Discuss the structure of the Rutherford model of the atom, the existence of the nucleus and electron orbits
The Rutherford atom consisted of a small positive nucleus with negatively charged electrons orbiting the nucleus. The Rutherford atom was devised following Rutherford’s experiments in which he fired positively charged alpha particles at thin gold foil. Rutherford found that while most alpha particles passed straight through the foil, a small proportion of them were reflected back. He hypothesised that they had encountered very dense areas of positive charge. The fact that most alpha particles passed through the gold foil led Rutherford to model the atom with a great deal of empty space. Rutherford modelled the atom with a dense, positively charged nucleus, negatively charged electrons that orbited the nucleus, and free space between the nucleus and the electrons. The model was essentially a simplified version of what we use today- it was groundbreaking at the time as it was a step in the right direction for other scientists to build on, but it lacked a description of where the electrons were and failed to address how atoms had stability without energy emission from accelerating electrons.
Remember- The Rutherford atom had a positive nucleus, negative electrons orbiting the nucleus, and empty space between the electrons and the nucleus.
Analyse the significance of the hydrogen spectrum in the development of Bohr’s model of the atom.
Bohr’s model of the atom was quite similar to Rutherford’s, but with two important differences- firstly, it assigned positions to the electrons, but secondly the electron energy levels were quantised. This was radically new, the idea that electrons had energy states and could absorb and emit energy to change states, and had no evidence. Bohr realised that if his model was correct, each atom would have a spectral fingerprint related to the differences between electron energy levels in that atom. The Rydberg equation, otherwise known as the Balmer equation, gave him evidence for the quantised emission of energy from the hydrogen atom, leading to him going on to further his model and define his postulates. So the hydrogen spectrum was very significant to the development of Bohr’s model of the atom, because without an understanding of it Bohr may not have continued to work on his model.
Remember- The hydrogen spectrum was extremely significant because it provided the only evidence at the time for an otherwise purely theoretical model.
Perform a first-hand investigation to observe the visible components of the hydrogen spectrum
In our experiment, we had a discharge tube (vacuum tube with a cathode and anode, powered by a high-voltage induction coil) with low-pressure hydrogen inside it. When high-voltage current was passed through the tube, the hydrogen fluoresced, emitting light that was visible in our darkened room. We observed the visible components of the spectrum with handheld spectrometers that used a diffraction grating to split the light. Using the spectrometer, we could clearly observe the red and blue/violet hydrogen emission lines, although the violet lines were very hard to observe. The red line was very clear and intense compared to the other observed lines.
Remember- Hydrogen discharge tube observed with a spectrometer.
Discuss Planck’s contribution to the concept of quantised energy
The concept of quantised energy is that energy can only occur in small packets of fixed amounts, and distinguished between energy increases due to increased intensity (bigger packets) and energy increases due to greater intensity (more packets). This was developed entirely by Planck in his work on black body radiation, and although Einstein significantly improved upon Planck’s ideas, the underlying idea was Planck’s alone, and so Planck made a huge contribution to the underlying concept of quantised energy. However, his involvement was limited to developing the mere concept- others developed it into a functional model.
Remember- Planck developed the concept of quantised energy, but not a functioning model.
Define Bohr’s postulates
Bohr had 3 postulates. The first was that electrons in an atom exist in stationary states of stability and emit no energy when in these states. The second was that energy is only lost or gained by an electron when it moves from state to state, and when it moves from a high energy state to a low energy state it releases a photon with energy equal to the difference between the states (and therefore a characteristic frequency). His third postulate was that electron angular momentum in a
Describe how Bohr’s postulates led to the development of a mathematical model to account for the existence of the hydrogen spectrum (the Rydberg equation)
Balmer originally devised the equation empirically by examining the lines in the hydrogen spectrum and creating a formula to fit them. Rydberg used Bohr’s postulates and manipulated them (especially the third) to create the same formula (derived from calculating differences in energy states). Essentially, there were two paths to the Rydberg equation and one of them used Bohr’s postulates to arrive at the equation, while the other didn’t.
Process and present diagrammatic information to illustrate Bohr’s explanation of the Balmer series
In this diagram, the energy levels described by Bohr are clearly marked. According to Bohr, the Balmer series (shown on the top of the diagram as the hydrogen spectrum) was caused by electrons changing energy levels. The electron makes a transition from a higher energy level to a lower energy level, in the process releasing light. As shown, larger energy changes produce more energetic photons, as seen in the Balmer series, and further, this diagram shows how the Balmer series is formed by successive electron transitions to the 2nd shell (transitions to other shells produce additional lines named after their discoverers).
Remember- Bohr explained the Balmer series as being the result of successive electron transitions down to the 2nd shell.
Discuss the limitations of the Bohr model of the hydrogen atom (including “Analyse secondary information to identify the difficulties with the Rutherford- Bohr model, including its inability to completely explain the spectra of larger atoms, the relative intensity of spectral lines, the existence of hyperfine spectral lines, and the Zeeman Effect”)
For all the questions the Bohr model answered, it posed still more. There was still no explanation for there being no energy emission from accelerating electrons as Maxwell predicted- instead it was simply an assumption. Further, there was no evidence for the Bohr model to give it scientific credibility. Finally, in terms of explaining spectral lines there were observed effects that simply could not be explained. These were
- Relative intensity of spectral lines- When observing spectra, some lines were much brighter than The Bohr model could not explain why some lines were more intense than others (i.e. why some electron transitions were preferred to others)
- Hyperfine splitting- When the spectral lines were examined closely, it was observed that each line actually consisted of many small lines, the existence of which the Bohr model could not explain as it only predicted one clear line for each transition
- Larger atoms- The Bohr model could not explain the spectra of larger atoms with more than one electron, a problem that Bohr tried unsuccessfully to
- Zeeman effect- The Zeeman effect occurs when a magnetic field is passed through the discharge The magnetic field increases the hyperfine splitting of spectral lines, further breaking them up. Again, the Bohr model was unable to explain the experimental evidence
Although the Bohr model lay down the framework for the quantum model of the atom, which ended in a scientific revolution out of which quantum mechanics (a vital part of modern physics) emerged, it was left to future scientists such as Pauli and Heisenberg to fully explain these phenomena.
Remember- Relative intensity, hyperfine splitting, larger atoms, Zeeman effect.