Atomic Size, Ionization Energy, and Electronegativity: Trends and Bonding, Lecture notes of Chemistry

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Announcements

 Lon-capa HW 6 – Type 1 due Monday

(11/26) and Type 2 due Wednesday

(11/28) by 7pm

 “Lab 5: Modern Atomic Theory” write-

up due tomorrow in discussion (and

text homework)

Announcements

 All previous lectures are posted on

course website (“Lectures” link in left

menu)

 Final exam is cumulative so you should

start reviewing material soon!

 Last textbook homework assignment

will be given this Friday (due November

30 th)

Ionization and atomic size

 When an element becomes an ion, how is

size and ionization energy affected?

 When electron is pulled off, there is still the same number of protons; larger positive charge pulls remaining electrons in tighter

Ionization and atomic size

 When an element becomes an ion, how is

size and ionization energy affected?

 When electron is added, there is still the same number of protons; larger negative charge means electrons are bound less tightly

Trends in Isoelectronic Species

 Which is larger, Na+^ or Ne?

 Isoelectronic (both contain 10 electrons)

 Have to look at number of protons:

 Na+: 11 protons; Ne: 10 protons

 Since Na+^ has 11 protons, it will pull in its

electrons tighter

 Therefore Ne will be larger

Trends in Isoelectronic Species

 Which has a higher ionization energy, Na+^ or

Ne?

 Isoelectronic (both contain 10 electrons)

 Already established that Ne is larger

 Since Na+^ is smaller, electrons are pulled in

tighter -> harder to remove an electron

 Therefore, Na+^ has the higher ionization

energy

Trends in Isoelectronic Species

 Which has a higher ionization energy, Cl-^ or

Ar?

 Isoelectronic (both contain 10 electrons)

 Already established that Cl-^ is larger

 Since Ar is smaller, electrons are pulled in

tighter -> harder to remove an electron

 Therefore, Ar has the higher ionization

energy

Clicker

 Rank the following isoelectronic species in order from highest to lowest ionization energy.

 Ba2+, Cs+, I-, Te2-, Xe

A) Xe > I-^ > Te2-^ > Ba2+^ > Cs+

B) Ba2+^ > Xe > Cs+^ > Te2-^ > I-

C) Cs+^ > Ba2+^ > Te2-^ > I-^ > Xe

D) Te2-^ > I-^ > Xe > Cs+^ > Ba2+

E) Ba2+^ > Cs+^ > Xe > I-^ > Te2-

Chemical Bonding: Ionic

 Contain ionic bonding between a metal and non-metal

 Electrons are transferred, and ions are

electrostatically held together

Chemical Bonding: Ionic

 Notice that valence shells are filled and that

there are no unpaired electrons in product

 Transferring electrons requires energy (endothermic)

 Example: ionization of magnesium

 Mg  Mg+^  Mg2+  First ionization energy: 737 kJ/mol  Second ionization energy: 1451 kJ/mol

 Takes more energy to remove second electron because atom is now more positively charged (so it holds electrons even tighter)

 If it requires so much energy, then how do ionic compounds form?

 Example: magnesium chloride

 Mg2+^ + 2Cl-^  MgCl 2  Heat released in forming MgCl 2 : 592 kJ/mol

 Energy is released when MgCl 2 forms because the electrostatic interaction between Mg2+^ and Cl-^ ions causes them to become more stable (lower energy)

Bonding: Covalent

 What about bonding in H 2? Is it ionic?

 Not likely; in order to do that, one H atom

would need to lose an electron and one

would need to gain one, then electrostatically

bond

H H H+^ H-

e-

Bonding: Covalent

 So why does H 2 form?  Because we get covalent bonding (sharing of valence electrons)

 Valence electrons spend majority of time in between the two atoms so both hydrogen atoms have a noble gas configuration ( not the cause of bonding, just a way to remember)

 This sharing of electrons lowers the energy between the two nuclei