Chemistry - Chapters 7 8 9 concepts

Explain the statement, Matter and radiation have a "dual nature."

• Both matter and radiation exhibit properties of both waves and particles

• This concept is known as particle-wave duality

• Radiation, like light, can behave as a wave (diffraction) and as a particle(photon in the photoelectric effect)

Schrödinger Equation, what are the different pieces for?

Ψ - is the wave function, describes the 3D shape of an electron, its magnitude squared equals the probability of an electrons location

H is the hamiltonian operator, it represents the total energy (kinetic + potential) of the system, written as an operator that acts on Ψ

Why is the shrodinger equation = Rydberg equation for 1 electron systems?

• One electron systems can be calculated using the Rydberg equation due to the absence of S(electron-electon repulsion) and Zeff(Effective nuclear charge)

Effective Nuclear Charge (Zeff)

• The net positive charge experienced by valence electrons (how strongly the nucleus attracts them)

• A relative Zeff can be found by Zeff = Z(# of protons) - S(shielding electrons, amount of electrons not in the last subshell) 

• This relative Zeff is only helpful in comparing to other electrons

• Increases across a period

• decreases down a group due to similar effective charge but further distance for valence electrons from new electrons due to the greater radius

• Explains why all the other trends exist (it’s the “master trend”)

Atomic Radius

• ½ the average distance between centers of neighboring touching atoms

• Atomic Radius decreases across a row/period

• This is because they are in the same orbital, but the number of protons is increasing as you get farther down the row. Since the amount of shielding electrons is constant, the effective nuclear charge increases, pulling the electrons closer, hence a smaller radius

• Atomic Radius increases down a column as the new principle energy levels increase the number of electron shells, thus increasing the atoms size

Ionization energy

• The energy required to remove an electron from an isolated neutral gaseous atom

• Ionization energy increases left to right in a period due the # of shielding electrons remaining fixed, while the the nuclear charge grows more positive, making each additional electron pulled in tighter, and harder to be removed

• exceptions are for elements already in stable situations, in which the ionization energy would be high, or electrons that are np^1, and could benefit from losing the lone electron. Would have a lower ionization energy compared to ns²

• Decreases down a collumn as the radius increases, following same period trends

Electronegativity

• The atom’s ability to attract electrons in a chemical bond

• Increases across a period along with Zeff, makes the radius smaller

Higher Zeff with a smaller radius can pull electrons more

• Decreases down a group, larger radius makes the nucleus less effective at attracting electrons

• Noble gases usually are not assigned EN value because they do not want to form bonds

• In bonds, if the difference between atoms is equal to or greater than 0.5, it is a polar bond, less than that would be a polar bond

Electron Affinity

• The electron change for the addition of an electron to an atom

• This value can be negative due to the release of energy due to the atom becoming more stable

• Therefore, the more negative the EA value, the more “favorable” the process/exchange is 

• Generally, the trend is that electron affinity increases from left to right(becomes more negative) because the rightmost electrons want electrons and towards the left would give them up. Stronger attraction for extra electrons (Zₑff ↑)

• electron affinity generally decreases down a group. As atoms get larger down a group, the added electron is further from the nucleus, resulting in a weaker attraction and a smaller release of energy

Exceptions for electron affinity

• Elements that have a filled shell, half filled shell, do not have a strong affinity

• Sodium has a higher electron affinity due to the nature that it would like an electron to become full, more stable, whereas magnesium right next to it, is essentially endothermic since it is stable as 3s², and adding an electron would be “unfavorable”

• Nitrogen is half filled, lower electron affinity compared to carbon on its left

Lattice energy

• energy required to completely separate one mole of an ionic solid into its gaseous ions - endothermic due to lattices being very stable

Why can solid metallic systems, such as Cu(s), conduct electricity and solid ionic compounds, such as NaCl(s), cannot? 

• Metal atoms release some of their valence electrons, which become delocalized — free to move throughout the entire crystal.

• These free electrons = electrical conductivity

• In ionic structures, the electrons cannot conduct electricity because of the ions in the lattice are fixed.