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Introducing the Elements

Introducing the Elements. The Element Song. 1869: Dmitri Mendeleev. Russian chemist Arranged elements in tabular form so that elements with similar properties were in the same column When listed in order by mass, elements generally repeat properties in groups of 8 (Law of Octaves).

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Introducing the Elements

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  1. Introducing the Elements The Element Song

  2. 1869: Dmitri Mendeleev • Russian chemist • Arranged elements in tabular form so that elements with similar properties were in the same column • When listed in order by mass, elements generally repeat properties in groups of 8 (Law of Octaves)

  3. The First Periodic Table • Most tables at the time listed elements by mass • Mendeleev also arranged elements by mass, but left several “holes” in his table and occasionally reversed the order of elements to fit the properties of others in that column • The “holes” were later filled in with newly discovered elements that had the properties predicted by Mendeleev’s table. • The reason for the reversal of elements was explained later by Henry Moseley, who noted that the elements were in order by atomic number (number of protons) rather than by mass

  4. Hydrogen • Most abundant element in the universe Why? • Makes up most the mass of stars • Can be H+ (hydrogen ion) or H- (hydride ion) • Used in Fuel Cells: How Stuff Works • In Fusion, H is converted to He

  5. Alkali Metals: Li, Na, K, Rb, Cs, Fr • Most reactive of the metals, +1 ions • Stored under kerosene or mineral oil • Na and K most important • Na2CO3 and NaHCO3 two important compounds • K is an important plant nutrient (macronutrient) • Fertilizers: N-P-K

  6. Total Molar Composition of Seawater (Salinity = 35)[7]

  7. Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra • Harder, more dense, and less reactive than alkali metals • Ca, Sr, and Ba most alike • Hard Water: Ca2+ and Mg2+ ions • Epsom salt: MgSO4 Boiler Scale

  8. …..more on the alkaline earths • CaCO3 is limestone becomes marble • Limestone is most abundant rock in the earth’s crust • CaO “Lime” or “quicklime” • CaSO4 “Plaster of Paris” (building material) Plaster of Paris footprints

  9. Aluminum Group: B, Al, Ga, In, Tl • Aluminum by far the most important • Third most abundant element in the earth’s crust • Important metal: abundant, light weight, strong • Al2O3 coating prevents corrosion

  10. Carbon Group: C, Si, Ge, Sn, Pb • Very diverse group of elements • C is the basis for organic compounds • CO2 and CO3-2 inorganic carbon • CO2 one of the earliest gases in the atmosphere • Carbon cycling one of the most important • Two basic parts: (1) photosynthesis (2) respiration

  11. Disrupting the carbon cycle • CO2 is a greenhouse gas (GWP=1) • Increasing concentration by: • Burning fossil fuels • Removing vegetation • Preindustrial 1800: 280 ppm • 1959: 316 ppm • 2010: 388 ppm • 2011: 391 ppm

  12. Growing CO2 Warms the Earth • Greenhouse Effect is essential for Life! • Earth’s radiative balance (solar input vs. IR output) leaves <TEarth> ~ – 20°C • Almost all water would be ice everywhere. • But Life requires ℓiquid water! • H2O(g) and CO2 absorb outbound IR and reradiate it omnidirectionally. • So Earth intercepts ~½ that absorbed IR and gains <T> to +15°C.  H2O(ℓ) & we exist.

  13. Venus, the Runaway Greenhouse • Being closer to the sun, Venus intercepts twice the solar flux of Earth. • But it is twice as reflective (albedo), so its temperature would be ~ –29°C. • But it’s surface T averages +435°C! • 90 atm at the surface, mostly CO2

  14. Keeling Curve

  15. Silicon - Si • Second most abundant element in the Earth’s crust • Found in clay, sand, sandstone, silica rock, quartz, other minerals • Many different bonding combinations • Is a semiconductor (Silicon Valley)

  16. Tin (Sn) and Lead (Pb) • Many Industrial Uses • Pb is a “heavy metal” and is toxic to many organs in the human body • Impedes the development of the nervous system • Taken out of gasoline in the late 1970’s and removed from most paints

  17. Ozone • Ozone absorbs much of the radiation between 240 and 310 nm. • It forms from reaction of molecular oxygen with the oxygen atoms produced in the upper atmosphere by photodissociation (< 242 nm). O + O2 O3

  18. Ozone Depletion In 1974 Sherwood Rowland and Mario Molina (Nobel Prize, 1995) discovered that chlorine from chlorofluorocarbons (CFCs) may be depleting the supply of ozone in the upper atmosphere.

  19. Chlorofluorocarbons CFCs were used for years as aerosol propellants and refrigerants. Mostly = CFCl3, CF2Cl2. They are not water soluble (so they do not get washed out of the atmosphere by rain) and are quite unreactive (so they are not degraded naturally).

  20. Chlorofluorocarbons • The C—Cl bond is easily broken, though, when the molecule absorbs radiation with a wavelength between 190 and 225 nm. • The chlorine atoms formed react with ozone: Cl + O3 ClO + O2

  21. Chlorofluorocarbons In spite of the fact that the use of CFCs in now banned in over 100 countries, ozone depletion will continue for some time because of the tremendously unreactive nature of CFCs.

  22. Sulfur • Sulfur dioxide is a by-product of the burning of coal or oil. • It reacts with moisture in the air to form sulfuric acid. • It is primarily responsible for acid rain.

  23. Sulfur • High acidity in rainfall causes corrosion in building materials. • Marble and limestone (calcium carbonate) react with the acid; structures made from them, erode.

  24. Sulfur • SO2 can be removed from flu gases by injecting powdered limestone which is converted to calcium oxide. • The CaO reacts with SO2 to form a precipitate of calcium sulfite. This process = “scrubbing”

  25. Carbon Monoxide • Carbon monoxide binds preferentially to the iron in red blood cells. • Exposure to CO can lower O2 levels to the point of causing loss of consciousness and death.

  26. Carbon Monoxide • Products that can produce carbon monoxide must contain warning labels. • Carbon monoxide is colorless and odorless, so detectors are a good idea.

  27. Nitrogen Oxides • What we recognize as smog, that brownish gas that hangs above large cities like Los Angeles, is primarily nitrogen dioxide, NO2. • It forms from the oxidation of nitric oxide, NO, a component of car exhaust.

  28. Photochemical Smog Nitrogen oxides react with water to form nitric acid, contributing to acid rain. Smog also contains ozone, carbon monoxide, hydrocarbons, and particles.

  29. Bonding: Influences • Valence Electrons • Nuclear Charge • Atomic Size/Radius • Distance between attractions • Screening or Shielding Effect

  30. Valence Electrons

  31. Core Configurations • Why? – Shows/focuses on the valence electrons. • Write the configuration for arsenic. • 1s22s22p63s23p64s23d104p3 • or • [Ar] 4s23d104p3 • How many valence electrons?

  32. Atomic Radius: How is it measured? • Half the distance between nuclei of two covalently bonded atoms of the same element. • Why not just measure from the nucleus to the outer edge of the atom?

  33. Atomic Radius

  34. Radius trends • Group trend? Radius increases down a group • Why? Adding new energy levels • Period trend? Radius decreases across a period • Why? Increasing nuclear charge has the effect of pulling electron cloud closer.

  35. Ions and their formation • Cations • Formed by the loss of electrons • Positively charged • Usually formed from metals • Are always smaller than the atom they are formed from • Anions • Formed by the gain of electrons • Negatively charged • Usually formed from nonmetallic elements • Are always larger than the atom they are formed from

  36. Ionization Energy • The energy required to remove an electron from an isolated, neutral, gaseous atom. • First ionization energy – energy required to remove a first electron from an atom. • Second ionization energy – energy required to remove a second electron from an atom. • Third ionization energy - ???? • Etc . . . .

  37. First Ionization Energies

  38. First Ionization energies • Group trend – IE1 decreases down a group. Why? • Valence electrons are further from the nucleus and the shielding effect is greater down a group. • Shielding effect – occurs when core electrons “shield” or interferes with the attraction that the nucleus has for the valence electrons.

  39. . . . IE cont. . . . • Period trend – IE1 is larger as you move across a period, left to right. Why? • Atoms are smaller so valence electrons are closer to the nucleus and . . . . . • . . . the nuclear charge is greater with no change in shielding effect (electrons are going in the same energy level)

  40. Ionization Energy Increasing Trend Periodic Table

  41. Successive Ionization Energies Where do the largest jumps occur for each Element and why do you think this happens?

  42. Electronegativity • A measure of the ability of an atom to attract electrons to itself when bonded to other atoms. • Trends in electronegativity are the same as ionization energy and the reasons why are essentially the same too.

  43. Electronegativities

  44. Electron Affinity • The amount of energy released or gained when an atom receives an electron. • When this happens a negatively charged ion, called an anion, forms.

  45. Electron Affinity Notice what groups have (-) negative affinities and what groups have (+) positive affinities (negative values are shown here as above zero.) A negative affinity means energy is released and a positive affinity means that much energy is gained when an atom acquires and electron.

  46. Place these elements in order of increasing: Ge, P, N, and Si (a) atomic radius (b) first ionization energy (c) electronegativity 2. Write the core configuration for the following elements: S, Ca, Sn • How many valence electrons does each element in #2 have?

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