Periodic Trends- Chemistry

Atomic Radius

half the distance between adjacent nuclei or bonded nuclei

*no defined edge of atom because of electron cloud- can't measure from edge

Atomic Radius- Period Trend

-left to right decrease

-increase in number of protons and valence electrons but same energy level; no further shielding so stronger nuclear pull

Atomic Radius- Group Trend

-going down atomic radius increases -there are more energy levels causing more shielding and less nuclear pull

Ion

a charged atom; formed when giving up or taking electrons; form ionic bonds

Cations

-positively charged ion due to electron loss

-usually metal

Anion

-negatively charged ion due to electron gain

-usually nonmetal

Ionic radius

-half the distance between adjacent charged nuclei or bonded charged nuclei (cations and anions)

Cation- Period Trend

-radius decreases left to right

-losing electrons, nuclear pull greater because there are more protons than electrons

Transition from Anion to Cation- crossing over stair step line

-radius increases

-gaining electrons, nuclear pull less because electrons outnumber protons

-more electron repulsion

Anion- Period Trend

-radius decreases left to right

-gaining fewer electrons with each element

-electrons don't outnumber protons as much

Atomic radius of Anion versus atomic radius of Cation

anionic radius is bigger than cationic radius if they're in the same period

Valence electrons

the electrons that are on the outermost energy level- are involved in a chemical reaction

*group number tells number of valence electrons

Reactivity

a chemical's ability to react- the ability to transfer electrons

Periodic Trends

apply mostly to the main groups/representative groups of elements (s and p blocks)

Reactivity- Period Trend

goes down as you go across a period left to right

Reactivity- Group Trend

-metals: goes up as you go down

-nonmetals: goes down as you go down; harder to gain electrons because of weaker nuclear pull caused by shielding

Electronegativity

-unit: Paulings

-the measure of the ability of an atom in a chemical bond to attract electrons

*atom with higher electronegativity will attract electrons of bond

Four most electronegative elements, greatest to least

-Fluorine

-Oxygen

-Chlorine

-Nitrogen

Electronegativity- Period Trend

-increases left to right across the period, but there are exceptions

-increase in protons, no change in shielding (no change in energy levels), greater nuclear pull, holds onto electrons tighter

Electronegativity- Group Trend

-decreases or stays the same a you go down a group

Noble Gases

no electronegativity values- already stable, don't need to react

Bottom left of Periodic Table

least electronegativity

Top right

greatest electronegativity

Alkalis and Alkaline Earth metals

least electronegative- want to lose, not gain/attract electrons

Ionization energy

the energy required to remove an electron from a neutral atom of an element in the gas phase

First ionization energy

the energy required to remove the first electron from an atom

Successive ionization energy

the energy required to remove more than one atom from an element

(second ionization energy: second electron, third ionization energy: third electron etc.)

Isoelectronic

-when multiple atoms, molecules, or ions have the same number of electrons or similar electron configurations

-Examples: Sulfide ion, Chloride ion, Argon, Potassium ion all have same electron configurations

Ionization

making an atom an ion

Nuclear pull- ionization energy

there is a greater nuclear pull with each electron removed because of the protons outnumbering the electrons- ionization energy becomes greater

Biggest jump in ionization energy

occurs when an element is in noble gas configuration

Alkali metals- ionization energy

have low ionization energy because they want to lose an electron

Noble gases- ionization energy

have highest ionization energy because they are the most stable and don't want to react/give up their eight valence electrons

Nonmetals- ionization energy

nonmetals have high ionization energy because they want to gain electrons, not lose electrons

Metals- ionization energy

low ionization energy because they want to lose electrons

Period Trend- ionization energy

-increases left to right across period

-increase in protons and no change in distance/shielding/energy levels (increase in nuclear pull- hold on to electrons tighter)

Group Trend- ionization energy

-goes down as you go down a group

-more electrons as you go down, more energy levels and more shielding causing a weaker nuclear pull