Unit 13

Chemical Periodicity

A Little Bit of History

In the early 1800s, German chemist J.W.

Dobereiner observed that several elements

could be classified into sets of three which he

called triads.

• (Li, Na, K), (Ca, Sr, Ba), (Cl, Br, I)

• The elements within each triad had similar

properties.

• Several properties of the middle elements

were averages of the other two.

In 1865, English chemist J.A.R. Newlands

observed that when the then 62 known elements

were arranged in order of increasing atomic mass,

every eighth element had similar properties.

• This was known as the Law of Octives

• The first was light the eighth; Na was like Li

• The second was like the ninth, Mg was like

Be

• He knew nothing of the noble gases

In 1869, Russian chemist Dmitri Mendeleev and

German chemist Lothar Meyer published nearly

identical schemes for classifying elements.

• Mendeleev got the credit because he published

his first and he was more successful at

demonstrating it value.

• Mendeleev eventually produces the first working

periodic table of the elements.

• He arranged his table so that elements in the

same column had similar properties.

• He switched the order of three pairs of

elements so as to keep elements in columns

with similar properties

• He switched the order of three pairs of

elements so as to keep elements in columns

with similar properties.

• This was contrary to all other chemists who

arranged elements according to increasing

atomic mass.

• He boldly pronounced that perhaps the

calculated atomic masses for those mixed up

elements should be recalculated

• Mendeleev also had the insight to predict the

existence and some of the properties of three

new elements (these elements were missing

from his periodic table.

In 1913, English chemist Henry Moseley

determined the atomic number of the elements.

• He already know that protons are positively

charged particles that exist inside of nuclei.

• He used x-rays to determine that each element

had a certain number of positive particles in its

nucleus.

• He said these positive particles must be the

protons discovered by Goldstein.

• This discovery led to the term atomic number.

The modern periodic law states that when

elements are arranged in order of

increasing atomic number, their physical

and chemical properties show a periodic

pattern.

Mosley’s discovery proved Mendeleev to be

correct.

Different kinds of periodic tables may present different

kinds of information or present information in different

ways.

• The shape of the periodic table comes from the

periodic law.

• Elements with similar properties are aligned in

vertical columns called groups or families.

• The horizontal rows are called periods.

• Some of the groups have family names. (Alkali

metals, alkaline earth metals, halogens, and the

noble gases.)

• Elements can be classified as metals, nonmetals, and

semimetals.

• Metals are typically solids at room temperature.

• characteristic shine, conduct heat and electricity,

malleable (hammered into sheets), ductile (drawn

into wires), lose e-‘s to become cations

• Nonmetals can be solids, liquids, or gases

• do not have luster (dull), poor conductors of heat

and electricity, brittle.

• quite a variation in physical properties (some are

colored, others are colorless; some are soft solids

while others form hard solids); gain e-‘s to become

anions

• Semimetals have some properties of metals and some

of nonmetals, or they have properties that are

intermediate.

The outermost “s” and “p” electrons, which are largely

responsible for an atom’s chemical behavior, are called

valence electrons.

• The elements in a group have similar properties

because they have the same number of valence

electrons (which means similar electron

configurations).

• To save space in writing electron configurations and

to focus attention on valence electrons, chemists

often use abbreviated electron configurations.

• An atom’s inner electrons are represented by the

symbol for the nearest noble gas with a lower

atomic number. (This is called the noble gas inner

core)

Octet rule is a very important concept to understanding

bonding in chemistry.

• Octet refers to the eight outer electrons of an atom.

• An atom with all of its “s” and “p” electrons is

extremely stable.

• The noble gases have those eight outer electrons and

will not bond with another atom.

• Metals will lose electrons to back up to have the same

number of electrons as the previous energy level’s

noble gas.(they become cations)

• Nonmetals will gain electrons to have the same number

of electrons as the noble gas at the end of the row.

(they become anions)

• Sometimes nonmetals share electrons to gain the

octet

The key to understanding the shape of the periodic table is

to examine the elements’ electron configuration. (The four

main sections of the periodic table correspond to the four

sublevels, s p d f.)

• The s-block elements are groups 1 and 2.

• The p-block elements are groups 3-8.

• The d-block elements are the elements whose last

electrons fill the d sublevel.

• The f-block elements are the elements whose last

electrons fill the f sublevel.

The shape of the periodic table is a result of the way the

electrons fill the s, p, d, and f orbitals of the different

energy levels.

• The four blocks have other names as well.

• The elements in the s- and p- blocks are called the

representative elements.

• The elements in the d-block are called the transition

metals.

• Those in the f-block are known as the inner transition

metals.

Periodic Trends

Many properties of the elements change in a

predictable way as you move through the

periodic table. These systematic variations are

called periodic trends. The periodic trends that

you must understand are…

• Atomic radius

• Ionization energy

• Electronegativity

• Shielding Effect

Atomic Radius

The atomic radius is the distance from the center of an atom’s nucleus to its outermost electron.

• Atoms get larger going down a group. (ie. size increases with increasing atomic #)

• Atoms get smaller moving from left to right across each period. Why?

– As a rule, atoms that have more positive charge in their nuclei exert a stronger pull on the electrons in a given principal quantum number.

– A stronger attractive force shrinks the electrons’ orbitals and makes the atom smaller.

• As you learned in Unit 5, an atom can gain or

lose electrons to form an ion.

• When an atom loses electrons it becomes

smaller.

• Loss of electrons not only vacates the atom’s

largest orbitals, it also reduces the repulsive

force between the remaining electrons,

allowing them to be pulled closer to the

nucleus.

• When an atom gains electrons it becomes larger.

• A fluoride ion is larger than the fluorine atom mainly because of the greater number of electrons, which increases the electric repulsive forces among them.

• The increased repulsions spread out the electrons, thus making the ion larger than the atom.

• The atoms of a particular element typically

form only certain ions.

• The elements on the left side of the table form

positive ions. (cations)

• The elements on the right side of the table

form negative ions. (anions)

• The elements in the eighth column do not

form ions. (noble gases)

• An atom’s ionization energy is the energy needed to remove one of its electrons.

• You can think of ionization energy as a reflection of how strongly an atom holds onto its outermost electron.

• Atoms with high ionization energies hold onto their electrons very tightly, whereas atoms with low ionization energies are more likely to lose one or more of their outermost electrons and gain a positive charge.

• Ionization energies generally increase as you

move from left to right across a period.

• Metals have low ionization energies while the

noble gases have the highest ionization

energies.

• There will be peaks and valleys due to filled

sublevels or partially-filled sublevels.

• Ionization energies generally decrease as you

move down a group.

• The energy required to remove the first electron from an isolated atom is called the first ionization energy.

• The successive ionization energies are the energies required to remove electrons beyond the first electron.

• For each element you can find one very large increase between a different pair of ionization energies.

• This happens whenever the inner core of electrons are attacked.

Electronegativity

• An atom’s electronegativity reflects its ability to

attract electrons in a chemical bond.

• Electronegativity increases as you go from left to

right across the period.

• Electronegativity decreases as you go down a

group.

– Metals have low electronegativity while nonmetals

have high electronegativity.

– Fluorine has the highest and francium has the lowest.

Shielding Effect

• Shielding effect is the result of full energy levels

separating the outermost electrons from the

nucleus.

• An element like lithium only has one energy level

separating its outer electron from the nucleus.

(therefore low shielding effect)

• But, an element like francium has many energy

levels separating its outer electrons from the

nucleus. (therefore high shielding effect)

• This shielding effect is the reason why

ionization energy and electronegativity

decreases as you go down a group.

• Shielding effect increases as you go down a

group.

• Shielding effect remains the same as you go

across a period.