Download Enriched Chemistry Chapter 5 * The Periodic Law

Enriched Chemistry
Chapter 5 – The Periodic Law
Geneva High School
Section 5-1 History of the Periodic Table
• By 1860, more than 60 elements had been discovered, yet chemists
had no way of organizing them or agreeing on basic properties of the
element.
Mendeleev’s periodic table grouped elements by their properties.
• Mendeleev placed the name of each known element on a card, with the
atomic mass and the observed phys/chem properties
• He then arranged them in various configurations, looking for patterns or
trends
• He noticed that when the elements were arranged in order of increasing
atomic mass, certain similarities appeared at regular intervals.
• He published his periodic table in 1869
• His P.T. left several empty spaces, but he
predicted the properties of the elements
that would fill three of these spaces.
• By 1886, all three elements had been
discovered.
• The success of his predictions persuaded
most chemists to accept his periodic table
and earned him credit as the discoverer of
the periodic table.
• periodic – a regular, repeating pattern
Eka-aluminum (now known as gallium)
Three truths and a lie about Mendeleev (You do
not need to copy this!)
Which of the following is NOT true?
A. Mendeleev only cut his hair once a year.
B. He had two cats, which he named after the first
two elements he discovered.
C. He was in charge of measuring the amount of
alcohol in vodka in Russia.
D. There is a crater named after him on the moon.
Henry Mosely arranged elements by their atomic numbers.
 Mosely discovered that the elements fit better when
arranged in increasing order by their atomic number
What elements does this effect?
 Mendeleev’s principle is known as the periodic law: the
physical and chemical properties of the elements are
periodic functions of their atomic numbers.
 This means that their properties repeat themselves based
on their atomic numbers
Modern periodic tables arrange the elements by both
atomic number and properties.
• The P.T. has changed over time, and new elements have been
discovered and synthesized
• synthesized – man made
• However, they can all be placed in a group of other elements
with similar properties
• The periodic table is an arrangement of the elements in order of
their atomic numbers so that elements with similar properties
fall in the same column, or group.
Periodic Table Structure
Rows are called
periods
Columns are called
groups or families
Significant periodic table discoveries
The Noble Gases
• In the late 1800s, the noble gases were discovered and placed in their own
group
• They had escaped detection because of their lack of reactivity
The Lanthanides and Actinides
• Lanthanides - 14 elements (58-71) that are all very similar in physical and
chemical properties which made them difficult to isolate
• Actinides – 14 elements (90 – 103)
• To save space, they are set off below the main portion of the periodic table
Periodicity
• The differences between the atomic numbers of successive elements
in a group are 8, 8, 18, 18, 32
• The reason for periodicity is explained by the arrangement of the
electrons around the nucleus.
Chapter 5
Section 2 – Electron Configuration and the Periodic Table
(aka Groups of the Periodic Table)
Generally, the electron configuration of an atom’s highest occupied
energy level governs the atom’s chemical properties.
Sodium atom – notice
This little guy is
the only
electron that
matters in
terms of
reactivity of
Francium!
how much closer the
outer shell electron is
to the positively
charged nucleus.
Main Idea – The period of an element is determined by its electron
configuration.
• Elements are arranged vertically in groups (or families) that share
similar chemical props
• They are also organized into rows, or periods
• The period of an element can be determined from the element’s
electron configuration
• ex: [Ar]3d104s24p3
• The P.T. can be divided into four blocks, the s, p, d, and f blocks
The s-Block Elements: Groups 1 and 2
• Elements of the s-block are chemically reactive
• Group 1 elements are known as the alkali metals
•
•
•
•
1 valence electron (v.e.)
silvery appearance and soft
never found in nature as free elements
combine vigorously with nonmetals
• Group 2 elements are known as the alkaline earth metals
•
•
•
•
2 v.e.
harder, denser, and stronger than the alkali metals
less reactive than alkali metals
never found as free elements
https://www.youtube.com/watch?v=cqeVEFFzz7E
Hydrogen and Helium
• despite having a 1s1 config, hydrogen does not share the same
properties as group 1 elements
• helium is non-reactive because its outer energy is filled with 2
electrons
The d-Block elements: groups 3-12 (aka the Transition Metals)
• The d-block elements are metals w/typical metallic props and are
often referred to as transition elements
• good conductors of electricity and have a high luster
• less reactive than groups 1 and 2
The p-Block elements: groups 13-18
• The p-block elements together with the s-block elements are called
the main-group elements
• props of p-block elements vary greatly
• includes all of the nonmetals except H and He, and the six metalloids
• The families are known by the first element in each family, so the
Boron Family, Carbon Family, Nitrogen Family and Oxygen Family.
• The elements of Group 17 are known as the halogens
• most reactive nonmetals, reacting vigorously with most metals to form salts
• this is due to their 7 v.e., one electron short of the stable noble-gas
configuration
• rarely found as free elements in nature
• Fluorine is the most reactive nonmetal.
Metalloids
• Also called semiconducting elements
• Mostly brittle solids with some properties of metals and some of
nonmetals
• They will conduct electricity, but a higher voltage is necessary than
with true conductors
• Used in transistors, LEDs, photovoltaic cells
Chapter 5
Section 3 – Electron Configuration and Periodic Properties
Main Idea – Atomic radii are related to electron configuration.
• The boundaries of an atom are fuzzy, making it difficult to describe its
size.
• One way to express an atom’s radius is to measure the distance
between the nuclei of two identical atoms that are bonded together
and then divide by two, giving the atomic radius.
Atomic radius – defined as one-half the distance between the nuclei of
identical atoms that are bonded together.
Atomic radii…( cont)
• Periodic trends
 Atomic radii decrease as you
go across a period.
 This is caused by the increasing
charge of the nucleus, which
gains a proton from one
element to the next.
 The attraction of the nucleus is
somewhat offset by repulsion
among the increased number
of electrons.
• Group trends
 The atomic radii of the maingroup elements increase down
a group.
• Main idea – Removing electrons from atoms to form ions requires
energy.
A + energy
A+ + eElement A+ represents an ion, which is an atom or group of bonded
atoms that has a positive or negative charge.
Any process that results in the formation of an ion is referred to as
ionization.
To compare the ease with which atoms of different elements give up
electrons, chemists compare ionization energies.
 the energy required to remove one electron from a neutral atom.
Ionization Energy
• Period Trends
In general, ionization energies of the main-group elements increase across
each period.
This is caused by an increasing nuclear charge.
• Group Trends
Among the main-group elements, ionization energies general decrease down
the groups.
This is because as you go down a group, there are more energy levels, and the
electrons are further from the nucleus, making them easier to remove.
Also, the electrons between the nucleus and the outer electrons have a
shielding effect.
• Removing Electrons from Positive Ions
Electrons can be removed from positive ions as well as neutral
atoms.
The energies for removal of additional electrons from an atom are
referred to as the second ionization energy, third ionization, and so
on.
The energy required is always greater as you remove more electrons.
Each successive electron removed from an ion feels an increasingly
stronger nuclear charge.
The biggest ionization energies occur when an ion assumes a noblegas configuration. (see figure 3.6 on page 147)
Main idea – Adding electrons to atoms to form ions also requires
energy.
• The energy change that occurs when an electron is acquired by a
neutral atom is called the atom’s electron affinity.
• Most atoms release energy when they acquire an electron.
A+ e-
A- + energy
• On the other hand, some atoms must be “forced” to gain an electron by the
addition of energy.
A + e -- + energy
A–
The vast majority of elements release energy when they acquire an electron.
Main Idea – When atoms become ions, their
radii change
• The formation of a cation by the loss of one or more electrons always
leads to a decrease in atomic radius.
• The formation of an anion by the addition of one or more electrons
always leads to an increase in atomic radius.
Main Idea – Only the outer electrons are involved in forming
compounds.
• Chemical compounds form because electrons are lost, gained, or shared between atoms.
• The electrons available to be lost, gained or shared in the formation of chemical
compounds are called valence electrons.
Group number
valence electrons
1
1
2
2
13
3
14
4
15
5
16
6
17
7
18
8
Main Idea – Atoms have different abilities to
capture electrons.
• Linus Pauling, one of America’s most famous chemists, devised a scale
that reflects the tendency of an atom to attract electrons.
• Electronegativity is a measure of the ability of an atom in a chemical
compound to attract electrons from another atom in the compound.
• The most electronegative element, fluorine, is arbitrarily assigned a value of
4.
• This concept is vitally important in many biochemical reactions, such as
photosynthesis and aerobic respiration.
• Period trends – Electronegativities tend to increase across each
period.
• Group trends – electronegativities tend to either decrease down a
group or remain about the same.