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Periodic Trends

Periodic Trends. Elemental Properties and Patterns. Open your Packet to page 10. Organizing Information Use Table 1 to make your cards (1set per group) You need 7 blue that lose e- 2 pink either gain or lose 3 green neither gain nor lose 6 yellow gain e-. 6 yellow gain e-.

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Periodic Trends

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  1. Periodic Trends Elemental Properties and Patterns

  2. Open your Packet to page 10 • Organizing Information • Use Table 1 to make your cards (1set per group) • You need • 7 blue that lose e- • 2 pink either gain or lose • 3 green neither gain nor lose • 6 yellow gain e- • 6 yellow gain e-

  3. Follow directions and Arrange Elements • Do not answer questions on page 10 • After # 12 and #13 Answer questions on PAGE 11

  4. Graphing Ionization Energy and Radius • There will be 2 graphs • Divide your paper into half (top and bottom) • You will graph both sets of Data • You will answer ALL questions.

  5. Homework 5.1 • Finish classwork • Read CH 5.1 for ½ sheet • Text page 137 (1-5) • Naming and Balancing Charges Paper

  6. Ch 5.1 Periodic Law Objectives: Explain Mendeleev and Moseley Describe the Periodic Table Explain how elements belong in Groups

  7. The Periodic Law • Dimitri Mendeleev was the first scientist to publish an organized periodic table of the known elements. • He was perpetually in trouble with the Russian government and the Russian Orthodox Church, but he was brilliant never-the-less.

  8. Mendeleev • Mendeleev saw that when atoms are arranged by atomic mass… • Similarities in chemical properties appeared at regular intervals. • These repeating patterns are called periodic trends. • First Periodic Table published in 1869

  9. The Periodic Law • Mendeleev even went out on a limb and predicted the properties of 2 at the time undiscovered elements. • He was very accurate in his predictions, which led the world to accept his ideas about periodicity and a logical periodic table.

  10. Moseley 1911 • Arranged the table by Atomic Number • Periodic Law= physical and chemical properties of the elements are periodic. • Or… when elements are arranged by their atomic number, similar properties appear regularly. (columns or Groups or families)

  11. The Periodic Law • ‘Periodic Law’ states: • When arranged by increasing atomicnumber, the chemical elements display a regular and repeating pattern of chemical and physical properties.

  12. The Periodic Law • Atoms with similar properties appear in 1groups or families (vertical columns) on the periodic table. • They are similar because they all have the same number of 2valence electrons(or Number of outer shell electrons), which governs their chemical behavior.

  13. Periodicity • Each Group of elements have the same number of protons added down each column. • This makes the properties similar, due to the arrangement of electrons around the nucleus.

  14. Atomic numbers change Predictably

  15. Valence Electrons • Do you remember how to tell the number of valence electrons for elements in the s- and p-blocks? • How many valence electrons will the atoms in the d-block (transition metals) and the f-block (inner transition metals) have? • Most have 2 valence e-, some only have 1.

  16. Bohr Diagrams showing Valence Electrons

  17. Ch 5.2 Electron Configuration and the Periodic Table • Objectives • Describe lengths of e- sublevels • Locate the 4 blocks of s,p,d,f • Compare configurations and group numbers • Know names of groups on table

  18. Ch 5.2 Electron Configuration and the Periodic Table • 7 periods on the chart • Horizontally organized • Length of the period is determined by the number of electrons that can occupy the sublevels being filled. (note: The number of orbitals that exist in a shell, or main energy level = n2 )

  19. Look for this in your book

  20. Blocks of the Periodic Table

  21. A Different Type of Grouping • Besides the 4 blocks of the table, there is another way of classifying element: • Metals • Nonmetals • Metalloids or Semi-metals. • The following slide shows where each group is found.

  22. Metals, Nonmetals, Metalloids

  23. Metals, Nonmetals, Metalloids • There is a zig-zag or staircase line that divides the table. • Metals are on the left of the line, in blue. • Nonmetals are on the right of the line, in orange.

  24. Metals, Nonmetals, Metalloids • Elements that border the stair case, shown in purple are the metalloids or semi-metals. • There is one important exception. • Aluminum is more metallic than not.

  25. Metals, Nonmetals, Metalloids Consider these questions….. • How can you identify a metal? • What are its properties? • What about the less common nonmetals? • What are their properties? • And what the heck is a metalloid?

  26. Metals • Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. • They are mostly solids at room temp. • What is one exception?

  27. Nonmetals • Nonmetals are the opposite. • They are dull, brittle, nonconductors (insulators). • Some are solid, but many are gases, and Bromine is a liquid.

  28. Metalloids • Metalloids, aka semi-metals are just that. • They have characteristics of both metals and nonmetals. • They are shiny but brittle. • And they are semiconductors. • What is our most important semiconductor?

  29. The s-Block ElementsGroup 1 • Known as Alkali Metals • Silvery in appearance • Most reactive of all groups and are NOT found as free elements in Nature • Extremely reactive metals are Group 1 • Outermost energy level contains 1e-

  30. The s-Block ElementsGroup 2 • Known as Alkaline-Earth metals • Less reactive than group 1 • Outermost energy level contains 2e- • Harder, denser, stronger than Group 1 • Still too reactive to exist as free elements • Have higher melting points than Group1

  31. The d-Block elements: Groups 3-12 • Known as the Transition Elements • Metals with traditional metal properties • Good conductors of heat and electricity • High luster (shiny) • Typically un-reactive, some do not form compounds easily, exist freely in nature • Platinum, gold and palladium are some of the least reactive of all elements

  32. The p-Block Elements: Groups 13-18 • Known as the Main Group Elements, (together with the s-block elements) • P-block includes -All non-metals, except Hydrogen and Helium -All six metalloids

  33. Group 17Halogens • The most reactive Non-metals • b/c they have 7 electrons in outer shell • One short of being stable with a full shell • React with metals to produce salts.

  34. The f-block elements:Lanthanides and Actinides • Wedged between Groups 3 and 4 in the 6th and 7th periods • Here b/c they fill the 4f sublevel for Lanthanides and the 5f sublevel for Actinides • 14 f-block elements in each series • Lanthanides as reactive as Group 2 • Actinides all radioactive and most manmade (First 4 are natural, Th>Np)

  35. Concept Check • How can you identify a metal? • What are its properties? • What about the less common nonmetals? • What are their properties?

  36. Homework 5.2 • Finish classwork • Bohr model practice • Read 5.2 • Text page 166 (1-9, 11-16) Test Thursday! Nov. 18th

  37. Ch 5.3 Electron Configuration and the Periodic Properties OR Trends in the Periodic Table Test Friday! Nov. 19th

  38. Objectives • Define atomic radii, ionization energy, electronegativity and electron affinity. • Compare trends • Define valence e- (in each main group) Open your books to page 150

  39. Trends in the Periodic Table • The arrangement of the periodic table reveals trends in the properties of elements. • A trend is a predictable change in a particular direction. There are several important atomic characteristics that show predictable trends that you should know.

  40. Trends in the Periodic Table • Understanding a trend among the elements helps us make predictions about their chemical behavior. • Periodic trends can be explained by the changes in atomic structure as we move across a period and down a group.

  41. Periodic Trends • The first and most important is atomic radius. • Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.

  42. Trends in the Periodic Table 1. Atomic Radius • Atomic radius is half the distance between the nuclei of identical atoms bonded together.

  43. *Increases down a Group and *Decreases across a period

  44. Atomic Radius • Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms. • Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10-10 m.

  45. 2.86Å 1.43 Å 1.43 Å Covalent Radius • Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å.

  46. Atomic Radius • The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family. • Why? • With each step down the family, we add an entirely new PEL to the electron cloud, making the atoms larger with each step.

  47. Atomic Radius • The trend across a horizontal period is less obvious. • What happens to atomic structure as we step from left to right? • Each step adds a proton and an electron (and 1 or 2 neutrons). • Electrons are added to existing PELs or sublevels.

  48. Atomic Radius • The effect is that the more positive nucleus has a greater pull on the electron cloud. • The nucleus is more positive and the electron cloud is more negative. • The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

  49. Effective Nuclear Charge • What keeps electrons from simply flying off into space? • Effective nuclear charge is the pull that an electron “feels” from the nucleus. • The closer an electron is to the nucleus, the more pull it feels. • As effective nuclear charge increases, the electron cloud is pulled in tighter.

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