1. Chapter 5 – The Periodic Law Chuck Norris destroyed the periodic table, because he only recognizes the element of surprise.

2. Some people are REALLY into the elements. geek blinggeek tattoo geek blinggeek tattoo

3. The Periodic Table…so far Displays the known elements in order of increasing atomic number Displays the known elements in order of increasing atomic number Arrangement is related to electron configurations of elements Arrangement is related to electron configurations of elements Horizontal rows are periods (energy levels) Horizontal rows are periods (energy levels) Vertical columns are groups or families (similar properties) Vertical columns are groups or families (similar properties)

4. The periodic table also… Allows us to predict the properties of elements. Allows us to predict the properties of elements. Helps us to understand the chemical behavior of elements. Helps us to understand the chemical behavior of elements. Allows us to detect trends in the properties of the elements. Allows us to detect trends in the properties of the elements. Shows us that chemical properties are predictable. Shows us that chemical properties are predictable.

5. 5-1 History of the Periodic Table Dmitri Mendeleev, Russian chemist – hoped to organize the known elements according to their properties Dmitri Mendeleev, Russian chemist – hoped to organize the known elements according to their properties At the time (1800s), nothing was known about the structure of the nucleus. At the time (1800s), nothing was known about the structure of the nucleus. Mendeleev put elements with their properties on cards, arranged them according to their properties, and looked for patterns Mendeleev put elements with their properties on cards, arranged them according to their properties, and looked for patterns

6. 5-1 Mendeleev and Chemical Periodicity Noticed that when elements placed in order of increasing atomic mass, certain similarities appeared at regular intervals Noticed that when elements placed in order of increasing atomic mass, certain similarities appeared at regular intervals Repeating pattern is called PERIODIC Repeating pattern is called PERIODIC Published the first periodic table in 1869 Published the first periodic table in 1869

7. 5-1 Mendeleev’s Periodic Table Left empty spaces in order to keep elements with similar properties together Left empty spaces in order to keep elements with similar properties together Suggested empty spaces represented elements not yet discovered AND predicted the properties of three undiscovered elements Suggested empty spaces represented elements not yet discovered AND predicted the properties of three undiscovered elements by 1886, all three had been discovered – predictions were very accurate by 1886, all three had been discovered – predictions were very accurate Other chemists accept his arrangement. Other chemists accept his arrangement.

8. 5-1 Mendeleev’s Periodic Table – A Problem A few elements are out of atomic mass order A few elements are out of atomic mass order This was necessary to keep elements with similar properties together This was necessary to keep elements with similar properties together Mendeleev thought atomic masses may have been incorrectly measured Mendeleev thought atomic masses may have been incorrectly measured

9. 5-1 Moseley and Atomic Number 1911, English physicist Henry Moseley was working with atomic spectra of metals 1911, English physicist Henry Moseley was working with atomic spectra of metals Noticed elements fit into patterns better when arranged in order of increasing NUCLEAR CHARGE Noticed elements fit into patterns better when arranged in order of increasing NUCLEAR CHARGE Led to modern definition of ATOMIC NUMBER and new basis for arrangement of periodic table Led to modern definition of ATOMIC NUMBER and new basis for arrangement of periodic table

10. 5-1 The Modern Periodic Table An arrangement of elements in order of their ATOMIC NUMBERS so that elements with similar properties fall in the same column or group An arrangement of elements in order of their ATOMIC NUMBERS so that elements with similar properties fall in the same column or group Properties repeat PERIODICALLY Properties repeat PERIODICALLY

11. 5-1 The Modern Periodic Table An interactive periodic table An interactive periodic table Even better interactive periodic table Even better interactive periodic table

12. 5-1 The Modern Periodic Table Noble Gases – earliest tables omitted them because they hadn’t been discovered yet (1894 – Ar, 1895 – He, 1898 – Kr and Xe, 1900 – Rn) Noble Gases – earliest tables omitted them because they hadn’t been discovered yet (1894 – Ar, 1895 – He, 1898 – Kr and Xe, 1900 – Rn) Lanthanides and Actinides – added in early 1900s when their properties were better understood – belong in periods 6 and 7, between s block and p block Lanthanides and Actinides – added in early 1900s when their properties were better understood – belong in periods 6 and 7, between s block and p block

13. 5-2 Electron Configuration and the Periodic Table The properties of an element are determined by the electron configuration of the atom’s highest occupied energy level (the valence electrons) The properties of an element are determined by the electron configuration of the atom’s highest occupied energy level (the valence electrons) A filled outer energy level (s & p sublevels filled) gives an atom special stability – this is why the noble gases are so unreactive. A filled outer energy level (s & p sublevels filled) gives an atom special stability – this is why the noble gases are so unreactive.

14. 5-1 The Periodic Table and Electron Configuration

15. 5-2 Using the Periodic Table to Determine Electron Configuration Ending For s and p blocks: period gives energy level, block gives sublevel, and column gives number or electrons For s and p blocks: period gives energy level, block gives sublevel, and column gives number or electrons Examples: K, Si, Rn, Cl, Sb, Sr, Mg, P Examples: K, Si, Rn, Cl, Sb, Sr, Mg, P

16. 5-2 The s-Block Elements: Groups 1 and 2 Chemically REACTIVE metals (group 1 > group 2) Chemically REACTIVE metals (group 1 > group 2) Group 1: ns 1, alkali metals, silvery, soft, can be cut with a knife, so reactive they are not found free in nature (only in compounds) Group 1: ns 1, alkali metals, silvery, soft, can be cut with a knife, so reactive they are not found free in nature (only in compounds) Group 2: ns 2, alkaline earth metals, harder, denser and stronger than alkali metals, higher melting points, less reactive than alkali metals but still too reactive to be found free in nature Group 2: ns 2, alkaline earth metals, harder, denser and stronger than alkali metals, higher melting points, less reactive than alkali metals but still too reactive to be found free in nature

17. 5-2 The Alkali Metals Alkali metal reactivity Alkali metal reactivity Alkali metal reactivity Alkali metal reactivity

18. 5-2 Alkaline Earth Metals

19. 5-2 Hydrogen and Helium Special cases Special cases Hydrogen: 1s 1, but is NOT an alkali metal – it is a nonmetal and a gas at room temperature Hydrogen: 1s 1, but is NOT an alkali metal – it is a nonmetal and a gas at room temperature Helium: 1s 2, placed with noble gases because of its properties (unreactive gas), has a filled outer energy level with only 2 electrons (1 st energy level only holds 2 electrons) Helium: 1s 2, placed with noble gases because of its properties (unreactive gas), has a filled outer energy level with only 2 electrons (1 st energy level only holds 2 electrons)

20. 5-2 The d-Block Elements: Groups 3-12 For energy level n, there are n possible sublevels, so n=3 has 3 sublevels For energy level n, there are n possible sublevels, so n=3 has 3 sublevels Because 3d has higher energy than 4s, the d-block first appears in the 4 th period Because 3d has higher energy than 4s, the d-block first appears in the 4 th period ns(n-1)d ns(n-1)d Transition metals – metals with typical metallic properties, good conductors of heat and electricity, high luster, less reactive than s-block metals, some so unreactive they exist in nature as free elements, palladium, platinum and gold are least reactive Transition metals – metals with typical metallic properties, good conductors of heat and electricity, high luster, less reactive than s-block metals, some so unreactive they exist in nature as free elements, palladium, platinum and gold are least reactive

21. 5-2 Exceptions to the Aufbau Principle Atoms with a filled or half-filled sublevel have special stability Atoms with a filled or half-filled sublevel have special stability Some d-block elements have unusual electron configurations that reflect this fact Some d-block elements have unusual electron configurations that reflect this fact Cr, Cu Cr, Cu

22. 5-2 The p-Block Elements – Groups 13-18 Electrons only add to p sublevel once s is filled, so all p-block elements have 2 electrons in ns. Electrons only add to p sublevel once s is filled, so all p-block elements have 2 electrons in ns. Along with s-block elements, p-block elements are called MAIN-GROUP ELEMENTS. Along with s-block elements, p-block elements are called MAIN-GROUP ELEMENTS. Number of valence electrons = group # - 10 Number of valence electrons = group # - 10 Properties vary greatly – metals (8), metalloids (6) and nonmetals (including halogens and noble gases) Properties vary greatly – metals (8), metalloids (6) and nonmetals (including halogens and noble gases)

23. 5-2 Noble Gases ns 2 [(n-1)d 10 ]np 6 ns 2 [(n-1)d 10 ]np 6 Have an octet (a filled s and p sublevel in their highest energy level) Have an octet (a filled s and p sublevel in their highest energy level) Have 8 valence electrons, filled outer energy level Have 8 valence electrons, filled outer energy level Helium is an exception – has a filled outer energy level with only 2 electrons Helium is an exception – has a filled outer energy level with only 2 electrons Stable and inert (unreactive) Stable and inert (unreactive)

24. 5-2 Halogens Group 17: fluorine, chlorine, bromine, iodine, astatine) Group 17: fluorine, chlorine, bromine, iodine, astatine) “Salt forming” “Salt forming” Most reactive nonmetals Most reactive nonmetals F, Cl – gases at RT, Br – liquid at RT, I – solid at RT F, Cl – gases at RT, Br – liquid at RT, I – solid at RT

25. 5-2 Metalloids Semiconducting elements Semiconducting elements Mostly brittle solids with some metallic properties and some nonmetallic properties Mostly brittle solids with some metallic properties and some nonmetallic properties Intermediate conductivity Intermediate conductivity Top – As, bottom - Si Top – As, bottom - Si

26. 5-2 p-Block Metals Harder and denser than s-block metals but softer and less dense than d-block metals Harder and denser than s-block metals but softer and less dense than d-block metals Reactive – only found in nature in compounds (except Bi) Reactive – only found in nature in compounds (except Bi) Once obtained as free metals, stable in air Once obtained as free metals, stable in air Top – Pb, bottom – Bi Top – Pb, bottom – Bi

27. 5-3 Electron Configuration and Periodic Properties Atomic Radius – distance from the center of the nucleus to the outer edge of the electron cloud Atomic Radius – distance from the center of the nucleus to the outer edge of the electron cloud OR ½ the distance between the nuclei of two identical atoms that are bonded together OR ½ the distance between the nuclei of two identical atoms that are bonded together

28. 5-3 Trends in Atomic Radii L to R across a period – atomic radius decreases due to increasing nuclear charge L to R across a period – atomic radius decreases due to increasing nuclear charge Elements in a period have outermost electrons in same energy level so might predict about same size, but increasing charge of nucleus pulls electron cloud in Elements in a period have outermost electrons in same energy level so might predict about same size, but increasing charge of nucleus pulls electron cloud in

29. 5-3 Trends in Atomic Radii Top to bottom down a group – atomic radius increases Top to bottom down a group – atomic radius increases Electrons occupy higher energy levels located farther from nucleus Electrons occupy higher energy levels located farther from nucleus

30. 5-3 Atomic Radii

31. 5-3 Atomic Radius as a Function of Atomic Number

32. 5-3 Valence Electrons and the Periodic Table

33. 5-3 Ions and Ion Formation Ion – an atom or group of bonded atoms that has a positive or negative charge Ion – an atom or group of bonded atoms that has a positive or negative charge Monatomic ions form when an atom gains or loses an electron or electrons Monatomic ions form when an atom gains or loses an electron or electrons Cations – positive ions formed by the loss of electrons Cations – positive ions formed by the loss of electrons Anions – negative ions formed by the gain of electrons Anions – negative ions formed by the gain of electrons

34. 5-3 The Octet Rule and Ion Formation Octet rule – atoms will gain, lose or share electrons in order to obtain a filled outer energy level (usually 8 electrons) Octet rule – atoms will gain, lose or share electrons in order to obtain a filled outer energy level (usually 8 electrons) Metals tend to lose electrons in order to gain an octet Metals tend to lose electrons in order to gain an octet Nonmetals tend to gain (or share) electrons in order to gain an octet Nonmetals tend to gain (or share) electrons in order to gain an octet

35. 5-3 The Octet Rule: Sodium Sodium has one valence electron. Sodium has one valence electron. It loses one valence electron to reveal a filled shell. It loses one valence electron to reveal a filled shell. A sodium ion has a charge of +1, Na +. A sodium ion has a charge of +1, Na +.

36. 5-3 The Octet Rule: Chlorine Chlorine has 7 valence electrons. Chlorine has 7 valence electrons. It is very close to having a filled outer energy level. It is very close to having a filled outer energy level. It will gain one electron to fill its outer energy level. It will gain one electron to fill its outer energy level. A chloride ion has a charge of -1, Cl -. A chloride ion has a charge of -1, Cl -.

37. 5-3 The Octet Rule and Ion Formation In general… In general… Atoms with 1, 2 or 3 valence electrons (metals) will lose their valence electrons, forming cations. Atoms with 1, 2 or 3 valence electrons (metals) will lose their valence electrons, forming cations. Atoms with 5, 6 or 7 valence electrons (nonmetals) will gain enough electrons to complete their octet, forming anions. Atoms with 5, 6 or 7 valence electrons (nonmetals) will gain enough electrons to complete their octet, forming anions.

38. 5-3 Ionization Energy Neutral atoms can lose electrons Neutral atoms can lose electrons Ionization energy, IE – energy required to remove one electron from a neutral atom of an element, measured in kJ/mol Ionization energy, IE – energy required to remove one electron from a neutral atom of an element, measured in kJ/mol A + energy  A + + e -

39. 5-3 Trends in Ionization Energy L to R across a period – ionization energy increases - it becomes harder to remove the valence electron(s) L to R across a period – ionization energy increases - it becomes harder to remove the valence electron(s) The closer the valence electrons are to the nucleus and the greater the nuclear charge, the harder they are to remove – held more tightly The closer the valence electrons are to the nucleus and the greater the nuclear charge, the harder they are to remove – held more tightly Nonmetals have higher ionization energies than metals Nonmetals have higher ionization energies than metals

40. 5-3 Trends in Ionization Energy Top to bottom down group – ionization energy decreases – it becomes easier to remove the valence electron(s) Top to bottom down group – ionization energy decreases – it becomes easier to remove the valence electron(s) Electrons removed from each successive element are farther from the nucleus so are held less tightly Electrons removed from each successive element are farther from the nucleus so are held less tightly More electrons lie between the nucleus and the valence electron(s), which partially shields the valence electron(s) from the pull of the nucleus More electrons lie between the nucleus and the valence electron(s), which partially shields the valence electron(s) from the pull of the nucleus

41. 5-3 Trends in Ionization Energy Where on the periodic table are the elements most likely to LOSE electron? Where on the periodic table are the elements most likely to LOSE electron?

42. 5-3 Multiple Ionization Energies Electrons can be removed from positive ions as well as neutral atoms Electrons can be removed from positive ions as well as neutral atoms These successive IEs are names 2 nd ionization energy, 3 rd ionization energy, etc These successive IEs are names 2 nd ionization energy, 3 rd ionization energy, etc 2 nd ionization energy is always higher than 1 st 2 nd ionization energy is always higher than 1 st 3 rd ionization energy is always higher than 2 nd 3 rd ionization energy is always higher than 2 nd As electrons are removed, fewer electrons remain to shield attractive force of nucleus, each successive electron removed feels greater effective nuclear charge As electrons are removed, fewer electrons remain to shield attractive force of nucleus, each successive electron removed feels greater effective nuclear charge

43. 5-3 Successive Ionization Energies of Selected Elements in kJ/mol LiBeNaMg IE 1 520900496738 IE 2 7298175745621451 IE 3 1181514,84969127733

44. 5-3 Ionization Energy Why do you think helium has the highest first ionization energy of any element? Why do you think helium has the highest first ionization energy of any element?

45. 5-3 Electron Affinity Neutral atoms can gain electrons. Neutral atoms can gain electrons. electron affinity – the energy change that occurs when an electron is acquired by a neutral atom electron affinity – the energy change that occurs when an electron is acquired by a neutral atom A + e -  A - + energy The quantity of energy released is represented by a negative number The more energy lost, the more likely an atom is to gain an electron A value of zero means the atom has NO tendency to gain an electron

46. 5-3 Trends in Electron Affinity Halogens gain electrons most readily Halogens gain electrons most readily L to R across p block, electron affinities get more negative L to R across p block, electron affinities get more negative Exception is between groups 14 and 15 because of stability of half-filled sublevel Exception is between groups 14 and 15 because of stability of half-filled sublevel Adding an electron to carbon is easier than adding an electron to nitrogen, which requires forcing an electron to pair up in an orbital of the already half-filled p sublevel Adding an electron to carbon is easier than adding an electron to nitrogen, which requires forcing an electron to pair up in an orbital of the already half-filled p sublevel Note noble gases have electron affinities of zero. Note noble gases have electron affinities of zero.

47. 5-3 Trends in Electron Affinity Top to bottom down group – generally decreases down a group, but there are some exceptions Top to bottom down group – generally decreases down a group, but there are some exceptions Generally, as atomic radius increases, attraction of an atom for an electron decreases Generally, as atomic radius increases, attraction of an atom for an electron decreases

48. 5-3 Ionization Energy, Electron Affinity and Reactivity

49. 5-3 Ionic Radii Cations – smaller than their respective atoms – outer energy level is lost and remaining electrons are drawn closer to nucleus Cations – smaller than their respective atoms – outer energy level is lost and remaining electrons are drawn closer to nucleus

50. 5-3 Ionic Radii Anions – larger than their respective atoms – electron cloud spreads out because of greater repulsion between increased number of electrons Anions – larger than their respective atoms – electron cloud spreads out because of greater repulsion between increased number of electrons

52. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons A measure of the ability of an atom in a chemical compound to attract electrons Most electronegative element, fluorine, is given a value of 4 Most electronegative element, fluorine, is given a value of 4 Values of other elements are assigned relative to fluorine Values of other elements are assigned relative to fluorine

53. Trends in Electronegativity L to R across a period – electronegativities tend to decrease L to R across a period – electronegativities tend to decrease The most electronegative elements are nitrogen, oxygen and the halogens The most electronegative elements are nitrogen, oxygen and the halogens The least electronegative elements are the alkali metals and the alkaline earth metals The least electronegative elements are the alkali metals and the alkaline earth metals

54. Trends in Electronegativity Top to bottom down a group – electronegativities either decrease or remain about the same Top to bottom down a group – electronegativities either decrease or remain about the same Noble gases are unusual – some do not form compounds so electronegativity cannot be measured Noble gases are unusual – some do not form compounds so electronegativity cannot be measured

55. Trends in Electronegativity

56. What does electronegativity mean?

57. 5-3 Polar Bonds If two atoms bonded to one another differ greatly in their electronegativities, the electron density will be drawn to one end of the bond. If two atoms bonded to one another differ greatly in their electronegativities, the electron density will be drawn to one end of the bond. The bond has a positive end and a negative end. The bond has a positive end and a negative end. This is a polar bond. This is a polar bond.

58. There are some polar bonds in a polar bear.