Download The Periodic Table - Ms. Simmons

The Periodic Table
The Picture Worth a Thousand Words
The Periodic Table
Function: organizational tool which classifies
the over 100 elements in a systematic way;
allows for easy comparison of the elements
and provides substantial information by
illustrating patterns
Imagine reading the content of the periodic
table in dictionary format
Organization of the Periodic Table
The periodic table visually classifies elements by
multiple traits
• number of protons (and electrons)
• groups and periods
• metals, nonmetals, and metalloids
• main-group elements (representative elements),
transition metals, inner transition metals
• s, p, d, f blocks
• valence electrons and electron configurations
• periodic trends
Key to the Periodic Table
Every element is represented by a box
containing the atomic number, the element
symbol, the element name, and the atomic
mass of the element
Overview of Periodic Table
Elements are arranged by increasing atomic
number in rows known as periods and
columns known as groups (or families)
Metals, metalloids, and nonmetals
Metals, Nonmetals, and Metalloids
• Classification as metals, nonmetals or
metalloids is based on specific properties
• Most important to this classification is
ability to conduct electricity or heat
• Metals are good conductors
• Nonmetals are poor conductors
• Metalloids are intermediate in nature
Metals, nonmetals, and metalloids
Property
1
Metals
Metalloids
Nonmetals
ability to conduct electric
good conductors of
poor conductors of
current varies with
electric current and heat
electric current and heat
temperature
often gases at room
temperature; when solids
typically brittle
2
usually solid at room
temperature (except Hg)
3
most are malleable
have properties of metals
and nonmetals; when
4
most metals are ductile
conductors they do not
conduct heat or
5
6
have low boiling points
reactivity varies; some electricity as well as the reactivity varies; some
highly reactive, others do
highly reactive, some are
metals
not react easily
extremely unreactive
usually have luster
(shiny)
when solids typically
dull-looking
Properties of Metals
Good conductors of electricity and heat
burner on stove is made of metal
handles on cookware are not metal
Usually solids at room temperature (often hard)
exception is Hg (mercury) – liquid
Most are malleable – can be formed into sheets/foils
Most are ductile – can be drawn into wires
Reactivity varies (most react with acids)
some are highly reactive
some do not react easily
Usually shiny (luster)
Au (gold)
Ag (silver)
Properties of Nonmetals
Poor conductors of electricity and heat
Often gases at room temperature
Have low boiling points
reason they are gases at room temperature
Reactivity varies
some are highly reactive – halogens
some are highly unreactive – noble gases
If solids they are typically dull-looking
Properties of Metalloids
Metalloids have properties of metals and
nonmetals; their properties tend to fall
between metals and nonmetals
• do not conduct heat or electricity as well
as the metals
• properties vary widely
Ability to conduct electricity varies with
temperature
semiconductors (used in electronics)
used in computer chips and solar cells
Metals
Shiny
Solid
Be
Nonmetals
Metalloids
dull solids
liquid/gas
varied
solid
C
Si
Ca
Ba
Bi
Os
P
Ge
S
Main Group Elements and Others
main group (representative elements): the s and p
blocks
transition metals: the d block
inner transition metals: the f block
Main Group Elements
• most clearly illustrate periodic patterns
• most abundant naturally occurring elements
• most important elements for living things,
including most common components of
organisms: O, C, N, H, P, Ca
Transition Metals
Metals (with the properties of metals)
Can have multiple oxidation states
d orbitals allow them to behave differently,
and to be less predictable
Form complexes with water
Inner Transition Metals
Display very similar reactivity to each other
Actinides are all radioactive
unstable
s, p, d, and f Blocks
Special Groups
Group = column
Group members typically have similar
properties
Four groups have special names:
Alkali metals
highly reactive
Alkaline earth metals
reactive
Halogens
highly reactive
Noble gases
highly unreactive
A
L
K
A
L
I
N
M E
E
T E
A A
L R
S T
H
A
L
K
A
L
I
Special Groups
N
O
B
L
H
E
A
L G
O A
G S
E E
N S
S
Versions of the Periodic Table
• Updated with new elements
• Updated with more accurate atomic masses
• Presented with labels and color-coding
• Different arrangements for different purposes
Versions of the Periodic Table
• Different versions emphasize different features
• They are all correct, and none is absolute
Origins of the Periodic Table
Mendeleev: created one of the earliest
versions of the periodic table
• Arranged elements by increasing mass
(right or wrong?)
• Grouped elements with similar properties
Mendeleev’s 1872 Periodic Table
Predicted undiscovered elements and their properties
Mendeleev’s Contribution
Mendeleev developed Periodic Law
(Modern) Periodic Law: the elements, when
listed in order of their atomic numbers, fall
into recurring groups, so that elements with
similar properties occur at regular intervals
Periodic law is observed by all members of a
column having the same ending to their
electron configuration
The Big Picture
• Periodic Table shows patterns
• Multiple types of patterns can be found
Periodic repetition of chemical and physical
properties of the elements as you move across
a period (row)
Or down a group (column)
This trait of periodic repetition of the
properties of elements when they are
arranged by increasing atomic number is
known as periodic law.
Electron Configurations
Pattern across a period and
pattern down a group
Periodic Trends
Certain properties show a systematic
variation across a period and down a group:
• metallic character
• valence electrons
• atomic radii (atomic size)
• ionic radii
• electronegativity
• ionization energy
Metallic Character
Valence Electrons
Alkali metals
Halogens
Valence Electrons
Equal number of valence electrons for all
members of a group
For main-group elements (IA-VIIIA) the
valence electrons are equal to the group number
Exception: He
(He only has 2 e-)
Atomic Radius
As protons are added across a period, the nucleus
has a larger positive charge
Na +11
Al +13
S +16
Ar +18
As electrons are added across a period, they
continue to fill the same energy level, n
The result?
Larger effective nuclear charge on the electrons
from left to right across a period
atomic radius increases
Atomic Radius
atomic radius increases
Ionic Radii
Na+
Mg2+
Al3+
P3-
S2-
Cl-
no of protons
11
12
13
15
16
17
electronic
structure of
ion**
2,8
2,8
2,8
2,8,8
2,8,8
2,8,8
ionic radius
(nm)
0.102
0.072
0.054
0.212
0.184
0.181
**How many electrons does each have?
Na+ has 10, Mg2+ has 10, Al3+ has ?
Cl- has 18, S2- has 18, P3- has ?
Ionic Radius
The positive ions (Cations)
Each ion has exactly the same electronic
structure, and is equal to the structure of Ne,
with n = 2.
However, the number of protons for each ion
is increasing. That causes a larger effective
nuclear charge that pulls the electrons towards
the nucleus and causes the ionic radii to
decrease.
Ionic Radius
The negative ions (Anions)
Each ion has the same electronic structure
again, but this time equal to the structure of
Ar with n = 3.
Compared to the cations, there is an extra
layer of electrons, so the ionic radii increase.
Ionic Radii
Cations become smaller than their atoms
Anions become larger than their atoms
Electronegativity
Electronegativity: a measure of the ability of
an atom in a compound to attract electrons
More electronegative elements are more able to
pull electrons from another atom towards their
own nucleus
Electronegativity
Fluorine is the MOST electronegative element
Electronegativity
Think: Are Noble Gases electronegative?
What is their electron configuration?
How does this impact tendency to gain or
lose electrons?
Noble gases DO NOT follow the trend for
electronegativity
Ionization Energy
Ionization energy: the energy required to remove
an electron from an atom in the gaseous state
Remember, atomic radius decreases across a period
due to greater effective nuclear charge
Electrons are closer to the nucleus and harder to
remove
From top to bottom within a group, the valence
electrons are in a higher energy, farther shell
This results in electron shielding, making it
easier to remove electrons
Ionization Energy
First Ionization Energies
Ionization Energy
Atoms can lose more than one electron
Always requires more energy to remove a
second electron (increased effective nuclear
charge)
Peaks change element
WHY?
Subsequent Ionization Energies
Each electron removed changes the peak
Electron Affinity
The energy given off when a neutral atom in
the gas phase gains an extra electron to form a
negatively charged ion.
F(g) + e-
F-(g)
Ho = -328.0 kJ/mol
The negative sign of Ho indicates that energy
is given off. This process is exothermic.
Electron Affinity
Electron affinity is more or less the opposite of
ionization energy
Halogens: greatest electron affinity
Noble gases: lowest electron affinity (requires E)
Learning Check
How do electron affinity and ionization energy
compare?
F
electron affinity
F
ionization energy
Learning Check
How do electron affinity and ionization energy
compare?
F + eF
FF+ + e-
electron affinity
ionization energy
Electron affinity is the energy required to gain
an electron. Ionization energy is the energy
required to lose an electron.