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The structure of the atom.07 The structure of the atom.07 Presentation Transcript

  • THE STRUCTURE OF THE ATOM 12 SACE PHYSICS-STAGE 2 SECTION 4 TOPIC 1 PRINCE ALFRED COLLEGE
  • THE STRUCTURE OF THE ATOM
    • As we find in physics, there are many objects and phenomena that are not easily seen as everyday experiences.
    • For this reason, models are used.
    • Light for example, is difficult to comprehend when sometimes it behaves as a particle and sometimes a wave.
    • Light is not ‘like’ either of these two but the model helps us to understand its behaviour.
  • THE STRUCTURE OF THE ATOM
    • Models are used to describe the atom as they are smaller than the wavelength of light and cannot be seen with the eye.
    • We will be exploring the patterns found in experiments, which are the basis, on which today’s model of the atom is built.
  • IN THE BEGINNING…
    • Around 430 BC, the first idea of the atom was developed by a Greek philosopher, Leucippus of Miletus.
  • IN THE BEGINNING…
    • He thought that if you keep dividing material again and again, there would come a time that you could not divide it any further.
    • This indivisible particle is called an atom ( atomos in Greek means ‘indivisible’).
  • IN THE BEGINNING…
    • His student Democritus of Abdera (c.460-371 BC) even went on to say that different substances were made up of different atoms or different combinations of atoms.
  • IN THE BEGINNING…
    • However, the two prominent philosophers of the time, Aristotle
  • IN THE BEGINNING…
    • and Plato
  • IN THE BEGINNING…
    • Believed in the idea that all matter was made up of air, fire, water and earth.
    • Because they were held in such high esteem, the atom concept was buried for over 2000 years.
  • IN THE BEGINNING…
    • John Dalton showed that reactions between elements followed predictable patterns during the 1700’s.
  • IN THE BEGINNING…
    • He also revived the concept of the atom, which was confirmed as more and more elements were discovered.
  • IN THE BEGINNING…
    • A Russian by the name of Dimitri Mendeleev arranged elements according to weight and chemical properties.
  • IN THE BEGINNING…
    • This table (known as the Periodic Table), gave an insight into the structure of the atom.
    • He even managed to predict the nature of some elements that were unknown at the time.
  • IN THE BEGINNING…
    • In 1911, Neils Bohr described the Structure of the Atom (nucleus and electron shells).
  • IN THE BEGINNING…
    • The atom is only a recent discovery-most physicists that have studied the atom are still alive today.
  • SPECTRA
    • Spectra provide a “picture or model” of what the atom looks like.
    • About midway through the 19th century, it was discovered that if a gas could be stimulated to produce light, the colour could indicate the nature of the gas.
    • The light was passed through a prism or diffraction grating to produce a spectrum.
  • SPECTRA
    • There are 3 types of spectrum:
    • CONTINUOUS SPECTRUM
    • Heating an object such as a filament in an incandescent globe produces this. When viewed through a spectroscope, all spectral colours are seen.
    • An object’s spectrum can extend beyond the visible part of the spectrum.
  • CONTINUOUS SPECTRA
    • A hotter object will produce more energy at all wavelengths than does a cooler object.
    • The hotter the object, the more energy is emitted at shorter wavelengths. The distribution of wavelengths depends on the temperature.
    • This is why the colour of objects change as the object is heated.
  • CONTINUOUS SPECTRA
    • A continuous spectrum is produced by the heating of solids and dense liquids.
    • The atoms in the solids and dense liquids vibrate and give off various frequencies of E-M energy.
    • They do not vibrate at the same rate and this leads to a wide range of emitted frequencies (continuous spectrum).
  • CONTINUOUS SPECTRA
  • CONTINUOUS SPECTRA
    • The inner core of a star could be considered a dense liquid and therefore the star will give off a continuous spectrum.
    • The diagram below shows how distributions of wavelengths change for stars. Notice how the distribution still gives a continuous spectrum for all temperatures.
  • CONTINUOUS SPECTRA
  • CONTINUOUS SPECTRA
    • Spectral Shift due to increasing temperature…
    • 20 ° C (Book on a table)-energy being emitted but at a low frequency. Energy is not visible to us.
    • 37 ° C (Humans)- emit infrared energy. Not visible to the naked eye. Spy satellites, night vision goggles.
  • CONTINUOUS SPECTRA
    • 3,000 ° C (Red Giant Star)- emits very little light. Mostly visible red (7,000 angstroms).
    • 6,000° C (Our sun)- ranges from Red to Violet Light. Most in yellow region. Makes white light.
  • CONTINUOUS SPECTRA
    • 24,000 ° C (Blue-White Dwarfs)- these are the hottest stars. They are faint because most of their energy is in the UV, Gamma, X-ray range of frequencies. Emit a little blue/violet light.
  • LINE EMISSION SPECTRA
    • Produced from light emitted by excited, or energised gases .
    • When gaseous atoms become excited, they emit characteristic colours (sodium - yellow and mercury - mauve).
  • LINE EMISSION SPECTRA
    • Looking through a spectrometer, the light is made up of specific frequencies and is referred to as spectral lines and the full spectrum seen is known as a line emission spectrum.
    • Unlike solids and dense liquids, gases do not produce a continuous spectrum.
  • LINE EMISSION SPECTRA
    • They are not vibrating back and forth at a variety of frequencies because they are not bonded together.
    • Instead, they emit SPECIFIC frequencies of light as seen in the spectroscope.
  • LINE EMISSION SPECTRA
    • These specific frequencies are produced by electrons falling from higher energy levels to lower energy levels.
    • A photon of a specific frequency is emitted in the process.
  • LINE EMISSION SPECTRA
    • A good example of line emission spectra is the burning of sodium.
    • The gaseous sodium’s electrons produce two distinct spectral lines in the yellow region of the E-M spectrum.
  • LINE EMISSION SPECTRA
  • LINE EMISSION SPECTRA
    • Only sodium has this spectral signature.
    • This allows you to identify elements through their spectral lines.
    • Atomic Emission Spectroscopy- we vaporise the mineral Quartz (SiO 2 ). The line emissions of Silicon and Oxygen are produced and detected by the spectrometer. This proves that the mineral is indeed quartz.
  • LINE EMISSION SPECTRA
    • Hydrogen gives off spectra that are in the very low frequency radio wave range.
    • Radio telescopes are used to study where hydrogen is present in the universe.
  • ABSORPTION SPECTRA
    • This is a continuous spectrum with dark lines. This means colours are removed due to their absorption by matter.
    • In 1814, a German scientist, Fraunhofer, focused the spectrum of light from the sun and found over 700 dark lines or gaps in the spectrum.
  • ABSORPTION SPECTRA
  • ABSORPTION SPECTRA
    • Today, astronomers have found over 30 000 ‘Fraunhofer lines.'
    • These gaps occurred because the cooler outer layers of the sun’s atmosphere remove some of the frequencies.
    • The frequencies absorbed are specific and depend not only on the atoms present, but also whether the atom has had electrons removed from it (ionised) or if it is neutral.
  • ABSORPTION SPECTRA
    • If we can match the absorption spectrum of the sun with the absorption spectra of an element like calcium, then we can say that calcium is in the sun’s outer atmosphere.
  • ABSORPTION SPECTRA
    • The absorbing material used in a lab is generally a gas or liquid but any state could be used. The dark lines depend on the nature of the absorbing material.
    • The absorption spectrum is as much a fingerprint of the element giving off the E-M energy as is the emission spectrum.
  • ABSORPTION SPECTRA
    • In all cases the absorption and the emission spectra will match perfectly.
  • ABSORPTION SPECTRA
  • ABSORPTION SPECTRA
    • Robert Bunsen
  • ABSORPTION SPECTRA
    • and Robert Kirchhoff,
  • ABSORPTION SPECTRA
    • Two German chemists, worked out a way in the late 1850’s, to identify elements using their line emission and absorption spectra.
    • They compared spectra they had from known elements with those from an unknown source and determined their chemical composition. This method was used to discover caesium and rubidium.
  • ABSORPTION SPECTRA
    • A Swedish astronomer in 1862
    • Anders Jonas Angstrom,
  • ABSORPTION SPECTRA
    • Used this technique to identify hydrogen in the sun.
    • Helium, although not found on earth for another 30 years, was found in the sun.
  • ABSORPTION SPECTRA
    • The spectra could be used as a ‘chemical fingerprint’ to identify elements.
    • However, questions such as how were the spectra produced and why were they different for each element could not be answered.
    • It did however suggest that there was some internal structure to the atom and not due to random vibrations.
  • THE BOHR MODEL OF THE ATOM
    • Bohr had just completed his doctorate in 1911 and had come to work in Cambridge under Rutherford in the Cavendish Laboratory (This is a must see on any trip to England).
    • By age 22, he had completed a mathematical treatment of the electronic structure of the hydrogen atom.
    • He predicted all the spectral lines of the atom.
  • THE BOHR MODEL OF THE ATOM
    • He ignored all the previous descriptions of the electronic structure as they were based on classical physics.
    • This allowed the electron to have any amount of energy.
    • Planck and Einstein used the idea of quanta for the energy carried by light.
  • THE BOHR MODEL OF THE ATOM
    • Bohr assumed that the energy carried by an electron was also quantized.
    • From this assumption, he formed three postulates (good intelligent guesses) from which he developed a mathematical description.
  • THE BOHR MODEL OF THE ATOM
    • 1. Electrons in the atom can only have fixed amounts of energy.
    • The electrons revolve around the nucleus only in certain allowed orbits called stationary states.
    • When in a stationary state, the electron cannot radiate any of its energy.
  • THE BOHR MODEL OF THE ATOM
    • The electron is only found in the stationary state so the energy is quantized inside the atom.
    • This is similar to rungs on a ladder. You can stand on the rungs but not in-between.
    • Stationary states are also referred to as energy levels or electron shells.
  • THE BOHR MODEL OF THE ATOM
  • THE BOHR MODEL OF THE ATOM
  • THE BOHR MODEL OF THE ATOM
    • 2. Electrons can jump from one stationary state to another by the absorption or emission of a photon
    • The energy of this photon will exactly equal the difference in energy of the two stationary states.
  • THE BOHR MODEL OF THE ATOM
    • Light energy is emitted or absorbed by atoms in fixed amounts called quanta. The quantum is equal to the difference in energy levels of the electrons.
    • This accounts for the existence of line emission and line absorption spectra.
  • THE BOHR MODEL OF THE ATOM
    • The frequencies of the bright lines of the line emission and dark lines of the absorption spectrum of an element will be identical.
    • Note how energy conservation is embodied in this postulate.
    • Bohr’s Third Postulate is not part of the examination.
  • THE BOHR MODEL OF THE ATOM
    • In summary, if the atom had electrons that varied in their energy levels, you would expect to get random frequencies emitted.
    • This is not the case.
    • Electrons give off photons of SPECIFIC frequencies.
    • More evidence for the Quantum Theory!
  • THE BOHR MODEL OF THE ATOM
    • This postulate laid the foundation for the mathematical treatment of the hydrogen atom.
  • THE BOHR MODEL OF THE ATOM
    • The number n is known as the principal quantum number and refers to the energy level under consideration.
    • n = 1 is the lowest energy level (called the K shell by chemists) or the ground state from the German grund meaning ‘fundamental’.
  • THE BOHR MODEL OF THE ATOM
    • n = 2 is the principal quantum number for the first energy level above the ground state or the first excited state (L shell) and so on.
    • Although Bohr allows the value of n to reach infinity, in reality only the first few are used.
  •  
  • THE BOHR MODEL OF THE ATOM
    • Transition - when an electron moves between energy levels.
    • Up would absorb energy (absorption).
    • An electron moving down would give off a photon of energy (emission).
  • THE BOHR MODEL OF THE ATOM
    • The electron n = infinity is not bound to the nucleus.
    • A larger nucleus will increase the eV (energy) values between the shells.
    • The Paschen series emits the least energy.
  • THE BOHR MODEL OF THE ATOM
    • Therefore, energy levels are not equally spaced.
    • The electron in the K shell (n=1) is closest to the nucleus. It has the LEAST amount of energy.
  • THE BOHR MODEL OF THE ATOM
    • An electron can be moved to a higher energy level by…
    • 1. INCOMING PHOTON - Must be of exactly the same energy as E 2 – E 1
    • 2. INCOMING ELECTRON - remaining energy stays with the incoming electron.
    • 3. HEAT- gives the electron vibrational energy.
  • THE BOHR MODEL OF THE ATOM
    • IONIZATION - energy required to remove an electron from the atom.
    • Example: the ionization energy required to remove an electron from its ground state (K=1) for Hydrogen is 13.6 eV.
  • THE BOHR MODEL OF THE ATOM
    • Notice that the Lyman Series would give off photons of the highest frequencies, followed by the Balmer Series then the Paschen Series.
  • THE BOHR MODEL OF THE ATOM
    • Let’s prove this using the following formula…
    • LYMAN SERIES
    • E n -E m = hf
    • (13.6 –0.85)(1.6X10 -19 )=hc/ λ
    • λ=9.7X10 -8 m(97nm)
    • This is UV Light
  • THE BOHR MODEL OF THE ATOM
    • BALMER SERIES
    • E n -E m = hf
    • (3.4 –0.85)(1.6X10 -19 )=hc/ λ
    • λ=4.87X10 -7 m
    • = 487 nm
    • This is Blue Light.
  • THE BOHR MODEL OF THE ATOM
    • SERIES LIMIT -The highest energy level transition for an electron in a particular series.
    • Example: The Lyman series limit is 13.6 eV.
  • THE BOHR MODEL OF THE ATOM
    • Below are the spectral lines that are seen for the emission spectra for Hydrogen.
    • Notice that the Lyman Series has the highest energy. The series limit is at the extreme right for each.
  • EXAMPLE 1
    • Use the first three energy levels for the electron in hydrogen to determine the energy and hence wavelength of the lines in its line emission spectrum.
  • EXAMPLE 1 SOLUTION
    • From the diagram, the atom can be excited to the first (n = 2) and second
    • (n = 3) excited states. From these, it will return to the ground state emitting a photon. The electron can make the following transitions:
  • EXAMPLE 1 SOLUTION
    • n = 3 n = 1, E photon = E 3 - E l
    • = -1.51 - (-13.6)
    • = 12.09eV
    •  
    • n = 2 n = l, E photon = E 2 - E 1
    • = -3.4 - (-13.6)
    • = 10.2eV
  • EXAMPLE 1 SOLUTION
    • n = 3 n = 2, E photon = E 3 - E 2
    • = -1.51 - (-3.4)
    • = 1.89 eV
  • EXAMPLE 1 SOLUTION
    • To find the wavelengths of the three photons, use
    • Note: convert eV to J
  • EXAMPLE 1 SOLUTION
    • n=3 n=l,
    • = 1.024 x 10 -7 m
    • = 102 nm (ultraviolet)
  • EXAMPLE 1 SOLUTION
    • n=2 n=l, λ=
    • = 1.21 x 10 -7 m
    • = 121nm (ultraviolet)
    • n= 3 n= 2 , λ=
    • = 6.55 X 10 -7 m
    • = 655 nm (visible-red)
  • EXAMPLE 1 SOLUTION
    • Note that we have two Lyman series lines (those ending in the ground state) and one Balmer line (ending in the first excited state).
  • EXAMPLE 1 SOLUTION
    • Those from the Lyman series produce lines in the ultra-violet part of the spectrum while the line in the Balmer series produces a line in the visible part of the spectrum.
  • EXAMPLE 2
    • The valence (outer) electron of sodium is involved in the emission of two photons that have almost the same energy. The wavelengths are 589.0 nm and 589.6 nm. It is known that the valence electron falling to the ground state from one or two higher levels will produce these photons. Determine the excitation energies of these two energy levels, and find the difference in their energy. (Work to 4 sig. Figs.)
  • EXAMPLE 2 SOLUTION
    • The diagram represents the two energy levels whose energies are E l and E 2 above the ground state.
  • EXAMPLE 2 SOLUTION
    • = 3.371 X 10 -19 J
    • = 3.374 X 10 -19 J
  • EXAMPLE 2 SOLUTION
    • Now it is more convenient to work in electron volts. E 1 = 2.107eV, and E 2 = 2.109eV, so the energy difference is 0.002 eV (3.2 x 10 -22 J). The energy level diagram can now be drawn.
  • FLOURESCENCE
    • If an electron is excited from one energy state, it may be able to make a jump of two or more energy levels.
    • When it returns to a lower level, it may do so in more than one jump.
    • The photons of light emitted will both have a lower energy, and hence frequency, than the photon that was absorbed.
  • FLOURESCENCE
    • If the absorbed photon comes from any high-energy part of the spectrum, and there are a larger number of emitted photons with less energy, we call this phenomena fluorescence .
  • FLOURESCENCE
    • We cannot see light that is in the UV part of the spectrum.
    • If the absorbed photon comes from this part of the spectrum, visible light has less energy than UV and so we see the object fluoresce.
    • An example of flourescence is shining UV light on the mineral FLOURITE.
  • FOURESCENCE
    • An example of this can be seen using the hydrogen energy levels.
  • FLOURESCENCE
    • Recall from Example 1 that the wavelength of the absorbed photon is 102 nm (ultraviolet).
    • The two emitted photons are at 121 nm (ultraviolet) and 655 nm (red light).
  • FLOURESCENCE
    • In this case, if we shone UV light of wavelength
    • 102 n m onto hydrogen, it would absorb the photon and we would see the hydrogen glow red with a wavelength of 655 nm.
    • Nothing would happen if we were to shine light of 105 nm as it does not match the energy level difference for hydrogen.
  • FLOURESCENCE
    • The wavelength at which fluorescence will occur will depend on the energy level of the atoms themselves.
    • This means that a substance will fluoresce with different colours.
    • Fluorescence is one way in which different compounds can be identified.
  • FLOURESCENCE
  • FLUORESCENCE
    • Fluorescent Diamonds
  • STIMULATED EMISSION
    • Stimulated emission is the principle behind LASER technology.
    • Normally, atoms that absorb energy and move to an excited state become unstable. The electron immediately (less than 10 -8 s) drops back to the ground state.
    • This may occur at any time and so two electrons that emit a photon of light do so at different times and so are not coherent.
  • STIMULATED EMISSION
    • If a photon with exactly the same energy required for the electron to jump down to the ground state interacts with an atom that has an electron in an excited state, it can stimulate the emission of a photon from the excited electron.
  • STIMULATED EMISSION
    • This causes two photons of light moving off in phase creating coherent light. It has the identical energy, direction and phase to the original photon.
  • STIMULATED EMISSION
    • This only happens when the electron is in an excited state.
    • As an electron is only in this state for 10 -8 s, or in the ground state, this does not happen under normal circumstances.
    • Stimulated Emission
  • STIMULATED EMISSION
    • To extend the time that the electron is in the excited state, the higher state must be METASTABLE.
    • In this way, the emission is less likely to be spontaneous but stimulated by other photons.
  • STIMULATED EMISSION
    • This metastable state is important, as it is a way of producing coherent light that is required for lasers.
    • The metastable state is used in luminous watch dials.
  • STIMULATED EMISSION
    • Phosphorescent materials are used as they have an excited state that will hold the electron in it for an extended period of time, perhaps as long as a second or two.
    • As there are many atoms, the electrons will take longer to drop down to the ground state and release their energy as light much more slowly.
  • STIMULATED EMISSION
    • This means the watch dial will glow in the dark for long periods after the light that originally stimulated the atoms has gone.
  • APPLICATION-LASERS
    • Laser is an abbreviation for “ Light Amplification of Stimulated Emission of Radiation ” .
    • It is attached to this topic because it utilizes the process of STIMULATED EMISSION to produce the LASER light.
    • What is a Laser?
  • APPLICATION-LASERS
    • It is used in many different pieces of equipment such as LADS, CD players and laser printers.
    • There are three main components of lasers that we will study in turn.
  • APPLICATION-LASERS
    • Laser Medium – Gas Medium (substance made from atoms in metastable state – Neon gas)
    • Imagine that a laser has two states, a ground and excited state. An electron can change state in three different ways.
    • 1.   Absorption (move to an outer shell).
    • 2.   Spontaneous emission (10 -8 s).
    • 3. Stimulated emission (Metastable).
  • APPLICATION-LASERS
    • As a laser amplifies light, the first two ways of changing state are not useful as no extra photons are created.
    • Electrons are elevated to an excited state and held there in a metastable state.
    • This causes a population inversion where there are more electrons in the stimulated state than in the ground state.
  • APPLICATION-LASERS
    • If an electron does decay by spontaneous emission, it can cause the spontaneous emission of another electron and so create another photon of light.
  • APPLICATION-LASERS
    • These two photons can now cause stimulated emission of another two, creating 4 photons and so on. This process amplifies the light.
    • Laser Principle
  • APPLICATION-LASERS
    • In the Helium Neon laser, Neon is the laser (gas) medium (hence the red colour characteristic of neon) but the helium makes the process easier and more efficient.
    • Helium is stimulated to the excited state. The excited state of helium is very close to the upper excited state of neon (it has two excited states).
  • APPLICATION-LASERS
    • If an electron in the excited state of helium decays, it can easily transfer the energy to the neon atom by exciting an electron in neon to the upper excited state.
    • We do not see the photon from helium as its energy is used to excite a neon atom. This particular state is metastable and so a population inversion occurs.
  • APPLICATION-LASERS
    • This leads to many stimulated emissions to neon’s lower excited state.
  • APPLICATION-LASERS
    • As this state is not metastable, many spontaneous emissions are made to the ground state, leaving space for more stimulated emissions.
    • This means the process can continue indefinitely resulting in a continuous beam of photons corresponding to the energy difference between the two excited states of neon.
    • The  of this beam is 6.328 x 10 -7 m.
  • APPLICATION-LASERS
    • The Pump (a way of exciting the electrons – high PD)
    • Previously it was stated that electrons need to be elevated, or ‘pumped’ to an excited level. Absorbing light from a source like a flash used in cameras can do this.
  • APPLICATION-LASERS
    • A helium neon laser uses an electrical discharge through a gas.
    • Electrons are excited across a P.D. and collide with electrons in the gas causing them to become elevated to an excited level.
    • This is more effective with helium than neon, which explains the presence of helium in the laser.
  • APPLICATION-LASERS
    • The Cavity (space in the tube)
    • To increase significantly the amplification of the light in a laser, it is made to cross the laser medium by reflecting it back and forth from mirrors. The mirrors are shaped so that they will focus the light and compensate for the spreading of the light due to diffraction.
  • APPLICATION-LASERS
    • One of the mirrors partially transmits light so when the intensity is great enough, it can pass through from the cavity and be seen as a beam.
  • APPLICATION-LASERS
    • Properties of Laser Light
    • As the beam is reflected back and forth between parallel mirrors, it diverges only marginally. The divergence is in the order of 1mm per 1m of travel. Laser light is said to be unidirectional .
  • APPLICATION-LASERS
    • The intensity can be varied to suit the use. In CD players, the intensity is very low.
    • It can be very high when used to weld metals.
    • Changing the partially reflecting mirror varies the intensity.
  • APPLICATION-LASERS
    • The wavelength of the light beam is determined by the energy levels of the atom being excited.
    • As they are fixed, the light is monochromatic.
    • There is, for reasons outside the course, a small range of wavelengths
    • (  10 -15 m).
  • APPLICATION-LASERS
    • The light from laser is extremely coherent.
    • As the photon emitted from each emission is identical in wavelength, phase and direction to the stimulating photon, all photons have the same properties.
  • APPLICATION-LASERS
    • Safe Handling
    • As laser light can have very high intensities, it is important that the beam doesn’t come in contact with the body.
    • The PAC laser is very weak, 0.95mW and so is only of concern if it is shone into the eyes.
    • The intensity is about the same as looking at the sun on a bright day.
  • APPLICATION-LASERS
    • If using the laser in a room that is darkened, the pupil is dilated and so the beam can cause more damage.
    • Even if not shone directly into the eye, reflections from other surfaces (or particles in the air) may also cause damage.
  • APPLICATION-LASERS
    • Some lasers produce radiation in the ultraviolet and infrared parts of the spectrum.
    • As this part of the spectrum is invisible to us, the laser may appear to have only a very low intensity beam.
    • The radiation from the UV or infrared beam could also cause serious damage.
  • APPLICATION-LASERS
    • As most lasers use a gas discharge, large P.D.’s may be involved.
    • If the case is tampered with, a person may receive a large electrical shock.
  • APPLICATION-LASERS
    • Uses of Lasers
    • Lasers are used in manufacturing for welding metals under computer control. Sails and clothing can also be cut exactly using lasers.
    • In the semiconductor industry, components such as resistors can be trimmed and integrated circuits can also be made.
  • APPLICATION-LASERS
    • Laser light is used in fibre optics for communications . This could take any form from cable T.V. to undersea telephone links.
    • In surveying , lasers are used to check the alignment of structures such as ceilings, walls, ribs or frames.
    • They are also used in the construction of tunnels such as the Hysen tunnel or bridges.
  • APPLICATION-LASERS
    • It can also measure distances. The first major use for this was to find the distance from the earth to the moon using a reflector placed on the moon by Apollo astronauts.
    • Nowadays, the military has found a use for it to determine the distance to targets and for aiming weapons.
  • APPLICATION-LASERS
    • Shops and libraries use lasers for barcode scanning . A beam is scanned across the barcode that has different width and spacing of bars to identify the product.
    • A computer decodes an electrical signal, looks up the price, then adds up the bill. Handheld scanners can also be used.
    • They employ semiconductor lasers as used in laser pointers.
  • APPLICATION-LASERS
    • Surgery
    • As lasers can be focussed very well onto small points, it can be used as a scalpel and burn target tissue without damaging the surrounding tissue.
    • Dermatologists use this to remove ‘portwine stains’ such as the one on Michael Gorbachev.
  • APPLICATION-LASERS
    • Eye surgeons such as at the Ashford Eye Centre can remove cataracts or alter the lens of a person so they no longer require glasses.
    • The heat from the laser also allows it to cauterise the wound and stop any bleeding which reduces the chance of infection.