Transcript Chapter 13

Chapter 13
Chemical Periodicity
Introduction

In the 19th century, chemists began to
categorize the elements according to similarities
in their physical and chemical properties


The end result of this was the modern periodic table
The periodic table is very useful for
understanding and predicting the properties of
elements
13.1 The Development of the
Periodic Table


Newland, an English chemist, published list of
elements arranged in order according to their
increasing atomic mass.
He stated that the elements properties
repeated when they were arranged according
to increasing atomic mass in groups of eight
He called this the arrangement the law of octaves
 Similar to musical scale that repeats every eighth note
 Law only works up to Ca

13.1 The Development of the
Periodic Table

Mendeleev, a Russian chemist, refined and added
to the arrangement of elements in a table
according to their atomic masses

With this arrangement he noticed a regular (periodic)
recurrence of their physical and chemical properties
13.1 The Development of the
Periodic Table

Mosely rearranged the periodic table according
to the atomic number of the elements, which is
how the modern periodic table is arranged today

The periodic table is a valuable organizational tool
for chemists
13.2 The Modern Periodic Table


The most commonly used modern periodic
table, sometimes called the long form (your
table)
The long form table lists many properties of the
elements so that the chemist can check them at a
glance
13.2 The Modern Periodic Table


The periodic law states that when elements are
arranged in order of increasing atomic number,
there is a periodic pattern in their physical and
chemical properties
The horizontal rows of the periodic table are
called periods – there are 7 periods in the
periodic table
13.2 The Modern Periodic Table

The vertical columns are called groups or
families identified by number and a letter
Groups 1A through 7A and group 0 make up the
representative elements (wide variety of
properties)
 Group B elements are the transition metals
 Two rows of elements below the periodic table are
the lanthanides and actinides

13.3 Electron Configurations and
Periodicity

The electron configuration of an element plays the
greatest part in determining it’s physical and
chemical properties


Most elements within the same group have the identical
electron configurations in their outer most energy level
(valence level)
Elements are classified into 4 different categories
according to their valence (outer) electron configuration

Noble Gases, Representative Elements, Transition Metals,
Inner Transition Metals
13.3 Noble Gases (Group 0)
1. Noble gases (group 0) are elements in which
the outermost s and p sublevels are filled
Also called inert gases because they do not react
with other elements – they are stable on their own
 Helium has 2 valence electrons (full 1s sublevel)
 The rest of the noble gases have 8 valence
electrons (full s and p sublevels): Ne, Ar, Kr, Xe, Rn

13.3 Representative Elements
2. Representative elements (Group A)
 Elements whose outermost s or p sublevels are
only partially filled
 Group 1A are known as the alkali metals
1 electron in outermost energy level
 Very reactive → only in compounds in nature

13.3 Representative Elements
2. Representative elements (Group A) continued
 Group 2A are known as the alkaline earth metals
2 electrons in outermost energy level
 Also reactive (but not as reactive as 1A) → only in
compounds in nature


Group 7A are known as the halogens
7 electrons in outermost energy level
 Nonmetals that are highly reactive
 Also called salt formers

13.3 Representative Elements

For any representative element, the group
number is equal to the number of electrons
in the outermost energy level

See periodic table 354 –355
13.3 Transition Elements
3. Transition Metals (Group B)
 Elements whose outermost s sublevel and
nearby d sublevel contain electrons
The d sublevels overlap with s sublevels – this is why
they are transition elements
 Characterized by having electrons added to the d
orbitals
 Not as reactive as Group A elements

13.3 Inner Transition Metals
4. Inner Transition Metals
 Elements whose outermost s sublevel and
nearby f sublevel generally contain electrons

Characterized by the filling of the f orbitals
13.3 Electron Configurations and
Periodicity


The periodic table can be divided into sections,
which correspond to the sublevels that are filled
with electrons (on your table) (blocks)
Group1A and 2A are in the s block (also Helium)


valence level = period #
Group 3A, 4A, 5A, 6A, 7A, and 0 belong to p
block

valence level = period #
13.3 Electron Configurations and
Periodicity

Transition belong to d block


Inner transition belong to f block


Exception – d sublevel is one less than period #
Exception – f sublevel is 2 less than period #
The valence electron configurations can be
determined by using the block diagram in figure
13.4 – on your periodic table
13.4 Periodic Trends in Atomic Size

Remember that, according to the quantum
mechanical model, an atom does not have a
specifically defined boundary that sets the limit
of its size.
13.4 Periodic Trends in Atomic Size

However, there are ways to estimate the relative
sizes of atoms.
X-ray diffraction – estimates the size of atoms in
crystalline solids
 The distance between the nuclei of diatomic
molecules (examples: O2 or Br2) can be used to
estimate the atomic radius of an atom.
atomic radius – half the distance between the nuclei
of two like atoms

13.4 Atomic Size – Group Trends


Atomic size generally increases as you move
down a group of the periodic table
The size increases because electrons are added
to higher principle energy levels

The added charge of nucleus pulls electrons inward,
but the net effect is an increase in size because
electrons are further from nucleus
13.4 Atomic Size – Periodic Trends


Atomic size generally decreases as you move
from left to right across a period
The size decreases because electrons are added
to the same principle energy level, but the added
charge of nucleus pulls electrons inward; the net
effect is a decrease in size

This trend is less pronounced in periods where there
are more electrons in the occupied principle energy
levels between the nucleus and the outermost
electrons; this is referred to as the shielding effect
13.5 Periodic Trends in Ionization Energy


When an atom gains or loses an electron it forms
an ion.
The energy that is required to overcome the attraction
of the nuclear charge and remove an electron from a
gaseous atom is called the ionization energy



The first ionization energy is the amount needed to remove
the first outermost electron
The second ionization energy is the amount needed to
remove the next outermost electron
The third ionization energy is the amount to remove the third
and so on (Table 13.1 page 362)
13.5 Periodic Trends in Ionization Energy

Ionization energies can be used to predict how
many electrons an atom will gain or lose in a
chemical reaction


1A vs. 2A – Table 13.1 page 362
Two factors affect ionization energy: nuclear
charge and distance from the nucleus
13.5 Ionization Energy – Group Trends

In general, the first ionization energy
decreases as you move down a group on the
periodic table.

The size of the atoms increases as you move down;
thus the outermost electron is farther from the
nucleus and more easily removed

This results in a lower ionization energy
13.5 Ionization Energy – Periodic Trends

For the representative elements, the first
ionization energy generally increases as you
move from left to right across a period.

The nuclear charge is increasing and the atomic size
is decreasing, therefore there is more of an
attraction between the nucleus and the outermost
electron

This results in a higher ionization energy
13.6 Trends in Ionic Size


When atoms lose electrons they become
positive ions (cations)
Cations are always smaller than the atoms from
which they are formed

There is a stronger attraction between the nucleus
(same number of protons) and the remaining
electrons (fewer)
13.6 Trends in Ionic Size


When atoms gain electrons they become
negative ions (anions)
Anions are always larger than the atoms from
which they are formed

There is less of an attraction between the nucleus
(same number of protons) and the resulting
electrons (more)
13.6 Trends in Ionic Size


Periodic Trend – There is a decrease in the size
of cations as you move across a period from left
to right – when you get to group 4A the anions
(which are much larger) start to decrease in size
Group Trend – Ionic size (both cations and
anions) increases as you go down each group.
13.7 Trends in Electronegativity

The electronegativity of an element is the
tendency for the atoms of the element to attract
electrons when they are chemically combined
with another element

The Pauling scale uses arbitrary units to express the
electronegativity of the all elements (except noble
gases)

The Pauling scale is based on a number of factors
including ionization energies and electron affinities
13.7 Trends in Electronegativity

Periodic Trend – As you go across a period from
left to right, the electronegativity of the
representative elements increases
Metallic elements far left have low electronegativities
 Nonmetallic elements far right have high
electronegativities


Group Trend - Electronegativity generally decreases
as you go down a group
*Transition metals do not show as regular trend of electronegativity