Chemistry Unit 5

Effective Nuclear Charge

The charge experienced by an electron on a many-electron atom, not the same as the charge on the nucleus because of the effect on the inner electrons — The attractive forced exerted by the nucleus on electrons

• increases left to right in a period

Shielding Effect

The repulsive force exerted by core electrons on valence electrons

• remains constant left to right across a period

• Increases top to bottom in a group

Atomic Radii

• As we move down a group the atoms become larger

  • As the energy level increases (move down a group), the distance of the valance electrons from the nucleus becomes larger - so the atomic radius increases

• As we move across a period atoms become smaller

  • As we move across the periodic table, the number of core electrons remains constant; however, the nuclear charge increases. Therefore, there is an increases attraction between the nucleus and the outermost electrons

Atomic Radii Graph

Peaks: Alkaline Metals - because that’s when a new period starts (Back to S) - they are the largest

Dips: Noble Gases - because they orbital is full

Electrons and Protons added as period increases, so the atomic radii gets smaller and the effect nuclear charge (attraction) gets stronger

Ionic Radii

Ion - an atom with a nonneutral charge

Cation and Antion

Cation:

• Electron has been removed(given to Antion) - Positive

• Metals

• Effective Nuclear Charge has increased (attraction increased)

• Smaller than atomic counterpart

• multiple charges

Antion

• Electron was added(from Cation) - Negative

• Nonmetals

• Repulsion has increased 

• Larger than atomic counterpart

• fixed charge 

For ions with the same charge, ionic size increases down a group

Ionic Radii Graph

• Noble gases are missing (don’t ionize naturally)

• Ion charged determined by the amount of valence electrons it can lose (how many it has)

• Cations are removing an energy level which is why they become smaller

• Anion - more electrons used the more it expands because you add repulsion

Ionization Energy

• _______ - the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion

• First ionization energy - the amount of energy required to remove an electron from a gaseous atom

• Second ionization energy - the energy required to remove the second electron from a gaseous ion

• The larger the ionization energy, the more difficult it is to remove the electron

• There is a sharp increase in ionization energy when a core electron is removed

• Ionization energies for an element increase in magnitude as successive electrons are removed

  • As each successive electron is removed, more energy is required to pull an electron away from an increasingly more positive ion

• A sharp increase in ionization energy occurs when an inner-shell electron is removed

Periodic Trends in First Ionization Energies

• Ionization energy generally increases across a period

• As we move across a period it is more difficult to remove an electron

• Ionization energy decreases down a group

• This means that the outermost electron is more readily removed as ew go down a group

• As the atom gets bigger, it becomes easier to remove an electron from the most extended orbital

First Ionization Energy Graph

First Ionization Energy: The amount of energy necessary to remove one mole of electrons from one mole of neutral atom (ground state to off an atom)

Peaks: Noble Gases - takes the most energy in a period to remove energy

Lowest: Alkaine Metals - lowest effective nuclear charge, little energy is used ro remove them

• Smaller radius, more attraction - more energy it takes to remove them

• As noble gases of down (He to Ne to Ar) less energy - atoms get bigger, effective nuclear charge is over a long distance so its easier to remove them (less energy)

• Dropoffs for single fill vs double fill (has to do with orbitals)

• 3,11,19 - easiest to remove (lowest level)(Alkali Metals)

• MAKES CATIONS (positive)(lose an electron)

Electron Affinities

_____ - the energy change when a gaseous atom gains an electron to form a gaseous ion

• Electron Affinity and ionization energy measure the energy changes of opposite processes

Electron Affinity Graph

Electron Affinity: The amount of energy necessary to ADD one mole of electrons to one mole of neutral atom

• MAKES ANIONS (negative) (add an electron)

• Ionization energy: about making cations, Electron Affinity: about making anions

• Negative energy: Energy that is going out of the system

• O, F, S, Cl - more stable in anion setting (lower energy the more stable)

• anion - more stable for negatively charged 

• release energy

Electronegative

• F - attraction force going beyond its own electron clouds, and can take them to become an ion (if it doesn’t go all the way it can become a covalent bond (shared))

• F (Fluorine) highest electronegativity

• Li (Lithium) one of the lowest

• Right to left period decreasing electronegativity

• Up to down family decreasing electronegativity

• Electronegative: Attraction the nucleus has for the electrons on another item

Alkali Metals (Group 1)

• In group 1

• soft

• loss of their single s electron

  • reactivity increases as we move down the group

Alkaline Earth Metals

• Harder and more dense than the alkali metals

• Reactivity increases down the group

Halogens

• Gaining an electron to form an anion

• Fluorine is one of the most reactive substances known

Noble Gases

• Group 8

• Nonmetals and monoatomic

• Unreactive because they have completely filled s and p orbitals

Periodic Law

When elements of the periodic table are arranged in order of ascending atomic number, patterns in their characteristics and behaviors can be observed

ENC

the attractive force exerted by the nucleus on electrons

• Based on number of protons in the nucleus

• we will focus in the effect ENC has on valence electrons

• increases left to right in a period

SHE

the repulsive force exerted by core electrons on valence electrons

• remains constant left to right across a period

• increases top to bottom in a group

• ENC in excess of SHE

Other

• core electrons found by last complete energy level

• Atomic Radii: the distance from the nucleus to the outermost portion of the electron cloud in an atom of an element

• Ionic Radii: the distance from the nucleus to the outermost portion of the electron cloud in an ionic form of an element

• Columb’s Law - as they get closer, increase in attraction; as they get further, decrease in attraction