Atomic Theory & Periodic Table Unit 3 Part 4 (Ch. 6) Trends in Periodic Table

Periodic Trends A trend is a predictable change in a particular direction A trend is a predictable change in a particular direction Trends occur in a group or period Trends occur in a group or period Trends are explained by electron configurations and variations in atomic structure Trends are explained by electron configurations and variations in atomic structure

Ionization Energy Ionization energy – the energy required to remove an electron from an atom or ion Ionization energy – the energy required to remove an electron from an atom or ion Ionization energy tends to decrease down a group Ionization energy tends to decrease down a group An element lower in a group has its outermost electrons further from the nucleus An element lower in a group has its outermost electrons further from the nucleus Electronic shielding – inner electrons shield outermost electrons from nucleus; outermost ones held less tightly Electronic shielding – inner electrons shield outermost electrons from nucleus; outermost ones held less tightly

Ionization Energy - 2 Ionization energy tends to increase across a period Ionization energy tends to increase across a period The number of electrons and protons both increase by one The number of electrons and protons both increase by one The additional electron is added to the same outer energy level The additional electron is added to the same outer energy level A higher nuclear charge more strongly attracts outer electrons in the same energy level A higher nuclear charge more strongly attracts outer electrons in the same energy level

Ionization Energy

Atomic Radius The exact size of an atom is hard to determine because electrons do not move in well-defined paths; the electrons occupy volume in an electron cloud The exact size of an atom is hard to determine because electrons do not move in well-defined paths; the electrons occupy volume in an electron cloud Scientists use bond radius to assign a size to an atom Scientists use bond radius to assign a size to an atom

Atomic Radius – bond radius Bond radius is half the distance of each line Bond radius is half the distance of each line

Atomic Radius Atomic radius increases down a group Atomic radius increases down a group As you move down another energy level is filled As you move down another energy level is filled The addition of another level of electrons increases the size, or atomic radius, of an atom The addition of another level of electrons increases the size, or atomic radius, of an atom Because of electron shielding, the attractive force of the nucleus is almost constant as you move down a group Because of electron shielding, the attractive force of the nucleus is almost constant as you move down a group

Atomic Radius - 2 Atomic radius decreases across a period Atomic radius decreases across a period Electron shielding does not play a role as you move across a period Electron shielding does not play a role as you move across a period As the nuclear charge increases across a period (one proton added moving one element to the next), the effective nuclear charge acting on the outer electrons also increases As the nuclear charge increases across a period (one proton added moving one element to the next), the effective nuclear charge acting on the outer electrons also increases

Atomic Radius

Electronegativity Not all atoms in a compound share electrons equally Not all atoms in a compound share electrons equally Knowing how strongly each atom attracts bonding electrons can help explain the properties of the compound Knowing how strongly each atom attracts bonding electrons can help explain the properties of the compound Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons

Electronegativity - 2 The atom with the higher electronegativity will pull on the electrons more strongly The atom with the higher electronegativity will pull on the electrons more strongly Fluorine is the element whose atoms most strongly attract shared electrons in a compound. Fluorine is arbitrarily given an electronegativity value of 4.0. Fluorine is the element whose atoms most strongly attract shared electrons in a compound. Fluorine is arbitrarily given an electronegativity value of 4.0. Values for the other elements were calculated in relation to this value Values for the other elements were calculated in relation to this value

Electronegativity - 3 Electronegativity decreases down a group Electronegativity decreases down a group The more protons an atom has, the more strongly it should attract electrons The more protons an atom has, the more strongly it should attract electrons BUT, electron shielding plays a role BUT, electron shielding plays a role Since effective nuclear attraction does not increase, but the electron is further from the nucleus, the ability of the atom to attract an electron decreases slightly Since effective nuclear attraction does not increase, but the electron is further from the nucleus, the ability of the atom to attract an electron decreases slightly

Electronegativity - 4 Electronegativity increases across a period Electronegativity increases across a period The effective nuclear charge increases across a period; electrons are more strongly attracted The effective nuclear charge increases across a period; electrons are more strongly attracted The increase in electronegativity across a period is much more dramatic than the decrease in electronegativity down a group. The increase in electronegativity across a period is much more dramatic than the decrease in electronegativity down a group.

Electronegativity

Other Periodic Trends Ionic Size Ionic Size Changes across/down the table just like atomic radius, but… Changes across/down the table just like atomic radius, but… Cations (+) always smaller than the neutral atom (pulls electrons in more) Cations (+) always smaller than the neutral atom (pulls electrons in more) Anions (-) always larger than the neutral atom (more electrons on the outside) Anions (-) always larger than the neutral atom (more electrons on the outside) Melting and boiling points Melting and boiling points

Example – Melting/Boiling points