Unit 6 The Periodic Table

1
Unit 6 The Periodic Table

• Addison-Wesley textbook End of Chapter 5 and
  end Chapter 14

2
Early History

• End of 1700s 30 elements known
• Technology led to many more elements needed a
  system to organize them
• Early 1800s Dobereiner (1780 1849)
• Proposed triads groups of 3 elements which had
  similar properties (Li, Na, K)

3
Newlands

• John Newlands (1837-1898)
• In 1865 he arranged 62 elements into a table by
  increasing atomic mass
• Called Law of Octaves
• Every 8th element repeated properties
• Compared to music so he was ignored
• Actually very close

4
Mendeleevs Work

• Dimitri Mendeleev (1834-1907)
• Published a table in 1869 (Meyer also, but not as
  much publicity)
• Lined cards of elements up on the floor by
  increasing atomic mass
• Arranged with elements in same columns having the
  same properties

5
Periodicity

• This is behavior of something that is
  predictable, acting in a repeating pattern
• Examples
• Tides, moon, seasons, physiological changes

6
Great Scientist

• Science should predict
• Mendeleev left spaces in his table when he
  thought an element was missing
• This predicted properties of missing elements and
  they were found
• Best science

7
One More Change

• One problem some elements were clearly wrong
  (Argon and Potassium)
• Henry Moseley (1887-1915) studied with Rutherford
• X-ray pattern directly related to charge on
  nucleus called the atomic number
• This straightens out all problems
• Sad story

8
Modern Periodic Law

• When elements are arranged in order of increasing
  atomic number, their physical and chemical
  properties show a periodic pattern

9
Electron Configuration and the Periodic Table

• Look at Li, Na, K, and Rb. Do their
  configurations.
• Underline the outer electrons (highest n value)
• Called valence electrons
• All elements in the same group have the same
  outer electron configuration and therefore the
  same properties

10
Lewis Dot Diagrams

• Show only valence electrons
• Arranged around the symbol
• Four orbitals around the symbol
• S and three p orbitals
• Na. K .
• Try others

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Periods

• Horizontal Row
• Numbered by s and p orbitals that are filling

12
Groups

• Vertical columns
• Numbered either 1 18 or with an A and B
  system

13
Family Names

• Learn (handout)
• http//onsager.bd.psu.edu/jircitano/names.html

14
Look at Hydrogen

• Separate unique
• Sometimes behaves like Group 1, sometimes like
  Group 17

15
Metals

• Brainstorm characteristics
• Most elements are metals
• They have luster or shine
• They are good conductors of heat and electricity
• They are solid
• They are malleable (can be pounded) and ductile
  (can be drawn into thin wire)

16
Chemical definition of Metal

• An element with three or fewer outer electrons
• Forms cations (ions with positive charge) by
  losing electrons
• Examples nickel, iron, copper, sodium

17
Nonmetals

• No luster
• Poor conductors of heat and electricity
• Solids, liquid, and gases
• If solid, brittle or soft

18
Chemical Definition of Nonmetal

• An element which has 5 or more outer electrons
• Forms anions by gaining electrons
• Examples chlorine, argon, oxygen, sulfur

19
Metalloids

• Properties of both
• Often called semiconductors
• Revolutionized the computer industry
• Examples silicon, arsenic

20
Position on Table
21
Element Song

• Tom Lehrer
• http//www.privatehand.com/flash/elements.html

22
Atomic Radius

• Def. distance from center of nucleus to
  outermost electrons
• Trend down a group
• What do you think?
• Size increases because n is increasing
• Trend across a period
• Size decreases because n is constant but
  positive charge of nucleus is increasing, pulling
  in tighter on those electrons

23
Ionization Energy

• Def. Energy needed to remove an electron from
  an atom (can be 1st, 2nd, etc.)
• Trends are opposite of size, lower energy for
  larger atoms
• Trend down a group first ionization energy
  decreases
• Trend across a period first ionization energy
  increases

24
Other Ionization Energies

• Complicated, depend upon which electron comes
  next.
• Example, electrons in carbon

25
Electron Affinity

• Amount of additional energy released when
  another electron is attracted to a neutral atom
• Same trend as ionization energy smaller atoms
  are better at attracting electrons and have
  greater electron affinity
• Decreases in a group
• Increases in a period

26
Electronegativity

• Rating system 0 to 4 which shows amount of
  attraction an element has for other electrons
  when it is already bonded.
• Again, same trends as ionization energy and
  electron affinity

27
Octet Rule

• Atoms lose or gain electrons to become stable (or
  have 8 outer electrons)

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Size of ions

• Atom loses electrons to form a cation, or
  positive ion.
• Why?
• Same nuclear charge, less electrons, smaller!
• Atoms gain electrons to form an anion. Or
  negative ion.
• Why?
• Same nuclear charge, more electrons, larger!

29
Riddle

• Ion Oh no! Ive lost my electron.
• Atom Are you sure?
• Ion Yes, Im positive!