1. 1 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. Development of the Periodic Table Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829). John Newlands Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863). Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

2. Dmitri Mendeleev 1834 – 1907 Russian chemist and teacher given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #) he even left empty spaces to be filled in later

3. At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He predicted their discovery and estimated their properties.

4. 4 IB Topic 3: Periodicity 3.1: The periodic table 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. 3.1.2 Distinguish between the terms group and period. 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20. 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

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6. Henry Moseley 1887 – 1915 arranged the elements in increasing atomic numbers (Z) –properties now recurred periodically

7. Design of the Table Groups are the vertical columns. –elements have similar, but not identical, properties most important property is that they have the same # of valence electrons

8. valence electrons- electrons in the highest occupied energy level all elements have 1,2,3,4,5,6,7, or 8 valence electrons

9. IB prefers this one.

10. 10 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. 3.1.2 Distinguish between the terms group and period. Development of the Periodic Table Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913) Glen Seaborg Discovered the transuranium elements (93- 102) and added the actinide and lanthanide series (1945) Elements arranged by increasing atomic number into periods (rows) and groups or families (columns), which share similar characteristics

11. 11 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. Metals Left side of the periodic table (except hydrogen) Good conductors of heat and electricity Malleable: capable of being hammered into thin sheets Ductile: capable of being drawn into wires Have luster: are shiny Typically lose electrons in chemical reactions

12. 12 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. Metals Alkali metals: Group 1 (1A) Alkaline earth metals: Group 2 (2A) Transition metals: Group B, lanthanide & actinide series

13. 13 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. Nonmetals Right side of the periodic table Poor conductors of heat and electricity Non-lusterous Typically gain electrons in chemical reactions Halogens: Group 17 (7A) Noble gases: Group 18 (0)

14. 14 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number. Metalloids Between metals and non-metals, along the stair step (except aluminum) Have properties of metals and non- metals Some are semi-conductors Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Astatine (At)

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16. 16 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Ground State Electron Configurations of the Elements

17. Electron Arrangement Core Electrons: electrons that are in the inner energy levels Valence Electrons: electrons that are in the outermost (highest) energy level Group = Sum of electrons in the highest occupied energy level (s + p) = Number of valence electrons 17

18. 18 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Arrangement of the Periodic Table Valence Electrons: electrons in the outermost (highest) energy level –Group 1 elements have 1 v.e.s –Group 2 elements have 2 v.e.s –Group 3 elements have 3 v.e.s –So on and so forth –Group 8 have 8 v.e. (except for helium, which has 2)

19. Lewis Dot-Diagrams/Structures valence electrons are represented as dots around the chemical symbol for the element Na Cl

22. 22 3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Electron dot diagrams Group 1A: 1 dotXGroup 5A: 5 dotsX Group 2A: 2 dotsXGroup 6A: 6 dots X Group 3A: 3 dotsXGroup 7A: 7 dotsX Group 4A: 4 dotsXGroup 0: 8 dots (except He)X

23. Look, they are following my rule!

24. Electron Dot Diagram Using the symbol for the element, place dots around the symbol corresponding to the outer energy level s & p electrons (valence electrons). Will have from one to eight dots in the dot diagram. Draw electron dot diagrams for the following atoms H Be O Al Ca Zr H Be O 24

25. Electron Dot Diagram Using the symbol for the element, place dots around the symbol corresponding to the outer energy level s & p electrons. Will have from one to eight dots in the dot diagram. Draw electron dot diagrams for the following atoms Al Ca Zr 25

26. 2.3.4 Deduce the electron arrangement for atoms and ions. Write electron configuration, orbital filling diagrams, and electron dot diagrams. Kr Tb 26

27. B is 1s 2 2s 2 2p 1 ; –2 is the outermost energy level –it contains 3 valence electrons, 2 in the 2s and 1 in the 2p Br is [Ar] 4s 2 3d 10 4p 5 How many valence electrons are present?

28. Periods are the horizontal rows –do NOT have similar properties –however, there is a pattern to their properties as you move across the table that is visible when they react with other elements

29. 29 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table Period = The highest occupied energy level = number of energy levels

30. 30 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table Na = 1s 2 2s 2 2p 6 3s 1 Sodium is in the 3 rd period because it has 3 energy levels  The highest occupied energy level is 3

31. 31 IB Topic 3: Periodicity 3.2: Physical properties 3.2.1 Define the terms first ionization energy and electronegativity. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li  Cs) and the halogens (F  I). 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.

32. 32 Periodic Trend Definitions Atomic Radius: half the internuclear distance between two atoms of the same element (pm) Ionic radius: the radius of an ion in the crystalline form of a compound (pm)

33. 33 Periodic Trend Definitions First ionization energy: The energy required to remove one electron from each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol -1 ) Electron Affinity: The energy released when one electron is added to each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol -1 )

34. 34 Periodic Trend Definitions Electronegativity: a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself Melting Point: the temperature at which a solid becomes a liquid at a fixed pressure (degrees Kelvin)

35. Trends in the table IB loves the alkali metals and the halogens

36. many trends are easier to understand if you comprehend the following the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces –the attraction between the electron and the nucleus –the repulsions between the electron in question and all the other electrons in the atom (often referred to the shielding effect) –the net resulting force of these two is referred to effective nuclear charge

37. This is a simple, yet very good picture. Do you understand it?

38. 38 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Group 1A: Alkali Metals Have 1 valence electron Shiny, silvery, soft metals React with water & halogens Oxidize easily (lose electrons) Reactivity increases down the group Group 7A: Halogens Have 7 valence electrons Colored gas (F 2, Cl 2 ); liquid (Br 2 ); Solid (I 2 ) Oxidizer (gain electrons) Reactivity decreases down the group

39. 39 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Atomic Radii The radius of an atom, measured in pm (picometers) Periodic trend (Period 3 Trend) –Atomic size decreases as you move across a period. –The increase in nuclear charge increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus Group trend for Alkali metals & Halogens –Atomic size increases as you move down a group of the periodic table. –Adding higher energy levels

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41. 41 Atomic Radii

42. 42 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Ionic Radii The radius of the ion form of atoms (cations and anions) Positive ions are smaller than their atoms. –Fewer electrons so nucleus attracts remaining electrons more strongly –One fewer energy level since valence electrons removed. Negative ions are larger than their atoms –More electrons so nucleus has less attraction for them –Greater electron-electron repulsion Periodic trend (Period 3 Trend) –Decrease as you move across a period, then spike and decrease again –This increase in nuclear charge increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus Group trend for Alkali metals & Halogens –Ions get larger down a group –More energy levels are added

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45. 45 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points First Ionization Energies The energy required to remove the first electron from a gaseous atom. Second ionization removes the second electron and so on. Can be used to predict ionic charges. Periodic Trend (Period 3 Trend) –Increases as you move from left to right across a period. –Effect of increasing nuclear charge makes it harder to remove an electron. Group trend for Alkali metals & Halogens –Generally decreases as you move down a group in the periodic table –Since size increases down a group, the outermost electron is farther away from the nucleus and is easier to remove.

46. 46 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

47. 47 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electronegativity Tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Helps predict the type of bonding (ionic/covalent). Periodic Trend (Period 3 Trend) –Increases as you move from left to right across a period. –Nonmetals have a greater attraction for electrons than metals & there is a greater nuclear charge that can attract electrons Group trend for Alkali metals & Halogens –Generally decreases as you move down a group in the periodic table. –For metals, the lower the number the more reactive. –For nonmetals, the higher the number the more reactive.

48. 48 Electronegativity

49. 49 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity The relative capacity of an atom, molecule or radical to undergo a chemical reaction with another atom, molecule or radical. Don’t worry about the periodic trend!!! Group trend for Alkali metals –Increases as you move down group 1 in the periodic table –Since alkali metals are more likely to lose an electron, the ones with the lowest 1 st ionization energy are the most reactive since they require the least amount of energy to lose a valence electron. Group trend for Halogens –Decreases as you move down group 7 in the periodic table –Since halogens are more likely to gain an electron, the ones with the greatest electronegativity are the most reactive since they are most effective at gaining a valence electron.

50. 50 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Melting Points The temperature at which a crystalline melts depends on the strength of the attractive forces and on the way the particles are packed in the solid state Don’t worry about the periodic trend!!! Alkali Metals: Melting point decreases down the group –Li (181 o C) to Cs (29 o C) –As the atoms get larger the forces of attraction between them decrease due to the type of bonding (metallic) Halogens: Melting point increases down the group –F 2 (-220 0 C) to I 2 (114 o C) –Weak attractive forces increase as the molecules get larger due to the type of bonding (non-polar covalent)

51. 51 IB Topic 3: Periodicity 3.3: Chemical properties Discuss the similarities and differences in the chemical properties of elements in the same group. Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

52. 52 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Alkali Metals React with water & react with many substances because… They have the same number of valence electrons

53. 53 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Alkali Metals 2Na(s) + 2H 2 O(l)  2NaOH (aq) + H 2 (g) In the reaction of alkali metals and water, all will: move around the surface of the water, give off hydrogen gas, create a basic solution.

54. 54 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Alkali Metals In the reaction of alkali metals and water, the reactivity will increase down the group because they get better at getting rid of their valence electron (the 1 st ionization energy decreases) So, alkali metals lower down will: React more vigorously React faster Give off a flame

55. 55 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Alkali Metals Reaction with halogens 2M (s) + X 2 (g)  2MX (s) where M represents Li,Na,K,Rb, or Cs Where X represents F,Cl,Br, or I 2Na (s) + Cl 2 (g)  2NaCl (s) Reactivity decreases down the group

56. 56 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. Halogens Halogens are diatomic as gases (two atoms bond together) and called halides when they form ions… These are BrINClHOF Halogens want to get one electron to fill its outer shell. Reactivity decreases down the group because electronegativity decreases Cl 2 reacts with Br - and I - Cl 2 (aq) + 2Br - (aq)  2Cl - (aq) + Br 2 (l) Cl 2 (aq) + 2I - (aq)  2Cl - (aq) + I 2 (s) Br 2 reacts with I- Br 2 (aq) + 2I - (aq)  2Br - (aq) + I 2 (s) I 2 non-reactive with halide ions

57. 57 Reactivity of Elements… in action Alkali Metals: http://www.youtube.com/watch?v=m55kgy ApYrY http://www.youtube.com/watch?v=m55kgy ApYrY Halogens: http://www.youtube.com/watch?v=tk5xwS5b ZMA&feature=related

58. 58 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Metallic Oxides in Period 3 Sodium oxide: Na 2 Oionic Magnesium oxide: MgOionic Aluminum oxide: Al 2 O 3 ionic Metalloid oxide in Period 3 Silicon dioxide: SiO 2 covalent Nonmetallic oxides in Period 3 Tetraphosphorus decoxide: P 4 O 10 covalent Sulfur trioxide: SO 3 covalent Dichlorine heptoxide: Cl 2 O 7 covalent

59. 59 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Acidic/Basic Metallic oxides in Period 3 are basic Sodium oxide: Na 2 O + H 2 O  2 NaOH basic Magnesium oxide: MgO + H 2 O  Mg(OH) 2 basic Aluminum oxide: Al 2 O 3 + H 2 O  2 Al(OH) 3 amphoteric Metalloid oxide in Period 3 is acidic Silicon dioxide:SiO 2 + H 2 O  H 2 SiO 3 acidic Nonmetallic oxides in Period 3 are acidic Tetraphosphorus decoxide: P 4 O 10 + 6H 2 O  4H 3 PO 4 acidic Sulfur trioxide: SO 3 + H 2 O  H 2 SO 4 acidic Dichlorine heptoxide: Cl 2 O 7 + H 2 O  2HClO 4 acidic Argon does not form an oxide

60. 60 Terms to Know Group Period Alkali metals Halogens Ionic radius Electronegativity First ionization energy

61. 61 Periodic Table of Video http://www.periodicvideos.com/