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1 Chapter 7 Atomic Structure. 2 Periodic Trends n Ionization energy the energy required to remove an electron form a gaseous atom n Highest energy electron.

Presentation on theme: "1 Chapter 7 Atomic Structure. 2 Periodic Trends n Ionization energy the energy required to remove an electron form a gaseous atom n Highest energy electron."— Presentation transcript:

1 Chapter 7 Atomic Structure

2 Periodic Trends n Ionization energy the energy required to remove an electron form a gaseous atom n Highest energy electron removed first. First ionization energy ( I 1 ) is that required to remove the first electron. First ionization energy ( I 1 ) is that required to remove the first electron. Second ionization energy ( I 2 ) - the second electron Second ionization energy ( I 2 ) - the second electron n etc. etc.

3 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding

4 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

5 Across a Period Generally from left to right, I 1 increases because Generally from left to right, I 1 increases because n there is a greater nuclear charge with the same shielding. As you go down a group I 1 decreases because electrons are further away and there is more shielding As you go down a group I 1 decreases because electrons are further away and there is more shielding

6 It is not that simple Z eff changes as you go across a period, so will I 1 Z eff changes as you go across a period, so will I 1 n Half-filled and filled orbitals are harder to remove electrons from n here’s what it looks like

7 First Ionization energy Atomic number

8 First Ionization energy Atomic number

9 First Ionization energy Atomic number

10 Atomic Size n First problem where do you start measuring n The electron cloud doesn’t have a definite edge. n They get around this by measuring more than 1 atom at a time

11 Atomic Size n Atomic Radius = half the distance between two nuclei of a diatomic molecule } Radius

12 Trends in Atomic Size n Influenced by two factors n Shielding n More shielding is further away n Charge on nucleus n More charge pulls electrons in closer

13 Group trends n As we go down a group n Each atom has another energy level n So the atoms get bigger H Li Na K Rb

14 Periodic Trends n As you go across a period the radius gets smaller. n Same energy level n More nuclear charge n Outermost electrons are closer NaMgAlSiPSClAr

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

16 Electron Affinity n The energy change associated with adding an electron to a gaseous atom n High electron affinity gives you energy- n exothermic n More negative n Increase (more - ) from left to right –greater nuclear charge. n Decrease as we go down a group –More shielding

17 Ionic Size n Cations form by losing electrons n Cations are smaller than the atom they come from n Metals form cations n Cations of representative elements have noble gas configuration.

18 Ionic size n Anions form by gaining electrons n Anions are bigger than the atom they come from n Nonmetals form anions n Anions of representative elements have noble gas configuration.

19 Configuration of Ions n Ions always have noble gas configuration n Na is 1s 2 2s 2 2p 6 3s 1 n Forms a 1+ ion - 1s 2 2s 2 2p 6 n Same configuration as neon n Metals form ions with the configuration of the noble gas before them - they lose electrons

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

21 Group trends n Adding energy level n Ions get bigger as you go down Li +1 Na +1 K +1 Rb +1 Cs +1

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

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

24 Size of Isoelectronic ions n Positive ions have more protons so they are smaller Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

25 Electronegativity

26 Electronegativity n The tendency for an atom to attract electrons to itself when it is chemically combined with another element. n How “greedy” n Big electronegativity means it pulls the electron toward itself. n Atoms with large negative electron affinity have larger electronegativity.

27 Group Trend n The further down a group more shielding n Less attracted (Z eff ) n Low electronegativity.

28 Periodic Trend n Metals are at the left end n Low ionization energy- low effective nuclear charge n Low electronegativity n At the right end are the nonmetals n More negative electron affinity n High electronegativity n Except noble gases

29 Ionization energy, electronegativity Electron affinity INCREASE

30 Atomic size increases, Ionic size increases

31 Parts of the Periodic Table

32 The information it hides n Know the special groups n It is the number and type of valence electrons that determine an atom’s chemistry. n You can get the electron configuration from it. n Metals lose electrons have the lowest IE n Non metals- gain electrons most negative electron affinities

33 The Alkali Metals n Doesn’t include hydrogen- it behaves as a non-metal n decrease in IE n increase in radius n Decrease in density n decrease in melting point n Behave as reducing agents

34 Reducing ability n Lower IE< better reducing agents n Cs>Rb>K>Na>Li n works for solids, but not in aqueous solutions. n In solution Li>K>Na n Why? n It’s the water -there is an energy change associated with dissolving

35 Hydration Energy n Li + (g) → Li + (aq) is exothermic n for Li + -510 kJ/mol n for Na + -402 kJ/mol n for K + -314 kJ/mol n Li is so big because of it has a high charge density, a lot of charge on a small atom. n Li loses its electron more easily because of this in aqueous solutions

36 The reaction with water n Na and K react explosively with water n Li doesn’t. Even though the reaction of Li has a more negative  H than that of Na and K Even though the reaction of Li has a more negative  H than that of Na and K n Na and K melt  H does not tell you speed of reaction  H does not tell you speed of reaction n More in Chapter 12.

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