Atom and Ev Atoms, Energy, and the Heisenberg Uncertainty Principle By Lee Wignall The idea that matter is made of small, indivisible “pieces” has been around for thousands of years. The word “atom” is derived from the Greek word “atomos”, meaning “that which cannot be divided”. Leucippus: Greek philosopher/scientist Democritus: student of Leucippus and considered “Father of Modern Science” Fast forward over 1000 years… In 1808, John Dalton came up with four postulates of his “atomic theory.” (there are actually five, but only the first four are wellknown) John Dalton Fast forward 100 years… In 1906, Thomson created a current of negatively charged particles in a cathode-ray tube that were much smaller than an atom. J.J. Thomson He called these tiny particles “corpuscles” and created the new model of the atom. Negatively charged “corpuscles” were spread around a positively charged sphere, like raisins in plum pudding. Plum pudding atomic model 2 years later… In 1908, Ernest Rutherford fired alpha particles (protons) at gold atoms and discovered that the center of the atom is densely packed and positively charged. Gold foil experiment He called the center of the atom the nucleus and proposed that the negatively charged electrons orbit around the nucleus like planets. 5 years later… In 1913, Neils Bohr introduced quantum mechanics into the atomic model. He proposed the idea that electrons were only allowed to orbit the nucleus at certain energy levels. Instead of an infinite number of possible orbits, only certain orbits were allowed. (“quantizing the atom”) Electrons could only jump to a higher energy orbit if they absorbed a photon with enough energy. Electrons jumping down energy levels would emit a photon with the exact energy difference between orbits. The Electron Volt (Ev) -13.6 Ev -3.4 Ev The energy of subatomic particles is measured in electron volts (Ev). An electron absorbs a photon with enough energy to jump to the next highest energy level. The Electron Volt (Ev) An electron jumps down to a lower energy level and emits a photon with the energy difference. -13.6 Ev -3.4 Ev The lowest “natural” energy state is the ground state. The energy of subatomic particles is measured in electron volts (Ev). Glow in the Dark Electrons in the paint absorb energy from ambient light and jump to higher energy orbits. However, this energetic state is unstable and as soon as the energy source (light) is removed, the electrons jump back down to their ground-state, emitting photons with energy equal to the difference between orbits. Louis de Broglie de Broglie proposed in 1924 that if light can act as both a particle and a wave, then shouldn’t all moving particles act as waves as well? Matter waves: the more massive an object, the smaller the associated wavelength. Therefore, large objects have wavelengths that are WAY too short to be noticeable on any level. 13 years later… In 1926, Erwin Schrodinger introduced probabilty into the structure of the atom. Borrowing de Broglie’s idea that all matter exhibited wave-like properties, Schrodinger created an equation that describes the way a wave evolves. The Schrodinger Wave Equation ? Don’t worry, I’ll break it down for you. The wavefunction: “psi” Called the how “del-squared number”: it’s operator”, the square this root ofover minus one. Describes Describes allthe thewavefunction forces acting changes on thequantity particle. time. This“imaginary is “h-bar”, Planck’s constant (6.6 *10^-34 J*s) The mass ofhow thepi. particle being described. describes the wavefunction, , changes from divided by 2 times one place to another. Applying the Schrodinger Equation to an Atom When you apply the Schrodinger Equation to an electron in an atom, you end up with a wavefunction that gives the probability of finding the electron in any given location. All the possible locations “smear” together to create an electron cloud. This computer analysis shows the results of the Schrodinger Equation under certain circumstances. You can clearly see the denser “cloud” where the probability of finding the electron is higher. ? ? ? ? ? ? ? nucleus ? ? ? ? ? ? ? Electron cloud Without direct observation, the position of the electron is given by the wavefunction--a probability wave that describes the chances of finding it at any given location. Observed location of electron ? ? ? ? ? ? ? nucleus ? ? ? ? ? ? ? Electron cloud As soon as observation we make andoes observation aboutbegin the precise Only upon the electron to so is Without But electrons direct observation, are both a wave the position and a particle, of the electron and location of the electron, wewith reduce the probability of act as a particle. Just like the double-slit given byitthe wavefunction--a probability wave by thatdirect we should be able to location determine it’s position finding in any other to zero, thereby experiment! describes observation thethe (with chances high-tech of finding equipment). it at any given location. “collapsing wavefunction.” The Uncertainty Principle An interesting feature of the Schrodinger wave equation is the more precise you calculate the position of the particle, the less you can determine the momentum (and vice versa). Werner Heisenberg This is called the Heisenberg Uncertainty Principle. (formulated in 1927) The combined uncertainty in boththe theless position So, the more you know about position, you The uncertainty ininthe particle’s momentum. The uncertainty the particle’s position. momentum has to be greater than (or know and about momentum! equal to) h-bar divided by two. Click below to begin youtube video on Heisenberg’s Uncertainty Principle. http://www.youtube.com/watch?v=KT7xJ0tjB4A I know it’s 1932 and I’m tardy to the party, but I just discovered the neutron. I’m kind of a big deal. James Chadwick Heisenberg and Bohr, two heavyweights in atomic theory for their quantum contributions, immediately adopted the theory and used the neutron to explain small discrepancies in their experiments.