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Periodic Trends Complete discussion in science 8.ppt
- 1. Periodic Trends Elemental Propertiesand Patterns
- 2. 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.
- 3. The Periodic Law •Mendeleev not only predicted the existence of 2 (at the time) undiscovered elements, but he predicted the properties too, and he was right! • He was very accurate in his predictions, which led the world to accept his ideas about periodicity and a logical periodic table.
- 4. The Periodic Law •Mendeleev ordered a periodic table according to atomic mass, which makes sense (look at a few in a row)
- 5. The Periodic Law •However this does not work for a 100% of the elements
- 6. • In 1913,Henry Moseley decided to arrange the periodic table by the number of protons, a.k.a The atomic number • This is the version that we use today The Periodic Law
- 7. The Periodic Law •Mendeleev and Moseley understood the ‘Periodic Law’ which states: • When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.
- 8. The Periodic Law •“Which elements act alike?” Lab • Review • Precipitate • Cation • Family names
- 9. The Periodic Law •“Which elements act alike?” Lab • What’s new? • Soluble (s) vs insoluble (i) • Polyatomic ions • Poly(many)atomic(atoms) ions(charged molecules) • Phosphate, Carbonate, Nitrate
- 10. The Periodic Law “Whichelements act alike?” Lab •Things to focus on • The Cations • Where are they on the periodic table? – Are they reacting similarly in the same group/family or the same period (row)?
- 11. The Periodic Law •“Which elements act alike?” Lab • Goggles on, Closed toed shoes on and… GO FORTH AND SCIENCE!
- 12. The Periodic Law •Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. • They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.
- 13. Valence Electrons • Doyou 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.
- 14. A Different Typeof Grouping • Besides the 4 blocks (s,p,d,f) of the table, there is another way of classifying element: • Metals • Nonmetals • Metalloids or Semi-metals.
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- 16. 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.
- 17. 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 a metal
- 18. 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 is a metalloid?
- 19. Metals • Metals arelustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. • They are mostly solids at room temp. • What is one exception?
- 20. Nonmetals • Nonmetals arethe opposite. • They are dull, brittle, nonconductors (insulators). • Some are solid, but many are gases, and Bromine is a liquid.
- 21. Metalloids • Metalloids, akasemi-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?
- 22. Periodic Trends • Thereare 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.
- 23. Atomic Radius • Sincea 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.
- 24. Covalent Radius • TwoBr 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
- 25. Atomic Radius • Thetrend 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 energy level to the electron cloud, making the atoms larger with each step.
- 26. Atomic Radius • Thetrend 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 sublevels. (They are not getting bigger)
- 27. Atomic Radius • Theeffect 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.
- 28. Effective Nuclear Charge •We call this effect… • 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.
- 29. Atomic Radius • Hereis an animation to explain the trend. • It might help to draw arrows like this:
- 30. Practice… 1. Explain whya magnesium atom is smaller than atoms of both sodium and calcium 2. 2. Predict the size of the astatine (At) atom compared to Predict the size of the astatine (At) atom compared to that of tellurium (Te). Explain your prediction. that of tellurium (Te). Explain your prediction. 3. Which effect on atomic size is more significant, an increase in nuclear charge across a period or an increase in occupied energy levels within a group? Explain.
- 31. Ionization Energy • Thisis the second important periodic trend. • If an electron is given enough energy 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.
- 32. Ionization Energy • Theenergy 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.
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- 35. Ionization Energy • Ifan atom as a low ionization energy, will it be easy or hard to remove an electron? • Think about the energy required to remove it, is it high or low? Do you have to put a large amount of energy into it or a small amount? • What if the atom has a high ionization energy?
- 36. Practice… 1. Compare thefirst ionization energy of sodium to that of potassium. 2. Compare the first ionization energy lithium to that of beryllium.
- 37. Metallic Character • Thisis simple a relative measure of how easily atoms lose or give up electrons. • Your help sheet should look like this: Which element is most metallic?
- 38. Electronegativity • Electronegativity isa measure of an atom’s attraction for another atom’s electrons. • It is an arbitrary scale that ranges from 0 to 4. • Generally, metals are electron donators and have low electronegativities. • Nonmetals are electron takers and have high electronegativities. • What about the noble gases? • Do they need any any electrons?
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- 40. Electronegativity • What isthe most electronegative atom? • Why? • Fluorine • It has a small atomic radius, and a high (relative) effective nuclear charge
- 41. Practice… 1. Is theelectronegativity of barium larger or smaller than that of strontium? Explain. 2. Arrange oxygen, fluorine, and sulfur in order of increasing electronegativity.
- 42. Overall Reactivity • Thisties all the previous trends together in one package. • However, we must treat metals and nonmetals separately, as one will gain electrons and one will lose electrons. • 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.
- 43. Overall Reactivity • Thinkcorners 0
- 44. Overall Reactivity • Takeout a piece of scrap paper or just find a spot in your notes… • We are going to watch a video that will react sodium, potassium, rubidium, and cesium with water. • Prediction: Order the 4 elements from least reactive to most reactive. • Justify your reasoning!
- 45. Alkali metal video •Extra link (slow motion): http://www.youtube.com/watch?v=iP6CRZdDu6o
- 46. Ionic Radius • Cationsare 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.
- 47. Cation Formation 11p+ Na atom 1valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na+
- 48. Anion Formation 17p+ Chlorine atom with7 valence e- 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.
- 49. Practice… 1. Explain whythe sulfide ion (S-2 ) is larger than the chloride ion (Cl- ). 2. Would you expect a S-2 ion to be larger or smaller than an K+ ion? Explain.
- 50. Apply… • Would youexpect a Cl- ion to be larger or smaller than an Mg+2 ion? Explain.
- 51. Miscellaneous • HONClBrIF’s/H7’s (diatomics) •Polyatomic Ions: • Phosphate • Carbonate • Nitrate