Download Periodic Table - Doral Academy Preparatory

Ch. 5:
Periodic
Table
C. Goodman, Doral Academy Preparatory High School, 2011-2013
Essential Question: Section 5.1
1. What is the history of the development of the Periodic Table?
2. What is the periodic law, and how can it be used to predict
physical and chemical properties of elements?
3. What is the overall organization of the modern Periodic Table?
1. Three types of elements
2. Named groups
3. Other families
Section 5.1 Vocabulary
• Mendeleev
• Moseley
• Periodic Law
• Period
• Group
• Main group elements
Periodic Table - Definition
Periodic Table
-- an arrangement of the elements in order of
their atomic numbers so that elements with
the same chemical properties are in the same
group (family). Examples: halogens, noble
gases, alkali metals.
Why is it cool?
• http://www.youtube.com/watch?v=u2ogMUD
Baf4&playnext=1&list=PLAC3A0775D813045F
&feature=results_main
History of the Periodic Table I
• Mendeleev: 1869
– Atomic mass
– Repeating (periodic)
patterns of reactivity
•In his favor: predicted
the discovery of Gallium,
which was isolated in his
lifetime
•Certain characteristic
properties of elements
can be foretold from their
atomic weights
•Problem:
Iodine and Tellurium
History of the Periodic Table II
Moseley: 1914
Atomic number
each element has a unique
atomic number; resequenced the
table by electronic charge
(=atomic #)
rather than atomic
weight.
Periodic Law
•In his favor:
solved the “Iodine and
Tellurium” problem
Moseley’s Periodic Law - Definition
• Periodic Law
The physical and chemical properties of the
elements are periodic functions of their atomic
numbers.
http://www.youtube.com/watch?v=OduTDU
GeAXEFind
How to use the periodic table…
Atomic number:
Symbol
Basically the abbreviation
for the element
# of protons in the nucleus of
an atom
Average atomic mass
# of Protons + # of Neutrons
(amu)
Weighted average mass of
isotopes of the element
Remember nuclear
notation for isotopes?
Notice that the atomic
mass is a whole number –
it’s not an average
Also notice the different
locations of the atomic
mass and atomic number.
Groups (families)
The Columns
Elements in
groups have
similar chemical
properties
Periods
The Rows
Elements properties vary
across periods
The length of each period is
determined by the number of
electrons that can occupy the
sublevels being filled in that
period
Important terms
• Main Group Elements
s-block + p-block elements
• Transition metals
d-block elements
• Lanthanides and actinides
f-block elements
• Metalloids, metals, non-metals (see below)
2 Main Sections in Periodic Table
Metals
- Majority of elements
-Good Electrical & Heat Conductors
- Room temperature = most solids
-Contain properties
- Malleability
- Ductility
- High tensile strength
Nonmetals
- Poor Electrical Conductors
- Poor Heat Conductors
- Room temperature = most gases
- One is a liquid at r.t. = Bromine
- Solid nonmetals generally brittle
Names of groups
• Group 1a – Alkali metals
• Group 2a – Alkali earth
metals
• Group 7a (17) – halogens
• Group 8a (18) –
Noble gases
Names of families–
Transition metals
Names of families–
Semiconductors
Section 2: Electron Configuration
1. What is the relationship between the
location of atoms in a group, their electron
configuration, and their chemical and
physical properties?
2. What are the s-, p-, d-, and f-blocks, and how
can their electron configurations of their
elements be determined?
Section 5.2 vocabulary
• Ion
• Valence
• Valence electrons
• s-, p-, d- and f-block elements
Valence Electrons
• Valence = outermost
energy level in which
contains electrons (in
unexcited state).
• Valence electrons are the
electrons on the outermost
energy level of the
element.
• The number of valence
electrons determines the
type of chemical reactions
available to the element!
What does this have to do with
groups?
• All main group elements in a particular group have
the same number of valence electrons.
• Prove it?
• Heh heh heh – that’s your job!
• Write electron configurations, noble gas notation,
of the s and p block elements in the first 5 rows.
Write valence electrons (s/b only) in contrasting
color
• Elements hydrogen – xenon, for columns #1a, 2a,
3a, 4a, 5a, 6a, 7a, 8a
• Huhhh? See next slide
Valence electrons &energy
levels
• Purpose: to determine the relationship between the
group number and number of valence electrons.
• Procedures
1. For each of the elements in the first 5 periods of the following
groups… group (1a, 2a, 3a, 4a, 5a, 6a, 7a, and 8a)
2. Write the name of the element
3. Write the type of element
4. Write the electron configuration, using noble gas notation
• Conclusion (Answer the following question)s:
– 1. How does the group # relate to the number of valence
electrons?
– 2. How do you think the chemical reactivity of the elements in a
particular group, relates to this number of valence electrons?
Valence electrons
& Energy Levels
• Conclusion (Answer the following questions):
– 1. How does the group # relate to the number of valence
electrons?
– 2. How do you think the chemical reactivity of the elements in a
particular group, relates to this number of valence electrons?
Group 14
Element
Electron configuration – Number of
noble gas notation
valence electrons
Carbon
[He]2s22p2
4
Silicon
[Ne]3s23p2
4
Germanium [Ar]4s23d104p2
4
[Kr]5s24d105p2
4
Tin
Valence Electrons
•
•
•
•
Main group elements have characteristic numbers of valence electrons.
Group 1 – 1 valence electron
Group 2 – 2 valence electrons
Groups 13-18
– # valence electrons = Group # - 10
– Example: Group 13 elements have 13-10 = 3 valence electrons
Valence Electrons
• Main group elements have characteristic
numbers of valence electrons.
• S block
– Group 1 – 1 valence electron
– Group 2 – 2 valence electrons
• P block
– Groups 13-18
• Group # - 10
• Example: Group 13 elements have 13-10 = 3 valence
electrons
Relationship Between Periodicity
and Electron Configurations
Sample problem A
1. Without looking at the periodic table,
identify the group, period, and block in which
the element that has the electron
configuration [Xe]6s2 is located.
2. Without looking at the periodic table, write
the electron configuration for the Group 1
element in the third period. Is this element
likely to be more reactive or less reactive
than the element described in (a)?
Sample problem B
An element has the electron configuration
[Kr] 5s24d5. Without looking at the periodic
table, identify the period, block, and group in
which this element is located. Then, consult the
periodic table to identify this element and the
others in its group.
Sample problem C
Without looking at the periodic table, write the
outer electron configuration for the Group 14
element in the second period. Then, use your
periodic table to name the element, and identify
it as a metal, nonmetal, or metalloid.
In book, p. 135
1: Identify period, block, group, element
[Kr]5s2
2. write configuration of…
a. Group 2 elements
b. the group 2 element in the fourth period.
c. the element in the 3rd period, group 15
Sample Problem C Solution
• p-block (group # >12
• 14-10 = 4 electrons in s, p
• 2 e- in s, 2e- in p
• The outer electron configuration is 2s22p2.
• The element is carbon, C, which is a nonmetal.
More practice problems!
Name the block and group in which each of the
following elements is located in the periodic table. Use
the periodic table to name each element. Identify each
element as a metal, nonmetal, or metalloid. Finally,
describe whether each element has high reactivity or
low reactivity.
1. [Xe] 6s24f145d8
2. [Ne]3s23p2
3. [Ne]3s23p5
4. [Xe]4f66s1
How do I know which groups are more
reactive than others?
For main group elements, look at the
number of valence (s/p) electrons
8 valence electrons (noble gases)
= not reactive
Less reactive 4<3<2<1 More reactive
Less reactive 4<5<6<7 More reactive
Chapter 5
Section 2 Electron Configuration
and the Periodic Table
Periods and Blocks of the Periodic Table, continued
• Sample Problem D Solution
a. The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially
filled with nine electrons. Therefore, this element is in the d block.
•
The element is the transition metal platinum, Pt, which is in
Group 10 and has a low reactivity.
•
•
•
b. The incompletely filled p sublevel shows that this element is in the
p block.
A total of seven electrons are in the ns and np sublevels, so this
element is in Group 17, the halogens.
The element is chlorine, Cl, and is highly reactive.
Chapter 5
Section 2 Electron Configuration
and the Periodic Table
Periods and Blocks of the Periodic Table, continued
• Sample Problem D Solution, continued
• c. This element has a noble-gas configuration and thus is in Group 18
in the p block.
•
The element is argon, Ar, which is an unreactive nonmetal and
a noble gas.
• d. The incomplete 4f sublevel shows that the element is in the f
block and is a lanthanide.
•
Group numbers are not assigned to the f block.
•
The element is samarium, Sm. All of the lanthanides are
reactive metals.
Section 3: Periodic Trends
Essential Questions
1. Compare the periodic trends of atomic radii,
ionization energy, electronegativity, and state the
reasons for these variations.
2. What are valence electrons, and how many are
present in atoms of each main-group element?
3. Compare the atomic radii, ionization energies, and
electronegativities of the d-block elements with
those of the main-group elements.
Section 5.3 Vocabulary
•
•
•
•
•
•
•
Atomic radius
Ion
Ionization energy
Cation
Anion
Electron affinity
Electronegativity
Atomic Radii: ½ the distance between
the nuclei of identical atoms bonded
Periodic Trends: Atomic Radii
DECREASES across periods
Because of increasing positive
charge of the nucleus
Holds electrons more tightly
INCREASES down groups
Higher principle quantum
number
Valence electrons in higher
main energy levels
Located farther from the
nucleus
Watch it kid,
I’ve got my ion you.
IONS
Ion – Definition
An Ion is…
An atom or group of bonded atoms,
which has a positive (+) or negative
(-) charge.
There are two kinds of ions…
Cation
•
•
•
•
CRUNCH (subtract) an electron
This results in a positive charge
When an electron is removed,
the atom loses bulk (like a
muscle which shrinks when it
atrophies)
So, the radius of a cation is
smaller than the atomic radius
Anions
•
•
•
•
ADD an electron
This results in a NEGATIVE charge
When an electron is added, the
atom gains bulk (like a muscle
which grows when you work out)
So, the radius of a anion is larger
than the atomic radius
Two sodium atoms bumped into
each other.
One said: "Why do you look
so sad?“
The first one asked "Are
you sure?“
The other responded:
"I lost an electron.“
The other replied "I'm
positive."
Which elements form
which type of ion?
Metal on left tend to form cations (+)
Nonmetals at the upper right tend to form
anions (-)
Hydrogen is a non-metal, but it forms a
cation (+)
Periodic Trends: Ionic Radii – very
similar to trends for atomic radii
Ionization Energy (IE) Definition
Ionization energy is…
The energy required to
remove one electron
from a neutral atom of
an element, forming a
cation.
Note: does not apply to
formation of anions!
Unit of measure: kJ/mol
It takes energy to “steal” electrons
Ionization energy is
(almost always)
positive (J)
Look p. 145, Figure 3.4

Low IE - lose electrons easily

High IE – it’s harder to lose electrons
Ionization Energy:
Atom (neutral) + energy 
Cation (A+)+ (e-) (removed)
Ionization Energy (IE)
 IE increases across periods
because of increasing nuclear charge.
Why?
Higher nuclear positive charge attracts e- in same energy level more
strongly
 Trend down groups
decrease because of more electrons between the
nucleus and furthest out electrons (lower nuclear charge)
Watch out!!!
Ionization energy only applies to the
formation of cations.
If you want to talk anion formation,
you need electron affinity.
Electron Affinity
The energy change that occurs when an
electron is acquired by a neutral atom
Most release energy when they acquire an
electron:
Take a neutral atom (A) add an electron (e-) 
get an anion (A-) + energy is released
Some must be “forced” to gain an electron
by adding energy: A + e- + energy  ACommon unit used is kJ/mol
Some positive affinities are difficult to
determine with any accuracy
Here is a mnemonic for electron
affinity:
If the ion is negatively charged
(anion), the electron affinity is more
strongly negative.
Why do some elements give up
their electrons more easily?
Q: Which group of elements do you think hold on to their
electrons the most strongly?
A: The noble gases! Their valence is full, so it’s very hard to
remove their electrons!
Electron Affinity
Period Trends: Halogens (VII A) gain
electrons the most readily
Group Trends: Electrons add with greater
difficulty down a group
Electron Affinity
• Question
Based on this trend what elements will most
likely form cations?
• Which will most likely form anions?
• Which will most likely have small ionic radii?
Adding Electrons to
Negative Ions
Difficult to add a second electron
to an already negatively charged ion
Second electron affinities are
therefore all positive
Halogens become negative ions by
adding one electron (i.e. Cl-)
Electronegativity
• A measure of the ability of an atom in a
chemical compound to attract
electrons from another atom in the
compound
• Most electronegative element is
Fluorine
Electronegativity
Trend across periods
increase (some exceptions)
More or less, trend down groups
decrease
Similar to IE
Additional Help Website
http://www.hcc.mnscu.edu/chem/stacks.php