1. Topic 3: Periodicity 1

2. 2 3.1: The periodic table Understandings 3.1.1 The periodic table is arranged into four blocks associated with the four sub-levels. 3.1.2 The periodic table consists of groups (vertical columns) and period (horizontal rows). 3.1.3 The period number (n) is the outer energy level that is occupied. 3.1.4 The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table. 3.1.4 The periodic table shows the positions of metals, non-metals and metalloids.

3. 3 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).

4. 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

5. 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.

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7. Henry Moseley 1887 – 1915 Arranged the elements in increasing atomic numbers (Z) –Properties now recurred periodically

9. IB prefers this one.

10. 10 Classification of the Elements

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13. Design of the Table A Group (aka family) is a vertical column –Elements have similar, but not identical, properties Most important property is that they have the same # of valence electrons

14. Electron Arrangement Core Electrons: electrons that are in the inner energy levels Valence Electrons: electrons that are in the outermost (highest) energy level 14

15. 15 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)

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

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

20. 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. 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

21. Look, they are following my rule!

22. 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 22

23. 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 23

24. 24 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

25. 2.3.4 Deduce the electron arrangement for atoms and ions. Write electron configuration, orbital filling diagrams, and electron dot diagrams. Kr Tb 25

26. 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?

27. 27 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

28. 28 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table Group = Sum of electrons in the highest occupied energy level (s + p) = Number of valence electrons

29. 29 2.3.4 Deduce the electron arrangement for atoms and ions. Valence electrons Valence electrons are electrons in the outermost energy level of an atom –The sum of electrons in the s & p orbitals in the highest energy level –Ex. Argon’s electron arrangement is 1s 2 2s 2 2p 6 3s 2 3p 6. Since the highest energy level is 3, we add the e-s in 3s 2 + 3p 6 = 8 –So, argon has 8 valence electrons The easy way is to look at its location on the periodic table (except for the transition metals)

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 Since the sum of electrons in the highest occupied energy level is 1, it will be in the 1 st group and have 1 valence electron

31. 31 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 It is in the 1 st group because it has 1 valence electron

32. 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?

33. 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

34. 34 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 (s and p) = number of energy levels

35. 35 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

36. 36 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table Write out the electron configuration for selenium State the relationship between the group that selenium is in and its electron configuration State the relationship between the period that selenium is in and its electron configuration

37. 37 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table Electron configuration for selenium: Se = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 Group # is 16 Since the highest energy level is 4, we add the e-s in 4s 2 + 4p 4 = 6 Therefore, Se has 6 valence e-s Period # is 4 The highest occupied s/p energy level is 4

38. 38 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

39. 39 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

40. 40 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)

41. 41 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)

42. 42 Green = MetalsBlue = MetalloidsYellow = Nonmetals http://www.windows2universe.org/earth/geology/metals.html

43. 43 3.2: Periodic trends Understandings 3.2.1 Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity and electronegativity. 3.2.2 Trends in metallic and non-metallic behavior are due to the trends above. 3.2.3 Oxides change from basic through amphoteric to acidic across a period.

44. 44 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)

45. 45 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 )

46. 46 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)

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

48. 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

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

50. Therefore… Periodic trends typically have to do with an increase in nuclear charge Group trends typically have to do with an increase in shielding effect (more energy levels)

51. 51 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

52. 52 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Atomic Radii Periodic trend (Period 3 Trend) –Atomic radius decreases as you move across a period. –Number of protons in the nucleus increases –Increase in nuclear charge increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus

53. 53 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Atomic Radii Group trend for Alkali metals & Halogens –Atomic radius increases as you move down a group of the periodic table. –More energy levels are added –More shielding H Li Na K Rb

55. 55 Atomic Radii

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57. 57 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 an ion in the crystalline form of a compound Atoms tend to gain or lose electrons in order to have the electron configuration of a noble gas Most want 8 valence electrons and take the easiest approach to obtaining that full “octet” Hydrogen and helium only want 2 There are some other weird exceptions that we’re not worried about yet Let’s assign charges!

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62. 62 Positive Ions (cations) Group 1: Lose 1 valence electron Charge of +1: Li +, Na +, K + Group 2 Lose 2 valence electrons Charge of +2: Mg 2+, Ca 2+ Group 3 Lose 3 valence electrons Charge of +3: Al 3+ Negative Ions (anions) Group 5: Gain 3 electrons Charge of -3: N 3-, P 3- Group 6: Gain 2 electrons Charge of -2: O 2-, S 2- Group 7 Gain 1 electron Charge of -1: F -, Cl -, Br -, I - Common Ion Charges

63. 63 Uncommon Ion Charges In most transition elements, d electrons can become involved in the reaction Iron can lose 2 electrons (the 2 in the 4s) (Fe 2+ ) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Fe 3+ ) The name of the Fe 2+ ion is iron(II) or ferrous The name of the Fe 3+ ion is iron(III) or ferric Chromium can lose 2 electrons (the 2 in the 4s) (Cr 2+ ) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Cr 3+ ) The name of the Cr 2+ ion is chromium(II) or chromous The name of the Cr 3+ ion is chromium(III) or chromic

64. 64 Uncommon Ion Charges These transition metals only form ONE ion: Ag +1, Zn +2 and Cd +2 Label these on your periodic table!!

65. 65 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Ionic Radii Look at the ions compared to their parent atoms Do atoms become smaller or larger when they do this?

67. Cations (+ ions) are smaller than the parent atom Have lost an electron and lost an entire energy level! Therefore, have fewer electrons than protons Li 0.152 nm Li+.078nm + Li forming a cation

68. Anions (– ions) are larger than parent atom Have gained an electron to achieve noble gas configuration Effective nuclear charge has decreased since same nucleus now holding on to more electrons Plus, the added electron repels the existing electrons farther apart (kind of “puffs it out”) F 0.064 nm 9e - and 9p + F - 0.133 nm 10 e - and 9 p + -

70. 70 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Ionic Radii Periodic trend (Period 3 Trend) –Decreases at first when losing electrons (+ ions) –Suddenly increase when gaining electrons ( – ions) –Decreases again due to increased nuclear charge Group trend for Alkali metals & Halogens (same as neutral atoms) –Increase down a group –More energy levels are added

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73. 73 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points 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 ) X(g)  X + (g) + e - Second ionization removes the second electron and so on.

74. 74 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points First Ionization Energy Periodic Trend (Period 3 Trend) –Increases as you move from left to right across a period. –Number of protons in the nucleus increases. –Effect of increasing nuclear charge makes it harder to remove an electron. Group trend for Alkali metals & Halogens –Decreases as you move down a group in the periodic table. –Number of energy levels increases. –Outer electrons are farther away from the nucleus and is easier to remove. –Inner core electrons “shield” the valence electrons from the pull of the positive nucleus and therefore easier to remove.

76. 76 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

77. 77 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electronegativity A measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself. Helps predict the type of bonding (ionic/covalent). Linus Pauling (1901 to 1994) came up with a scale where a value of 4.0 is arbitrarily given to the most electronegative element, fluorine, and the other electronegativities are scaled relative to this value.

79. 79 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electronegativity Periodic Trend (Period 3 Trend) –Increases as you move from left to right across a period. –Number of protons in the nucleus increases. –Increasing nuclear charge makes it more likely to want an electron.

80. 80 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electronegativity Group trend for Alkali metals & Halogens –Generally decreases as you move down a group in the periodic table. –Number of energy levels increases. –Outer electrons are farther away from the nucleus and aren’t as attracted to one another. –Inner core electrons “shield” the valence electrons from the pull of the positive nucleus and therefore less attracted.

81. 81 Electronegativity

82. 82 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points 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 ) In other words, the neutral atom’s likelihood of gaining an electron. Example F (g) + e -  F - (g) will release 328 kJ mol -1 of energy The more negative the value, the greater the attraction for the electron, the more affinity the atom has

83. 83 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electron Affinity Periodic Trend (Period 3 Trend) –Values decrease (become more negative) as you move from left to right across a period…. –Energy released increases… –Meaning affinity for electrons INCREASES –Number of protons in the nucleus increases  increasing nuclear charge makes it more likely to add an electron. Group trend for Alkali metals & Halogens –Generally increase (become less negative) as you move down a group in the periodic table… –Meaning affinity for electrons DECREASES –Number of energy levels increases  outer electrons are farther away from the nucleus, adding to the shielding effect.

84. 84 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electron Affinity Periodic Trend (Period 3 Trend) –Values decrease (become more negative) as you move from left to right across a period…. –Energy released increases… –Meaning affinity for electrons INCREASES –Number of protons in the nucleus increases  increasing nuclear charge makes it more likely to add an electron. Group trend for Alkali metals & Halogens –Generally increase (become less negative) as you move down a group in the periodic table… –Meaning affinity for electrons DECREASES –Number of energy levels increases  outer electrons are farther away from the nucleus, adding to the shielding effect.

85. This one same as “IB textbook”

87. 87 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 solid becomes a liquid at a fixed pressure (degrees Kelvin) 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 Requires understanding of concepts covered in later topics (this year and next year) Know the type of bonding

88. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Melting Points Don’t worry about the periodic trend!!! You will need to know the group trends for Alkali metals and halogens… they’re different! 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) –The “sea of electrons” is further away from the metal ionsElement Melting Point (K) Li453 Na370 K336 Rb312 Cs301 Fr295

89. 89 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Melting Points Halogens: Melting point increases down the group –F 2 (-220 0 C) to I 2 (114 o C) –Halogens molecules are held together with weak van der Waals’ attractive forces due to their non-polar covalent nature –Larger molecules have more electrons, increasing the strength of the IMF

91. increases

92. 92 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 to undergo a chemical reaction with another atom, molecule or radical. Why do atoms react? Atoms that are good at becoming stable (getting a full valence shell) are the most reactive Don’t worry about the periodic trend because elements on opposite sides of the periodic table can be equally reactive… but for different reasons!!!

93. 93 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity of Alkali metals The relative capacity of an atom to undergo a chemical reaction with another atom, molecule or radical. How many valence electrons do the alkali metals have? Are alkali metals more likely to give up or get electrons? What is the name of the property responsible for doing that?

94. 94 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity of Alkali metals It all has to do with ionization energy since they want to lose electrons! Lower ionization energy = more reactive Group trend for Alkali metals –Increases as you move down group 1 –Since alkali metals are more likely to lose an electron, the ones with the lowest ionization energy are the most reactive since they require the least amount of energy to lose a valence electron. Which alkali metal is the best at losing electrons? That’s the most reactive alkali metal!

95. 95 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity of halogens The relative capacity of an atom to undergo a chemical reaction with another atom, molecule or radical. How many valence electrons do the halogens have? Are halogens more likely to give up or get electrons? What is the name of the property responsible for doing that?

96. 96 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity of halogens It all has to do with electronegativity (or electron affinity) since they want to gain electrons! Higher electronegativity = lower electron affinity = more reactive Group trend for Alkali metals –Decreases as you move down group 17 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. Which alkali metal is best at getting electrons? That’s the most reactive halogen!

97. most reactive least reactive

98. 98 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.

99. 99 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

100. 100 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.

101. 101 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

102. 102 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

103. 103 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

104. 104 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

105. 105 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

106. 106 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

107. 107 Terms to Know Group Period Alkali metals Halogens Ionic radius Electronegativity First ionization energy

108. 108 Periodic Table of Video http://www.periodicvideos.com/