Draw the periodic table and label the electron blocks and areas
of non-metals, metals, and metalloids. Relate the Lewis dot
structure to its place in the periodic table. Explain periodic
trends as one moves along periods and down groups in the periodic
table 2
Slide 3
Chapter 6.1-6.2 Periodic Law Group Period Representative
Element Transition Element Metal Alkali Metal Alkaline Earth Metal
Transition Metal Inner Transition Metal Lanthanide Series Actinide
Series Nonmetal Halogen Noble Gas Metalloid 3
Slide 4
Dmitri Mendeleev noticed in his table that there were
repetitions of physical and chemical properties when the elements
were arranged by atomic mass. 4
Slide 5
5
Slide 6
PropertyPredicted (1869) Atomic Mass72 u ColorDark gray
Density5.5 g/mL Melting PointHigh Density of Oxide4.7 g/mL Oxide
solubility in HCl Slightly dissolved by HCl Formula of chloride
EsCl 4 Properties of Germanium (Ge) 6 Actual (1886)
Slide 7
Periodic Law states that chemical and physical properties
repeat in regular cyclic patterns when they are arranged by
increasing atomic number. Starts with metals at left and goes to
non-metal (noble gas) on right Properties change in orderly
progression across a period. 7
Slide 8
8
Slide 9
Periodic Table Periods 9 Transition Elements Inner Transition
Elements Halogens Noble Gases Alkali Metals Alkaline Earth Metals
Representative Elements Columns, Groups or Families Metals
MetalloidsNonmetals
Slide 10
What are some of the elemental properties that make the
periodic table, well, periodic? Classification by metals, nonmetals
and metalloids Metals - shiny ductile, malleable solids, good
conductors of heat and electricity Nonmetals - dull, brittle
solids; or gas, poor conductors of heat and electricity Metalloids
- have chemical and physical properties of both metals and
nonmetals 10
Slide 11
Representative Elements (Sometimes called A Group) Group # =
number of valence electrons Means similar Lewis dot structure and
similar properties. s-block s-block elements have 1-2 electrons in
s-orbital p-block p-block elements have 1-6 electrons in p-orbitals
Noble gases have filled valence shells Energy level of valence
electrons is at energy level given by period (row) number 11
Slide 12
Transition Elements (Sometimes called B Group) d-block d-block
elements have 1-10 electrons in d- orbitals Columns 3-12 in
periodic table Energy level of valence electrons at n and partially
filled n-1 d orbitals (example: 4s and 3d) f-block f-block
(Lanthanides and Actinides) have 1-14 electrons in f-orbitals
12
Slide 13
Fill in the missing info for the following elements: Identify
the element fitting the description. a) Group 2 (2A) element in 4
th period: b) Noble gas in 5 th period: c) Group 12 (2B) element in
4 th period: d) Group 16 (6A) element in 2 nd period: 13
ConfigurationGroupPeriodBlock [Ne]3s 2 [He]2s 1 [Kr]5s 2 4d 10 5p 5
7 (7B)4
Slide 14
14
Slide 15
Effective Nuclear Charge (Z*) Not in book! Shielding (Not in
book) Ion Ionization Energy Octet Rule Metallic Character (Not in
book) Electronegativity 15
Slide 16
Atomic and ionic size Ionization energy Electronegativity
Metallic Character 16 Higher effective nuclear charge Electrons
held more tightly Larger orbitals. Electrons held less
tightly.
Slide 17
Z* is the nuclear charge experienced by the outermost
electrons. (Note: not in book!) Z* increases across a period owing
to shielding by inner electrons. Shielding is blocking by inner
electrons. For a period (row), the number of shielding electrons
remain the same, but the number of protons in the nucleus
increases. Example: All elements in the second period have the same
underlying [He] noble gas configuration. However, the number of
protons increase from left to right. 17
Slide 18
So we can estimate as Z* = [ Z - (no. inner electrons) ] or or
Z* = Z S (inner electrons) Z is total number of electrons S is the
number of electrons blocking the valence shell electrons, the
underlying noble gas electrons. Charge felt by 2s e - in Li Z* = 3
- 2 = 1 Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3and so on! 18
Slide 19
19 Orbital energies drop as Z* increases
Slide 20
Atomic size is a periodic trend influenced by electron
configuration. For metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the element. 20
Slide 21
For other elements, the atomic radius is half the distance
between nuclei of identical atoms that are bonded together. 21
Slide 22
22
Slide 23
Size (radius) goes UP on going down a group. See previous
slide. Because electrons are further from the nucleus, there is
less attraction. Size (radius) goes DOWN on going across a period.
Size (radius) goes UP on going down a group. See previous slide.
Because electrons are further from the nucleus, there is less
attraction. Size (radius) goes DOWN on going across a period.
23
Slide 24
Size (radius) decreases across a period owing to increase in
Z*. Each added electron feels a greater and greater positive
charge. Note: Electrons in the same energy level dont shield each
other too much. 24 Large Small Increase in Z*
Slide 25
25
Slide 26
The radius of an atom when it has become an ion. An ion is an
atom or bonded group of atoms that has an overall positive or
negative charge. An atom acquires a positive charge by losing
electrons or negative charge by gaining electrons!! 26
Slide 27
To form positive ions from elements remove 1 or more e- from
subshell of highest n [or highest (n + l)]. Al: [Ne] 3s 2 3p 1 - 3e
- Al 3+ : [Ne] 3s 0 3p 0 27
Slide 28
28 Atoms tend to gain, lose, or share electrons to get 8
valence electrons (except small atoms up to Boron)
Slide 29
1.Write the electron configuration and orbital box diagram for
Mg when it is an ion. Hints: What is its noble gas configuration?
What will they do to get an octet? 2.Write the electron
configuration and orbital box diagram for O when it is an ion.
29
Slide 30
Positive ions are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.
Electron Configuration as ion is: [He] 2s 0 30 Li,152 pm 3e and 3p
Li +, 78 pm 2e and 3 p + Forming a positive ion.
Slide 31
Negative ions are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes. Electron
configuration as ion: 1s 2 2s 2 2p 6 (just like neon.) 31 Forming a
negative ion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -
Slide 32
32 See Figure 6-14
Slide 33
Why do metals lose electrons in their reactions? Why does Mg
form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons?
Why do metals lose electrons in their reactions? Why does Mg form
Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons?
33
Slide 34
IE = energy required to remove an electron from an atom in the
gas phase. 34 Mg (g) + 738 kJ Mg + (g) + e-
Slide 35
35 Mg + (g) + 1451 kJ Mg 2+ (g) + e- Mg + has 12 protons and
only 11 electrons. Therefore, IE for Mg + > Mg. IE = energy
required to remove an electron from an atom in the gas phase.
Slide 36
1 st : Mg (g) + 735 kJ Mg + (g) + e- 2 nd : Mg + (g) + 1451 kJ
Mg 2+ (g) + e- 36 3 rd : Mg 2+ (g) + 7733 kJ Mg 3+ (g) + e- Energy
cost is very high to dip into a shell of lower n.
Slide 37
37
Slide 38
38
Slide 39
39 As Z* increases, orbital energies drop and IE
increases.
Slide 40
40
Slide 41
IE increases across a period because Z* increases. Metals lose
electrons more easily than nonmetals. Nonmetals lose electrons with
difficulty. 41 High ionization energy: atoms want to hold on to
electrons; likely to form negative ion Low ionization energy: atom
gives up electron easily; likely to form positive ion
Slide 42
IE decreases down a group Because size increases. Ability to
lose electrons generally increases down the periodic table. See
reactions of Li, Na, K 42
Slide 43
Which element in each pair has the larger 1 st ionization
energy? A. Na or Al B. Ar or Xe C. Ba or Mg 43
Slide 44
44 Lithium SodiumPotassium
Slide 45
*Note: metallic character not in book. An element with metallic
character is one that loses electrons easily. Metallic character:
is more prevalent in metals on left side of periodic table is less
for nonmetals on right side of periodic table that do not lose
electrons easily 45
Slide 46
46
Slide 47
Relative ability of an element to attract electrons in a
chemical bond. Ionization energy reflects ability of atom to
attract electrons in an isolated atom Generally, the higher the
ionization energy of an atom, the more electronegative the atom
will be in a molecule There are many electro negativity scales well
use the one by Linus Pauling (values dimensionless) Will be used to
determine things like polarity of a chemical bond. 47
Slide 48
48
Slide 49
Decreases down a group Why? Due to greater atomic radius
Increases across a period Why? Increased positive charge in nucleus
(Greater Z*) Same trend as for ionization energy. Surprised?
49
Slide 50
Moving Left Right (periods) Z * Increases Atomic & ionic
Radius Decrease Ionization Energy Increases Electronegativity
Increases Metallic Character Decreases Moving Top Bottom (groups)
Z* is roughly constant, but val e - distance increases Atomic &
Ionic Radius Increase Ionization Energy Decreases Electronegativity
Decreases Metallic Character Increases 50
Slide 51
a)Electronegativity b)Ionic Radius c)Atomic Radius d)Ionization
Energy e)Metallic character 51