keep it simple science
From QUANTA to QUARKSWhat is this topic about?To keep it as simple as possible, (K.I.S.S.) this topic involves the study of: 1. RUTHERFORD & BOHR MODELS OF THE ATOM 2. DE BROGLIE & MATTER WAVES 3. INTO THE NUCLEUS 4. APPLICATIONS OF NUCLEAR PHYSICS ...all in the context of the history, nature and practice of Physics.
HSC Physics Option Topic
1. RUTHERFORD & BOHR MODELS OF THE ATOMWhat Has Gone Before...The entire Science of Chemistry and much of Physics is built on the foundation of Atomic Theory... the concept that all matter is composed of atoms. Initially conceived as tiny, unbreakable particles of matter, by the beginning of the 20th century it became apparent that the atom was composed of smaller parts.+ + In his famous experiment with cathode rays, J.J.Thomson had discovered the (negatively charged) electrons in all atoms. This meant that there also had to be a positive part of each atom.
The Rutherford Model of the AtomIn 1911, Ernest Rutherford carried out an experiment which indicated that the positively charged part of an atom must be concentrated into a tiny nucleus, with the electrons orbiting around it.Rutherfords ATOM Electrons in orbit around central nucleus Atom mostly empty space Nucleus of positively charged matter, possibly made up of of particles
In 1900, Max Plank had proposed the Quantum Theory to explain the details of the Black Body Radiation Curves. In 1905, Einstein then explained the strange phenomenon of the Photoelectric Effect by using Planks quantum idea. He proposed that light is not just a wave, nor a stream of particles, but made up of wave packets.Light is NOT a stream of particles...
Rutherfords model proposed that: At the centre is a tiny, dense nucleus with a positive electrical charge. The negatively charged electrons orbit around the nucleus. The distance from nucleus to the electron orbits is very large compared to the size of the particles, so the atom is mostly empty space.
Light is NOT a wave...
Since negative charge was carried by particles (the electrons) Rutherford thought it likely that the nucleus was made of positive particles. These were soon called protons and their existence was confirmed a few years later. The electrons were too light to account for much of the mass of an atom, so he thought the protons must be relatively heavy. Even at this early stage there was speculation that there might be another massive particle in the nucleus as well, but its discovery had to wait 20 years. 1Usage & copying is permitted according to the Site Licence Conditions only
Light is a stream of wave packets... PHOTONS
Each photon is both a particle AND a wave!
Einstein also proposed his Theory of Relativity in 1905. Classical Physics was being turned upside-down by this sequence of new, fundamental discoveries.HSC Physics Option Topic From Quanta to Quarks Copyright 2006-2 2009 keep it simple science www.keepitsimplescience.com.au
keep it simple science
Problems with Rutherfords AtomEven as he proposed his atomic model, Rutherford knew there was a problem with it. The existing theory of Electromagnetic Radiation (EMR) contained the concept that if an electrically charged particle was accelerating, then it must emit EMR, in the form of light waves. Since Rutherfords electrons were imagined to be in circular orbits around the nucleus, and since circular motion involves constant (centripital) acceleration, then it follows that each electron should be constantly emitting light. Trouble is... they obviously dont! Existing accepted theory required that an orbiting electron should emit light energy continuously. Obviously they dont, or all matter would constantly glow with light. However, atoms DO emit light if stimulated with energy, such as in a high-v voltage discharge tube.
Emission Spectrum of HydrogenYou will have observed the emission spectrum for hydrogen by using a spectrometer to view the light from a discharge tube filled with lowpressure hydrogen gas.from induction coil Slit & lens
SpectroscopePrism Optical viewing system
Tube filled with Hydrogen gas
Tube glows with emitted light
Telescope can be rotated to view the different lines of the emission spectrum
light emission from electrons
You will have seen that the light from a hydrogen discharge tube is composed of 4 visible bright lines of light. Each line is one single wavelength of light.
You should be familiar with the idea of a spectrum of light. For example, if white light is passed through a prism, the different wavelengths are separated, and the familiar rainbow colours appear.white light is a mixture of wavelengths different wavelengths spread out to form a spectrum
The Balmer Series & Rhydberg EquationThe lines in the emission spectrum of hydrogen had been discovered some 20 years before Rutherfords work, and were known as the Balmer Series. Each line was given a name (H, H, H & H) and the precise wavelength of each had been measured. Other similar series of lines were known to exist in the invisible infra-red and ultra violet parts of the EMR spectrum. No-one could explain them, but mathematicians Balmer and (later) Rhydberg had worked out that the exact wavelengths of the hydrogen spectrum lines could be calculated from an empirical equation:
Red Orange Yellow Green Blue Violet
(use your imagination... we cant print colours)
If the light emitted by atoms of a particular element is put through a prism, the spectrum shows very narrow bright lines on a dark background because only certain wavelengths are given out. The pattern of lines is characteristic for each element.Element A Element B Element C
The Rhydberg Equation 1 = RH( 1/nf - 1/ni ) = wavelength of the spectral line (in metres) RH = the Rhydberg constant = 1.097 x 107 nf = an integer number. For the Balmer series nf = 2 ni = an integer number. To calculate the wavelengths of the 4 lines of the Balmer series, ni takes the values 3, 4, 5 or 6.The fact that the Rhydberg equation worked was strong evidence that there was an underlying law controlling the hydrogen spectral lines. The fact that a series of integer numbers were involved was a clue that connected the whole thing to Planks Quantum Theory... 2 2
Each line is light of one exact wavelength. Light is only emitted at certain precise wavelengths
Each element has its own unique set of spectral lines
HSC Physics Option Topic From Quanta to Quarks Copyright 2006-2 2009 keep it simple science www.keepitsimplescience.com.au
Usage & copying is permitted according to the Site Licence Conditions only
keep it simple science
Planks Quantum TheoryA quick revision of what you learned previously...
Neils Bohr Puts It All TogetherBohr used Planks Quantum Theory to modify the Rutherford model of the atom in such a way that: the problem of radiation that should be emitted constantly from accelerating electrons was overcome. the underlying reasons for emission spectra were explained. the empirical nature of the Rhydberg Equation was given theoretical backing and mathematical validity. the reasons for the valency of different atoms, and how and why they combine in fixed ratios became clearer. Not bad for an afternoons work! (The last point above is fundamental to Chemistry and understanding chemical bonding and formulas. It will not be pursued any further in this topic)
In 1900, Max Plank proposed a radical new theory to explain the black body radiation. He found that the only way to explain the exact details coming from the experiments, was that the energy was quantised: emitted or absorbed in little packets called quanta (singular quantum). The existing theories of classical Physics assumed that the amount of energy carried (say) by a light wave could have any value, on a continuous scale. Planks theory was that the energy could only take certain values, based on units or quanta of energy. Plank proposed that the amount of energy carried by a quantum of light is related to the frequency of the light, and can be calculated as follows:
Bohrs Postulates Electrons revolve only in certain allowed orbits. Bohr theorised that there are a series of orbits, at fixed distances from the nucleus, in which an electron will not constantly emit radiation as demanded by classical theory. (Why was explained later by de Broglie) Allowed orbit positions. Electrons cannot orbit anywhere else. Electrons can jump from one orbit to another, but must absorb energy to jump higher, or emit energy to drop lower. 3 Quantum numbers of the orbits.
E = h.fE = energy of a quantum, in joules ( J)h = Planks constant, value 6.63x10-34
f = frequency of the wave, in hertz (Hz) You are reminded also, of the wave equation:
V = .f (or, for light) c = .fc = velocity of light (in vacuum) = 3.00x10 ms .8 -1
= wavelength, in metres (m). f = frequency, in hertz (Hz)
Example Calculation a) Use the Rhydberg Equation to find the wavelength of the H line of the hydrogen spectrum, given that nf= 2 and ni = 6.
1 = RH( 1/nf - 1/ni ) = 1.097x107( 1/22 - 1/62 )(410 nm nanometres)
Electrons gain or lose energy to jump between orbits. To jump up to a higher orbit, an electron must gain a certain quantity of energy. If it drops back to lower orbit, it must emit that exact same amount of energy. These quantities of energy are quantised, so each orbit is really a quantum energy level within the atom. The amount of energy absorbed or emitted during a jump is defined by Planks Equation E = hf, and the corresponding wavelengths of light