3. 1 & 3. 2 Chemistry Notes Dalton * In 1805, John Dalton reintroduced the idea to explain 3 fundamental principles Experimental Work * Atoms of different elements have different properties * Law of definite proportion and multiple compositions: atoms of 2 or more elements can combine in a fixed ratio to form new substances depending on their combining capacities (eg. H2O vs H2O2) * Law of conservation of mass: atoms cannot be created or destroyed during a chemical reaction Conclusions * All matter is composed of atoms Atoms are the smallest pieces of matter and cannot be broken down further * All atoms of one element have identical properties Problem * Development of a cathode ray tube (by William Crookes) Thompson (1897) Experimental Work * Used a cathode ray tube a vacuum tube with electrodes at both ends * Found that there were charged particles that were travelling from one end of the tube to the other (from negative end to positive end) Conclusion * Proposed that an atom was a positively charged empty sphere containing negatively charged electrons raisin scone analogy What Thompson left us with? Atoms consist of negative electrons embedded within a positively charged sphere * Analogy of raisin bun often used Milikin’s Famous Oil Drop Experiment * Determined size and charge on electron * Discovered charge on single electron was 1. 6 x 10^19 C How it worked? * Knew mass of single drop of oil, calculated gravity on one drop * Charge was applied to falling drops by illuminating bottom chamber with x-rays, exciting electrons, causing them to attach to oil. * Using a battery, electric voltage was applied to the plates.
When just right, the electromagnetic force would balance out the force of gravity, suspending particles in midair. * Noticed charge was always multiples of 1. 6 x 10^19 * Q= mg/E Gold Foil Experiment * Radioactive particles (alpha radiation) were fired at thin gold sheets * Screens coated with zinc sulfide detected the presence of the alpha radiation * Vast majority of alpha particles passed straight through gold sheet, however, approximately 1 in 8000 particles were deflected Chadwick and the Neutron When calculating the mass of specific nuclei, the calculated mass did not correlate with the associated charge of the nucleus * Chadwick proposed that neutral particles must be present to make up for the missing mass * Chadwick proposed a positive nucleus containing neutral particles Isotopes * Mass spectrometers were used to discover that all atoms of the same element were not the same * Elements contained several different forms of isotopes (atoms with the same number of protons, but different numbers of neutrons) Problem with the Rutherford Model Physics – bodies are accelerating when they change speed and/or direction * And electron travelling in a circular orbit is constantly changing its direction and therefore accelerating * This acceleration would result in the electrons emitting electromagnetic radiation, lose electrons, and collapsing the atom as it continuously spirals inward because it is losing electrons Enter Max Plank * Her was studying the emission of light from hot objects * What is visible light? When objects are heated, they emit various colors of light depending on how hot the object is * Ex. “white hot” objects are emitting the whole range of the visible spectrum * “red hot” objects emit light with wavelength of the infrared – longest wavelength * “blue hot” objects are the hottest as they emit light of shortest wavelength * Hot objects emit radiation. The hotter they are, the more energetic the radiation emitted is. The electromagnetic radiation changes as the object gets hotter. * The color of light emitted reveals temperature Explaining Intensity vs.
Energy – The staircase which changed physics * Planck suggested that energies of the vibrating atoms in the heated solid were multiples of small quantities of energy (energy was not continuous) * Introduced the term “quantum” * The slope is actually more like a staircase * Each step represents a ‘quanta’ of energy * A quanta is derived from quantity and refers to the smallest possible unit of energy that can be associated with a specific sub-microscopic even * An atom has to absorb or release an entire package (quanta) of energy or none at all.
There is no ‘in between’ Heinrich Hertz: the photoelectric effect * Photoelectric effect when light is shone on a metal surface, electrons are released from the surface of the metal. The number of electrons released per second can be measured by a connected ammeter * Frequency is different from intensity. Electrons will only jump off if the frequency is right, however, how many electrons jump off will depend on the intensity of the light. How fast they jump off will also increase with higher frequency * The amount of energy in a light wave is proportionally related to its frequency.
High frequency light has high energy, low frequency light has low energy (violet has the most energy and red has the least) Einstein puts 2 and 2 together * In 1905, Einstein received the Nobel Prize for applying Planck’s idea to the photoelectric effect * When light strikes metal, some of the energy is used to allow the electron to break free from the metal, the rest of the energy is left over as the kinetic energy of the ejected electron * If one electron absorbs one photon (quanta of energy), it must be great enough or the electron to be able to escape * No electrons escape at low photon energies because the energy of the single photon was insufficient for the electron to escape the metal Energy of Quanta of Energy – Photons * E = h x f, where E is the amount of energy in joules (J), h is Planck’s constant 6. 6 x 10^-34, and f is the frequency in hertz * A photon is a packet of energy, with energy values corresponding to the frequency of the electromagnetic wave Einstein’s Proposals Light is quantized like a particle (photon) * Light exist as bundles of photons, with each photon independent of each other * This means that light has certain particle properties as well * The energy of a photon is proportional to its frequency and nothing else. * Therefore, a phonon is a small packet of energy corresponding to a specific frequency of light (E=hf) Spectroscopy The spectroscope was invented by Robert Bunsen and Gustav Kirchhoff in the 1850s to study light * When white light passes through spectroscope (containing a prism or diffraction grating), the light is divided into a continuous rainbow of colors (continuous spectrum) Bunsen and Kirchhoff (1859): invented the spectroscope * When elements were heated in a Bunsen burner flame, each element produced a flam color and a bright line spectrum that was characteristic of the element * Continuous Spectrum – a display of all colors.
It comes from the “dispersion” (refraction) of white light passing through a prism * Dark Line spectrum (absorption spectrum) – certain colors are missing from a display of colors produced by white light passing through a gas and then through a prism. These missing lines enable scientists to identify the gas that the light passed through * Bright line spectrum (emission spectrum) – when a gas is “energized” by electricity or heat or light, the gas emits light of a specific color (not white light). When this light is passed through a prism it is refracted into a pattern of a few bright lines of color.
Each substance has a unique, bright line “signature”. This pattern of colored lines represents the same pattern of dark lines of missing color in the dark line spectrum] Bohr’s theory was needed to explain the bright/dark spectrum and Einstein’s photons 1. Electrons travel in an atom in circular orbits. Each orbit represents a specific energy level. All electrons in one orbit/energy level will have the same amount of energy, which is quantized (discrete packet) 2. There is maximum number of electrons allowed in each orbit 3. When electrons absorb a photon of light, they jump from a lower energy level to a higer energy level.
This absorption of a photon of light energy results in a dark line in the absorption spectrum 4. When electrons jump from a higher energy level to a lower energy level, energy is released as a photon of light. This release of photon from the atom results in the bright line in the emission spectrum 5. When electrons are at the lowest energy level, they are in “ground state” How does Bohr’s Energy levels of electrons relate to the periodic table? * Each period represents one energy level – Period 1 1 Energy level, Period 2 2 energy levels, etc. There is a maximum number of electrons in each lever (level 1 2 electrons, level 2 8 electrons, level 3 8 electrons) Power Point 2 Problems with Planetary Model * If electrons were accelerating, photons of electromagnetic radiation should be emitted * Obviously this is not the case * The Rutherford planetary model is insufficient as a model to explain matter Quantum Theory * All electrons in all atoms can be described by 4 unique quantum numbers * Quantum numbers are used to describe the approximate location and characteristics of electrons surrounding an atom based on the energy levels of an atom * There are 4 quantum numbers Principle quantum number (n) * Designates main E level of electron * Secondary quantum number (l) * Describes E sublevels of electrons * Magnetic Quantum Number (ml) * Relates to direction of electron orbit * Spin Quantum number (ms) * Relates to the spin of an electron Principle Quantum Number (n) * n=1, 2, 3, 4 etc. * n=1 means Energy level 1 and so on Secondary Quantum number, l * (l) describes shapes of sublevels (subshells) of the main energy level * Sommerfeld looked more closely at the H line spectrum. Found that main lines of bright line spectrum split into more lines. The number of sublevels equals the value of the principle quantum number * Has integral values from 0 to (n-1) for each value of n * If n=3, then there are three sublevels. L = 0, 1, 2 * Each l number represents a possible shape of the orbital. (hence if l=0, 1, 2, then there are 3 possible shapes) Third Quantum Number: Magnetic Quantum number, ml * describes the orientation of electron orbital in space (therefore orbitals could exist at different angles to each other in 3-d) * For each value of l, ml, can vary from -1 to 1 Shapes of Orbitals ‘s’ (l=0) orbital is spherical, ml = 0 * ‘p’ (l=1) , ml = -1, 0, 1 * ‘d’ (l=2) , ml = -2, -1, 0, 1, 2 * ‘f’ (l=3) orbitals are much more complex, ml = -3, -2, -1, 0, 1, 2, 3 Classification of Energy Subshells * Each distinct sublevel has specific number of orbitals. * Each orbital has a different orientation The spin quantum number, ms * Pauli – each electron spins on its axis in one of 2 ways clockwise or counterclockwise * The spin quantum possesses only two values; either +1/2 (clockwise) or -1/2 (counter-clockwise) New Orbital Way Orbitals are 3 dimensional probability distribution graphs which help chemists visualize where electrons are most likely to be found Electron Orbitals * An electron orbital is described as the region of space where an electron may be found * Orbits are rings surrounding the nucleus, whereas orbitals are probability clouds or clouds of electron density * More than one orbital can be found within an energy level Pauli’s Exclusion Principle * No two electrons in an atom can have the same 4 quantum numbers!
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