Following the Geiger-Marsden experiment, Rutherford proposed a model of the atom.
Bohr modified this model to explain the spectrum of hydrogen observed in experiments.
The Bohr-Rutherford model of the atom consists of electrons in energy levels around a positive nucleus.
How do features of this model account for all the experimental evidence above? Support your answer with a sample calculation and a diagram, and refer to energy, forces and photons. (9 marks)
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The Geiger-Marsden experiment, which involved firing alpha particles at a thin sheet of gold foil produced results which can be explained by the Bohr-Rutherford model:
→ The majority of fired alpha particles passed through the gold foil undeflected. Rutherford concluded from this that the atom had a small, central nucleus.
→ Some alpha particles were deflected and some of these were deflected at very large angles. Rutherford concluded from this that the nucleus was dense and positively charged exerting a repulsive electromagnetic force on the fired alpha particles.
→ The model accounts for Rutherford’s conclusions, placing electrons in orbits around a small positive nucleus.
Rutherford’s model alone could not explain the emission spectra of elements such as hydrogen. Bohr’s contribution to the Bohr-Rutherford model amended this:
→ Bohr proposed that electrons orbited the atomic nucleus in quantised orbits at fixed energies. He proposed that electrons could move from a higher energy orbit (eg. n=1) to a lower energy orbit (n=3) by emitting a photon with energy `E=hf` equal to the energy difference between the two orbits.
→ Additionally, he proposed that electrons could move from a lower energy orbit to a higher energy orbit by absorbing a photon with energy `E=hf` equal to the energy difference between the two orbits.
→ This is able to account for the given emission spectra of hydrogen, where emission lines correspond to electron transitions from higher energy orbits to the second energy orbit which produce photons within the spectrum of visible light.
Using Rydberg’s equation it is possible to predict the emission lines of hydrogen, using an electron moving from the sixth to the second Bohr energy orbit as an example:
`(1)/(lambda)` | `=R((1)/(n_(f)^(2))-(1)/(n_(i)^(2)))` | |
`=(1.097 xx10^7)((1)/(2^(2))-(1)/(6^(2)))` | ||
`=(2 xx1.097 xx10^7)/(9)` | ||
`=2.438 xx10^6` | ||
`lambda` | `=410 text{nm}` |
→ This value corresponds to the leftmost line on the given spectrum, reflecting how Bohr’s model can account for the emission spectra of hydrogen.
The Geiger-Marsden experiment, which involved firing alpha particles at a thin sheet of gold foil produced results which can be explained by the Bohr-Rutherford model:
→ The majority of fired alpha particles passed through the gold foil undeflected. Rutherford concluded from this that the atom had a small, central nucleus.
→ Some alpha particles were deflected and some of these were deflected at very large angles. Rutherford concluded from this that the nucleus was dense and positively charged exerting a repulsive electromagnetic force on the fired alpha particles.
→ The model accounts for Rutherford’s conclusions, placing electrons in orbits around a small positive nucleus.
Rutherford’s model alone could not explain the emission spectra of elements such as hydrogen. Bohr’s contribution to the Bohr-Rutherford model amended this:
→ Bohr proposed that electrons orbited the atomic nucleus in quantised orbits at fixed energies. He proposed that electrons could move from a higher energy orbit (eg. n=1) to a lower energy orbit (n=3) by emitting a photon with energy `E=hf` equal to the energy difference between the two orbits.
→ Additionally, he proposed that electrons could move from a lower energy orbit to a higher energy orbit by absorbing a photon with energy `E=hf` equal to the energy difference between the two orbits.
→ This is able to account for the given emission spectra of hydrogen, where emission lines correspond to electron transitions from higher energy orbits to the second energy orbit which produce photons within the spectrum of visible light.
Using Rydberg’s equation it is possible to predict the emission lines of hydrogen, using an electron moving from the sixth to the second Bohr energy orbit as an example:
`(1)/(lambda)` | `=R((1)/(n_(f)^(2))-(1)/(n_(i)^(2)))` | |
`=(1.097 xx10^7)((1)/(2^(2))-(1)/(6^(2)))` | ||
`=(2 xx1.097 xx10^7)/(9)` | ||
`=2.438 xx10^6` | ||
`lambda` | `=410 text{nm}` |
→ This value corresponds to the leftmost line on the given spectrum, reflecting how Bohr’s model can account for the emission spectra of hydrogen.