Chem Exam 4 ch 6.3-8.8

Pauli Exclusion Principle

No two electrons can have the same set of four quantum numbers

S orbitals hold _ electrons

2

P orbitals hold _ electrons

6

D orbitals hold _ electrons

10

F orbitals hold _ electrons

14

Degenerate (single electron atom)

subshells with the same principal quantum number (meaning the same energy)

Higher principal quantum number = _ energy

higher

Energy of N =

-R_yhc/

n²

Degenerate (multielectron atom)

orbitals in the same subshell (l) have the same energy

Energy from least to greatest in terms of subshell due to penetration (multielectron atom)

S < P < D < F

Shielding

When electrons that are close to the nucleus protect the further away ones from being pulled in to the nucleus

Shielding causes _ effective nuclear charge (Z_eff)

lower

When n increases, energy _

increases due to less attraction to the nucleus

Effective nuclear charge (Z_eff)

the net charge experienced by one electron in a multielectron atom

Net charge

The balance of the attractive forces of the nucleus and the repulsive forces of other electrons

Smaller Z_eff = _ energy

Greater

Penetration

The ability of core electrons to get closer to the nucleus rather than valence electrons

Closer proximity = _ attraction to the nucleus

greater

Bohr Model of Hydrogen

Electron in a hydrogen atom moves around in one fixed set of circular orbitals

Neils Bohr orbits

Energy levels

n=1 (bohr model)

ground state

n>1 (bohr model)

excited states

Where are electrons located (Bohr Model)

They must occupy one of the energy levels

Is there energy emitted when electrons are in an energy level (Bohr model)

No

Electrons can move from orbitals by what (Bohr Model)

Photon of light emission or absorption

Photon Emission Motion of Electron (Bohr Model)

Higher to lower

Photon Absorption Motion of Electron (Bohr Model)

lower to higher

As n increases, distance between energy levels _

decreases

Energy is 0 when

The electron is free of the nucleus

As n increases, the distance between the electron from the nucleus _ , and the energy becomes _ (less negative)

increases, higher

Energy of a transition

-R_y (1/n_f² - 1/n_i²)

Energy (if given wavelength)

hc

/λ

Energy (if given frequency)

hv

De Broglie Equation

λ = h/ m (kg) x v (velo m/s)

Heisenberg Uncertainty Principle

We will never know for certain the position and velocity of electrons at any time

quantum numbers

describe orbitals and their properties within different atoms

Principle quantum number (n)

determines distance of electron from the nucleus and the energy of the orbital as well as the shell where it is located.

n (possible values)

any positive whole number 1 or greater

Angular momentum quantum number (l)

determines the shape of the orbital where the electron is located

l (possible values)

n-1

l value of 0 subshell letter

s

l value of 1 subshell letter

p

l value of 2 subshell letter

d

l value of 3 subshell letter

f

Magnetic quantum number (m_l)

the orientations of the orbital

m_l (possible values)

-l to +l

spin quantum number (m_s)

spin behavior of an electron

m_s (possible values)

+1/2 or -1/2

A single orbital can hold a max of _ electrons

2

Number of orbitals in a subshell

2l + 1

Orbital

the area around a nucleus where electrons can be found

S orbitals shape

sphere

p orbitals shape

dumbbell

d orbitals shape

clover

Nodes

area where there are no electrons

total number of nodes

n-1

Order of electron filling

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p

Condensed Electron Configuration

[closest noble gas] and then the rest of configuration

Valence electrons (main group elements)

electrons in highest n

Core electrons (main group elements)

any electrons not in highest n

Valence electrons (transition metals)

s and d electrons with the highest n

Core electrons (tranistion metals)

any electrons not in the highest s or d

Transition metal configuration exceptions

Cr, Mo, Cu, Ag, Au

During transition metal configuration exceptions

One s electron is in the d orbitla

Excited state electron configurations

electron jumps from lower energy subshell to higher energy subshell

Pauli exclusion principle

no two electrons can have same set of quantum numbers

hund’s rule

in a shared orbital, all of the positive spinning electrons will be added before the negative ones

Electron configuration of cations

lose electrons from highest energy level

Electron configuration of anions

add electrons to the partially filled orbitals

Paramagnetic

having unpaired electrons in the electron configuration of ions or atoms

Diamagnetic

have a full electron configuration

effective nuclear charge(Z_eff) (equation)

atomic number - core electrons

Atomic radius (periodic trend)

decreases left to right, increases up to down

effective nuclear charge (Z_eff) trend

increases left to right

atomic radius trend reasoning (decrease left to right)

increase in number of protons without any more core electrons increasing (increasing Z_eff), pull electrons in tighter

atomic radius trend reasoning (increase top to bottom)

energy level increases so they are further from the nucleus

Stronger attraction to the nucleus = _ the radius

smaller

trends in ion sizes (for an atom)

anion > neutral atom > cation

Gaining electrons = _ nuclear attractions (larger size)

less

Isoelectronic ions

ions with an identical number of electrons

If isoelectronic, more _ means smaller radius

protons

Ionization energy

amount of energy required to remove an electron from an ion or atom in gas phase

ionization energy is always _ (in terms of heat)

endothermic

first ionization energy

energy required to remove the first electron from an atom or ion in gas phase

1st Ionization Energy (generic equation)

X (g) → X⁺ (g) + e^-

2nd Ionization Energy (generic equation)

X⁺ (g) → X²⁺ (g) + e^-

second ionization energy

energy required to remove the second electron from an atom or ion in gas phase

2nd and higher ionization energies are always _ than first ionization energies

larger, more endothermic, more positive

Removing a core electron requires a much _ amount of energy

larger

First ionization energy (periodic trend)

increases left to right, decreases top to bottom

Reason for first ionization energy trend (left to right)

greater Z_eff = stronger nuclear attraction = harder to remove electrons

Reason for first ionization energy trend (top to bottom)

radius increases = less attraction = electrons easier to remove

Ionization energy exceptions

If shell is close to full or half full (group 13,16), it takes less energy to make it full. If it is already full or half full, it takes more energy to make it not full (2,15).

Electron attachment enthalpy (ΔH_EA)

the energy change associated with adding an electron to a neutral atom in gas phase

Electron attachment enthalpy (ΔH_EA) (generic equation)

A (g) + e^- → A^- (g)

Electron attachment enthalpy (ΔH_EA) is always _ in terms of heat

exothermic 

Electron attachment enthalpy (ΔH_EA) (period trend)

more negative left to right (more likely), less negative top to bottom (less likely)

Electron attachment enthalpy (ΔH_EA) exceptions

More likely to go towards complete shell and hard to add to an already full shell.

Electron Affinity

energy required to remove an electron from an anion with a -1 charge

Electron Affinity (generic equation)

A^- (g) → A (g) + e^-