Periodic Table- PT

1
Periodic Table- PT
2
History of the Periodic Table

• one of the major milestones in the history of
  chemistry
• helped to predict the existence of elements that
  had yet to be discovered

3
JW Dobereiner

• 1829- about 30 known elements
• Noticed that elements with similar atomic weights
  and their chemical properties
• Classified some elements in triads (sets of
  three)
• Li, Na, K
• Ca, Sr, Ba
• Cl, Br, I
• Look on the current periodic table- notice a
  pattern?

http//www.juntadeandalucia.es
4
JAR Newlands

• 1865 - about 56 elements were known
• Noticed that if elements were arranged in
  increasing atomic mass, the properties of the 8th
  element were like those of the 1st, the 9th like
  those of the 2nd, the 10th like those of the 3rd,
  and so on
• Called this the Law of Octaves (nobody believed
  him) again check the current perodic table notice
  anything?

http//falcon.sbuniv.edu
5
Dimitri Mendeleev

• 1869
• Published his periodic table arranging elements
  by their atomic mass
• Also arranged the table so that elements in the
  same column have similar properties
• Left blank spots in his periodic table
• Predicted the existence of elements yet to be
  discovered

http//www.rit.ac.th
6
Mendeleevs Periodic Table
I began to look about and write down the
elements with their atomic weights and typical
properties, analogous elements and like atomic
weights on separate cards, and this soon
convinced me that the properties of elements are
in periodic dependence upon their atomic
weights. --Mendeleev, Principles of Chemistry,
1905, Vol. II
7

• one of the predicted elements he named ekasilicon
• eka - first
• located below silicon on the periodic table
• 1886 ? germanium was discovered very close to
  the properties that Mendeleev predicted for
  ekasilicon

8
HGJ Moseley

• 1913 HGJ Moseley developed the concept of
  atomic number
• ( he was working in Rutherfords lab who
  discovered protons)
• the correct way to arrange the elements is by
  atomic number NOT atomic mass

http//dbhs.wvusd.k12.ca.us
9

• Mendeleevs table is close because generally
  atomic mass increases with atomic number
• Can you find the exceptions to the rule?
• Periodic Law the physical and chemical
  properties of the elements are periodic functions
  of their atomic numbers
• Today we give credit for the periodic table to
  Mendeleev and Moseley (think MM)

10
The Modern Periodic Table
11

• Elements are found in squares
• Different periodic tables have different
  information in the squares but they may include
  Symbol, name, atomic number, either atomic mass
  or mass number.
• Some periodic tables even include electron
  configurations and oxidation numbers (charge of
  element)
• The squares are then arranged into rows columns
• The rows are called periods
• The columns are called groups or families

12

• 7 periods ?
• 18 groups ?

13
Periods

• Horizontal rows
• Arranged by atomic number
• As you go down the table each period has more and
  more elements
• 1st row has only 2 elements- Hydrogen Helium
• 2nd and 3rd rows have 8 elements- Li?Ne, Na?Ar
• 4th and 5th rows have 18 elements
• How many does the 6th row have?
• 32

14
Groups/Families

• labeling and naming of the groups
• IUPAC method just numbers each group 1-18
• American method- uses numbers 1-8 and the letters
  A B (some PTs use roman numbers I-VII)
• 1-8 A? elements in the S P blocks
• 1-8 B? elements in the D F blocks
• This method is more useful to identifying the
  charges of elements

15

• European Method uses the numbers 1-8 and the
  letters A B only A is for metals B is for non
  metals. (Not very helpful method)
• Hydrogen- can be found in different locations
• is usually not connected to the table because in
  terms of reactivity, it is more like the halogens
  (group 17), but its electron configuration is
  like that of the alkali metals (group 1)

16
Metal

• characteristic luster or shine
• good conductors of heat and electricity
• typically solids at room temperature
• malleable can be hammered into thin sheets
• many are ductile can be pulled into thin wires

http//www.webelements.com
17

• PT also seperates the metals from the non-metals
  this can be done by colors

18
Non-metal

• no luster
• poor conductors
• not malleable or ductile
• many are gases at room temperature, others are
  solids
• Br - liquid
• varied properties
• colored, colorless
• soft solids, hard solids

http//www.webelements.com
19
Semi-metals or metalloids

• some properties of metals and non-metals or
  intermediate
• Si - principle component of computer chips

http//www.webelements.com
20
Electron Configuration and the Periodic table

• electron configuration for the first 3 elements
  in group 1A
• H 1s1
• Li 1s22s1
• Na 1s22s22p63s1
• the highest energy electron is in the s orbital

21

• electrons occupying the highest principle energy
  level are the atoms outermost electrons
• Valence electrons
• located in s and p orbitals
• these valence electrons are the electrons that
  are going to interact with other elements

22
Abbreviated Electron Configuration

• the atoms inner electrons are represented by the
  symbol for the nearest noble
• gas with a lower atomic number
• K - 1s22s22p63s23p64s1
• Ar 4s1
• argons electron configuration
• configuration of electrons after that

23

• H 1s1
• Li He 2s1
• Na Ne 3s1
• K Ar 4s1
• Rb Kr 5s1
• Cs Xe 6s1
• all elements in group 1A have a single valence
  electron in the s orbital
• the principal quantum number of this s orbital
  is the same as the elements period or row number

24
Shape and the Periodic Table

• the shape of the table is related to the electron
  configuration of the elements

http//www.webelements.com
25
s block

• H, He and the elements in group 1A and 2A
• valence electrons are in s orbitals
• the s block is 2 elements wide
• an s orbital can only hold 2 electrons
• the elements in group 1A have 1 valence electron
• the elements in group 2A have 2 valence electrons

26
p block

• group 3A 8A
• the p block is 6 elements wide
• p orbitals can hold a maximum of 6 electrons

27
d block

• takes up most of the middle of the table
• 10 elements wide

28
f block

• the 28 elements below the main body of the
  periodic table

29
http//www.webelements.com
30

• the s and p blocks are also called the
  representative elements or the main block
  elements
• d block is also called the transition metals
• f block is also called the inner transition metals

31
Periodic Trends
32
Periodic Trends

• the properties of elements change in a
  predictable way as you move through the periodic
  table
• nuclear charge
• atomic radius
• shielding
• ionization energy
• electronegativity

33
Nuclear Charge

• as more protons are added to the nucleus the
  charge on the nucleus increases

34
Atomic radius

• radius
• center of nucleus to outermost electrons
• as you move down a group the principal quantum
  number of the outermost electrons increases
• 1A ? 1s1-2s1-3s1-4s1-5s1
• electrons with a larger principal quantum number
  are found in orbitals that extend farther and
  farther from the nucleus which make the atomic
  radius larger

35

• as you move left to right across the table the
  atomic radius decreases
• in any period as you move from left to right, the
  atoms nuclei gain more protons
• more protons
• more positive charge
• a stronger pull is exerted on the electrons in a
  given principal quantum level
• electrons are pulled in closer to the nucleus
• the atom becomes smaller

36

• http//images.encarta.msn.com/xrefmedia/aencmed/ta
  rgets/illus/ilt/1e67a7ad.gif

37
Shielding effect

• comes into play as you go down the group
• as electrons are added to successively higher
  principal energy levels, the innermost electrons
  shield the outer electrons from the pull of the
  positive charge from the nucleus, making the atom
  larger

38
Ionic Size

• when an atom gains or loses an electron, it forms
  an ion
• if it loses an electron
• positive ion (cation)
• it becomes smaller
• 1 less electron reduces the electron electron
  repulsion, allowing the electrons to be pulled
  closer to the nucleus

39

• if it gains an electron
• negative ion
• it becomes larger
• greater number of electrons increases the
  repulsion forces among the electrons
• Group 1A ? generally form 1 ions
• Group 2A ? generally form 2 ions
• Group 7A ? generally form 1- ions
• Group 6A ? generally form 2- ions
• Noble gases ? ?
• do not form ions

40
Ionization Energy

• energy needed to remove an electron
• measured in kJ/mole
• reflects how strongly an atom holds onto its
  outermost electrons

41

• high ionization energy ? holds onto electrons
  very tightly
• low ionization energy ? more likely to lose one
  or more of its outermost electrons
• to ionize a mole of magnesium atoms it takes 738
  kJ of energy

42

• ionization energies decrease as you move down a
  group
• larger atoms ? electrons are held less strongly
  less energy needed to remove an electron

43

• ionization energies increase as you move left to
  right across a period
• smaller atoms ? hold electrons more strongly -
  so more energy is required to remove an electron
• opposite to atomic radius trends

44
(No Transcript)
45
Equations

• Mg (g) ? Mg1(g) e-
• measured in the gas state electrons must be far
  apart for an accurate measurement to be made

46
Successive ionization energies

• energies required to remove electrons beyond the
  first electron
• 1st
• Mg (g) ? Mg1(g) e- 738 kJ/mol
• 2nd
• Mg1 (g) ? Mg2(g) e- 1450 kJ/mol
• 3rd
• Mg2 (g) ? Mg3(g) e- 7730 kJ/mol

47

• the large jump in energy between the 2nd and 3rd
  ionization energy is due to the fact that the
  third electron removed from the magnesium atom is
  part of the noble gas inner core
• these electrons are very difficult to remove

48

• the increase in ionization energy is also due in
  part to the reduced
• electron electron repulsion
• as each electron is removed from the atom the
  remaining electrons are pulled closer and tighter
  to the nucleus

49
Electron Affinity

• the energy change that occurs when an electron is
  gained by an atom
• represented in kJ/mole
• Equation
• F (g) e- ? F- (g) -328 kJ/mol
• fluorine has a negative affinity because energy
  is released when a mole of F atoms gains
  electrons
• exothermic

50

• in general, non-metals have more negative
  electron affinities than do metals
• except noble gases
• positive electron affinities
• electron affinity is related to the number of
  electrons needed to fill its outer energy level

51

• group 7A elements have a very strong electron
  affinity
• by looking at ionization energies and electron
  affinities we can derive an important principle
  about atoms
• Octet rule
• atoms tend to gain, lose or share electrons in
  order to acquire a full set of valence electrons

52
http//tptc.iit.edu/Center/research/PhaseDiagram/C
ontent/periodic20table/electron20affinity.jpg
53
Electronegativity

• a measure of the ability of an atom in a chemical
  compound to attract electrons
• fluorine is the most electronegative element

54

• ?

• atomic radius decreases
• ionization energy increases
• electronegativity increases
• nuclear charge increases
• shielding is constant

55

• ?

• atomic radius increases
• ionization energy decreases
• electronegativity decreases
• nuclear charge increases
• shielding increases

56
(No Transcript)