Chapter 18 Electronic Structure of the Hydrogen Atom

Total Page:16

File Type:pdf, Size:1020Kb

Chapter 18 Electronic Structure of the Hydrogen Atom Chapter 18 Electronic Structure of the Hydrogen Atom P. J. Grandinetti Chem. 4300 P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Electronic Structure of the Hydrogen Atom Hydrogen atom is simplest atomic system where Schrödinger equation can be solved analytically and compared to experimental measurements. Analytical solution serve as basis for obtaining approximate solutions for multi-electron atoms and molecules, where no analytical solution exists. Warning: Working through analytical solution of H-atom may cause drowsiness. Do not operate heavy machinery during this derivation. P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Two Particle Problem ⃗ ⃗ 2 particles with mass mp and me at rp and re. z Total mass, center of mass, and inter-particle distance vector is ⃗ ⃗ mprp + mere M = m + m ; R⃗ = ; and ⃗r = ⃗r * ⃗r p e M e p y ⃗ ⃗ ⃗ ⃗ Express rp and re in terms of M, R and r as m mp ⃗r = R⃗ * e ⃗r and ⃗r = R⃗ + ⃗r x p M e M Express individual momenta of 2 particles as ⃗ drp d⃗r p⃗ = m and p⃗ = m e p p dt e e dt then in terms of M, R⃗ and ⃗r as 0 1 0 1 ⃗ ⃗ dR m d⃗r dR mp d⃗r p⃗ = m * e and p⃗ = m + p p dt M dt e e dt M dt P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Two Particle Problem Next, consider total energy ⃗ 2 2 pp p⃗ E = + e + V.⃗r/ 2mp 2me V.⃗r/ depends only on distance between 2 particles. Write energy in terms of M, R⃗ and ⃗r, and obtain 0 1 0 1 2 2 1 dR⃗ 1 d⃗r E = M + 휇 + V.⃗r/ 2 dt 2 dt 휇 is reduced mass, given by 1 1 1 휇 = + mp me Define 2 new momenta associated with center of mass and reduced mass, dR⃗ d⃗r p⃗ = M ; and p⃗ = 휇 R dt r dt P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Two Particle Problem Total energy becomes p⃗ 2 p⃗ 2 E = R + r + V.⃗r/ 2M 2휇 1st term is translational energy of center of mass 2nd term is kinetic energy due to relative motion of 2 particles Translating to quantum mechanics, we write time independent Schrödinger equation for 2 particle system as L M 4 5 p⃗ 2 p⃗ 2 `2 `2 R + r + V.⃗r/ f .R⃗; ⃗r/ = * (⃗ 2 * (⃗ 2 + V.⃗r/ f .R⃗; ⃗r/ = Ef .R⃗; ⃗r/ 2M 2휇 2M R 2휇 r P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Two Particle Problem 4 5 `2 `2 * (⃗ 2 * (⃗ 2 + V.⃗r/ f .R⃗; ⃗r/ = Ef .R⃗; ⃗r/ 2M R 2휇 r Wave function can be separated into product of two wave functions f .R⃗; ⃗r/ = 휒.R⃗/ .⃗r/ 휒.R⃗/ depending only on center of mass .⃗r/ depending only on relative motion of 2 particles Substitute product into Schrödinger Eq above, we obtain 2 wave equations 4 5 `2 `2 * (⃗ 2 휒.R⃗/ = E 휒.R⃗/ and * (⃗ 2 + V.⃗r/ .⃗r/ = E .⃗r/ 2M R R 2휇 r r where E = ER + Er On left is wave equation for translational motion of free particle of mass M On right is wave equation for particle with mass 휇 in potential V.⃗r/ For electron bound positively charged nucleus, we focus on PDE for .⃗r/ P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Schrödinger Equation in Spherical Coordinates P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Schrödinger Equation in Spherical Coordinates Focus on PDE for .⃗r/: Electron bound to positively charged nucleus 4 5 `2 * (2 + V̂ .r/ .⃗r/ = E .⃗r/ 2휇 Coulomb potential is 2 ̂ Zqe V.r/ = * 휋휖 4 0r Central potential with 1_r dependence. V → −∞ as r → 0 V → 0 as r → ∞ P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Schrödinger Equation in Spherical Coordinates Since V̂ .r/ only depends on r, we adopt spherical coordinates 4 0 ( ) 15 `2 )2 2 ) 1 1 ) ) 1 )2 * + + sin 휃 + + V̂ .r/ = E 2휇 )r2 r )r r2 sin 휃 )휃 )휃 sin2 휃 )휙2 «­­­­­­­­­­­­­­­­­­­­­­­­­­¯­­­­­­­­­­­­­­­­­­­­­­­­­­¬ *L̂ 2_`2 «­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­¯­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­­¬ (2 Take term in parentheses as *L̂ 2_`2, rearrange, and simplify to )2 .r; 휃; 휙/ ) .r; 휃; 휙/ 2휇r2 ( ) 1 r2 + 2r + E * V̂ .r/ .r; 휃; 휙/ = L̂ 2 .r; 휃; 휙/ )r2 )r `2 `2 To solve PDE use separation of variables .r; 휃; 휙/ = R.r/Y.휃; 휙/ P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Schrödinger Equation in Spherical Coordinates Substitute .r; 휃; 휙/ into PDE, and dividing both sides by .r; 휃; 휙/ r2 )2R.r/ 2r )R.r/ 2휇r2 ( ) 1 1 + + E * V̂ .r/ = L̂ 2Y.휃; 휙/ R.r/ )r2 R.r/ )r `2 Y.휃; 휙/ `2 Left side depends only on r Right side depends only on 휃 and 휙. We have turned one PDE into two ODEs. ̂ 1 We know eigenfunctions and eigenvalues of L2 on right side, ̂ 2 휃; 휙 l l `2 휃; 휙 L Yl;m. / = . + 1/ Yl;m. / 2 Define l.l + 1/ as separation constant and obtain ODE for radial part r2 d2R.r/ 2r dR.r/ 2휇r2 ( ) + + E * V̂ .r/ = l.l + 1/ R.r/ dr2 R.r/ dr `2 P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Schrödinger Equation in Spherical Coordinates Expanding and rearranging ODE for radial part, we obtain 4 0 1 5 `2 d2 2 d `2 l.l + 1/ * + + + V̂ .r/ R.r/ = ER.r/ 2휇 )r2 r dr 2휇 r2 «­­­­­­­­­­­¯­­­­­­­­­­­¬ Veff.r/ Looks like 1D time independent Schrödinger Eq with effective potential `2 l.l + 1/ Zq2 V̂ .r/ = * e eff 2휇 r2 4휋휖 r «­­­­¯­­­­¬ 0 Centrifugal Term ù Angular momentum, l.l + 1/`, creates centrifugal force that pushes electron away from nucleus. ̂ Notice 2 terms in Veff.r/ always have opposite signs. P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom l Plot of Veff for = 0, 1, 2, and 3. 0.2 0.1 0.0 -0.1 -0.2 -0.3 -0.4 0 2 4 6 8 10 12 14 Note: potential minimum shifts to higher radii with increasing angular momentum (i.e., centrifugal force). P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom ̂ Veff.r/: Effective Potential of H-Like Atom Switching to Atomic Units: Convenient units for atomic physics Atomic unit of length, also called Bohr radius, is defined as 4휋휖 `2 a ≡ 0 = 52:9177210526763 pm 0 2 qeme Atomic unit of energy is `2 q2 1E ≡ = e = 27:21138602818051 eV h 2 4휋휖 a mea0 0 0 P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Solving the Wave Equation for the Radial Part L 0 1 M 2 `2 d2 2 d `2 l.l + 1/ Zq * + + * e R.r/ = ER.r/ 휇 2 휇 2 휋휖 2 )r r dr 2 r 4 0r pour yourself a fifth cup of coffee... P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Solving the Radial Wave Equation To obtain radial wave function, R.r/, we must solve L 0 1 M 2 `2 d2 2 d `2 l.l + 1/ Zq * + + * e R.r/ = ER.r/ 휇 2 휇 2 휋휖 2 )r r dr 2 r 4 0r Begin by dividing and both sides by E, and obtain 4 5 H I 2 `2 d2 2 d `2 l.l + 1/ Zq * + R.r/ + * e * 1 R.r/ = 0 휇 2 휇 2 휋휖 2 E )r r dr 2 E r 4 0rE Define 2휇E `2휅2 휅2 ≡ * or E = * `2 2휇 휅 is in wave numbers. Rearrange potential energy expression to 2 2 2 Zq Zq 2휇 Zq 휇 1 2휌 * e = e = e = 0 휋휖 휋휖 `2휅2 휋휖 `2휅 휅 휅 4 0rE 4 0r 2 0 r r Zq2휇 휌 is a dimensionless quantity: 휌 ≡ e 0 0 휋휖 `2휅 2 0 P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Solving the Radial Wave Equation Substitute expression for potential energy into ODE 4 5 0 1 1 d2 2 d l.l + 1/ 2휌 + R.r/* 1 + + 0 R.r/ = 0 휅2 )r2 r dr .휅r/2 휅r Define dimensionless variable: 휌 ≡ 2휅r Rewrite radial part as function of 휌 as 4 5 0 1 d2 2 d l.l + 1/ 2휌 + R.휌/* 1 + + 0 R.휌/ = 0 d휌2 휌 d휌 휌2 휌 Further simplify ODE by defining du.휌/ dR.휌/ d2u.휌/ d2R.휌/ dR.휌/ u.휌/ = 휌R.휌/; = 휌 + R.휌/; = 휌 + 2 d휌 d휌 d휌2 d휌2 d휌 Re-express ODE as 0 1 d2u.휌/ 1 휌 l.l + 1/ * * 0 + u.휌/ = 0 d휌2 4 2휌 휌2 휌 varies from 0 to ∞. 1st look at asymptotic solutions: (1) 휌 → ∞ and (2) 휌 → 0 P. J. Grandinetti Chapter 18: Electronic Structure of the Hydrogen Atom Solving the Radial Wave Equation, 휌 → ∞ Starting with 0 1 d2u.휌/ 1 휌 l.l + 1/ * * 0 + u.휌/ = 0 d휌2 4 2휌 휌2 (1) 휌 → ∞. At large values of 휌 approximate ODE as d2u.휌/ u.휌/ * ù 0 d휌2 4 Solutions to this ODE have form u.휌/ ∼ Ae*휌_2 + Be휌_2 Reject positive exponent since require that solution be finite everywhere, u.휌/ ∼ Ae*휌_2 in limit that 휌 → ∞ P.
Recommended publications
  • Quantum Theory of the Hydrogen Atom
    Quantum Theory of the Hydrogen Atom Chemistry 35 Fall 2000 Balmer and the Hydrogen Spectrum n 1885: Johann Balmer, a Swiss schoolteacher, empirically deduced a formula which predicted the wavelengths of emission for Hydrogen: l (in Å) = 3645.6 x n2 for n = 3, 4, 5, 6 n2 -4 •Predicts the wavelengths of the 4 visible emission lines from Hydrogen (which are called the Balmer Series) •Implies that there is some underlying order in the atom that results in this deceptively simple equation. 2 1 The Bohr Atom n 1913: Niels Bohr uses quantum theory to explain the origin of the line spectrum of hydrogen 1. The electron in a hydrogen atom can exist only in discrete orbits 2. The orbits are circular paths about the nucleus at varying radii 3. Each orbit corresponds to a particular energy 4. Orbit energies increase with increasing radii 5. The lowest energy orbit is called the ground state 6. After absorbing energy, the e- jumps to a higher energy orbit (an excited state) 7. When the e- drops down to a lower energy orbit, the energy lost can be given off as a quantum of light 8. The energy of the photon emitted is equal to the difference in energies of the two orbits involved 3 Mohr Bohr n Mathematically, Bohr equated the two forces acting on the orbiting electron: coulombic attraction = centrifugal accelleration 2 2 2 -(Z/4peo)(e /r ) = m(v /r) n Rearranging and making the wild assumption: mvr = n(h/2p) n e- angular momentum can only have certain quantified values in whole multiples of h/2p 4 2 Hydrogen Energy Levels n Based on this model, Bohr arrived at a simple equation to calculate the electron energy levels in hydrogen: 2 En = -RH(1/n ) for n = 1, 2, 3, 4, .
    [Show full text]
  • Principal, Azimuthal and Magnetic Quantum Numbers and the Magnitude of Their Values
    268 A Textbook of Physical Chemistry – Volume I Principal, Azimuthal and Magnetic Quantum Numbers and the Magnitude of Their Values The Schrodinger wave equation for hydrogen and hydrogen-like species in the polar coordinates can be written as: 1 휕 휕휓 1 휕 휕휓 1 휕2휓 8휋2휇 푍푒2 (406) [ (푟2 ) + (푆푖푛휃 ) + ] + (퐸 + ) 휓 = 0 푟2 휕푟 휕푟 푆푖푛휃 휕휃 휕휃 푆푖푛2휃 휕휙2 ℎ2 푟 After separating the variables present in the equation given above, the solution of the differential equation was found to be 휓푛,푙,푚(푟, 휃, 휙) = 푅푛,푙. 훩푙,푚. 훷푚 (407) 2푍푟 푘 (408) 3 푙 푘=푛−푙−1 (−1)푘+1[(푛 + 푙)!]2 ( ) 2푍 (푛 − 푙 − 1)! 푍푟 2푍푟 푛푎 √ 0 = ( ) [ 3] . exp (− ) . ( ) . ∑ 푛푎0 2푛{(푛 + 푙)!} 푛푎0 푛푎0 (푛 − 푙 − 1 − 푘)! (2푙 + 1 + 푘)! 푘! 푘=0 (2푙 + 1)(푙 − 푚)! 1 × √ . 푃푚(퐶표푠 휃) × √ 푒푖푚휙 2(푙 + 푚)! 푙 2휋 It is obvious that the solution of equation (406) contains three discrete (n, l, m) and three continuous (r, θ, ϕ) variables. In order to be a well-behaved function, there are some conditions over the values of discrete variables that must be followed i.e. boundary conditions. Therefore, we can conclude that principal (n), azimuthal (l) and magnetic (m) quantum numbers are obtained as a solution of the Schrodinger wave equation for hydrogen atom; and these quantum numbers are used to define various quantum mechanical states. In this section, we will discuss the properties and significance of all these three quantum numbers one by one. Principal Quantum Number The principal quantum number is denoted by the symbol n; and can have value 1, 2, 3, 4, 5…..∞.
    [Show full text]
  • Further Quantum Physics
    Further Quantum Physics Concepts in quantum physics and the structure of hydrogen and helium atoms Prof Andrew Steane January 18, 2005 2 Contents 1 Introduction 7 1.1 Quantum physics and atoms . 7 1.1.1 The role of classical and quantum mechanics . 9 1.2 Atomic physics—some preliminaries . .... 9 1.2.1 Textbooks...................................... 10 2 The 1-dimensional projectile: an example for revision 11 2.1 Classicaltreatment................................. ..... 11 2.2 Quantum treatment . 13 2.2.1 Mainfeatures..................................... 13 2.2.2 Precise quantum analysis . 13 3 Hydrogen 17 3.1 Some semi-classical estimates . 17 3.2 2-body system: reduced mass . 18 3.2.1 Reduced mass in quantum 2-body problem . 19 3.3 Solution of Schr¨odinger equation for hydrogen . ..... 20 3.3.1 General features of the radial solution . 21 3.3.2 Precisesolution.................................. 21 3.3.3 Meanradius...................................... 25 3.3.4 How to remember hydrogen . 25 3.3.5 Mainpoints.................................... 25 3.3.6 Appendix on series solution of hydrogen equation, off syllabus . 26 3 4 CONTENTS 4 Hydrogen-like systems and spectra 27 4.1 Hydrogen-like systems . 27 4.2 Spectroscopy ........................................ 29 4.2.1 Main points for use of grating spectrograph . ...... 29 4.2.2 Resolution...................................... 30 4.2.3 Usefulness of both emission and absorption methods . 30 4.3 The spectrum for hydrogen . 31 5 Introduction to fine structure and spin 33 5.1 Experimental observation of fine structure . ..... 33 5.2 TheDiracresult ..................................... 34 5.3 Schr¨odinger method to account for fine structure . 35 5.4 Physical nature of orbital and spin angular momenta .
    [Show full text]
  • 1.2 Formulation of the Schrödinger Equation for the Hydrogen Atom
    2 1 Hydrogen orbitals 1.2 Formulation of the Schrodinger¨ equation for the hydrogen atom In this initial treatment, we will make some practical approximations and simplifi- cations. Since we are for the moment only trying to establish the general form of the hydrogenic wave functions, this will suffice. To start with, we will assume that the nucleus has zero extension. We place the origin at its position, and we ignore the centre-of-mass motion. This reduces the two-body problem to a single particle, the electron, moving in a central-field potential. To take the finite mass of the nu- cleus into account, we replace the electron mass with the reduced mass, m, of the two-body problem. Moreover, we will in this chapter ignore the effect on the wave function of relativistic effects, which automatically implies that we ignore the spins of the electron and of the nucleus. This makes us ready to formulate the Hamilto- nian. The potential is the classical Coulomb interaction between two particles of op- posite charges. With spherical coordinates, and with r as the radial distance of the electron from the origin, this is: Ze2 V(r) = − ; (1.1) 4pe0 r with Z being the charge state of the nucleus. The Schrodinger¨ equation is: h¯ 2 − ∇2y(r) +V(r)y = Ey(r) ; (1.2) 2m where the Laplacian in spherical coordinates is: 1 ¶ ¶ 1 ¶ ¶ 1 ¶ 2 ∇2 = r2 + sinq + : (1.3) r2 ¶r ¶r r2 sinq ¶q ¶q r2 sin2 q ¶j2 Since the potential is purely central, the solution to (1.2) can be factorised into a radial and an angular part, y(r;q;j) = R(r)Y(q;j).
    [Show full text]
  • The Principal Quantum Number the Azimuthal Quantum Number The
    To completely describe an electron in an atom, four quantum numbers are needed: energy (n), angular momentum (ℓ), magnetic moment (mℓ), and spin (ms). The Principal Quantum Number This quantum number describes the electron shell or energy level of an atom. The value of n ranges from 1 to the shell containing the outermost electron of that atom. For example, in caesium (Cs), the outermost valence electron is in the shell with energy level 6, so an electron incaesium can have an n value from 1 to 6. For particles in a time-independent potential, as per the Schrödinger equation, it also labels the nth eigen value of Hamiltonian (H). This number has a dependence only on the distance between the electron and the nucleus (i.e. the radial coordinate r). The average distance increases with n, thus quantum states with different principal quantum numbers are said to belong to different shells. The Azimuthal Quantum Number The angular or orbital quantum number, describes the sub-shell and gives the magnitude of the orbital angular momentum through the relation. ℓ = 0 is called an s orbital, ℓ = 1 a p orbital, ℓ = 2 a d orbital, and ℓ = 3 an f orbital. The value of ℓ ranges from 0 to n − 1 because the first p orbital (ℓ = 1) appears in the second electron shell (n = 2), the first d orbital (ℓ = 2) appears in the third shell (n = 3), and so on. This quantum number specifies the shape of an atomic orbital and strongly influences chemical bonds and bond angles.
    [Show full text]
  • Importance of Hydrogen Atom
    Hydrogen Atom Dragica Vasileska Arizona State University Importance of Hydrogen Atom • Hydrogen is the simplest atom • The quantum numbers used to characterize the allowed states of hydrogen can also be used to describe (approximately) the allowed states of more complex atoms – This enables us to understand the periodic table • The hydrogen atom is an ideal system for performing precise comparisons of theory and experiment – Also for improving our understanding of atomic structure • Much of what we know about the hydrogen atom can be extended to other single-electron ions – For example, He+ and Li2+ Early Models of the Atom • ’ J.J. Thomson s model of the atom – A volume of positive charge – Electrons embedded throughout the volume • ’ A change from Newton s model of the atom as a tiny, hard, indestructible sphere “ ” watermelon model Experimental tests Expect: 1. Mostly small angle scattering 2. No backward scattering events Results: 1. Mostly small scattering events 2. Several backward scatterings!!! Early Models of the Atom • ’ Rutherford s model – Planetary model – Based on results of thin foil experiments – Positive charge is concentrated in the center of the atom, called the nucleus – Electrons orbit the nucleus like planets orbit the sun ’ Problem: Rutherford s model “ ” ’ × – The size of the atom in Rutherford s model is about 1.0 10 10 m. (a) Determine the attractive electrical force between an electron and a proton separated by this distance. (b) Determine (in eV) the electrical potential energy of the atom. “ ” ’ × – The size of the atom in Rutherford s model is about 1.0 10 10 m. (a) Determine the attractive electrical force between an electron and a proton separated by this distance.
    [Show full text]
  • The Hydrogen Atom: Consideration of the Electron Self-Field Arxiv
    The hydrogen atom: consideration of the electron self-field Biguaa L.V ∗ Kassandrov V.V † ∗ Quantum technology center Moscow State University y Institute of Gravitation and Cosmology, Peoples’ Friendship University of Russia January 8, 2021 Keywords: Spinor electrodynamics. Dirac-Maxwell system of equations. Solitonlike solu- tions. Spectrum of characteristics. “Bohrian” law for binding energies PACS: 03.50.-z; 03.65.-w; 03.65.-Pm 1 THE HYDROGEN ATOM AND CLASSICAL FIELD MODELS OF ELEMENTARY PARTICLES It is well known that the successful description of the hydrogen atom, first in Bohr’s theory, and then via solutions to the Schr¨odinger equation was one of the main motivations for the development of quantum theory. Later, Dirac described almost perfectly the observed hydrogen spectrum using his relativistic generalization of the Schr¨odingerequation. A slight deviation from experiment (the Lamb shift between 2s1=2 and 2p1=2 levels in the hydrogen atom) arXiv:2101.02202v1 [quant-ph] 5 Jan 2021 discovered later was interpreted as the effect of electron interaction with vacuum fluctuations of electromagnetic field and explained in the framework of quantum electrodynamics (QED) using second quantization. Nonetheless, the problem of the description of the hydrogen atom still attracts attention of researchers, and, in some sense, is a “touchstone” for many new theoretical developments. In particular, it is possible to obtain the energy spectrum via purely algebraic methods [1]. Theories of the ”supersymmetric” hydrogen atom are also widespread [2]. In the framework of the classical ideas, the spectrum can be explained by the absence of radiation for the electron ∗E-mail: [email protected] †E-mail: [email protected] 1 motion along certain “stable” orbits.
    [Show full text]
  • The Hydrogen Atom
    Quantum Mechanics: The Hydrogen Atom 13th April 2011 I. The Hydrogen Atom In this next section, we will tie together the elements of the last several sections to arrive at a complete description of the hydrogen atom. This will cul- minate in the definition of the hydrogen-atom orbitals and associated energies. From these functions, taken as a complete basis, we will be able to construct approximations to more complex wave functions for more complex molecules. Thus, the work of the last few lectures has fundamentally been aimed at estab- lishing a foundation for more complex problems in terms of exact solutions for smaller, model problems. II. The Radial Function We will start by reiterating the Schrodinger equation in 3D spherical coordi- nates as (refer to any standard text to get the transformation from Cartesian to spherical coordinate reference systems). Here, we have not placed the constraint of a constant distance separting the masses of the rigid rotor (refer to last lec- ture); furthermore, we will keep in the formulation the potential V (r; θ; φ) for generality. Thus, in spherical polar coordinates, H^ (r; θ; φ) (r; θ; φ) = E (r; θ; φ) be- comes: E (r; θ; φ) = −h¯2 1 @ @ 1 @ @ 1 @2 r2 + sinθ + + V (r; θ; φ) (r; θ; φ) 2µ r2 @r @r r2sinθ @θ @θ r2sin2θ @φ2 Now, for the hydrogen atom, with one electron found in "orbits" (note the quotes!) around the nucleus of charge +1, we can include an electrostatic poten- tial which is essentially the Coulomb potential between a positive and negative charge: −Ze2 V (jrj) = V (r) = 4πor where Z is the nuclear charge (i.e, +1 for the nucleus of a hydrogen atom).
    [Show full text]
  • The Hydrogen Atom Atomic Emission Spectroscopy
    The Hydrogen atom Atomic Emission Spectroscopy the atomic emission spectrum is the result of many atoms emitting light simultaneously. a spectrometer Absorption and Emission Spectroscopy Light will be absorbed or emitted only if ∆E = hν The spectrum is QUANTIZED! ∆E51=24G ∆E41=15G ∆E31=8G ∆E21=3G G= h2/8mL2 Atomic Emission Spectroscopy The emission spectra of H, He and Hg are all quantized! Energy eigenvectors/eigenstates is the eigenstate vector solution to the Schrödinger equation. is the energy eigenvalue associated with this eigenstate vector. Hydrogen eigenstates/wavefunctions State Vectors Three Dimensional Wavefunction (Hydrogen Atom) The spatial component of the eigenstate vector can be represented by a three dimensional wavefunction ψ(r) = ψ(x,y,z). The Hydrogen Atom Hamiltonian V(r) is called the Coulomb potential. WHEN we solve the Schrödinger equation, we always obtain two things: 1. a set of eigenstate energies, En. Example: PIAB n = 1, 2, 3... WHEN we solve the Schrödinger equation, we always obtain two things: 1. a set of eigenstate energies, En. 2. a set of eigenstate vectors |ψn>. Example: PIAB n = 1, 2, 3... Eigenstate vectors are labeled with the quantum number n. WHEN we solve the Schrödinger equation, we always obtain two things: 1. a set of eigenstate energies, En. Hydrogen atom solution: n = 1, 2, 3... WHEN we solve the Schrödinger equation, we always obtain two things: 1. a set of eigenstate energies, En. 2. a set of eigenstate vectors |ψn>. Hydrogen atom solution: n = 1, 2, 3... State vectors are labeled with the quantum number n... AND three more: l, ml, m s WHEN we solve the Schrödinger equation, we always obtain two things: 1.
    [Show full text]
  • Name___Making Models of Atoms & Isotopes
    Hydrogen Models 1 Name______________________ Making Models of Atoms & Isotopes Problem: To build 3 dimensional models of the hydrogen atom. Background Information: An atom is defined as Re-write the main ideas of the background a small particle that makes up most types of matter. Atoms are so small it would take about 1 information here: million of them lined up in a row to equal the thickness of a human hair. Atoms are made up ____________________________________ of even smaller particles. The largest of these ____________________________________ particles are protons, neutrons and electrons. The identity of a type of matter depends on the ____________________________________ number of protons in the nucleus of an atom of that type of matter. All atoms of the same type of ____________________________________ matter have the same number of protons in the ____________________________________ nucleus. For example, all carbon atoms have six protons. ____________________________________ Not all atoms of the same type of matter have the ____________________________________ same number of neutrons. Most carbon atoms ____________________________________ have 6 neutrons, although some have more and some have less. Atoms of the same type of ____________________________________ matter that have different numbers of neutrons are called isotopes. Most types of matter have ____________________________________ isotopes. The atomic mass, or mass number, can be used to identify isotopes. The mass number of a type of matter is the number of protons and neutrons in one atom of that matter. In our example of carbon, 6 protons + 6 neutrons = 12 particles in the nucleus; the mass number = 12. The name for this isotope is Carbon 12. The symbol for Carbon 12 is: Mass Number 12 C A tomic Number 6 Diagrams of carbon isotopes: M.
    [Show full text]
  • Hydrogen Atomic and Molecular Hydrogen
    Hydrogen Hydrogen, the simplest atom, combines with itself to make the simplest molecule, H2. Most other elements in the periodic table combine with hydrogen to make compounds. The chemical bonds between hydrogen and other elements result from sharing electrons. The electrons are shared equally in non-polar covalent bonds and unequally in polar covalent bonds. The chemical energy stored in covalent bonds can be transformed into heat or electrical energy. Outline • Atomic and Molecular Hydrogen • Hydrogen Compounds • Energy from Hydrogen • Homework Atomic and Molecular Hydrogen Electronic Structure of Atomic Hydrogen We can represent a hydrogen atom with the symbol H and a dot that stands for the single electron in the neutral atom. In its lowest energy form, the electron occupies the 1s orbital, spherically symmetrical around the hydrogen nucleus. As you saw in the section on electronic configuration, all atoms have many possible energy levels for their electrons. The lowest energy state (ground state) of an atom has its electrons filling orbitals beginning from lowest energy, with 2 electrons per orbital. When energy is added to the ground state atom, an electron can be promoted to one of the higher energy orbitals. The energy added must be exactly the same as the energy between the orbital occupied by the electron and one of the higher energy, unoccupied orbitals. Below you can see an orbital energy diagram showing the ground state hydrogen atom on the left. When hydrogen absorbs a quantity of energy exactly equal to ΔE1, the electron goes from the orbital in the first shell (n = 1) to an orbital in the second shell (n = 2).
    [Show full text]
  • 28-1 Line Spectra and the Hydrogen Atom Figure 28.1 Gives Some Examples of the Line Spectra Emitted by Atoms of Gas
    28-1 Line Spectra and the Hydrogen Atom Figure 28.1 gives some examples of the line spectra emitted by atoms of gas. The atoms are typically excited by applying a high voltage across a glass tube that contains a particular gas. By observing the light through a diffraction grating, the light is separated into a set of wavelengths that characterizes the element. The spectrum of light emitted by excited hydrogen atoms is shown Figure 28.1: Line spectra from hydrogen (top) and helium in Figure 28.1(a). The decoding of the (bottom). A line spectrum is like the fingerprint of an element. hydrogen spectrum represents one of Astronomers, for instance, can determine what a star is made of by the great scientific mystery stories. carefully examining the spectrum of light emitted by the star. First on the scene was the Swiss mathematician and schoolteacher, Johann Jakob Balmer (1825 – 1898), who published an equation in 1885 giving the wavelengths in the visible spectrum emitted by hydrogen. The Swedish physicist Johannes Rydberg (1854 – 1919) followed up on Balmer’s work in 1888 with a more general equation that predicted all the wavelengths of light emitted by hydrogen: , (Eq. 28.1: The Rydberg equation for the hydrogen spectrum) 7 –1 where R = 1.097 ! 10 m is the Rydberg constant, and the two n’s are integers, with n2 greater than n1. Neither Balmer nor Rydberg had a physical explanation to justify their equations, however, so the search was on for such a physical explanation. The breakthrough was made by the Danish physicist Niels Bohr (1885 – 1962), who showed that if the angular momentum of the electron in a hydrogen atom was quantized in a particular way (related to Planck’s constant, in fact), that the energy levels for an electron within the hydrogen atom were also quantized, with the energies of the electrons being given by: , (Eq.
    [Show full text]