lyman Tutorial (j) State limitation of Bohr’s atomic model. Lecture (k) State the dual nature of electron using de Broglie’s Postulate and Heisenberg’s Uncertainty Principle. Lecture 2.2 Quantum mechanical model (a) Define the term orbital. Lecture (b) Explain all four quantum numbers of an electron in an orbital: i. principal quantum number, n ii. angular momentum quantum number iii. magnetic quantum number, m iv. electron spin quantum number, s. Lecture (c) Sketch the 3-D shapes of s, p and d orbitals. Lecture 2.3 Electronic Configuration (a) Explain Aufbau principle, Hund’s rule and Pauli’s Exclusion Principle . Lecture & Tutorial (b) Predict the electronic configuration of atoms and monoatomic ions using spdf notation and orbital diagram Tutorial (c) Justify the anomalous electronic configuration of chromium and copper. Tutorial
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp 2.1 BOHR’S ATOMIC MODEL BOHR’S ATOMIC MODEL POSTULATES Emission series of Hydrogen Spectrum Series Spectrum region Electron drop to Level (GROUND STATE) Electron drops from Lyman Ultraviolet n=1 n=2,3,4,5… Balmer Visible n=2 n=3,4,5,6… Paschen Infrared n=3 n=4,5,6,7… Brackett Infrared n=4 n=5,6,7,8… Pfund Infrared n=5 n=6,7,8,9… Leman BAwa PAkcik Beli Pizza In the specific energy level, the energy of electron is fixed in value or is quantized. Electron moves in a circular orbit of certain radii with specific energy around the nucleus. At ordinary condition, the electron is at the ground state (lowest energy state). If energy is supplied, electron absorbed the energy and is promoted from lower energy level to a higher energy level (electron is excited) Electron at excited state is unstable. Electron will fall back to lower energy level and released a specific amount of energy in the form of light or photon. The light emitted has a specific wavelength. Thus, the line spectrum is formed. Formation of Line Spectrum of Hydrogen Spectrum
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp Energy level diagram of the transition of electrons Example: 1. How is the second line of Brackett series produced? 2. How is the third line of Balmer series produced? 3. How is the fourth line of Lyman series produced? 4. How is the second line of Paschen series produced? n Intial n Final (series) Line 1 st 2 nd 3 rd 4 th 5 th ENERGY LEVEL DIAGRAM show the transition of electron from higher energy to lower energy. LINE SPECTRUM Series of lines produced during the transition of electron from higher energy level to lower energy level From 1st to 5 th line: ∆ ↑, ↓, ↑ 1 st 2 nd 3 rd 4 th 5 th
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp RYDBERG EQUATION Calculate the energy of an electrons 2 1 n En RH RH(Rydberg constant) = 2.1810-18J Calculate the photon of energy emitted by an electron that produces a particular wavelength during transition ΔE hv h Planck’s constant=6.63x10-34 Js v frequency Calculate the energy change of an electron during transition 2 2 1 1 i f H n n E R RH(Rydberg constant) = 2.1810-18J c v c speed of light =3.0x108ms-1 v frequency wavelength Calculate wavelength of an electron during transition 2 2 2 1 1 1 1 n n RH n1<n2 RH (Rydberg constant) =1.097x107m-1 hc ΔE h Planck’s constant=6.63x10-34 Js c speed of light =3.0x108ms-1 wavelength Weakness (limitation) of Bohr’s atomic model 1.Unable to explain the line spectrum of atoms or ions containing more than one electron (such as helium). 2.Electron is restricted to move in a certain distance around the nucleus of an atom. 3.Unable to explain the extra lines formed in the hydrogen spectrum 4.Unable to explain the dual nature of electrons
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp 2.2 Quantum mechanical model n: Principle quantum no Determine the energy (or energy level) of the electron and the size of orbital (n=1,2,3,4) l: Angular momentum quantum no @azimuthal/subsidiary/orbital quantum no Determines the shape of orbital (integers from 0 to (n-1) l Subshell Shape of orbital 0 s Spherical 1 p Dumb-bell 2 d Cloverleaf m: Magnetic quantum no Determine the orientation of orbital in space l m Orbital 0 0 s 1 +1,0,-1 px, py, pz 2 +2,+1,0,-1,-2 dxy, dyz, dxz, dx 2 - y 2 , dz 2 s: Electron spin quantum no Determine the direction of spinning motion (clockwise or anti clockwise) of an electron n (n= 1,2,3,4) orbital size l (l=0,1,2,3) orbital shape m (m=-l,0, +l) orbital orientation S (s=-1/2, +1/2) e - spin direction Sets of quantum no 1 0 1s 0 + ½ @ - ½ (1,0,0, + ½) (1,0,0, - ½) 2 2 0 2 1 6 3 0 2 1 6 2 14
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp 2.3 ELECTRONIC CONFIGURATION Aufbau principle Hund’s rule Pauli’s Exclusion Principle Electrons fill the lowest energy orbitals first and other orbitals in order of ascending energy. Electrons are added to the orbital of equivalent energy (or degenerate orbitals), each orbital are filled singly with electron of the parallel spins (same) first before it is paired No two electrons in an atom can have the same four quantum numbers (n, , m, s). Set Quantum No: (n,l,m,s) (2,0,0, + ½) and (2,0,0, + ½) Electronic configuration Example: Write the electronic configuration Na Na+ Mn Mn2+ Fe Mn3+ O O2- Orbital diagram: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2……..
2.0 ATOMIC STRUCTURE Special Gift 2 SKO15_CHEMISTRY_2021/22_nfp Anomalous Electronic Configuration of Chromium and Copper Chromium Copper Proton no 24 29 Expected Electronic Configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d4 1s2 2s2 2p6 3s2 3p6 4s2 3d9 Actual Electronic Configuration 1s2 2s2 2p6 3s2 3p6 4s1 3d5 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Explanation The reason for these anomalous are unusual stability in half-filled orbitals of chromium. One electron from 4s orbital is promoted to 3d orbital. Hence, it has a half-filled orbital arrangement Half-filled orbital is very stable compare to partially-filled orbital. The reason for these anomalous are unusual stability in full-filled orbitals of copper. One electron from 4s orbital is promoted to 3d orbital. Hence, it has a full-filled orbital arrangement Full-filled orbital is very stable compare to partially-filled orbital.