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The Electron

The Electron

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The Electron

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  1. The Electron 6.0 Chemistry

  2. Development of the Periodic Table • History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C). • By the mid 1800’s, about 60 elements had been identified. • Sept 1860 – chemists assembled at the First International Congress of Chemists in Germany to settle the controversial issues such as atomic mass. Standard values set for atomic mass and improved communication for research.

  3. Johann Dobereiner: 1817 Organized the elements into sets of three with similar properties. He called these groups triads. The middle element is often the averageof the other two. Ex) Cl – 35.5 Br – 79.9 I – 126.9 Ca Avg Sr Ba

  4. Triads on the Periodic Table

  5. B. John Newlands: 1866 • Arranged elements in order of increasing atomic mass. • Noticed repeating patterns in the elements’ properties every 8th element. • Law of Octaves - properties of elements repeated every 8th element. • There were 62 known elements at the time.

  6. C. Dmitri Mendeleev: 1869 • Arranged elements in order of increasing atomic mass. • Similar properties occurred after periods (horizontal rows) of varying lengths. • Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column. • Periodic – repeating properties or patterns • Noticed inconsistencies in the arrangement.

  7. Mendeleev’s 1st Periodic Table

  8. He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I) • He also left several blanks in his table. • In 1871, he correctly predicted the existence and properties of 3 unidentified elements – Sc, Ga and Ge • These elements were later identified and matched his predictions.

  9. 1st Periodic Law Properties of the elements repeat periodically when the elements are arranged in increasing order by atomic mass Mendeleev is known as the Father of Chemistry #101 honors Mendeleev

  10. D. Henry Moseley: 1911 • Studied X-ray spectral lines of 38 metals. Each element had a certain amount of positive charge in the nucleus which are called protons. • Analyzed data and found that the elements in the PT fit into patterns better when arranged in increasing nuclear charge, which is the Atomic Number. • The Modern Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

  11. Glenn Seaborg “Seaborgium” Sg #106 • Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII. • He suggested pulling the “f-block” elements out to the bottom of the table. • He was the principle or co-discoverer of 10 transuranium elements. • He was awarded the Noble prize in 1951 and died in 1999.

  12. Seaborgium is the exception… • After some argument between the USA and the rest of the world, element 106 was named Seaborgium shortly before he died. This was a matter of some controversy because the International Union of Pure and Applied Chemistry, IUPAC, the body that deals with naming in chemistry, had previously ruled that elements should not be named after living people.

  13. Parts of the Periodic Table • Horizontal Rows – PERIODS • There are 7 periods in the periodic table • Elements in a period do NOT have similar properties. • Vertical Columns – GROUPS or FAMILIES • Labeled 1-18 • IA-VIIIA are the Main-group or representative elements. • Elements in a group have similar properties. • Why?

  14. Hydrogen (1) Alkali metals (1) – most reactive metals; reactivity increases down the group Alkaline earth metals (2) Boron family (13) Carbon family (14) Nitrogen family (15) Oxygen or Chalcogen family (16) Halogens (17) Noble gases (18) - inert Transition elements or metals (3-12): d-block Inner transition elements or metals (f-block) Lanthanides or lanthanide series Actinides or actinide series Transuranium elements Family Nameswrite these on your P.T.

  15. Parts of the Periodic Table

  16. E. Metal, Nonmetals and Metalloids (Semimetals): - Good conductors of heat & electricity - High melting points most solids at room temperature • Metals Found on the LEFT side of the PT 2. Nonmetals Located on the RIGHT side of the PT 3. Metalloids - Properties of both metals & nonmetals - High luster (shiny) - Ductile (can be drawn into thin wire) - Malleable (bends without breaking) - High densities - Reacts with acids - Brittle (easy to break) - No luster (dull) - Insulators nonconductors - Neither ductile nor malleable - Nonreactive with acids (Semimetals)

  17. Review of Early Atomic Theories • Dalton  • Thomson  (plum pudding)

  18. Atomic Theories

  19. Bohr – electrons in a particular path have a fixed energy called energy levels • Rungs of a ladder • Quantum Mechanical (Schrödinger) Model • Electrons better understood as WAVES • Does not tell where the electrons are located • Electrons have a certain amount of energy - QUANTIZED

  20. Parts of a Wave

  21. Light as a Wave Characteristics of a Wave • Amplitude:Height of the wave from the baseline. The higher the wave the greater the intensity. • Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m). Distance between similar points on 2 consecutive waves. • Frequency: (ν , “nu”) The number of waves that pass a fixed point per unit of time. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz) ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 106Hz (cycles/sec)

  22. Wavelength Trough Baseline Amplitude Crest 3 Amplitude Wavelength

  23. D. Electromagnetic Radiation- a form of energy that exhibits wavelike behavior as it travels through space- all forms of EM radiation move at the speed of light

  24. Speed of Light (c) E. 3.00 x 108 m/s or 186,000 miles/sec. The relationship between wavelength and frequency can be shown with the following equation: This is an indirect relationship. If λ then ν . c = λ ν

  25. Visible Light Microwaves Radio/TV Ultraviolet Gamma Rays Radar Infrared X-Rays Low High Long Short High Low Energy Energy Red Orange Yellow Green Blue Violet

  26. Quantum Theory • Planck’s Hypothesis: (Max Planck 1900) • Studied emission of light from hot objects • Observed color of light varied with temperature • Suggested the objects do not continuously emit E, but emit E in small specific amounts • Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural). • Quantum = minimum amount of E that can be lost or gained by an atom • Energies are quantized. (Think steps not a ramp) e- e- X e- e-

  27. Incandescent light bulbs give off most of their energy in the form of heat-carrying infrared light photons -- only about 10 percent of the light produced is in the visible spectrum. This wastes a lot of electricity. Cool light sources, such as fluorescent lamps and LEDs, don't waste a lot of energy generating heat -- they give off mostly visible light. For this reason, they are slowly edging out the old reliable light bulb.

  28. Max Planck’s Energy Equation • Proposed that energy is directly proportional to frequency. E = h Planck’s equation for each quantum h = Plank’s constant = 6.626 x 10-34 J.s This is a direct relationship. As energy increases, frequency increases.

  29. Albert Einstein While well-known for the equation E=mc2 , Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics. (1879 – 1955) German Physicist

  30. Albert Einstein and the Photoelectric Effect Refers to the emission of electrons from a metal when light shines on the metal Observations: • Electrons are ejected by light of sufficient energy. Energy minimum is different for different metals. • The current (# of electrons emitted/s) increases with brightness of the light.

  31. Albert Einstein and the Photoelectric Effect Conclusions: • Proposed that light consists of quanta of energy that behaves like particles. • Quantum of light = photon = massless particle that carries a quantum of energy. • Proposed the Dual Nature of Light: its wave and particle nature. • Light travels through space as waves • Light acts as a stream of particles when it interacts with matter.

  32. Light (Electromagnetic Radiation)

  33. Spectroscopy Definition: a method of studying substances that are exposed to some sort of continuous exciting energy. A. Emission Line Spectra: contains only certain colors or wavelengths (  ) of light. 1. Every element has its own linespectrum (fingerprint).

  34. Continuous Spectrum – White Light Line Spectrum – Excited Elements

  35. B. White light gives off a Continuous Spectruma blending of every possible wavelength

  36. Gas Discharge Tubes • Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.

  37. Flame Tests • used to test qualitatively for the presence of certain metals in chemical compounds.  • the heat of the Bunsen flame excites electrons that emit visible light. Copper(II) sulfate Lithium chloride Potassium chloride Barium nitrate

  38. Spectroscope • Uses a diffraction grating to diffract the light into particular wavelengths of light.

  39. A Line Spectra result from excited elements - as electrons of an element gain energy and rise to an excited state they then fall back to their ground state in the same pattern producing the same energy drop each time which we see as individual wavelengths of light.

  40. III. Atomic Spectra and the Bohr Model of Hydrogen (1913) Neils Bohr - Danish Scientist Explained the bright-line spectrum of hydrogen Study: • Added E as electricity to H gas at low pressure in a tube. • Emitted E as visible light, was observed through a prism Result: Hydrogen emitted 4 distinct bright lines of color, aka bright line spectrum

  41. Electrons release energy as they fall back to a lower energy level

  42. Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their ground states.

  43. Path of an excited electron as it “falls” back to the Ground State • When electrons gain energy, they jump to a higher energy level (excited state). • Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state. • As they fall, they emit electromagnetic radiation. • Depending on how far they fall determines the type of radiation (light) released.

  44. Bohr Model of Hydrogen Conclusion: • *Unique line spectrum is due to quantized electron energies. • *Electrons are in specific orbits related to certain amounts of energy known as stationary states. • *Orbits are related to energy levels. • *Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …) • *Lowest energy level = ground state • *Electrons absorb certain amounts of energy to move to a higher energy level farther away from the nucleus = excited state • *Electrons return to the more stable ground state and release a photon that has energy equal to the difference in energy between the energy levels. • from E2 to E1: Ephoton = E2 – E1 (difference in energy)

  45. The Bohr Atom for Hydrogen -a Model • Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum. • Successful in calculating the energy needed to remove hydrogen’s electron H(g) + energy  H+1(g) + 1e- Calculated ionization E = observed ionization E = 1312.1 kJ/mol

  46. Lyman, Balmer and Paschen series of the Hydrogen Atom • Lyman series: electrons fall to n = 1 and give off UV light. • Balmer series: electrons fall to n = 2 and give off visible light. • Paschen series: electrons fall to n = 3 and give off infrared light.