1 Electromagnetic Radiation and Energy Electromagnetic Radiation: –Energy traveling through space Three Characteristics of Waves: 1.Wavelength: (symbolized  1.Distance between two consecutive peaks or troughs in a wave 2.Frequency: (symbolized  1.How many waves pass a given point per second 3.Speed: (symbolized c) 1.How fast a given peak moves through space

2

3 Electromagnetic Radiation and Energy c = λ x ν C = speed of light = 3 x 10 8 m/s ν = frequency (s -1 or Hz) λ = wavelength (m)

4

5 Standing Waves Tie down a string at both ends  pluck Has 2 or more nodes Distance between nodes is λ/2 Only certain wavelengths are allowed (n x λ/2)  as is found in atomic theory

6 Planck Scientists tried to explain relationship between intensity and wavelength for radiation given off by heated objects –Electromagnetic radiation color depends on temperature –Wrongly surmised that the shorter the wavelength the greater the radiation intensity –Planck solved the riddle He came up with term quantized –Only certain vibrations with specific frequencies allowed Planck’s equation –Vibrating system energy proportional to frequency of vibration –E = h ; E in joules, h = (J  s) = x J  s = Planck’s constant

7 As temperature increases… –Maximum energy released in visible spectrum goes towards UV “white hot”

8 Einstein’s Photon Photoelectric effect –Electron ejection after light strikes metal surface Must be the right frequency Automatic door openers Einstein –Light has particle-like properties Photons

9 Einstein + Planck What happens to energy as frequency increases? What happens to energy as wavelength increases?

10 Spectra Sunlight yields continuous spectrum Energized gaseous elements yield certain wavelengths –Line emission spectrum

11 Rydberg Why did gaseous atoms emit certain wavelengths? Rydberg equation –N=3, red line –N=4, green line –N=5, blue line Balmer series –N=6,7,8

12 The Bohr Model of the Atom 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)

13 Bohr’s Model Explains emission spectrum of H –Movement of electrons from one quantized energy level to a lower one Balmer series –n > 2 to n = 2 Visible wavelengths Lyman series –n > 1 to n = 1 UV (invisible) Model only good for one electron system

14 Others De Broglie –Electron’s properties (mass & velocity) related to wave property ( ) Schrödinger Wave Equation –“quantum mechanics” quantum numbers e - has wave-particle duality Heisenberg Uncertainty Principle –Probability of e - presence Orbital pathways

15 Quantum numbers Used to solve Schrödinger Wave Equation n = principle quantum number –Principle energy level Energy shell –n  1 (the letter L) = angular momentum quantum number –Subshells Orbitals –0,1,2,…n-1 (s,p,d,f) m = magnetic quantum number –Orientation of orbitals (p x, p y, p z ) (how many orbitals within subshell) m s = magnetic spin number =  1/2 –Spin direction of electron in orbital

16

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

18 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

19 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 2p x 2p y 2p z

20 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 3d yz 3d xz 3d xy 3d x 2 -y 2 3d z 2

21 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

22 Electron configuration Electron must be identified as to where it is located –Hydrogen: One electron in first energy level and s-subshell –Thus, 1s 1 (= 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

23 Pauli Exclusion Principle An atomic orbital holds a maximum of two electrons Both electrons must have opposite spins m s = +1/2 & -1/2

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

25 Subshell filling order – not what one expected

26 Using the Periodic Table to advantage

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

28 Exercises Give me the Aufbau electron configurations for: –Y –Te –Hf –Tl –112

29 Sundry matters pertaining to d-block metals Stability is increased when: –d-subshell is half-filled (d 5 ) –d-subshell is completely filled (d 10 ) 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

30 Sundry matters pertaining to f-block metals Stability is increased when: –f-subshell is half-filled (f 7 ) –f-subshell is completely filled (f 14 ) Electron will be taken from the d-subshell to fill the f-subshell –Eu & Yb –Am & No