1. Unit 5: Periodicity

2. History of the Periodic Table  Dimitri Mendeleev (1836-1907) –Russian chemistry professor  Noticed that as he was arranging elements according to their properties that they generally were organized in order of increasing atomic mass.

3. Genius of Mendeleev’s Work  Left spaces for elements not yet discovered.  He predicted that some still-unknown elements must exist to fit in the holes.

4. Mendeleev’s Periodic Table  Contained 1 inconsistency. –He placed the elements in order of atomic mass  Forced to break pattern a couple times to preserve the patterns he had discovered.

5. Henry Moseley Once the proton was discovered by Rutherford, Moseley worked with X-rays to determine how many protons elements had.Once the proton was discovered by Rutherford, Moseley worked with X-rays to determine how many protons elements had. He arranged the elements in order of increasing atomic number.He arranged the elements in order of increasing atomic number. The eliminated the inconsistencies seen in Mendeleev’s table.The eliminated the inconsistencies seen in Mendeleev’s table. He was killed at the age of 28 in WWI.He was killed at the age of 28 in WWI.

6. Periodic Law  Periodic Law: The properties of the elements are periodic functions of their atomic numbers.  What this means is that if we arrange the elements in order of increasing atomic number, we will periodically encounter elements that have similar chemical and physical properties.  These elements appear in the same vertical column (group).

7. Elements  Science has come along way since Aristotle’s theory of Air, Water, Fire, and Earth.  Scientists have identified 92 naturally occurring elements, and created about 28 others.

8. Periodic Table  We can learn a lot about an element based on where it is located.  You can predict the physical and chemical properties and how an element will react.

9. Metals Nonmetals Metalloids LABEL THESE ON YOUR PERIODIC TABLE

10. Properties of Metals  Good conductors of heat and electricity.  Shiny  Ductile (can be stretched into thin wires).  Malleable (can be pounded into thin sheets).  Reacts with water which results in corrosion.

11. Properties of Non-Metals  poor conductors of heat and electricity.  not ductile or malleable.  Solid non-metals are brittle and break easily.  dull  Many non-metals are gases. Sulfur Bromine

12. Properties of Metalloids  Metalloids (metal-like) have properties of both metals and non-metals.  They are solids that can be shiny or dull.  They conduct heat and electricity better than non- metals but not as well as metals.  They are ductile and malleable.

13.  Columns of elements are called groups or families.  Elements in each family have similar but not identical properties.  For example, lithium (Li), sodium (Na), potassium (K), and other members of family 1 are all soft, white, shiny metals.  All elements in a family have the same number of valence electrons.

14.  Each horizontal row of elements is called a period.  The elements in a period are not alike in properties.  In fact, the properties change greatly across even given row.  The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas.

15. Hydrogen  The hydrogen square sits atop Family 1, but it is not a member of that family. Hydrogen is in a class of its own.  It’s a gas at room temperature.  It has one proton and one electron in its one and only energy level.  Hydrogen only needs 2 electrons to fill up its valence shell.

16. Alkali Metals  first column of the periodic table  1 valence electron  They are shiny, have the consistency of clay, and are easily cut with a knife That is Sodium!

17. Alkali Metals  most reactive metals  react violently with water  never found as free elements in nature. They are always bonded with another element. That is sodium on top of the water.

18. Alkaline Earth Metals  never found alone in nature, always bonded with something else.  two valence electrons

19. Transition Metals  elements in the d-block (Groups 3-12)  These are the metals you are probably most familiar

20. Transition Elements  Transition elements have properties similar to one another and to other metals, but their properties do not fit in with those of any other family.  Many transition metals combine chemically with oxygen to form compounds called oxides.

21. Boron Family  The Boron Family is named after the first element in the family.  3 valence electrons  This family includes a metalloid (boron), and the rest are metals.  This family includes the most abundant metal in the earth’s crust (aluminum).

22. Carbon Family  4 valence electrons  This family includes a non- metal (carbon), metalloids, and metals.  The element carbon is called the “basis of life.” There is an entire branch of chemistry devoted to carbon compounds called organic chemistry.

23. Nitrogen Family  The nitrogen family is named after the element that makes up 78% of our atmosphere.  This family includes non- metals, metalloids, and metals.  5 valence electrons. They tend to share electrons when they bond  Other elements in this family are phosphorus, arsenic, antimony, and bismuth.

24. Oxygen Family  6 valence electrons  Most elements in this family share electrons when forming compounds.  Oxygen is the most abundant element in the earth’s crust. It is extremely active and combines with almost all elements.

25. Halogen Family  7 valence electrons, which explains why they are the most active non-metals. They are never found free in nature.  Halogen atoms only need to gain 1 electron to fill their outermost energy level.  They react with alkali metals to form salts.

26. Noble Gases  Noble Gases are colorless gases that are extremely un- reactive.  One important property of the noble gases is their inactivity. They are inactive because their outermost energy level is full.  Because they do not readily combine with other elements to form compounds, the noble gases are called inert.  All the noble gases are found in small amounts in the earth's atmosphere.

27. Rare Earth Elements  The thirty rare earth elements are composed of the lanthanide and actinide series.  One element of the lanthanide series and most of the elements in the actinide series are called trans-uranium, which means synthetic or man-made.

28. Octet Rule   Octet means to have 8 valence electrons.   Is associated with the stability of the noble gases.   Helium (He) is stable with 2 valence electrons   Valence Electrons – –He 1s 2 2 – –Ne 1s 2 2s 2 2p 6 8 – –Ar 1s 2 2s 2 2p 6 3s 2 3p 6 8 – –Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 8

29. Electron Shielding Shielding electrons are those electrons in the energy levels between the nucleus and the valence electrons. They are called "shielding" electrons because they "shield" the valence electrons from being pulled closer to the nucleus. Remember, the nucleus has a positive charge!

31. Periodic Trends  Trends in properties of the elements that follow a pattern down a group and across a period in the periodic table. down a group across a period

32. Trends in Atomic Radius (size)  atomic radius distance from center of nucleus to edge of electron cloud  atomic radius: distance from center of nucleus to edge of electron cloud

33. Atomic Radius  group trend: increases going down a group. –As you move down a group, you add energy levels, thus increasing the size of the electron cloud, so the atoms get larger.

34. Atomic Radius  periodic trend: decreases going left to right across a period.  Why?: You are not adding energy levels, you are adding protons. As you do this, the nucleus gets a greater positive charge which pulls the electrons closer to it, decreasing the radius.

35.  Atomic Radius –Increases to the LEFT and DOWN Atomic Radius

36. Ionization Energy  ionization energy: the energy required to remove a valence electron.

37. Trends in Ionization Energy group trend: decreases going down a group This is due to the shielding effect - an electron in the outer energy level of a large atom is easier to remove because it is well-shielded from the pull of the nucleus by the inner electrons. periodic trend: increases going across a period This is due to nuclear charge - across a period, nuclear charge increases, so they hold onto their electrons tighter. metals have a much greater tendency to lose electrons than nonmetals do.

38.  Ionization Energy –Increases UP and to the RIGHT Ionization Energy

39. Trends in Electron Affinity  the energy change that accompanies the addition of an electron to an atom.  electron affinity (electron-liking): the energy change that accompanies the addition of an electron to an atom.  group trend: EA decreases going down a group. –Why? It is harder to add an electron when it is so far away from the nucleus. The nucleus cannot “grab onto it”.  periodic trend: EA increases going across a period. –Why? As the nuclear charge becomes more positive, it is able to hold onto electrons better. –Note that this periodic trend supports the idea that nonmetals have a much greater tendency to gain electrons than metals do.

40.  Electron Affinity –Increases UP and to the RIGHT Electron Affinity

41. Trends in Electronegativity  the tendency of an atom to attract electrons to itself when it is chemically bonded with another element.  electronegativity: the tendency of an atom to attract electrons to itself when it is chemically bonded with another element.  Diagram of water molecule: Diagram of water molecule: Diagram of water molecule: –In H 2 O, oxygen is more electronegative than hydrogen, so it pulls the electrons closer, and thus obtains a partially negative charge.

42. Trends in Electronegativity group trend: decreases down a group. Larger atoms have more energy levels, so it is harder for them to attract electrons to the nucleus (shielding effect). periodic trend: increases across a period. The nuclear charge is greater with more protons and can hold onto electrons closer to the nucleus.

43.  Electronegativity –Increases UP and to the RIGHT Electronegativity

44. Ionic Radius  Ion: charged atom  When atoms lose electrons, they become positive ions –More p+ than e-  When atoms gain electrons, they become negative ions. -more e- than p+

45. Positive Ions  Example…  Sodium has 11 p + and 11 e -. To become the sodium ion it loses an e -. Now it has 11 p + and 10 e - Sodium atom NaSodium ion Na +1

46. Negative Ions  Example…  The oxygen atom has 8 p+ and 8 e-. The oxygen ion gains 2 e- to have a total of 8 p+ and 10e-.

47. Ionic Radius Positive ions are smaller than the neutral atom. Positive ions are smaller than the neutral atom. This is because you remove valence electrons and they no longer have the outer shell.

48. Ionic Radius  Negative ions are larger than the neutral atom.  They get bigger because adding electrons to the outer shell causes them to want to be further apart (like charges repel).

49. Trends in Ionic Radius  Ionic radius increases going down and to the left.

50.  Ionization Energy, Electron Affinity, and Electronegativity –Increases UP and to the RIGHT Summary of Periodic Trends  Atomic Radius –Increases to the LEFT and DOWN IE EA EN AR