Atomic Structure and Periodic Trends Flashcards
(45 cards)
1
Q
proton charge
A
1+
2
Q
proton mass
A
1 amu
3
Q
atomic number
A
- Z
- number of protons
- nuclear charge
- determines the element
4
Q
neutrons charge
A
0
5
Q
neutrons mass
A
1 amu
6
Q
mass number
A
- A
- number of protons and neutrons
- weighted average all of isotopes
7
Q
isotope
A
- differs in number of neutrons
8
Q
electron charge
A
-1
9
Q
electron mass
A
0 amu
10
Q
charge
A
- number of protons and electrons
11
Q
C>0
A
cation
12
Q
C=0
A
atom
13
Q
C<0
A
anion
14
Q
Bohr model of the atom
A
- electrons orbit at fixed distances from the nucleus
- distance between orbits decreases with distance from the nucleus
15
Q
quantization of electron energies
A
- electron energies are quantized are are related to their fixed radius orbits
- energy increases with distance from the nucleus
16
Q
principle quantum number, n
A
- describes the radial distance of an electron’s orbit from the nucleus.
17
Q
electron transitions excitation
A
- electrons absorb only specific, allowed quantities of energy
- allowed energies match the energy difference between an electron’s ground state and excited state.
- initial promotion most difficult
18
Q
electronic transitions relaxation
A
- electrons in an excited state can return to a lower energy orbit, emitting a photon equal in energy to the energy difference between the energy levels.
- electrons can return to the ground state in a single transition, or in multiple transitions.
19
Q
hydrogen absorption spectrum
A
- dark bands on a light background
20
Q
hydrogen emission spectrum
A
- bright bands on a dark background
21
Q
lowest to highest energy of waves in electromagnetic spectrum
A
- radio
- micro
- IR
- ROYGIV
- UV
- Xray
- Gamma
22
Q
energy of a photon
A
E = hf = (hc)/λ
23
Q
quantum model of the atom
A
- electrons exist within 3-D orbitals of various sizes and shapes
- electron energies are quantized and are related to their specific orbital
- four quantum numbers fully describe the electronic structure of the quantum model
24
Q
quantization of electron energies
A
- energy increases with distance from the nucleus
- energy increases with complexity of the orbital shape
25
energy shell and subshell
- each period corresponds to a different energy shell
- each shell is higher energy and larger than the last.
- each block corresponds to a different energy shell
- each subshell is more complex and higher energy than the last.
26
magnetism and spin
- each subshell higher than s has multiple orbital orientations
- each orbital orientation has the same energy
- each orbital can only hold up to 2 electrons
- each electron can be spin up or spin down
27
Aufbau principle
- electron added to orbitals from lowest to highest energy
- valence electrons are in the highest energy shell
- electrons first removed from valence orbitals (outermost) from highest to lowest energy
28
Hund's rule
- electrons fill degenerate orbitals one per orbital before pairing
- occupy singly before pairing
29
paramagnetic
- at least one electron is unpaired
| - attracted to a magnetic field
30
diamagnetic
- all electrons are paired
| - doesn't like magnetic field
31
Pauli principle
- no two electrons may be identical
| - this limits the occupancy of an orbital to a maximum of 2 electrons.
32
column 4 and 9 of d-block
- remove the first electron from the highest s subshell first then add to d subshell.
- then remove the rest from the d subshell.
33
excited state of electron configuration
- any number of configurations that have higher energy than the lowest energy electron configuration
- make sure the configuration has the correct total number of electrons
- the electrons can be in ANY orbital as long as that orbital exists.
34
effective nuclear charge (Zeff)
- the nuclear charge experienced by valence electrons
- the pull of protons on valence electrons
- reduced due to shielding
- increases from left to right and from bottom to top on the periodic table. Points to NY
35
atomic radius
- as the force increases, valence electrons are pulled more strongly toward the nucleus, decreasing atomic radius
- increases from right to left and top to bottom on the periodic table. atomically hot in death valley
36
ionic radius
- valence electron repulsion is slightly increased in anions
- valence electron repulsion is slightly decreased in cations.
- ionic radius increases with increasing negative charge
37
shielding
- core electrons shield the valence electrons from the full nuclear charge
- the nuclear charge experienced by a valence electron is Zeff
38
ionization energy
- the minimum amount of energy required to remove the outermost electron from an atom
- increases from left to right and bottom to top
- less energy to pull of an electron as they are further from the nucleus with increasing atomic radius
- hard to do from for nobel gases.
39
electron affinity
- energy change to add an electron to an atom.
- most elements release energy upon the addition of an electron
- the larger the energy change, the more stable the resulting ion
- becomes more negative from left to right and bottom to top
40
multiple ionization
- as the charge on a given ion increases, so too does its ionization energy
- as the extent of ionization increases, valence electron repulsion decreases.
- the second ionization energy is greater than the first.
41
electronegativity
- the ability of an atom to attract electrons to itself in a covalent bond
- increases from left to right and from bottom to top
- F>O>N>C>O>L>Br>I>S>C=H
42
Acidity
- the measure of a compound's ability to lower the pH of a solution, donate protons or accept electrons
- depends on the relative stability of the acid and its CONJUGATE BASE!
- as the size of the anion increases, its stability increases
- increases from left to right and from top to bottom on the periodic table. Points to Florida where there is a lot of acidic orange juice.
43
d block
- transition metals
44
s and p blocks
- representative elements
45
f block
- rare earth metals
