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Chapter 6 1

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Draw the periodic table and label the electron blocks and areas of non-metals, metals, and metalloids. Relate the Lewis dot structure to its place in the periodic table. Explain periodic trends as one moves along periods and down groups in the periodic table 2

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Chapter 6.1-6.2 Periodic Law Group Period Representative Element Transition Element Metal Alkali Metal Alkaline Earth Metal Transition Metal Inner Transition Metal Lanthanide Series Actinide Series Nonmetal Halogen Noble Gas Metalloid 3

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Dmitri Mendeleev noticed in his table that there were repetitions of physical and chemical properties when the elements were arranged by atomic mass. 4

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PropertyPredicted (1869) Atomic Mass72 u ColorDark gray Density5.5 g/mL Melting PointHigh Density of Oxide4.7 g/mL Oxide solubility in HCl Slightly dissolved by HCl Formula of chloride EsCl 4 Properties of Germanium (Ge) 6 Actual (1886)

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Periodic Law states that chemical and physical properties repeat in regular cyclic patterns when they are arranged by increasing atomic number. Starts with metals at left and goes to non-metal (noble gas) on right Properties change in orderly progression across a period. 7

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Periodic Table Periods 9 Transition Elements Inner Transition Elements Halogens Noble Gases Alkali Metals Alkaline Earth Metals Representative Elements Columns, Groups or Families Metals MetalloidsNonmetals

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What are some of the elemental properties that make the periodic table, well, periodic? Classification by metals, nonmetals and metalloids Metals - shiny ductile, malleable solids, good conductors of heat and electricity Nonmetals - dull, brittle solids; or gas, poor conductors of heat and electricity Metalloids - have chemical and physical properties of both metals and nonmetals 10

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Representative Elements (Sometimes called A Group) Group # = number of valence electrons Means similar Lewis dot structure and similar properties. s-block s-block elements have 1-2 electrons in s-orbital p-block p-block elements have 1-6 electrons in p-orbitals Noble gases have filled valence shells Energy level of valence electrons is at energy level given by period (row) number 11

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Transition Elements (Sometimes called B Group) d-block d-block elements have 1-10 electrons in d- orbitals Columns 3-12 in periodic table Energy level of valence electrons at n and partially filled n-1 d orbitals (example: 4s and 3d) f-block f-block (Lanthanides and Actinides) have 1-14 electrons in f-orbitals 12

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Fill in the missing info for the following elements: Identify the element fitting the description. a) Group 2 (2A) element in 4 th period: b) Noble gas in 5 th period: c) Group 12 (2B) element in 4 th period: d) Group 16 (6A) element in 2 nd period: 13 ConfigurationGroupPeriodBlock [Ne]3s 2 [He]2s 1 [Kr]5s 2 4d 10 5p 5 7 (7B)4

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Effective Nuclear Charge (Z*) Not in book! Shielding (Not in book) Ion Ionization Energy Octet Rule Metallic Character (Not in book) Electronegativity 15

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Atomic and ionic size Ionization energy Electronegativity Metallic Character 16 Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.

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Z* is the nuclear charge experienced by the outermost electrons. (Note: not in book!) Z* increases across a period owing to shielding by inner electrons. Shielding is blocking by inner electrons. For a period (row), the number of shielding electrons remain the same, but the number of protons in the nucleus increases. Example: All elements in the second period have the same underlying [He] noble gas configuration. However, the number of protons increase from left to right. 17

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So we can estimate as Z* = [ Z - (no. inner electrons) ] or or Z* = Z S (inner electrons) Z is total number of electrons S is the number of electrons blocking the valence shell electrons, the underlying noble gas electrons. Charge felt by 2s e - in Li Z* = 3 - 2 = 1 Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3and so on! 18

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19 Orbital energies drop as Z* increases

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Atomic size is a periodic trend influenced by electron configuration. For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element. 20

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For other elements, the atomic radius is half the distance between nuclei of identical atoms that are bonded together. 21

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Size (radius) goes UP on going down a group. See previous slide. Because electrons are further from the nucleus, there is less attraction. Size (radius) goes DOWN on going across a period. Size (radius) goes UP on going down a group. See previous slide. Because electrons are further from the nucleus, there is less attraction. Size (radius) goes DOWN on going across a period. 23

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Size (radius) decreases across a period owing to increase in Z*. Each added electron feels a greater and greater positive charge. Note: Electrons in the same energy level dont shield each other too much. 24 Large Small Increase in Z*

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The radius of an atom when it has become an ion. An ion is an atom or bonded group of atoms that has an overall positive or negative charge. An atom acquires a positive charge by losing electrons or negative charge by gaining electrons!! 26

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To form positive ions from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]. Al: [Ne] 3s 2 3p 1 - 3e - Al 3+ : [Ne] 3s 0 3p 0 27

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28 Atoms tend to gain, lose, or share electrons to get 8 valence electrons (except small atoms up to Boron)

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1.Write the electron configuration and orbital box diagram for Mg when it is an ion. Hints: What is its noble gas configuration? What will they do to get an octet? 2.Write the electron configuration and orbital box diagram for O when it is an ion. 29

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Positive ions are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. Electron Configuration as ion is: [He] 2s 0 30 Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a positive ion.

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Negative ions are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes. Electron configuration as ion: 1s 2 2s 2 2p 6 (just like neon.) 31 Forming a negative ion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

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32 See Figure 6-14

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Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? 33

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IE = energy required to remove an electron from an atom in the gas phase. 34 Mg (g) + 738 kJ Mg + (g) + e-

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35 Mg + (g) + 1451 kJ Mg 2+ (g) + e- Mg + has 12 protons and only 11 electrons. Therefore, IE for Mg + > Mg. IE = energy required to remove an electron from an atom in the gas phase.

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1 st : Mg (g) + 735 kJ Mg + (g) + e- 2 nd : Mg + (g) + 1451 kJ Mg 2+ (g) + e- 36 3 rd : Mg 2+ (g) + 7733 kJ Mg 3+ (g) + e- Energy cost is very high to dip into a shell of lower n.

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39 As Z* increases, orbital energies drop and IE increases.

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IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals. Nonmetals lose electrons with difficulty. 41 High ionization energy: atoms want to hold on to electrons; likely to form negative ion Low ionization energy: atom gives up electron easily; likely to form positive ion

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IE decreases down a group Because size increases. Ability to lose electrons generally increases down the periodic table. See reactions of Li, Na, K 42

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Which element in each pair has the larger 1 st ionization energy? A. Na or Al B. Ar or Xe C. Ba or Mg 43

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44 Lithium SodiumPotassium

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*Note: metallic character not in book. An element with metallic character is one that loses electrons easily. Metallic character: is more prevalent in metals on left side of periodic table is less for nonmetals on right side of periodic table that do not lose electrons easily 45

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Relative ability of an element to attract electrons in a chemical bond. Ionization energy reflects ability of atom to attract electrons in an isolated atom Generally, the higher the ionization energy of an atom, the more electronegative the atom will be in a molecule There are many electro negativity scales well use the one by Linus Pauling (values dimensionless) Will be used to determine things like polarity of a chemical bond. 47

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Decreases down a group Why? Due to greater atomic radius Increases across a period Why? Increased positive charge in nucleus (Greater Z*) Same trend as for ionization energy. Surprised? 49

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Moving Left Right (periods) Z * Increases Atomic & ionic Radius Decrease Ionization Energy Increases Electronegativity Increases Metallic Character Decreases Moving Top Bottom (groups) Z* is roughly constant, but val e - distance increases Atomic & Ionic Radius Increase Ionization Energy Decreases Electronegativity Decreases Metallic Character Increases 50

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a)Electronegativity b)Ionic Radius c)Atomic Radius d)Ionization Energy e)Metallic character 51