1 / 22

Electromagnetic Radiation and Energy

Electromagnetic Radiation and Energy. Electromagnetic Radiation: Energy traveling through space Three Characteristics of Waves: Wavelength : (symbolized l) Distance between two consecutive peaks or troughs in a wave Frequency : (symbolized n) How many waves pass a given point per second

koren
Download Presentation

Electromagnetic Radiation and Energy

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Electromagnetic Radiation and Energy • Electromagnetic Radiation: • Energy traveling through space Three Characteristics of Waves: • Wavelength: (symbolized l) • Distance between two consecutive peaks or troughs in a wave • Frequency: (symbolized n) • How many waves pass a given point per second • Speed: (symbolized c) • How fast a given peak moves through space

  2. Electromagnetic Radiation and Energy c = λ x ν c = speed of light = 2.9979 x 108 m/s ν = frequency (s-1 or Hz) λ = wavelength (m)

  3. Spectra • Sunlight yields continuous spectrum • Energized gaseous elements yield certain wavelengths • Line emission spectrum

  4. Rydberg • Why did gaseous atoms emit certain wavelengths? • Didn’t find out why, but came up with an equation • Rydberg equation • N=3, red line • N=4, green line • N=5, blue line • Lyman series • n > 1 to n = 1 • UV (invisible) • Balmer series • n > 2 to n = 2 • Visible wavelengths

  5. The Bohr Model of the Atom • Explained Rydberg • Electron energy quantized • Electron only occupies certain energy levels or orbitals • If it didn’t, electron would crash into protons in nucleus • As “n” increases energy becomes less negative • Increases • Only certain amts of E may be absorbed/emitted • If electron in lowest possible energy level • Ground state • If electron in excited energy level • Excited state • One can calculate energy needed to raise H electron per atom from ground state (n=1) to first excited state (n=2)

  6. Bohr’s Model • Explains emission spectrum of H • Movement of electrons from one quantized energy level to a lower one gave distinct emission wavelengths • Model only good for one electron system

  7. Atomic orbital The probability function that defines the distribution of electron density in space around the atomic nucleus.

  8. The s-orbital • The simplest orbital • The only orbital in the s-subshell • Occurs in every principal energy level • “s” stands for “sharp” • The first energy level only houses this orbital • Can house up to 2 electrons

  9. The p-orbitals • Start in second principle energy level (n=2) • There are three p-orbitals in the p-subshell (see below) • And one s-orbital • “p” stands for “principle” • Can house up to 6 electrons • Has one nodal surface • Nodal plane = a planar surface in which there’s zero probability of find an electron 2px 2py 2pz

  10. The d-orbitals • Start in third principle energy level (n=3) • There are five d-orbitals in the d-subshell • And one s-orbital • And three p-orbitals • Can house up to 10 electrons • “d” stands for “diffuse” • Has two nodal surfaces 3dyz 3dxz 3dxy 3dx2-y2 3dz2

  11. The f-orbitals • Start in fourth principle energy level (n=4) • There are seven f-orbitals in the f-subshell • And one s-orbital • And three p-orbitals • And five d-orbitals • Can house up to 14 electrons • “f” stands for “fundamental” • Has 3 nodal surfaces

  12. Electron configuration • Electron must be identified as to where it is located • Hydrogen: • One electron in first energy level and s-subshell • Thus, 1s1 (= Aufbau electron configuration) • 1 states energy level (n) • s designates subshell • Superscript 1 tells how many electrons are in the s-subshell • Can also use orbital box or line diagrams • Let’s take a look

  13. Pauli Exclusion Principle • An atomic orbital holds a maximum of two electrons • Both electrons must have opposite spins • ms = +1/2 & -1/2

  14. Hund’s Rule • Electron configuration most stable with electrons in half-filled orbitals before coupling

  15. Subshell filling order – not what one expected

  16. Using the Periodic Table to advantage

  17. Short-hand vs. long-hand Aufbau electron configuration • F • Al • Ca • Br

  18. Exercises • Give me the Aufbau electron configurations for: • Y • Te • Hf • Tl • 112

  19. Sundry matters pertaining to d-block metals • Stability is increased when: • d-subshell is half-filled (d5) • d-subshell is completely filled (d10) • Electrons will be taken from the s-subshell to fill the d-subshell • But there is a limit • No more than 2 electrons taken from s-subshell • Given the above, which subshell electrons will d-block metals lose first when they ionize? • So what are the correct electron configurations of Cr and Ag? • Caveat • Not all metals follow the above; i.e., take from s-subshell and give to d-subshell • Ni & Pt, for example

  20. Sundry matters pertaining to f-block metals • Stability is increased when: • f-subshell is half-filled (f7) • f-subshell is completely filled (f14) • Electron will be taken from the d-subshell to fill the f-subshell • Eu & Yb • Am & No

More Related