The Periodic Table Chapters 6 & 7.1

1. 1 The Periodic Table Chapters 6 & 7.1

2. 2 Periodic Table Goals • SC4. Students will use the organization of the Periodic Table to predict properties of elements. • Use the Periodic Table to predict periodic trends including atomic radii, ionic radii, ionization energy, and electronegativity of various elements. • Compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table

3. 3 Periodic Table Organizing the Elements

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5. 5 • Grouped elements on the basis of similar chemical properties.  • Left blank spaces open to add new elements where he predicted they would occur.  • Accepted minor inversions when placing the elements in order of increasing atomic mass.  • Predicted properties for undiscovered elements. Dmitri IvanovitchMendeléev

6. Dmitri IvanovitchMendeléev 6 • Organized by increasing atomic mass Clip

7. 7 Modern Periodic Table • Now arranged by increasing atomic number • Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical properties

8. 8 Metals, Nonmetals, and Metalloids

9. Metals vs. Nonmetals • Lose their valence electrons easily. • Good electrical conductors and heat conductors. • High luster (sheen) • Malleable & Ductile • Solid at room temperature 9 • Gain or share valence electrons easily. • Poor conductors of heat and electricity. (Carbon is an exception) • Brittle - if a solid. • Nonductile. • Usually gas @ room temp (some are Solids, liquids @ room temp. ) Clip

10. 10 Metalloids • Properties similar to metals and non-metals…depending on the conditions -like mixing it. • Boron, silicon and germanium behave as semiconductors • Ex: Silicon is a poor conductor, but when mixed with boron is a good conductor. • They can be shiny or dull in appearance and brittle in nature. • They behave as non metals when they react with metals and behave as metals when they react with non metals. • They are ductile and can be bent into pipes. These seven metalloids:Boron (B)Silicon (Si)Germanium (Ge)Arsenic (As)Antimony (Sb)Tellurium (Te)Polonium (Po)

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12. 12 Trend

13. Groups or Families increasing atomic number periods 12 13

14. Groups/Families Have Similar Properties Examples of Physical Properties • Density • Boiling Point • Melting Point • Conductivity • Heat Capacity Examples of Chemical Properties - Valence - Reactivity - Radioactivity

15. Alkali Metals 14 13 Clip • Soft, silvery metals, solids @room temp • Soft enough to be cut with a dull knife • low melting and boiling temperatures • Very reactive metals that do not occur freely in nature. • Only one valence electron • Give up valence electron easily • Malleable, ductile, and are good conductors of heat and electricity. • Cesium and francium are the most reactive elements in this group. • Alkali metals can explode if they are exposed to water.

16. Alkaline Earth Metals 15 • Called “alkaline” because when these compounds are mixed in solutions, they form solutions with a pH greater than 7. • Two valence electrons • Give up valence electrons easily • Not as reactive as the alkali metals, but still highly reactive • Because of their reactivity, the alkaline metals are not found free in nature. • Mg & Ca are very important in animal and plant physiology. (Calcium-bones. Mg-chlorophyll) Clip radioactive

17. Transition Metals 16 • All solids at room temp. • Are both ductile and malleable, and conduct electricity and heat. • The interesting thing about transition metals is that their valence electrons change. • There are three noteworthy elements in the transition metals family. -iron, cobalt, and nickel, and they are the only elements known to produce a magnetic field. Clip

18. Rare Earth Metals • Composed of the lanthanide and actinide series. • The lanthanide series can be found naturally on Earth. • Only 1 element in the series is radioactive. • The actinide series is much different. They are all radioactive and some are not found in nature. • One element of the lanthanide series and most of the elements in the actinide series are calledtransuranium, which means synthetic or man-made. • All found in group 3 of the periodic table, and the 6th & 7th periods. • All have 3 valence e- & are solid at room temp. Youtube Clip Actinide Clip 17

19. Metalloids 18 • Metalloids have properties of both metals and non-metals. • Some of the metalloids, such as silicon and germanium, are semi-conductors. • This means that they can carry an electrical charge under special conditions. This property makes metalloids useful in computers and calculators • All solids at room temperature Other Metals Clip Metalloids Clip

20. 19 Halogens most reactive • “Halogen" means "salt-former" and compounds containing halogens are called "salts“ or “halides”. • All have 7 valence electrons. • Highly reactive. Often boning w/group IA (alkali metals). • exist, at room temperature, in all three states of matter: Solid- Iodine, Astatine Liquid- Bromine Gas- Fluorine, Chlorine Clip

21. Noble Gases 20 • All have 8 valence electrons • Chemically inert • All gases at room temp. • They all have very low boiling and melting points. • They all put out a color in the visible wavelengths when a low pressure of the gas is put into a tube and a high voltage current is run through the tube. • Neon-advertising signs. Argon-light bulbs. Helium-balloons. Xenon-headlights for new cars. • When you move down the periodic table, as the atomic numbers increase, the elements become rarer Clip Strange Clip

22. Review: Periodic Table

23. Some Reminders…. • Chapter 6 Vocab Quiz tomorrow • Project should be complete by tomorrow • You still have 2 video clips to watch by Thur • ChemThink: “Ions” should be complete by Thur. • Read Chapter 6 & Ch 7 (sec 1) by Thur • Test on Thur

24. 21 Electron Dot Diagrams Diagrams that show valence electrons as dots.

25. 22 Becoming stable…….Ions • In order to be come stable, atoms gain, lose or share electrons(we will discuss this more later). • When they gain or lose electrons, they are then called IONS. • Why gain or lose? • To fill their valence shell. Ions Clip

26. 23 Cation Anion

27. 24 24 ION = Charged atom The Octet Rule • In forming compounds atoms tend to achieve the electron configuration of a noble gas. • An octet is a set of 8. • Can you think of any exceptions?

28. 25 Ions and Energy Ionization Energy: the energy required to remove an electron

29. Ion Notation Charge is determined by the number of electrons gained or lost. Examples:

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31. 27 Periodic Trends

32. 28 Atomic Size(Atomic Radius) • Atomic Radius: ½ the distance between the nuclei of two atoms of the same element when the atoms are joined. • the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons.

33. Atomic Radius TREND 29 The greater the number of protons present, the stronger the attraction that holds the electrons closer to the nucleus, and the smaller the size of the shells.

34. 30 Atomic Radius TREND Answer the question

35. 31 Ionization Energy Removing an electron from an atom requires enough energy to overcome the magnetic pull of the positive charge of the nucleus. Ionization energy (I.E. or I) is the energy required to completely remove an electron from a gaseous atom or ion. The Ionization Energy is always positive.

36. 32 TREND

37. 33 Factors Governing Ionization Energy • Size of the atom/ distance of the valence electron to the nucleus • the nuclear charge • Screening effect of the inner electrons • Screening effect • outermost electrons are shielded or “screened” from the nucleus by the inner electrons. • The outer most electrons do not feel the complete charge of the nucleus. • When there are more inner electrons • the screening effect will be large, the nuclear attraction will be less. • Thus when the inner electrons increase the ionization energy will decrease.

38. 34 Ionic Size • Cations are always smaller than the atom from which they form • Anions are always larger than the atoms from which they form

39. 35 Electronegativity • Electronegativity: The ability of an atom to attract electrons when the atom is in a compound. (electron attraction) • Electronegativity is basically the property of an element that helps in attracting a pair of electrons or a single electron towards it.

40. 36 • Ionization energy is related to electronegativity as low ionization electrons exhibit low electronegativity. This is because their nuclei does not have a strong attractive force on electrons. When you refer to a decreasing electronegativity chart, you will find, electronegativity decreases as atomic number increases. This is because of the distance between valance electrons and nucleus

41. 37 Mr. Andersen’s Review of Periodic Trends

43. I I Br Br • Where are these found on the Periodic table? • Never found in nature alone. 38

44. 39 Periodic Table Review