Chapter 3 (Test 2)

Question 1

When an electron moves away from a nucleus, is energy absorbed or emitted by the atom?

Answer 1

Absorbed - need energy to separate positive and negative charges

Question 2

When an electron moves closer to a nucleus, is energy absorbed or emitted by the atom?

Answer 2

Emitted - lower potential energy as it gets closer (energy has to go somewhere)

Question 3

Energy Level Diagram

Answer 3

depicts the allowed energy states for an electron

Question 4

Ground state

Answer 4

lowest (MOST STABLE) energy state for an atom (n=1 for H)

Question 5

Excited state

Answer 5

electrons move to higher energy states when an atom absorbs a photon of sufficient energy

Question 6

Absorbed and Emitted photon energy equations

Answer 6

△E = +E (photon)
△E = -E (photon)

Question 7

The potential energy of an electron in a specific energy level can be calculated with:

Answer 7

E = -2.178 x 10^-18(1/n^2)

HYDROGEN ONLY

Question 8

H Atom Change in Energy between energy levels Equation

Answer 8

△E = E (final) - E (initial)

Question 9

how to calculate the electron energy when it moves energy levels for HYDROGEN

Answer 9

△E = -2.18x10^-18J(1/n(f)^2 - 1/n(i)^2)

Question 10

Change in energy vs energy of photon

Answer 10

△E(atom) = E(absorbed) = hv(absorbed)

△E(atom) = -E(emitted) = -hv(emitted)

Question 11

Delocalized

Answer 11

the position of an electron cannot be precisely defined (because a particle occupies a particular location, but a wave has no exact position)

Question 12

Heisenberg’s uncertainty principle

Answer 12

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

Question 13

How is an orbital’s energy related to it’s distance from the nucleus?

Answer 13

the farther away, the higher the energy (inversely proportional)

Question 14

The 2p orbitals

Answer 14

2Px, 2Py, 2Pz orbitals (infinity signs) combine to make the P subshell (n=2 and above)

Question 15

3d orbitals

Answer 15

3Dxy, 3Dxz, 3Dyz, 3D(x^2-y^2), 3Dz^2 orbitals (flowers) combined to make the D subshell (n = 3 and above)

Question 16

Degenerate

Answer 16

for HYDROGEN ONLY all orbitals in the same energy level have the same energy (degenerate)
2s energy = 2p energy
2s energy < 3s energy

Question 17

Why do the line spectra of all other atoms have more emission lines than H?

Answer 17

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)

Question 18

Aufbau Principle

Answer 18

electrons occupy the lowest energy orbital (most stable) available and an orbital can hold a max of 2 electrons with opposite spins

Question 19

Valence electron rules

Answer 19

(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

Question 20

Where can irregularities in electron configurations come from?

Answer 20

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

Question 21

Why might elements in the same group exhibit similar chemical behavior?

Answer 21

Because they have the same # of valence electrons - they’re the reactive ones!

Question 22

How to do Electron configurations for ions

Answer 22

Anions: add electrons following Hund’s Rule
Cations: remove electrons from the electron shell of the HIGHEST OCCUPIED ENERGY LEVEL

Question 23

Characteristics of metals

Answer 23

-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

Question 24

Sheilding

Answer 24

• 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

Question 25

Effective nuclear charge

Answer 25

the net attraction of an electron to a nucleus after screening Zeff = Z - core electrons Z: actual nuclear charge: # of protons

Question 26

Effective nuclear charge trends w consequences

Answer 26

- 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))

Question 27

Atomic radius trends

Answer 27

- decreases left to right (from increasing Zeff and constant n) - increases top to bottom (from increasing n - gets bigger)(EVEN IONS FOLLOW THIS)

Question 28

How ions change atomic radius

Answer 28

Cations: SMALLER Anions: BIGGER (more electrons added on)

Question 29

Isoelectronic series

Answer 29

- same number of electrons | - ionic size decreases with an increasing nuclear charge (pulls in electrons, more stable)

Question 30

Ionization energy

Answer 30

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

Question 31

Why is the IE 2 for He+ 4 times greater than the IE 1 for H?

Answer 31

The electron is pulled to the +2 nucleus' charge, so you need more energy to pull it off

Question 32

Trends in first ionization energies

Answer 32

- 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)

Question 33

Discontinuities in the 1st ionization energies

Answer 33

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)

Question 34

What is the trend in successive ionization energies?

Answer 34

(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

Question 35

Which would you expect to have a higher 2nd ionization energy, Na or Mg?

Answer 35

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.

Question 36

Electron affinity

Answer 36

the energy change accompanying the ADDITION of a mole of electrons to a mole of gaseous atoms

Question 37

Why are the ionization energies positive, but the electron affinity for Cl is negative?

Answer 37

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)

Question 38

Positive affinity (discontinuities)

Answer 38

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!

Question 39

Trends in electron affinity

Answer 39

-gets more negative left to right (larger Zeff, electron ends up closer to the nucleus - more stable!)