History of Periodic Table and Trends

• Dmitri Mendeleev created the periodic table

  • He organized it by increasing atomic mass

  • He made sure elements in rows had similar properties

• Henry Moseley rearranged the periodic table

  • Rearranged it by increasing atomic number

  • Modern periodic table is based on his idea

• It is better to organize the periodic table by increasing atomic number as

  • The Atomic number (# of protons) determines a certain element

    • (for example, ONLY lithium has 3 protons)

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

  • Reflects arrangement of electrons in atoms

    • For example, energy levels can be seen in periods

  • Easier to predict an element property based on position.

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

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

• Period: Horizontal rows of the periodic table

  • Elements on each period (row) have the same number of occupied energy levels

  • Do not have similar properties

• Group: Vertical columns of the periodic table

  • Elements on the same group have similar properties

  • Numbered from 1 to 18

• Representative metals: S & P Blocks

  • Include Alkali Metals, alkaline earth metals

• Transition Metals: D blocks

• Inner Transition Metals: F Blocks

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

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

Hallogens: Group 7, highly reactive with 7 valence electrons making them prone to forming compounds with other elements.

Noble Gases: Group 18, gaseous elements, very stable and least reactive.

Transition Metals:

• Filled in d orbitals.

• Found between alkaline earth metals and nonmetals.

• Have high melting and boiling points, good conductivity, and can form colored compounds.

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

• Post-transition metals: Less metallic properties, lower melting points

Metalloids:

• Share similar properties to both metals and nonmetals

What are inner transition metals?

• Lanthanides and actinides

• 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.

Atoms do not have fixed radius..

• To measure atoms, we measure by atom radius.

Atomic Radius: ½ of the radius between two nuclei of two like atoms

Atomic Size Trends:

• Increases from top to bottom.

• Decreases from left to right

• Measured in pico-meters (.pm)

• Group trend of atomic size,

  • Increases as you go down due to more occupied energy levels

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

• Period trend of atomic size

  • Decreases from left to right

    • Shielding effect is constant between periods, so all atoms will have the same number of orbitals

    • Yet increased protons the farther you go right means that there is an increased positive charge thus electrons are more attached to center protons which reduces the atomic size

Shielding Effect: 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

Ionization Energy & Trends

Ionization Energy: Energy required to remove an electron from a gaseous atom

• Measured in kJ

• depends on the distance between electrons and nucleus.

• Nuclear charge (# of protons)

  • Example: the more protons there are, the more electrons will be attracted to nucleus thus it will be harder to remove an electron

• Group trend: The ionization energy decreases as you go down

  • due to greater orbital level the further you go down

    • Less effort to remove electron when it is farther away from proton results in lower I.E.

• Period Trend: 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

• Ionization energy is higher when removing the 2nd or 3rd electron from the atom

  • 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 (1st electron that was pulled)

• There will be a very large increase of ionization energy when an electron is isoelectronic with a noble gas

  • 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 to remove electron

Ionic Size of Cations and Anions

• Cations Ionic Size: Smaller than neutral atom from which they were made from

  • Smaller due to 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.

• Anions Ionic Size: Anions are always larger than their neutral atoms from which they were made from

• The more electrons gained, the bigger the ion becomes

  • Larger size due to more electrons than protons

    • Resulting in other electrons having less pull/force to proton

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

    • Increase in electrons = more and farther away electron orbitals

Electronegativity & Trends

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

Electronegativity Trends in Groups and Periods

• Electronegativity levels decrease as you go down a group

  • More electrons = more electrons farther away from nucleus

  • The farther away from nucleus which electrons are, the less attracted they are to the nucleus.

  • Less electrons attracted to nucleus = less electronegativity

• Electronegativity levels increase from left to right across a period.

  • More nuclear charge due to increased atomic size the farther right means that increased protons result in electrons more attracted to nucleus