Honors Chemistry, Chapter 5

1
Chapter 5 Periodic Law
2
Mendeleevs Periodic Table

• In 1869, a Russian chemist, Dmitri Mendeleev
  published the first periodic table
• Mendeleev arranged the elements by properties
  rather than by atomic mass
• His procedure left several empty spaces where
  elements were predicted to fill in when they were
  discovered

3
Mendeleevs Periodic Table
4
Moseley and the Periodic Law

• Henry Moseley, working with Ernest Rutherford,
  examined the x-ray spectra of 38 elements and
  discovered that Mendeleevs order was by charge
  in the nucleus rather than atomic mass.
• This confirmed Mendeleevs principle of chemical
  periodicity.

5
Periodic Law

• The physical and chemical properties of the
  elements are periodic functions of their atomic
  numbers.
• This lead to the formation of our periodic table.

6
Periodic Table

• The periodic table is arranged with elements
  appearing in order of their atomic number so that
  elements with similar properties fall in the same
  column or group.

7
Families of Elements

• The noble gases, group 18.
• The halogens (F, Cl, Br, I, At), group 17.
• The lanthanides (Ce through Lu)
• The actinindes (Th through Lr)

8
Periodicity

• Atomic Difference
  Atomic
• He 2
  3 Li
• Ne 10
  11 Na
• Ar 18
  19 K
• Kr 36
  37 Rb
• Xe 54
  55 Cs
• Rn 86
  87 Fr

8 8 18 18 32
9
Chapter 5, Section 1 Review

• How were Mendeleev and Moseley involved in the
  development of the periodic table?
• Describe the modern periodic table.
• Explain how the periodic law can be used to
  explain the physical and chemical properties of
  elements.
• Explain how the elements belonging to a group are
  interrelated in terms of atomic number.

10
s-Block Elements

• The group 1 and group 2 elements are the s-block.
• The group 1 elements are called the alkali metals
  (H, Li, Na, K, Rb, Cs, Fr).
• The group 2 elements are called the alkaline
  earth metals (Be, Mg, Ca, Sr, Ba, Ra) (note that
  He is NOT in group 2)
• Note that H is a special case and really doesnt
  fit in either group 1 or 17.

11
Sample Problem 5-1

• What is the element with the electron
  configuration Xe6s2 ?
• The highest level with electrons is n6. The
  second element is Ba.
• What is the electron configuration of the group 1
  element in the 3rd period?
• Ne3s1 is sodium.

12
d-Block Elements

• d-Block elements are located in group 3 to group
  12.
• These have properties of metals are called the
  transition elements.
• Although the s level is not always filled, the
  sum of the s and d electrons equals the group
  number.
• Exceptions Pd is Kr 4d10 5s0
• Pt is Xe 4f14 5d9 6s1

13
Sample Problem 5-2

• An element has the electron configuration Kr
  4d5 5s1. What is the period? Block? Group?
  What is the element?
• period 5
• block d
• group 6
• element Mo

14
p-Block Elements

• The p-block elements are in groups 13 through 18
  where the p orbitals are filled.
• The p-block and the s-block are called the
  main-group elements.
• The p-block contains all non-metals, all
  metalloids (B, Si, Ge, As, Sb, and Te) and a few
  metals.
• Group 17 is called the halogens
• Group 18 is called the Noble gases.

15
Group Number, Blocks, Electron Configurations
16
Sample Problem 5-3

• What is the outer electron configuration of the
  period 2 element in group 14?
• He2s2 2p2
• Carbon

17
Sample Problem 5-4

• For the electron configurations given
• Name the block and group
• Identify the element as metal, non-metal or
  metalloid
• Describe it as likely to be high or low
  reactivity
• Xe 4f14 5d9 6s1 c. Ne3s2 3p6
• Ne 3s2 3p5 d. Xe 4f6 6s2

18
Chapter 5, Section 2 Review

• Describe the relationship between electrons in
  the sub-levels and the length of each period of
  the periodic table.
• Locate and name the four blocks of the periodic
  table. Explain the reasons for these names.
• Discuss the relationship between group
  configurations and group numbers.

19
Review continued

• Describe the locations in the periodic table and
  the general properties of the alkali metals, the
  alkaline-earth metals, the halogens, and the
  noble gases.

20
Atomic Radii

• The atomic radius is defined as one-half the
  distance between the nuclei of identical atoms
  that are bonded together.
• The trend to smaller atoms across a period is
  caused by the increasing positive charge.
• In general, the atomic radii of main group
  elements increase down a group.

21
Sample Problem 5-5

• Of the elements Mg, Cl, Na, and P which has the
  largest atomic radius? Why?
• Na Sodium
• Of the elements Ca, Be, Ba, and Sr which has the
  largest atomic radius? Why?
• Ba Barium

22
Ionization Energy

• An ion is an atom or group of bonded atoms that
  has a positive or negative charge.
• Ionization is any process which results in the
  formation of an ion.
• Ionization energy is the energy required to
  remove one electron from a neutral atom of an
  element.
• A energy ? A e-

23
Trends in Ionization Energy

• In general, in main-group elements, ionization
  energy increases across a period due to the
  increase in nuclear charge.
• Among main-group elements, ionization generally
  decreases down the group. Moving down a group
  the electrons are in higher orbitals are
  partially shielded from the nucleus making them
  easier to remove.

24
Removing Electrons from Positive Ions

• The energy to remove the first electron from a
  neutral atom is the first ionization potential,
  IE1
• The energy to remove the second electron is the
  second ionization potential, IE2
• The energy to remove the third electron is the
  third ionization potential, IE3
• Etc.

25
Trends In Ionization Energy

• In general, it is harder to remove the second
  electron than the first because of the stronger
  effective nuclear charge (the nuclear charge
  minus the electron shielding)
• Once an ion reaches a noble gas configuration,
  removal of the next electron requires a much
  higher input of energy.

26
Ionization Energies for Period 3
27
Sample Problem 5-6

• Element A has an ionization energy of 419 KJ/mol.
  Element B has an ionization energy of 1000
  KJ/mol. For each element, is it likely to be in
  the s-block or the p-block? Which element is
  more likely to form a positive ion?
• Element A s-block, likely to form a ion
• Element B p-block

28
Electron Affinity

• Electron affinity is energy change that occurs
  when an electron is acquired by a neutral atom.
• A e- ? A- energy
• In general, electron affinities increase (become
  more negative) from left to right across the
  periodic table.
• With less regularity, electron affinity generally
  decreases down the periodic table.

29
Ionic Radii

• A positive ion is called a cation.
• A negative ion is called an anion.
• Cationic and anionic radii decrease from left to
  right across the periodic table due to the
  increase in nuclear charge.
• In general ionic radii increase down a group
  since the outer electrons are further from the
  nucleus.

30
Valence Electrons

• The electrons in an atom which are available to
  be lost, gained, or shared are called valence
  electrons.
• Valence electrons are often located in
  incompletely filled main-energy levels (s and p
  electrons)

31
Electronegativity

• Electronegativity is an arbitrary scale invented
  by Linus Pauling
• Electronegativity is the measure of how strongly
  an atom attracts the electrons in a molecule.
• Electronegativity increases from left to right on
  the periodic table
• Electronegativity decreases from top to bottom on
  the periodic table

32
Sample Problem 5-7

• Among Ga, Br, and Ca, which has the highest
  elecronegativity? Explain in terms of periodic
  trends.
• All are in the 4th period.
• Br because Br is farthest right.
  Electronegativity increases from left to right.

33
Properties of d-Block Elements

• The properties of d-block elements (all metals)
  vary less and with less regularity than those of
  the main-group elements.
• Atomic radii generally decrease from left to
  right across the periods of the d-block because
  of the increase in nuclear charge.
• d-Block Ionization energies and
  electronegativities generally increase from left
  to right across the periods.

34
Chapter 5, Section 3 Review

• Define atomic and ionic radii, ionization energy,
  electron affinity, and electronegativity.
• Compare the periodic trends of atomic radii,
  ionization energy, and electronegativity, and
  state the reasons for these variations.

35
Review Continued

• Define valence electrons, and state how many are
  present in atoms of each main group element.
• Compare the atomic radii, ionization energies and
  electronegativities of the d-block elements
  with those of the main group elements.