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Periodic Trends (SL) u OBJECTIVES: Interpret group and period trends in atomic radii, ionic radii, ionization energies and electronegativity (m.p., b.p.,

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

1 Periodic Trends (SL) u OBJECTIVES: Interpret group and period trends in atomic radii, ionic radii, ionization energies and electronegativity (m.p., b.p., and chemical properties must also be studied)

2 Trends in Atomic Size u First problem: Where do you start measuring from? u The electron cloud doesn’t 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 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

11 SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

12 What determines IE u The greater the nuclear charge, the greater IE. u Greater distance from nucleus decreases IE u Shielding effect (inter-electron repulsion)

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.

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 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.

20 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.

21 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 ‘main’ groups elements have noble gas configuration.

22 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.

23 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.

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

25 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-

26 Size of Isoelectronic ions u Iso- means the same u Iso electronic ions have the same number 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

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

28 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.

29 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.

30 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.

31 Ionization energy, Electronegativity, and Electron Affinity INCREASE

32 Atomic size increases, shielding constant Ionic size increases


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