Chapter 5 History of the Periodic Table

1
Chapter 5 History of the Periodic Table
Periodic Trends
2
HISTORY OF THE PERIODIC TABLE

• There are a number of scientists who worked to
  develop our modern periodic table.
• You should thank them, as this tool makes
  chemistry MUCH easier!

3
Some History

• Antoine Lavoisier (c.1989) made the first major
  contribution to the periodic table in his book
  Traité Élémentaire de chimie, and began a
  journal, and Annales de Chimie by suggesting that
  elements were things other than earth, air, water
  and fire
• He listed 33 known elements at that time.

4
Some History

• John Dalton (c. 1803) suggested the first model
  of the atom in his atomic theory (recall chapter
  3)

5
Some History

• Johann Döbereiner (c. 1829) described the idea of
  the triad
• Triads were groups of three elements that shared
  similar properties
• He listed three triads
• Ca, Sr, Ba
• S, Se, Te
• Cl, Br, I

6
Some History

• John Newlands (c. 1864) arranged the existing
  elements by increasing atomic mass
• Noticed that properties reemerge after 8 elements
• Called this the law of octaves
• The Octave states that the properties of the
  elements seem to repeat in a periodic fashion
  after every 8.

7
Some History
8
Some History

• Russien teacher Dimitri Mendeleev (c. 1869)
  proposed the first true version of the periodic
  table
• Mendeleevs table was not complete in that he
  left holes where he thought elements would
  later be discovered.

9
Some History

• For example, Mendeleev proposed that there was an
  element with similar properties as silicon that
  should exist, though it had not been discovered
• This element was called ekasilicon
• Later experiments isolated ekasilicon. It is now
  known as Germanium today.
• Properties of germanium were almost the same as
  predicted

10
(No Transcript)
11
Some History

• English scientist Henry Moseley (c. 1913)
  proposed the modern periodic table
• He arranged the periodic table in order of
  increasing atomic number as opposed to atomic
  mass
• Came up with periodic law
• (def) the physical and chemical properties of the
  elements are periodic functions of their atomic
  numbers

12
Our Periodic Table
Our periodic table can be arranged into three
general sections
13
Metals Nonmetals

• Metals Na, Ca, Li, Cu, Zn, etc.
• Ductile can be drawn into wires
• Malleable Can be hammered into sheets
• Metallic luster
• Good conductors of electricity
• Nonmetals Cl, O, N, S, He
• Brittle
• No luster
• Do not conduct electricity
• Semimetals B, Si, As, Sb
• Have both metallic and nonmetallic properties

14
Important Groups
15
Periodic Trends

• Several trends exist within the periodic table
  which will allow us to predict reactivity of
  various elements
• Atomic size
• Ionic size
• Ionization energy
• Electron affinity
• Electronegativity

16
Atomic Size

• What is the electron configuration of H?
• 1s1
• What is the electron configuration of Li?
• 1s22s1
• What is the electron configuration of Na?
• 1s22s22p63s1
• Notice that with each new row, we add another
  energy level
• This causes atoms to GET BIGGER as we move down

17
Atomic Size/Atomic Radius

• (def) half the distance between the nuclei of
  identical atoms that are bonded together (called
  the Van der Waals radius)

18
Atomic Size (cont.)

• As we move across a period, what are we adding?
• Protons, neutrons and electrons
• Think of the protons and electrons as magnets
  with attractive force
• More protons and electrons more attractive
  force ? more attractive force means that atoms
  get smaller as we move across a period (left to
  right)

19
(No Transcript)
20
Ionic Radius

• Recall, an ion is an atom with either a positive
  or negative charge
• Positively charged ions are called cations
• Negatively charged ions are called anions

• To determine the type of ion an atom will form,
  look at its electron configuration and keep in
  mind the octet rule for valence e-
• Ex. Li1s22s1
• Li has 1 valence electron
• Would it be easier for Li to lose 1 e- or gain 7?
• Li will likely lose 1 electron and become Li
  1s2 (same as He thus stable)

21
Ionic Radius

• How about Magnesium?
• 1s22s22p63s2
• 2 valence electrons
• Easier to lose 2 rather than gain 8
• Mg2 (cation)
• What about F?
• 1s22s22p5
• 7 valence electrons
• Easier to gain 1 than lose 7
• F- (fluoride anion)

• Mg becomes 1s22s22p6 (same as Ne)
• Said to be isoelectric with Neon (thus stable)
• What happened to its valence shell?
• Went from 3rd to 2nd
• Formation of a cation makes a SMALLER ion
• F becomes 1s22s22p6 (isoelectric with Nestable)
• No change in valence shell
• Added another negatively charged particle
• Formation of an anion makes a LARGER ion

22
(No Transcript)
23
Ionization Energies

• (def) the energy required to remove one valence
  electron from a neutral atom is the first
  ionization energy (IE1)
• A energy ? A e-
• (def) the energy required to remove a second
  electron from a 1 cation is the second
  ionization energy (IE2)
• etc
• Usually represented in kJ/mol

24
Ionization Energy Trend

• This trend exists for IE1
• As we move DOWN the periodic table, what happens
  to size?
• Increases
• Is it easier or harder to hold on to an electron
  thats far away?
• Harder, thus the energy required to remove it
  gets lower
• As we move ACROSS (left?right) the periodic
  table, atoms come closer to filling their valence
• Dont want to give up their electrons as easily

25
(No Transcript)
26
Successive Ionization Energies

• Consider Na Would you expect high or low IE?
• Relatively low
• Upon removal of an electron, Na is isoelectric
  with Ne
• Would you expect IE2 for Na to be high or low?
• VERY high because Na is stable

27
Electron Affinity

• (def) the energy change that occurs when an
  electron is acquired by a neutral atom
• A e- ? A- energy
• OR
• A e- energy ? A-
• Values are in kJ/mol (negative because energy is
  released)
• Basically, some atoms want electrons in order
  to fill their valence shells
• F, O, Cl, Br, etc
• Other things dont want electrons because theyd
  rather get rid of them
• Na, Li, Ca, Mg
• So why do the alkali metals have an electron
  affinity whereas the alkaline earth metals dont?
• What about the electron affinities of noble
  gases?

28
(No Transcript)
29
Electronegativity REALLY IMPORTANT!!!

• (def) a measure of the ability of an atom in a
  chemical compounds to attract electrons from
  another atom in the compound
• Electronegativity is abbreviated by c (chi)
• F is the most electronegative of the non-noble
  gas elements
• Fr is the least electronegative atom
• Follows same trend as electron affinity (almost)
• A scale of electronegativity was proposed by
  Linus Pauling (though it is not the only series
  of values)

30
(No Transcript)
31
Electron Sharing
32
In Summary All Trends