Electrons and the Periodic Table

1. Electrons and the Periodic Table Honors Chemistry

2. The Periodic Table

3. 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.

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

5. Triads on the Periodic Table

6. 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.

7. 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

8. Mendeleev’s 1st Periodic Table

9. Mendeleeve Noticed inconsistencies in the arrangement of increasing atomic mass. • 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.

10. 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 Element 101 (Md) honors Mendeleev

11. What’s wrong with Mendeleev’s PT? • We know that Mendeleev's periodic table was underpinned by false reasoning. • Mendeleev believed, incorrectly, that chemical properties were determined by atomic weight. • In 1869 the electron itself had not been discovered until 1896 - 27 years later. • It took 44 years for the correct explanation of the regular patterns in Mendeleev's periodic table to be found...

12. Henry Moseley: 1913 • Fired electrons at atoms, resulting in emission of x-rays. Each element emitted x-rays at a unique frequency. • He examined the pattern that was best explained if the positive charge in the nucleus increased by exactly one unit from element to element. • Element are different from one another because their atoms have different number of positive charge (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 Table: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

13. 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.

14. 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.

15. 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?

16. 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.

17. 1A 1 8A 18 3A 4A 5A 6A 7A 13 14 15 16 17 2A 2 Transition Metals 3 4 5 6 7 8 9 10 11 12 Inner Transition Metals Lanthanide Series Actinide Series 15= Nitrogen Group 1 = Alkali Metals and Hydrogen Group 16 = Oxygen or Chalcogen Group 2 = Alkaline Earth Metals 17 = Halogen Group 13 = Boron Group 14 = Carbon Group 18 = Noble Gas Group

18. Parts of the Periodic Table

19. The ELECTRON How are the electrons arranged around the nucleus?

20. To Review • Dalton  Thomson  Rutherford  Bohr  Quantum Mechanical (Schrödinger) Model

21. 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

22. Parts of a Wave

23. 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)

24. Wavelength Trough Baseline Amplitude Crest 3 Amplitude Wavelength

25. 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

26. 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

27. In the US, AM stations all have longer wavelengths than the FM stations. • AM broadcast stations are licensed to operate only in the band 550-1700 kHz • FM broadcast stations are licensed to operate only in the band 88-108 MHz

29. 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

31. 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. http://home.howstuffworks.com/light-bulb2.htm

32. 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-

33. 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.

34. 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

35. 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. my.hrw.com

36. 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. my.hrw.com

37. Light (Electromagnetic Radiation)

38. 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).

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

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

41. 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.

42. 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

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

45. 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.

46. 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

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

50. 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.