Orbits and Line Spectra The Bohr Model: Orbits and Line Spectra

C12-2-04 ENERGY AND ATOMIC MODELS OUTCOME QUESTION(S): C12-2-04 ENERGY AND ATOMIC MODELS Explain average atomic mass using isotopes and their relative abundance. Include: radioisotopes Describe the electromagnetic spectrum in terms of frequency, wavelength, and energy. Include: Plank’s Equation, quantum, photon Understand the historical development of the Quantum Mechanical Model of the atom. Describe how unique line spectra are created for each element. Include: Bohr model Vocabulary & Concepts  Ground State Absorption Spectra

How do scientists explain the very predictable, replicable structure of line spectra with the disorganized Rutherford model?

Bohr Model (1913) – proposed that spectral lines are light from excited electrons. Restricted electrons to fixed orbits (n) of different quantized energy levels Created an equation for energy of an electron at each orbit ΔE = -2.18 x 10-18 J Z2 nf2 ni2 _ His created equations correctly predicted the structured spectral lines of Hydrogen…

Electron fall to ground state Electrons absorb energy and jump from ground state (its resting state) to a higher unstable energy level (excited state). Electron fall to ground state – releasing a photon of energy equivalent to the distance moved. “unstable” is the KEY - electrons are attracted to the nucleus and can’t stay away for long Free Atom e− EMR Ground State Excited State Ionization Absorption nucleus > Threshold Energy < Threshold Energy

ΔE = E higher-energy orbit - E lower-energy orbit = Ephoton emitted = hf The difference in energy requirements between orbits determines the “colour” of photon released by the electron when it drops back towards the nucleus

3. Levels are discrete (like quanta) – No in-between 4. Every jump/drop has a specific energy requirement - same transition, same photon.

_ ΔE = -2.18 x 10-18 J Z2 nf2 ni2 = Ephoton emitted = hf Predicts the type of EMR released for any transition ΔE = Enf - Eni = Ephoton emitted = hf Change in energy is released as a photon of EMR

What is the energy of the photons released when an electron drops from n=5 to n=2 in a hydrogen atom? ΔE = -2.18 x 10-18 J Z2 nf2 ni2 _ ΔE = -2.18 x 10-18 J 12 2f2 5i2 _ ΔE = -2.18 x 10-18 J 1 4 25 _ ΔE = - 4.58 x 10 -19J

Visible light – blue to be exact What type of energy is released when an electron drops from n=5 to n=2 in a hydrogen atom? ΔE = - 4.58 x 10 -19J c = λƒ λ 6.91 x 1014 Hz = 3.00 x 108 m/s E = hf ƒ 6.626 x 10-34 J·s = -4.58 x 10-19 J λ = 4.34 x 10-7 m λ = 434 nm ƒ = 6.91 x 1014 Hz Visible light – blue to be exact

The size of the nucleus will affect electron position around the atom – and the energy requirements Each element has a unique line spectrum as each element has a unique atomic configuration – making it jump/drop unique 17 p+

We only “see” the emitted energy of an transitioning electron if it is a photon of energy in our visible spectrum – some emissions lines are invisible to us

Notice energy absorbed is the same as energy released Absorption spectrum – portion of visible light absorbed by an element – heating up. Emission spectrum – portion of visible light emitted by that element – cooling down.

C12-2-04 ENERGY AND ATOMIC MODELS CAN YOU / HAVE YOU? C12-2-04 ENERGY AND ATOMIC MODELS Explain average atomic mass using isotopes and their relative abundance. Include: radioisotopes Describe the electromagnetic spectrum in terms of frequency, wavelength, and energy. Include: Plank’s Equation, quantum, photon Understand the historical development of the Quantum Mechanical Model of the atom. Describe how unique line spectra are created for each element. Include: Bohr model Vocabulary & Concepts  Ground State Absorption Spectra