Ch. 5- Electrons in Atoms Unit 7 Targets: The Electronic Structure of Atoms (Chap 5) I CAN Utilize appropriate scientific vocabulary to explain scientific concepts. I CAN Perform calculations involving the energy, wavelength and frequency of electromagnetic waves. I CAN Perform calculations to determine the de Broglie wavelength of any object. I CAN Compare and contrast Bohr’s solar system model with Schrodinger’s wave mechanical model. I CAN Predict the movement of electrons from the ground state to an excited state, back to the ground state. I CAN Recognize elements based on their emission spectra. (Flame lab). I CAN Generate electron configurations for elements: full electron configuration, noble gas configuration, orbital diagram. I CAN Differentiate relationships between orbitals, sublevels and energy levels. I CAN Correlate the relationship between electrons in a single orbital or single sublevel.

 Electromagnetic radiation- a form of energy that exhibits wavelike behavior. Wavelength ( ) - length of one complete wave Frequency ( ) - # of waves that pass a point during a certain time period  SI UNIT for frequency is hertz (Hz) = 1 wave / s Amplitude (A) - distance from the origin to the trough or crest

A greater amplitude (intensity) greater frequency (color) crest origin trough A

To understand the electronic structure of atoms we must understand light and how it is emitted or absorbed by substances. We will examine visible light a type of Electromagnetic Radiation (EM) which carries (radiant) energy through space (speed of light) and exhibits wavelike behavior. Also need to think of light as particle, to help understand how EM radiation and atoms interact

LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

yMove through a vacuum at the ‘speed of light’ 3.00 x 10 8 m/s yBehaves like waves that move through water, which are the result of a transfer of energy to the water (from a stone), expressed as up and down movement of water yBoth electric and magnetic properties

LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY ROYG.BIV redorangeyellowgreenblueindigoviolet

yWave Speed = (distance between wave peaks) x (frequency) = (wavelength) x (frequency) EM radiation moves through a vacuum at the “speed of light” 3.00 x 10 8 m/s also called c. yA lower energy wave (infrared and red) has a longer wavelength( ) and lower frequency(f) yA higher energy wave (blue - violet) has a shorter wavelength( ) and higher frequency(f).

 Frequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

 EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN: = ? = 434 nm = 4.34  m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  m = 6.91  Hz

 Planck (1900) Observed - emission of light from hot objects Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta) Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

E:energy (J, joules) h:Planck’s constant (  J·s) :frequency (Hz) E = h zPlanck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency zThe energy of a photon is proportional to its frequency.

 EX: Find the energy of a red photon with a frequency of 4.57  Hz. GIVEN: E = ? = 4.57  Hz h =  J·s WORK : E = h E = (  J·s ) ( 4.57  Hz ) E = 3.03  J

 Planck (1900) Observed - emission of light from hot objects Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta) Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

 Planck (1900) vs. Classical Theory Quantum Theory

 Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, …. The allowed energies are quantized, that is their values are restricted to certain quantities.  The notion of quantized rather than continuous energies is strange. Consider a ramp and a staircase, on a ramp you can vary the length your steps and energy used on the walk up. When walking up steps you must exert exactly the specific amount of energy needed to reach the next step. Your steps on steps are quantized, you cannot step between them.

 Einstein (1905) Observed - photoelectric effect

 Einstein (1905)  Concluded - light has properties of both waves and particles (photons) “wave-particle duality”  Photon - particle of light that carries a quantum of energy  Used planck’s quantum theory to deduced that: E photon = hv

ground state excited state ENERGY IN PHOTON OUT yDEFINITION: A set of frequencies of EM waves emitted by atoms of the element. EX. neon light absorb electrical energy, e - get excited, become somewhat unstable and release energy in the form of light, a prism separates the light into an atomic emission spectrum.

 e - exist only in orbits with specific amounts of energy called energy levels  Therefore… e - can only gain or lose certain amounts of energy only certain photons are produced

 Energy of photon depends on the difference in energy levels  Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom

 Each element has a unique bright-line emission spectrum. “Atomic Fingerprint” Helium zBohr’s calculations only worked for hydrogen! 

 Examples: Iron  Now, we can calculate for all elements and their electrons – next section

 Louis de Broglie (1924) Applied wave-particle theory to e - e - exhibit wave properties EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT

 Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time

 Schrödinger Wave Equation (1926) finite # of solutions  quantized energy levels defines probability of finding an e - Wavelength = plank’s constant/mass X velocity

 Orbital (“electron cloud”) Region in space where there is 90% probability of finding an e - Radial Distribution Curve Orbital

 Four Quantum Numbers: Specify the “address” of each electron in an atom UPPER LEVEL

1. Principal Quantum Number ( n ) Main energy level occupied the e- Size of the orbital n 2 = # of orbitals in the energy level

2. Angular Momentum Quantum # ( l ) Energy sublevel Shape of the orbital s p d f

zn=# of sublevels per level zn 2 =# of orbitals per level zSublevel sets: 1 s, 3 p, 5 d, 7 f

3. Magnetic Quantum Number ( m l ) Orientation of orbital around the nucleus Specifies the exact orbital within each sublevel

pxpx pypy pzpz

 Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x

4. Spin Quantum Number ( m s ) Electron spin  +½ or -½ An orbital can hold 2 electrons that spin in opposite directions.

 Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers. Each e - has a unique “address”: 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orientation electron

Read Section 4-2!