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Periodic Trends u OBJECTIVES: Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties.

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Presentation on theme: "Periodic Trends u OBJECTIVES: Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties."— Presentation transcript:

1 Periodic Trends u OBJECTIVES: Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties

2 Trends in Atomic Size u First problem: Where do you start measuring from? u The electron cloud doesnt have a definite edge. u They get around this by measuring more than 1 atom at a time.

3 Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

4 Trends in Atomic Size u Influenced by three factors: 1. Energy Level Higher energy level is further away. 2. Charge on nucleus More charge pulls electrons in closer. u 3. Shielding effect e e repulsion

5 Group trends u As we go down a group... u each atom has another energy level, u so the atoms get bigger. H Li Na K Rb

6 Periodic Trends u As you go across a period, the radius gets smaller. u Electrons are in same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

7 Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

8 Trends in Ionization Energy u The amount of energy required to completely remove a mole of electrons from a mole of gaseous atoms. u Removing an electron makes a +1 ion. u The energy required to remove (1 mole of) the first electron is called the first ionization energy.

9 Ionization Energy u The second ionization energy is the energy required to remove (1 mole of) the second electron(s). u Always greater than first IE. u The third IE is the energy required to remove a third electron. u Greater than 1st or 2nd IE.

10 SymbolFirstSecond Third H He Li Be B C N O F Ne

11 SymbolFirstSecond Third H He Li Be B C N O F Ne

12 What determines IE u The greater the nuclear charge, the greater IE. u Greater distance from nucleus decreases IE u Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. u Shielding effect

13 Shielding u The electron in the outermost energy level experiences more inter-electron repulsion (shielding). u Second electron has same shielding, if it is in the same period

14 Group trends u As you go down a group, first IE decreases because... u The electron is further away. u More shielding.

15 Periodic trends u All the atoms in the same period have the same energy level. u Same shielding. u But, increasing nuclear charge u So IE generally increases from left to right. u Exceptions at full and 1/2 full orbitals.

16 First Ionization energy Atomic number He u He has a greater IE than H. u same shielding u greater nuclear charge H

17 First Ionization energy Atomic number H He Li has lower IE than H Outer electron further away l outweighs greater nuclear charge Li

18 First Ionization energy Atomic number H He Be has higher IE than Li same shielding l greater nuclear charge Li Be

19 First Ionization energy Atomic number H He B has lower IE than Be same shielding greater nuclear charge l p orbital is slightly more diffuse and its electron easier to remove Li Be B

20 First Ionization energy Atomic number H He Li Be B C

21 First Ionization energy Atomic number H He Li Be B C N

22 First Ionization energy Atomic number H He Li Be B C N O u Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter- electron repulsion.

23 First Ionization energy Atomic number H He Li Be B C N O F

24 First Ionization energy Atomic number H He Li Be B C N O F Ne u Ne has a lower IE than He u Both are full, u Ne has more shielding u Greater distance

25 First Ionization energy Atomic number H He Li Be B C N O F Ne Na has a lower IE than Li Both are s 1 Na has more shielding l Greater distance Na

26 First Ionization energy Atomic number

27 Driving Force u Full Energy Levels require lots of energy to remove their electrons. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

28 Trends in Electron Affinity u The energy change associated with adding an electron (mole of electrons) to a (mole of) gaseous atom(s). u Easiest to add to group 7A. u Gets them to full energy level. u Increase from left to right: atoms become smaller, with greater nuclear charge. u Decrease as we go down a group.

29 Trends in Ionic Size u Cations form by losing electrons. u Cations are smaller that the atom they come from. u Metals form cations. u Cations of representative elements have noble gas configuration.

30 Ionic size u Anions form by gaining electrons. u Anions are bigger that the atom they come from. u Nonmetals form anions. u Anions of A groups elements have noble gas configuration.

31 Configuration of Ions u Ions have noble gas configurations (not transition metals). u Na is: 1s 2 2s 2 2p 6 3s 1 u Forms a 1+ ion: 1s 2 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

32 Configuration of Ions u Non-metals form ions by gaining electrons to achieve noble gas configuration. u They end up with the configuration of the noble gas after them.

33 Group trends u Adding energy level u Ions get bigger as you go down. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

34 Periodic Trends u Across the period, nuclear charge increases so they get smaller. u Energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

35 Size of Isoelectronic ions u Iso- means the same u Iso electronic ions have the same # of electrons u Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- and N 3- u all have 10 electrons u all have the configuration: 1s 2 2s 2 2p 6

36 Size of Isoelectronic ions u Positive ions that have more protons would be smaller. Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- N 3-

37 Electronegativity u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u High electronegativity means it pulls the electron toward it. u Atoms with large negative electron affinity have larger electronegativity.

38 Group Trend u The further down a group, the farther the electron is away, and the more electrons an atom has. u More willing to share. u Low electronegativity.

39 Periodic Trend u Metals are at the left of the table. u They let their electrons go easily u Low electronegativity u At the right end are the nonmetals. u They want more electrons. u Try to take them away from others u High electronegativity.

40 Ionization energy, Electronegativity, and Electron Affinity INCREASE

41 Atomic size increases, shielding constant Ionic size increases


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