Atomic Trends in Periodic Tables

Dmitri Mendeleev

• By increasing atomic mass

• So that elements in the same row have similar properties

Henry Moseley

• By increasing atomic number

• What we use today

• Atomic number (protons) determine the element

• Helps to identify trends and patterns in properties (periodic law)

• Reflects the arrangement of electrons in atoms.

• Allows for easy comparison of elements in the same group.

• Enables prediction of an element's properties based on its position.

When elements are arranged in order of increasing atomic number, there is a pattern in their physical and chemical properties

The horizontal rows of the periodic table

Vertical columns of the periodic table

Have similar properties

1 to 18

• Do not have similar properties.

• But they have the same number of occupied energy levels

S & P Blocks

D Block

F Block

• Alkali metals: Group 1 elements, highly reactive, soft metals

• Alkaline earth metals: Group 2 elements, reactive but less than alkali metals

• D-block elements, good conductors of heat and electricity

Share properties between metals and nonmetals

• also known as lanthanides and actinides

• are located at the bottom of the periodic table.

• occupy f-orbitals in their electron configurations.

• These elements possess unique properties and find applications in magnets, catalysts, and nuclear reactors.

• filled in d orbitals.

• They are found between alkaline earth metals and nonmetals.

• These metals have high melting and boiling points, good conductivity, and can form colored compounds.

• They are also known for multiple oxidation states and catalytic activity.

• Examples include iron, copper, zinc, silver, and gold.

• Atoms do not have fixed radius

• The radius of an atom is found by measuring the nuclei in between two touching atoms of the same element and then halving that distance.

½ of the radius between two nuclei of two like atoms

Increases as you go down due to more occupied energy levels

• More occupied energy levels = more orbits = greater atomic size

Decreases from left to right

• Shielding effect is constant between periods

• Increased protons = increased nuclear charge = electrons are more attracted to center protons = less atomic size

• The more electrons which are closer to the proton results in outer electrons being repelled due to increase in negative charge.

• As negatives and negatives repel —> Outer electrons are repelled and move into farther away energy orbitals.

• Thus outer electrons have less attraction to nucleus

• Increases from top to bottom

• Decreases from left to right

Energy required to remove an electron from a gaseous atom

• Distance between electrons and nucleus

• Nuclear charge (# of protons)

• By increasing atomic number

• What we use today

Increases as you go to the right as there is an increased nuclear charge (higher proton number)

• More protons = electrons are more attracted to proton

• More attraction between electrons and protons = harder to remove electron thus higher ionization energy

• Yes, the ionization energy is much higher

• As the 2nd or 3rd electron removed from an atom tend to be closer to the proton

• Meaning that they will have a higher pull which is harder to break than with the outer and farthest electron

There will be a very large increase of ionization energy when an electron is very

Generally tend to be the opposite of metals

• Brittle (breaks easily)

• Dull

• Poor conductors of heat or electricity

• Good insulator

• Isoelectronic with a noble gas means having the same number of electrons as a noble gas.

• Noble gases have full electron shells, making them stable.

• Removing an electron from an isoelectronic species disrupts this stability, requiring a significant amount of energy

• Ionization energy increases substantially.

• Smaller than neutral atom from which they were made from

• Loss of energy levels (loss of orbitals = less rings around atom)

• More protons than electrons means that more electrons will be pulled closer to nucleus

The more electrons lost, the smaller the ion becomes

Anions are always larger than their neutral atoms

• More electrons than protons result in less attractive force to proton

• Electrons that are less attracted to protons will be farther away

• Increase in electrons = more electron orbitals

• Lustrous (shiny)

• Good conductors of heat and electricity

• Malleable (Example: Aluminium can be split into thin sheets)

• Ductile (Can be turned into metals)

The more electrons gained, the bigger the ion becomes

Ability for an element to attract other electrons when a compound (when it is chemically combined with another element)

4.0

Decreases as you go down

• More electrons = more electrons farther away from nucleus

• Less electrons attracted to nucleus = less electronegativity

• Increases from left to right

• More nuclear charge (more protons) = electrons will be more attracted to nucleus