Chapter 3 (Test 2) Flashcards
When an electron moves away from a nucleus, is energy absorbed or emitted by the atom?
Absorbed - need energy to separate positive and negative charges
When an electron moves closer to a nucleus, is energy absorbed or emitted by the atom?
Emitted - lower potential energy as it gets closer (energy has to go somewhere)
Energy Level Diagram
depicts the allowed energy states for an electron
Ground state
lowest (MOST STABLE) energy state for an atom (n=1 for H)
Excited state
electrons move to higher energy states when an atom absorbs a photon of sufficient energy
Absorbed and Emitted photon energy equations
△E = +E (photon) △E = -E (photon)
The potential energy of an electron in a specific energy level can be calculated with:
E = -2.178 x 10^-18(1/n^2)
HYDROGEN ONLY
H Atom Change in Energy between energy levels Equation
△E = E (final) - E (initial)
how to calculate the electron energy when it moves energy levels for HYDROGEN
△E = -2.18x10^-18J(1/n(f)^2 - 1/n(i)^2)
Change in energy vs energy of photon
△E(atom) = E(absorbed) = hv(absorbed)
△E(atom) = -E(emitted) = -hv(emitted)
Delocalized
the position of an electron cannot be precisely defined (because a particle occupies a particular location, but a wave has no exact position)
Heisenberg’s uncertainty principle
showed that the more precisely we know the momentum (mass x velocity) of a particle is known, the less precisely its position is known
- thus, the uncertainty of the calculated position of an electron within an atom is greater than the diameter of the atom
How is an orbital’s energy related to it’s distance from the nucleus?
the farther away, the higher the energy (inversely proportional)
The 2p orbitals
2Px, 2Py, 2Pz orbitals (infinity signs) combine to make the P subshell (n=2 and above)
3d orbitals
3Dxy, 3Dxz, 3Dyz, 3D(x^2-y^2), 3Dz^2 orbitals (flowers) combined to make the D subshell (n = 3 and above)
Degenerate
for HYDROGEN ONLY all orbitals in the same energy level have the same energy (degenerate)
2s energy = 2p energy
2s energy < 3s energy
Why do the line spectra of all other atoms have more emission lines than H?
there are less possible energy levels in hydrogen (for degenerate, you can only go n=1 to n=2 etc and for non-degenerate you can go from 1s to 2s to 2p etc)
Aufbau Principle
electrons occupy the lowest energy orbital (most stable) available and an orbital can hold a max of 2 electrons with opposite spins
Valence electron rules
(the reactive electrons in an atom) they’re the electrons in ALL partially filled d and f subshells and the ones in the highest occupied energy level
Where can irregularities in electron configurations come from?
Ex: [Ar] 4s^1 3d^5 - because full and half full are more stable. when one moves up and creates a half full d subshell it lowers the energy in d and makes it more stable
Ex2: [Ar] 4s^1 3d^10
Why might elements in the same group exhibit similar chemical behavior?
Because they have the same # of valence electrons - they’re the reactive ones!
How to do Electron configurations for ions
Anions: add electrons following Hund’s Rule
Cations: remove electrons from the electron shell of the HIGHEST OCCUPIED ENERGY LEVEL
Characteristics of metals
-good conductors of heat and electricity
-malleable (can be made into thin sheets)
-shiny solids (except mercury)
-ductile (can be made into wire)
metallic character decreases as you go to the upper right hand side of the periodic table
Sheilding
- the partial cancelling of the nuclear attractive force on an electron due to the electron-electron repulsion
- it influences periodic trends!
- core electrons are the most effective at sheilding
Effective nuclear charge
the net attraction of an electron to a nucleus after screening
Zeff = Z - core electrons
Z: actual nuclear charge: # of protons
Effective nuclear charge trends w consequences
- increases left to right (orbitals shrink (stronger charge pulls in electrons) and stability increases (electrons closer to nucleus))
- increases (SLIGHTLY) top to bottom (orbitals expand (more electrons) and stability decreases (more electrons farther from nucleus))
Atomic radius trends
- decreases left to right (from increasing Zeff and constant n)
- increases top to bottom (from increasing n - gets bigger)(EVEN IONS FOLLOW THIS)
How ions change atomic radius
Cations: SMALLER
Anions: BIGGER (more electrons added on)
Isoelectronic series
- same number of electrons
- ionic size decreases with an increasing nuclear charge (pulls in electrons, more stable)
Ionization energy
the energy required to remove one mole of electrons from one mole of an atom or ion
Ex: H -> H+ + e- IE(1H) = 2.18x10^-18J
Why is the IE 2 for He+ 4 times greater than the IE 1 for H?
The electron is pulled to the +2 nucleus’ charge, so you need more energy to pull it off
Trends in first ionization energies
- increases left to right (Zeff increases, orbitals get smaller, stronger force for electrons)
- decreases w increasing atomic # (in an energy level further from the nucleus, easier to pull off electron - also more electrons to block the valence electrons from the nucleus’ charge)
Discontinuities in the 1st ionization energies
Groups 2 and 13 (group 13 goes down bc it’s easier to pull off an electron from a partially filled p subshell than a full s subshell - farther away, too)
Groups 15 and 16 (group 16 goes down bc group 15 has a half-full p subshell which is really stable and hard to break apart, group 16 can get more stable when you pull off an electron)
What is the trend in successive ionization energies?
(removing each successive electron - IE1 (first), IE2 (second), IE3 (third), etc)
- removing core electrons takes way more energy (removing valence electrons is way easier)
- ionization energy jumps significantly when going from removing valence to core electrons
Which would you expect to have a higher 2nd ionization energy, Na or Mg?
Na because it only has 1 valence electron. After that, it’s trying to remove an electron from the noble gas core. Mg is going to be lower bc it’ll remove another 2s electron.
Electron affinity
the energy change accompanying the ADDITION of a mole of electrons to a mole of gaseous atoms
Why are the ionization energies positive, but the electron affinity for Cl is negative?
Ionization: add energy bc we have to pull the electron from the positive nucleus (usually positive)
Affinity: energy released - electron attracted to positive nucleus (usually negative)
Positive affinity (discontinuities)
Be, Mg, and N
- adding electrons takes energy when putting it in half full or empty subshells (also close to nucleus) - don’t want it!
Trends in electron affinity
-gets more negative left to right (larger Zeff, electron ends up closer to the nucleus - more stable!)
