# Atomic Structure - Academic Computer Center Atomic Structure I Its not about Dalton anymore http://plus.maths.org/latestnews/may-aug07/strings/atoms.jpg First To understand the electronic structure of the atom we need to review the properties of electromagnetic radiation.

The Wave Nature of Light Figure 7.1 Frequency and Wavelength c = wavelength

frequency C = speed of light Amplitude (intensity) of a wave. he waveheight or amplitude determines radiation intensity. The wavelength is related to the energy of the radiation. , , and Energy As decreases and increases, what happened to the energy of the

radiation? hc E =h = where h = Plancks constant (6.626 10-34 m2 kg/s) The infinite number of wavelengths of electromagnetic radiation have been classified into groups as shown below.

Regions of the electromagnetic spectrum. Interconverting Wavelength and Frequency o PROBLEM: A dental hygienist uses x-rays (= 1.00A) to take a series of dental radiographs while the patient listens to a radio station ( = 325 cm) and looks out the window at the blue sky (= 473 nm). What is the frequency (in s-1) of the electromagnetic radiation from each source? (Assume that

the radiation travels at the speed of light, 3.00x108 m/s.) Use c = SOLUTION: o -10 -10 1.00A 10 o m = 1.00x10 m 1A 3x108 m/

= 3x1018 s= s 1.00x10-10 1 10-2 m 325 -2 = 325x10 m m 1 cm

cm 3x108 m/ = = 9.23x107 ss 325x10-2 m1 10-9 473nm = 473x10-9 m 1 nm m 3x108 m/ = 6.34x1014

= s 473x10-9 ms-1 Calculating the Energy of Radiation from Its Wavelength PROBLEM: A cook uses a microwave oven to heat a meal. The wavelength of the radiation is 1.20cm. What is the energy of one photon of this microwave radiation? After converting cm to m, we can use the energy equation, E = h combined with = c/ to find the energy.

SOLUTION: E= E = hc/ 6.626X10-34J*s x 3x108m/ = 1.66x10-23J 10- s 1.20c 2 m m

cm Particle or Wave? Different behaviors of waves and particles. The diffraction pattern caused by light

passing through two adjacent slits. Light is a waveright? Light falling on alkali metals causes electrons to be released from the metal. The # of electrons depends on the intensity of light. There are specific

wavelengths of light that cause the release of e-. This is called the photoelectric effect. Light is a waveright? Einsteins interpretation of the photoelectric effect (1905) was that light is quantized in packets of set energy called photons. (He won the Nobel Prize for this.)

This meant that light had characteristics of particles! Electrons are particles right? In 1925, de Broglie stated that all particles have a wavelength described by the equation: = h/p where p= momentum Electrons show diffraction pattern when passing through a slit So light and particles have a dual

nature. Back to atomic structure We already know an atom contains a nucleus with p+ and no. Electrons orbit the nucleus. It was known that atoms emit a unique spectrum of lines when excited. Rydberg derived an equation that related the lines. Rydberg equation

1 = R 1 n1 2 R is the Rydberg constant = 1.096776x107 m-1

1 n2 2 Flame test colors derive from electrons changing energy levels. Atomic emission spectra Clockwise from lower left: neon, helium, hydrogen,

mercury, nitrogen http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/atspect2.html Spectra Site http:// jersey.uoregon.edu/vlab/elements/El ements.html Absorption and emission spectra for element arranged on the periodic table

Back to atomic structure Bohr theorized that the emission spectra of atoms described by Rydbergs equation were caused by the transition of electrons between specific energy levels (orbits). http://www.upscale.utoronto.ca/Gene ralInterest/Harrison/BohrModel/Flash/ BohrModel.html

Electron locations When an electron occupies its usual energy level it is in the ground state. When an electron absorbs a photon and moves to a higher energy level it is in an excited state. The energy levels are quantized. Atoms can only transition between set levels. Why are the levels set where they are?

More on electrons as waves Since electrons have wave motion Schrdinger applied the classic wave equations to the motion of a hydrogen electron. Certain wavelengths reinforced each other and were allowed. This generated regions occupied by an electron of set energy termed orbitals.

Wave motion in restricted systems. More on electrons as waves Heisenberg stated that in measuring the electron there is uncertainty so we can only calculate a probable location for the electron. This is called the Heisenberg Uncertainty Principle.

Electron probability in the ground-state H atom. The 2p orbitals. The 3d orbitals. F orbitals

CLASSICAL THEORY Matter Energy particulat continuou e, s, massive wavelike

Summary of the major observations and theories leading from classical theory to quantum theory. Since matter is discontinuous and particulate perhaps energy is discontinuous and particulate. Observatio Theory n

blackbody Planck: Energy is quantized; only certain radiation values allowed photoelectric Einstein: Light has particulate behavior effect (photons) atomic line

Bohr: Energy of atoms is quantized; spectra photon emitted when electron changes orbit. Since energy is wavelike perhaps matter is wavelike Observatio Theory n

Davisson/Germer: deBroglie: All matter travels in waves; energy of electron atom is quantized due to wave diffraction motion of electrons by metal crystal Since

matter has mass perhaps energy has mass Observatio Theory n Compton: photon Einstein/deBroglie: Mass and energy are wavelength equivalent; particles have increases wavelength and

(momentum photons have momentum. decreases) after QUANTUM THEORY colliding with electron Energy same as Matter particulate, massive, wavelike