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The Electron 6.0 Chemistry. Development of the Periodic Table 1)History of the Periodic Table – By the end of the 1700’s, scientists had identified.

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Presentation on theme: "The Electron 6.0 Chemistry. Development of the Periodic Table 1)History of the Periodic Table – By the end of the 1700’s, scientists had identified."— Presentation transcript:

1 The Electron 6.0 Chemistry


3 Development of the Periodic Table 1)History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H 2, N 2, O 2, C). 2)By the mid 1800’s, about 60 elements had been identified. 3)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 – Ca AvgSr Ba

5 Triads on the Periodic Table

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

7 C. Dmitri Mendeleev: Arranged elements in order of increasing atomic mass. 2.Similar properties occurred after periods (horizontal rows) of varying lengths. 3.Organized the 1 st periodic table according to increasing atomic mass and put elements with similar properties in the same column. 4.Periodic – repeating properties or patterns 5.Noticed inconsistencies in the arrangement.

8 Mendeleev’s 1 st Periodic Table

9 6.He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I) 7.He also left several blanks in his table. 8.In 1871, he correctly predicted the existence and properties of 3 unidentified elements – Sc, Ga and Ge 9.These elements were later identified and matched his predictions.

10 1 st 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

11 D. Henry Moseley: Studied X-ray spectral lines of 38 metals. Each element had a certain amount of positive charge in the nucleus which are called protons. 2.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. 3.The Modern Periodic Law : When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.

12 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.

13 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. Atomic # IUPAC Unnilquadeum Unq Unnilpentium Unp Unnilhexium Unh Unnilseptium Uns Agreed in 1995 Dubnium (Dubna, Russia) Joliotium (Frederic Joliot) Rutherfordium (Earnest Rutherford) Bohrium (Neils Bohr) Agreed in 1996 RutherfordiumDubniumSeaborgiumBohrium

14 Parts of the Periodic Table A.Horizontal Rows – PERIODS – There are 7 periods in the periodic table – Elements in a period do NOT have similar properties. B.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?

15 Family Names write these on your P.T. 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

16 Parts of the Periodic Table

17 E. Metal, Nonmetals and Metalloids (Semimetals): 1.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 - Good conductors of heat & electricity - High luster (shiny) - Ductile (can be drawn into thin wire) - Malleable (bends without breaking) - High melting points most solids at room temperature - High densities - Reacts with acids - Brittle (easy to break) - No luster (dull) - Neither ductile nor malleable - Nonreactive with acids - Insulators nonconductors (Semimetals)

18 Review of Early Atomic Theories Dalton  Thomson  (plum pudding)

19 Atomic Theories

20 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

21 Parts of a Wave

22 Light as a Wave Characteristics of a Wave A.Amplitude: Height of the wave from the baseline. The higher the wave the greater the intensity. B.Wavelength: ( λ, “lambda”) in nanometers (1 x m). Distance between similar points on 2 consecutive waves. C.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 10 6 Hz (cycles/sec)

23 Baseline Crest Wavelength Amplitude WavelengthTrough 3

24 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

25 Speed of Light (c) E x 10 8 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 Visible Light Radar Microwaves Infrared Radio/TV Ultraviolet X-Rays Gamma Rays Low Long High Short RedOrangeYellowGreenBlueViolet Energy Low High


29 Quantum Theory A.Planck’s Hypothesis: (Max Planck 1900) 1.Studied emission of light from hot objects 2.Observed color of light varied with temperature 3.Suggested the objects do not continuously emit E, but emit E in small specific amounts a.Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural). b.Quantum = minimum amount of E that can be lost or gained by an atom c.Energies are quantized. (Think steps not a ramp) e-e- e-e- X e-e- e-e-

30 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.fluorescent lampsLEDs

31 Max Planck’s Energy Equation 4.Proposed that energy is directly proportional to frequency. E = h Planck’s equation for each quantum h = Plank’s constant = x J. s This is a direct relationship. As energy increases, frequency increases.

32 Albert Einstein (1879 – 1955) German Physicist While well-known for the equation E=mc 2, Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics.

33 Refers to the emission of electrons from a metal when light shines on the metal Observations : 1.Electrons are ejected by light of sufficient energy. Energy minimum is different for different metals. 2.The current (# of electrons emitted/s) increases with brightness of the light. Albert Einstein and the Photoelectric Effect

34 Conclusions: 1.Proposed that light consists of quanta of energy that behaves like particles. 2.Quantum of light = photon = massless particle that carries a quantum of energy. 3.Proposed the Dual Nature of Light : its wave and particle nature. a)Light travels through space as waves b)Light acts as a stream of particles when it interacts with matter. Albert Einstein and the Photoelectric Effect

35 Light (Electromagnetic Radiation)

36 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 line spectrum (fingerprint).

37 Continuous Spectrum – White Light Line Spectrum – Excited Elements

38 B. White light gives off a Continuous Spectrum a blending of every possible wavelength

39 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.

40 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) sulfateLithium chloride Potassium chlorideBarium nitrate


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

43 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.gain energy and rise to an excited state they then fall back to their ground state

44 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

45 Electrons release energy as they fall back to a lower energy level

46 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.

47 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.

48 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 E 1, E 2, E 3, … (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 E 2 to E 1 : E photon = E 2 – E 1 (difference in energy)

49 The Bohr Atom for Hydrogen - a Model 1.Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum. 2.Successful in calculating the energy needed to remove hydrogen’s electron H (g) + energy  H +1 (g) + 1e -  Calculated ionization E = observed ionization E = kJ/mol

50 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.

51 When electrons absorb energy they jump to a higher ( excited ) state. n=2 n=3 n=4 n=5 n=6 n=7 Electrons are not stable. Radiation ( light ) is emitted when an electron falls back from a higher level to a lower level.


53 Ultraviolet Light Visible Light Infrared Light

54 Atomic Spectra Hydrogen Lithium Mercury Helium Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.

55 Limitations of the Bohr Model a.Model could not calculate the wavelengths of observed spectra of multi-electron atoms. b.Model could not explain the chemical behavior of atoms. c.Bohr used classical mechanics to understand the behaviors of small particles. d.The Bohr model is also known as the planetary, solar system, or satellite model.

56 Quantum Mechanical Model of the Atom A.Louie De Broglie (1924-5) 1.Took Einstein’s idea that light can exhibit both wave and particle properties 2.Very small particles (like electrons) display properties of waves. 3.Behavior of electrons in Bohr’s quantized orbits was similar to behavior of waves Known: any wave confined to a space can only have specific frequencies De Broglie suggested electrons are waves confined to the space around the atomic nucleus. Electrons could exist only at specific frequencies which correspond to specific energies (E = h quantized E of Bohr) French scientist

57 4.Experimentally proven in 1927 by diffraction of electrons by Davisson & Germer (showed diffraction of electrons by a crystal of Ni)diffraction of electrons B.Wave-Particle Duality of Nature a.Light and electrons (very small particles like electrons, atoms, molecules) have properties of waves and particles  QUANTUM MECHANICS (based on WAVE properties) **Large objects obey the laws of classical mechanics** Quantum Mechanical Model of the Atom

58 C. Werner Heisenberg: (1927) 1.Heisenberg’s Uncertainty Principle : states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. 2.You cannot predict future locations of particles. 3.He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.

59 D. Erwin Schrödinger Wave Equation (1926) 1.Wave nature of an electron is described by a mathematical equation. 2.Four quantum numbers in the equation are used to describe an electron’s behavior – location and energy. 3.Electron is treated as a wave with quantized energy. 4.Describes the probability of the electrons found in certain locations around the nucleus. (1887 – 1961) Austrian Physicist

60 Electron Density An orbital is a region in which an electron with a particular energy is likely to be found. Where the density of an electron cloud is high there is a high probability that is where the electron is located. If the electron density is low then there is a low probability.




64 E. Atomic Orbitals - region around the nucleus where an electron with a particular energy is likely to be found (not the same as Bohr’s orbits!) 1.Orbitals have characteristic shapes, sizes, & energies. 2. Orbitals do not describe how the electron moves. 3. The drawing of an orbital represents the 3-dimentional surface within which the electron is found 90% of the time.

65 4. Sublevels can have 4 different shapes s – orbital spherical

66 1s, 2s & 3s orbitals Superimposed on one another

67 Electron-Cloud Models

68 p-orbital – dumbbell shaped

69 p-orbital - dumbbell shaped

70 d-orbital - double dumbbell or fan blades

71 s,p and d orbitals For a more complete representation and presentation of atomic orbitals go to s orbital p orbitals d orbitals

72 Models of d-orbitals

73 f-orbital – more complex!

74 f orbitals

75 f – orbitals (3D)

76 Quantum Numbers Each quantum number provides more specific information on the probable location of an electron. Each electron within an atom can be described by a unique set of 4 quantum numbers.

77 Quantum Numbers - Finding an address for each electron: 1.“state” Principle Quantum Number ( n ) or the energy level ; a.Describes the relative size of the electron cloud. b.Positive integer values (n = 1 to n = 7) 2.“city” Sublevel ( l ) a.Describes the shape of the electron cloud. b.The maximum number of sublevels within a level = n c.Shapes are s, p, d,or f. d.Lowest energy = s Highest energy = f

78 Quantum numbers cont. 3.“street” Orbital (m l ) odd # of orbitals 1.Describes the orientation or direction in space a)s – 1 orbital b)p – 3 orbitals (x, y, z) c)d – 5 orbitals (xy, yz, xz, x 2 – y 2, z 2 ) d)f – 7 orbitals (y 3 – 3yx 2, 5yz 2 -yr 2, x 3 -3xy 2, zx 2 -zy 2, xyz, 5xz 2 -3xr 2, 5z 3 -3zr 2 ) 2.Orbitals within the same sublevel have the same energy are called degenerate orbitals 3.An orbital can hold a maximum of 2 electrons

79 4.“house” Spin (m s ) 1.Describes the direction of electron spin in an orbital. 2.The clockwise or counterclockwise motion of electrons. 3.Only electrons with opposite spins can occupy the same orbital. 4.The opposite spin is written as +1/2 or -1/2 or  or Quantum numbers cont.

80 E. Electron Configurations: 1.Shorthand notation for indicating the number of electrons in each level, sublevel, and orbital. 1s 2 2.Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.

81 Electron Configuration Rules 1.The Aufbau Principle: electrons are added one at a time to the lowest energy orbital available.

82 Pauli Exclusion Principle: 1. Each orbital can only hold 2 electrons. 2. The electrons must have opposite spins. s-sublevel = max 2 electrons p-sublevel = max 6 electrons d-sublevel = max 10 electrons f-sublevel = max 14 electrons incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓

83 Hund’s Rule: Electrons will remain unpaired in degenerate orbitals before they pair up. incorrect ↑↓ ↑ __ correct ↑ ↑ ↑


85 Electron Blocks on the Periodic Table


87 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin. Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel) Aufbau Principle: Electrons fill lowest energy levels first (n=1)

88 Electron Configuration Examples: Ex) electron configuration for Na: 1s 2 2s 2 2p 6 3s 1 Ex) orbital filling box diagram for Na:

89 Electron Dot Diagrams: Write the symbol for the element. Place dots around the symbol to represent the s & p valence s & p electrons only. Do NOT include d & f orbitals in diagram. p orbital electronss orbital electrons

90 Electron Configuration Orbital Box Diagram Electron-dot Diagram 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 4 1s 2 2s 2 2p 4 What does the Tellurium electron-dot resemble???

91 Mark your Periodic Tables

92 Unpaired vs. Paired Electrons Filled and Half-filled orbitals Atoms with unpaired electrons are said to be paramagnetic. These are weakly attracted to a magnetic field. Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field. ½ filled and filled orbitals have special stability

93 Noble Gas or Shorthand Electron Configurations Rb Se At Draw the Dot Diagrams for these elements

94 Exceptions to the Rules Max stability - ½ filled and filled orbitals – Cr – Mo – Cu – Ag – Au

95 Exceptions to the Rules Max stability - ½ filled and filled orbitals – Cr – Mo – Cu – Ag – Au

96 Electron Configuration for Ions K P Al Se K +1 P -3 Al +3 Se -2

97 Excited vs. Ground State If an electron absorbs energy, it is in an EXCITED state Ne: 1s 2 2s 2 2p 5 3s 1 How is this different from the ground state configuration? Ne: 1s 2 2s 2 2p 6

98 C. Electron Configuration & Families 1.Valence electrons – outermost electrons (s and p); responsible for bonding and chemical behavior. 2.Elements in the same group have the same number of valence electrons. Carbon has 4 valence electrons

99 Electron Configurations s, p, d, f blocks

100 Electron Configurations 8 electrons Stable Octet : 8 electrons in the outer level is very stable (includes He) Ions – gain/lose electrons to achieve a stable octet Isoelectronic – same electron configuration Examples: N, O, F, Na, Mg, Al are isoelectronic with Ne – this is called an isoelectronic series Pseudoisoelectronic – same electron configuration but includes the d orbitals Fe +2 is pseudoisoelectronic with Ar

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