Transcript Chapter 5: Electrons

Chapter 5
The Periodic Law
Objectives:
 Describe the periodic tables of Moseley and
Mendeleev.
 Identify the various families of elements on the
periodic table.
 State the trends in atomic radius, ionization
energy, electron affinity and ion size with a group
or period on the periodic table.
 Identify the relationship between these trends and
the structure of the atom.
Chapter 5
Section 1
History of the
Periodic Table
By 1860, more than 60 elements had
been discovered.
Chemists needed to learn the properties
of these elements as well as those of the
compounds they formed.
At the time there was no method for
accurately determining an element’s
atomic mass.
Things were disorganized and confusing
among chemists.
Dmitri Mendeleev
Mendeleev hoped to organize the elements
according to their properties.
He arranged the elements in order of
increasing atomic mass and noticed certain
similarities in their chemical properties would
appear at regular intervals.
In 1869 Mendeleev created the first periodic
table grouping the elements with similar
properties.
Henry Moseley
In 1911 Moseley revised the periodic table
of Mendeleev by arranging the elements in
order of increasing atomic number (number
of protons) and not atomic mass.
The elements were still grouped by similar
physical and chemical properties.
Modern Periodic Table
The periodic table has undergone
extensive changes since Mendeleev’s
time.
More than 40 new elements have been
discovered or synthesized.
Periodic Table – an arrangement of the
elements in order of their atomic numbers
so that elements with similar properties fall
in the same column or group.
The elements arranged vertically in the
periodic table share chemical properties.
Example:
F, Cl, Br and I all react in a similar manner.
Later additions to the periodic table
include:
Noble Gases (1868-1900) – group 18
elements that are characterized by their
relative un-reactivity.
The Lanthanides (early 1900’s) – the 14
elements with atomic numbers from 58
(cerium, Ce) to 71 (lutetium, Lu)
The Actinides – the 14 elements with
atomic numbers from 90 (thorium, Th) to
103 (lawrencium, Lr).
The lanthanides and actinides belong in
periods 6 and 7 of the periodic table.
To save space and to group them
together they are set off below the main
portion of the periodic table.
Periodic Table

Elements are classified in three major
groups:




Metals
Nonmetals
Metalloids
Use the periodic table to distinguish the
classes of elements
Physical Properties of Metals






Malleability
Ductility
Luster
Heat conductors and
electrical conductors
Solids
Ex. Iron (Fe), tin (Sn),
zinc (Zn), and copper
(Cu)
Properties of Nonmetals
Dull in appearance
 Brittle
 Do not conduct electricity
 Ex. Carbon (C), Oxygen (O)
and Sulfur (S)
 Solids, liquids or gases

The metalloids divide the metals from the
nonmetals.
They are mostly brittle solids with some
properties of metals and some of nonmetals.
The electrical conductivity falls between the
metals and nonmetals.
The metals of the p block are reactive enough
to be found in nature only as compounds and
not free elements (except bismuth).
Metalloids

Properties of metals and nonmetals

Ex. Silicon (Si) and Germanium (Ge)

Common in the computer
industry
Summary – Section 5.1
Recognize the work of Mendeleev and
Moseley.
How the modern periodic table is
arranged with respect to the elements.
Know the three sets of elements added
to the periodic table after Mendeleev.
Chapter 5
Section 2
Electron Configuration
and the
Periodic Table
In the first period, the 1s sublevel is filled.
The 1s sub level can hold a total of two
electrons.
Therefore, the first period consists of two
elements - hydrogen and helium.
The second period can held 8 electrons and
therefore consists of 8 elements.
The third period is similar to the second period
– 8 elements.
When you get to the fourth and fifth rows
you have to include the transition metals.
There are 10 transition metals in each of
these rows.
The total number of elements in the fourth
and fifth rows are therefore 18 elements.
When you get to the sixth and seventh rows
you have to include the lanthanides and
actinides.
There are 14 elements in each of these
groups.
The total number of elements in the six and
seventh rows can therefore include 32
elements.
The periodic table is divided into 4 blocks
The period of an element can be
determined from the element’s electron
configuration.
For example, arsenic – As, has the
electron configuration of:
[Ar]3d104s24p3
The 4 in 4p3 indicates that arsenic is in
the fourth period of the periodic table.
Without looking at the periodic table give the
period number for the following electron
configuration:
[Xe]6s2
The s-Block Elements
The elements of the s block are chemically
reactive metals.
The Group 1 metals are more reactive than
the Group 2 metals.
The outermost energy level in an atom of the
Group 1 elements contain a single electron.
The ease with which the single electron is lost
make the Group 1 metals extremely reactive.
Alkali Metals – The elements in Group 1 of
the periodic table (lithium, sodium, potassium,
rubidium, cesium and francium).
Because the alkali metals are so reactive they
are not found in nature as free elements.
They combine vigorously with most
nonmetals.
They react strongly with water to produce
hydrogen gas.
Alkaline-Earth Metals – The elements in
Group 2 of the periodic table (beryllium,
magnesium, calcium, strontium, barium and
radium).
Atoms of the alkaline-earth metals contain a
pair of electrons in their outermost shell.
Slightly less reactive than the alkali metals.
Still too reactive to be found in nature as free
elements.
Classwork
Review the practice problems 1 and
2a-d on page 133.
The d-Block Elements
The elements of the d block are known as the
transition metals.
They are good conductors of electricity and
have a high luster.
They are typically less reactive than the alkali
and alkaline-earth metals.
Palladium, platinum and gold are among the
least reactive of all elements.
Classwork
Review the practice problems 1 and
2a-b on page 136.
The p-Block Elements
The elements of the p-block and s-block are
known as the main-group elements.
The p-block consists of the elements in Groups
13-18 (except helium).
The properties of the p-block vary greatly.
The p-block contains metals (Al), metalloids
(Si) and nonmetals (Br).
The electron configuration of bromine is:
[Ar]3d104s24p5
Halogens – the elements of Group 17
(fluorine, chlorine, bromine, iodine and
astatine).
The halogens are the most reactive
nonmetals.
They react vigorously with most metals to
form compounds known as salts (NaCl).
Fluorine and chlorine are gases. Bromine is
a liquid and iodine is a solid.
The Group 18 elements (Noble Gases)
undergo few chemical reactions.
This stability results from the gases
electron configuration.
Their highest occupied levels are
completely filled with electrons (octet).
An atom’s electron configuration governs
the atom’s chemical properties.
Classwork
Review the practice problems 1a-b
and 2a-b on page 138.
Homework
Section Review – page 139
Questions 1, 2, 4 and 5
End of chapter problems – page 156-157
Questions 27, 28 and 29
Due:
Chapter 5
Section 3
Electron Configuration
and the
Periodic Properties
So far you have learned that the elements
are arranged in the periodic table according
to their atomic number.
There is also a correlation between the
arrangement of the elements and their
electronic configuration.
We will look at the relationship between the
electron configurations and the periodic
trends of the elements.
Atomic Radii
Ideally, the size of an atom is defined by
the edge of its orbital.
However, this boundary varies under
different conditions.
One way to express the atomic radius is to
measure the distance between the nuclei
of two identical atoms that are bonded
together, then divide this distance by two.
Atomic radius may be defined as one-half the
distance between the nuclei of identical atoms
that are bonded together.
Period Trends
There is a gradual decrease across a row.
There is a gradual decrease across a row.
The trend to smaller atoms across a period is
caused by increasing positive charge of the
nucleus. Adding of protons or increasing
atomic number.
The electrons are pulled closer to the
nucleus.
The increase pull results in a smaller atomic
radius.
Group Trends
There is an increase down a group.
There is an increase down a group.
As electrons are added to sublevels in
higher energy levels located further from the
nucleus, the size of the atoms increase.
An exception is between aluminum (radius –
143 pm) and gallium (radius – 135 pm).
This is due to the gallium being proceeded
by the 10 d-block elements.
The nuclear charge is considerably higher.
Problem
Of the elements magnesium-Mg, chlorineCl, sodium-Na, and phosphorus-P, which
has the largest atomic radius and why?
Of the elements calcium-Ca, beryllium-Be,
barium-Ba and strontium-Sr, which as the
largest atomic radius and why?
Ionization Energy
Ionization energy (IE) – the energy
required to remove one electron from a
neutral atom of an element.
A + energy
A+ + e-
An ion is an atom or group of bonded
atoms that has a positive or negative
charge.
Period Trends
In general, ionization energies of the
elements increase across a period.
Group 1 elements – have the lowest
ionization energies. Therefore they lose
electrons most easily. Very reactive.
Group 18 elements - have the highest
ionization energies . They do not lose
electrons easily. Very low reactivity.
The increase is due to increasing nuclear
charge (more protons going across a period).
A higher positive charge more strongly
attracts electrons in the same energy level.
Therefore, it is tougher to remove an electron
from an atom.
Increasing nuclear charge is responsible for
both an increasing ionization energy and
decreasing atomic radius across a period.
Group Trends
Ionization energies generally decrease down
a group.
Electrons removed from atoms of the
elements down a group are farther from the
nucleus.
Also, the electrons from the lower energy
levels shield the outer electrons.
Therefore, they are removed more easily.
Classwork
Review practice problems 1and 2 on
page 142.
Review practice problems 1 (a-c) on
page 146.
Electron Affinity
Neutral atoms can also acquire electrons.
Electron Affinity – the energy change that
occurs when an electron is acquired by a
neutral atom.
A + e-
A- + energy
The quantity of energy released is
represented by a negative number.
Period Trends
Among the elements of each period, the
halogens (Group 17) gain electrons most
readily.
The ease with which halogen atoms gain
electrons is a major reason for the high
reactivities of the Group 17 elements.
As electrons add to the same p sublevel of
atoms, electrons affinities increase.
Period Trends
An exception to this trend occurs between
Group 14 and 15.
Compare the electron affinities of carbon
([He] 2s2 2p2 ; -126.3) and nitrogen
[He] 2s2 2p3; 0).
Adding an electron to nitrogen is more
difficult because it forces the electron to pair
with another electron.
Group Trends
Trends for electron affinities within groups
are not as regular as trends for ionization
energies and atomic radii.
Cations and Anions
Cation (A+) – a positive ion that results with
loss of an electron.
Anion (A-) – a negative ion that results with
the gain of an electron.
Valence Electrons
Chemical compounds form because
electrons are lost, gained or shared
between atoms.
The electrons that interact are those in the
highest energy levels.
Valence electrons – the electrons available
to be lost, gained or shared in the
formation of chemical compounds
Valence electrons are often located in
incompletely filled main-energy levels.
For example, the electron lost in Na to form
Na+ is from the 3s sublevel.
For the main group elements, the valence
electrons are the electrons in the outermost
s and p sublevels.
The inner electrons are in filled energy
levels and are held to tightly to be involved.
Electronegativity
Electronegativity is a measure of the ability
of an atom in a chemical compound to
attract electrons.
In many compounds the negative charge of
the valence electrons are not shared
evenly between atoms.
The uneven concentration of charge has
an effect on the chemical properties of a
compound.
Period Trends
Electronegativities tend to increase across a
period.
The Group 1 and 2 elements have the lowest
electronegativities.
Groups 16 and 17 have the highest
electronegativities.
Group Trends
Electronegativities tend to decrease down a
group.
The combination of these trends in
electronegativities results in the highest
values belonging to the elements in the
upper right of the periodic table (fluorine).
Problem
Of the elements gallium-Ga, bromine-Br,
and calcium-Ca, which has the highest
electronegativities and why?
Homework
End of chapter problems – page 157-158
Questions 32, 34, 36 a-b, 37 and 46.
Due: