Periodic Properties of the Elements

1
Periodic Properties of the Elements

• Chapter 7 AP Chemistry

2
I History and Development

• Elements in pure form like Gold and Silver were
  known thousands of years ago
• 31 elements were known in 1800 by 1865 the number
  had more than doubled to 63.
• The increase lead to the grouping and
  classification of the elements

3
First Periodic Law

• Properties of elements are a periodic function of
  atomic weight
• Although Meyer also grouped elements like
  Mendeleev
• Mendeleev was given more credit because he left
  blanks for un-discovered elements

• Mendeleev

4
Modern periodic law

• Henry Mosley

• In 1913, by using x-ray spectra obtained by
  diffraction in crystals, he found a systematic
  relation between wavelength and atomic number,
• Using Atomic number Mosley rearranged the
  elements noting that the elements were a periodic
  function of atomic number

5
Design of the Table

• Rows and Columns
• Horizontal rows series or periods numbered
  1-7 Notes the shell or principle QN
• Vertical columns group or a family
  numbered 1-18 or 1A-8A, 1B-8B
• Review electron configuration notation

6
Periodic Trends

• Size of Atoms
• Within each group the atomic radius tends to
  increase as you go down the group the principle
  QN increases size of the cloud gets larger
• Within a period the atomic radius decreases as
  you go from left to right the effective nuclear
  charge increases the number of protons
  increases while the number of core electrons
  stays the same the outer electrons dont shied
  each other well

7
Ionization Energy (I)

• I the energy required to remove the outermost
  electron from a gaseous atom or ion
• I1 the energy to remove one electron I2 the
  second electron
• Metals tend to give up e- to obtain a Nobel gas
  configuration

8
Ionization Energy Trends

• Ist Ionization energy generally follows effective
  nuclear charge increases from left to right in
  periods and from bottom to top in groups
• Except from group 2A to 3A there is a slight
  increase the second e- in the S sub-shell is
  harder to remove that the 1st P e-
• GR 5A to 6A there is a slight increase due to
  repulsion of paired e-s in the P4 configuration
• Every element shows large increase in IE when es
  are removed from a Nobel gas core

9
Electron Affinities

• Is the energy change of the reaction of adding an
  electron to a gaseous atom or ion.
• Tends to be an exothermic processes (Delta H is
  neg) although in some cases it is positive or
  endothermic
• In general electron affinity tends to ( become
  more negative) from left to right in periods.
  There is very little change going down a group in
  the value.

10
Properties of Metal, Non-Metals Metalloids

• Metals conduct heat and electricity are
  lustrous, malleable, and ductile
• Tend to lose es to become cations
• Form basic oxides metal oxides react with water
  to produce basic or alkaline solutions
• Ex 7.7,7.8 in text

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Metal Properties Cont.

• Metal oxides an acid yield a salt and water
• Ex 7.9,7.10 in the text

12
Non-Metal Properties

• Poor conductors of heat and electricity are often
  dull in color and shatter if forged
• Nonmetals tend to gain electrons form other sub
  and become anions
• Non-metals from acidic oxides non-metal oxide
  react with water to form acidic solutions
• Ex 7-12, 7-13 in text
• Non-metal oxides a base yield a salt plus water
  ex 7.14, 715 in text.

13
Mettaloids

• Boundary between metals and non-metals
• Mettaloids can either gain or lose electrons
• Al and Po are metals not mettaloids
• Properties between metals and non-metals
• Several are semiconductors and re important
  because of their use in circuits and computer
  chips
• Metallic character increases from left to right
  and from top to bottom in a group

14
Group Trends

• Alkali Metals
• Soft, gray metals
• Low Ionization energies lose es easily form
  1 cations very reactive
• Electrolysis is a technique to force an electron
  back on a cation to produce a neutral metal the
  metals can be obtained

15
Alkali Metals

• by passing an electrical current through a
• molten salt.
• The metals combine with hydrogen to from hydrides
  ex. 7.19, or with Sulfur to form Sulfides ex.
  7.20 or with chlorine to form chlorides. ex 7.21
• The alkali metals react with water to produce
  hydrogen gas and hydroxides ex 7.22.
• See examples of reactions with oxygen ex 7.23-7.25

16
Alkaline Earth Metals

• Slightly harder and more dense than Alkali metals
• 1st ionization energies are low but not a s low
  as alkali metals form 2 cations
• Reactivity tends to increase as you progress down
  a group.
• See examples of reactions 7.26 7-29

• Got Milk?

17
Trends in Selected Non-metals

• Hydrogen
• Usually placed in grp 1A because of its 1s
  electron config. Although a non-metal
• It reacts with other non-metals to form molecular
  cmps ex 7.30-32
• Reacts with metals to form hydrides 7.31-32

18
Oxygen Family

• Non-metallic top of family metallic at bottom
• O, S, Se are typical non-metals O and S are
  allotropes molecules of elements with different
  structures ( O2, O3, S2, S4, S6, S8)
• Polonium is a rare and radioactive element
• See examples of reactions 7.33-36

19
The Halogens

• All non-metals (astatine is radioactive seldom
  discussed)
• All are diatomic in gas phase
• Bromine normally a liquid iodine is normally a
  solid
• See reactions 7.37-42

20
Nobel Gases

• Mostly non-reactive
• Some compounds of Xenon and Krypton are known
• All monatomic gases
• Highest ionization energies of any family

• xenon tetrafluoride