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Parts of periodic table

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Periodic Table of Elements
Periodic Table of Elements
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Parts of periodic table

  1. 1. Periodic Table Geography
  2. 2. The horizontal rows of the periodic table are called PERIODS.
  3. 3. The elements in any group of the periodic table have similar physical and chemical properties! The elements in a group have the same number of electrons in their outer orbital. The vertical columns of the periodic table are called GROUPS, or FAMILIES.
  4. 4. Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties. When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.
  5. 5. The Periodic Table
  6. 6. Examples of Families -Alkali Metals - Alkaline Earth Metals - Transition Metals - Halogen Gases - Noble Gases Examples of Physical Properties -Density - Boiling Point - Melting Point - Conductivity - Heat Capacity Examples of Chemical Properties - Valence - Reactivity - Radioactivity
  7. 7. HaLiNa Korina, Ruby, Ces, Francisco They are very reactive. Why? They all have one electrons in their outer shell. That's one electron away from being happy (full shells). Alkali metals are not the type of metals you would use for coins or houses. Group I: Alkali Metals
  8. 8. Group 1- Alkali Metals Lithium Sodium Potassium Rubidium Cesium Francium
  9. 9. Group 1- Alkali Metals • 1 valence electron (ns1) • Form a 1+ ion. • Note: Hydrogen, a nonmetal, is located in the first column because it has one valence electron.
  10. 10. ALKALI METALS •very reactive metals that do not occur freely in nature •malleable, ductile, good conductors of heat and electricity. •can explode if they are exposed to water
  11. 11. This is the second most reactive family of elements in the periodic table. Each of them has two electrons in their outer shell. They are ready to give up those two electrons in electrovalent/ionic bonds. Group II : Alkaline Earth Metals
  12. 12. Group 2- Alkali Earth Metals Beryllium Magnesium Calcium Strontium Barium Radium
  13. 13. Group 2- Alkali Earth Metals • Have 2 valence electrons. • Form 2+ ions.
  14. 14. ALKALINE EARTH METALS •metals •very reactive •not found free in nature From
  15. 15. Transition metals are able to put more than eight electrons in the shell that is one in from the outermost shell. Transition Metals
  16. 16. •ductile and malleable, and conduct electricity and heat •iron, cobalt, and nickel, are the only elements known to produce a magnetic field.
  17. 17. These elements are also called the rare-earth elements. Inner Transition Metals
  18. 18. RARE EARTH ELEMENTS •many are man-made From
  19. 19. HALOGENS •"halogen" means "salt-former" and compounds containing halogens are called "salts" •exist in all three states of matter
  20. 20. (Group XVII). Halogens What Makes Them Similar? They have seven electrons in their outer shell. Because they are so close to being happy, they have the trait of combining with many different elements.
  21. 21. This group is composed of stable gases otherwise known as the nonreactive or inert elements. They have 8 electrons in the sub shell. Noble Gases
  22. 22. The s and p block elements are called REPRESENTATIVE ELEMENTS.
  23. 23. The periodic table is the most important tool in the chemist’s toolbox!
  24. 24. Other Types of Grouping • Metals • Nonmetals • Metalloids or Semimetals.
  25. 25. Metals, Nonmetals, Metalloids
  26. 26. 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.
  27. 27. 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.
  28. 28. Metals, Nonmetals, Metalloids • 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?
  29. 29. OTHER METALS •are ductile and malleable •are solid, have a high density
  30. 30. 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?
  31. 31. NON-METALS •not able to conduct electricity or heat very well •very brittle •Do not reflect light. From
  32. 32. Nonmetals • Nonmetals are the opposite. • They are dull, brittle, nonconductors (insulators). • Some are solid, but many are gases, and Bromine is a liquid.
  33. 33. METALLOIDS •have properties of both metals and nonmetals •some of the metalloids are semi-conductors. This means that they can carry an electrical charge under special conditions. This property makes metalloids useful in computers and calculators
  34. 34. Metalloids • Metalloids, aka semimetals are just that. • They have characteristics of both metals and nonmetals. • They are shiny but brittle. • And they are semiconductors.
  35. 35. Periodic Trends • There are several important atomic characteristics that show predictable trends that you should know. • 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.
  36. 36. 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.
  37. 37. Covalent Radius • Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å. 2.86 Å 1.43 Å 1.43 Å
  38. 38. 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.
  39. 39. 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.
  40. 40. 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.
  41. 41. 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.
  42. 42. Atomic Radius • The overall trend in atomic radius looks like this.
  43. 43. Atomic Radius • Here is an animation to explain the trend. • On your help sheet, draw arrows like this:
  44. 44. Shielding • As more PELs are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. • The effective nuclear charge (enc) on those outer electrons is less, and so the outer electrons are less tightly held.
  45. 45. Ionization Energy • This is the second important periodic trend. • If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely. • The atom has been “ionized” or charged. • The number of protons and electrons is no longer equal.
  46. 46. Ionization Energy • The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ) • The larger the atom is, the easier its electrons are to remove. • Ionization energy and atomic radius are inversely proportional. • Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.
  47. 47. Ionization Energy
  48. 48. Ionization Energy (Potential) • Draw arrows on your help sheet like this:
  49. 49. Electron Affinity • What does the word ‘affinity’ mean? • Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). • Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.
  50. 50. Electron Affinity • Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy. • If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic. • This is true for the alkaline earth metals and the noble gases.
  51. 51. Electron Affinity • Your help sheet should look like this: + +
  52. 52. Metallic Character • This is simple a relative measure of how easily atoms lose or give up electrons. • Your help sheet should look like this:
  53. 53. Electronegativity • Electronegativity is a measure of an atom’s attraction for another atom’s electrons. • It is an arbitrary scale that ranges from 0 to 4. • The units of electronegativity are Paulings. • Generally, metals are electron givers and have low electronegativities. • Nonmetals are are electron takers and have high electronegativities. • What about the noble gases?
  54. 54. Electronegativity • Your help sheet should look like this: 0
  55. 55. Overall Reactivity • This ties all the previous trends together in one package. • However, we must treat metals and nonmetals separately. • The most reactive metals are the largest since they are the best electron givers. • The most reactive nonmetals are the smallest ones, the best electron takers.
  56. 56. Overall Reactivity • Your help sheet will look like this: 0
  57. 57. The Octet Rule • The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level. • They may accomplish this by either giving electrons away or taking them. • Metals generally give electrons, nonmetals take them from other atoms. • Atoms that have gained or lost electrons are called ions.
  58. 58. Ions • When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion. • In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons. • They become positively charged cations.
  59. 59. Ions • Here is a simple way to remember which is the cation and which the anion: + + This is Ann Ion. This is a cat-ion. She’s unhappy and negative. He’s a “plussy” cat!
  60. 60. Ionic Radius • Cations are always smaller than the original atom. • The entire outer PEL is removed during ionization. • Conversely, anions are always larger than the original atom. • Electrons are added to the outer PEL.
  61. 61. Cation Formation Effective nuclear charge on remaining electrons increases. Na atom 1 valence electron 11p+ Valence elost in ion formation Result: a smaller sodium cation, Na+ Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
  62. 62. Anion Formation Chlorine atom with 7 valence e17p+ One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.