Unit 3 Quantum Numbers Flashcards
Quantum numbers
Describe probable location of where an electron is likely to be found in its ground state. (Lowest energy state)
Valence electrons
1-8 the s and p electrons in outermost energy level.
Ground state vs excited state of electron
Ground is lowest energy state, when not in ground state atom is excited.
Space orbital
Region of space around nucleus where there is the greatest probability that an electron will be found.
Principal quantum number-1st
Horizontal rows "periods" on table. Energy level or distance from nucleus where n= a whole number. 1=1s 2=2s, 2p 3=3s,3p,3d 4=4s, 4p, 4d, 4f
Momentum quantum number-L 2nd
Shape of the orbital. Specified by 2nd quantum number. L, where L can have values 0-3 (number of shapes in that sub shell)
L=SPDF
0,1,2,3
L=0 if The electron is in the S sub shell
L=1 if the electron is in the P sub shell
L=2 If the electron is in the D subshell
L=3 If the electron is in the F subshell
Magnetic quantum number-3rd
Describes position with respect to three axes on space (x,y,z).
SPDF=1,3,5,7
Ml=0
If Electron is in the S subshell. 1 shape
Ml=-1,0,1
The electron is in the P subshell. Px=-1 Py=0 Pz=+1 3 shapes
Ml=-2, -1, 0, +1, +2
in D subshell. 5 shapes
Ml=-3,-2,-1,0,+1,+2,+3
F subshell 7 shapes
Number of positions for each type
S=1 P=3 D=5 F=7
Spin Quantum number-4th
Direction of electron spin. 4th quantum number. Ms. -1/2 to +1/2
Electron configuration
How electrons are arranged
Determines reactivity
Aufbau principle
Orbitals with the lowest energy are occupied furst
Hunds rule (spins)
Electrons with parallel spins will enter unoccupied orbitals of the same energy level one at a time BEFORE pairing up.
Pauli exclusion principle
No two electrons can be described by the same set of quantum numbers. Cannot have the same n, L, Ml, and Ms
Representative elements
S and P electrons. 1-2 are S. 3-8 (13-18) are P.
Transition elements
D area
Actinide and lanthanides
Inner transition-F area
Order of filling orbitals-D, F
D n value is one lower than row. F n value is 2 lower than row.
Electron dot notation
Outermost (valence) electrons only. S and P main group elements only. Not atomic number.
Group number = number of
Valence electrons
P electron configuration
.
.P:
.
n=
Row number
Periodic law
Properties of elements show up periodically when elements are arranged in increasing order of atomic number.
Groups (families)
Vertical columns
Group 1
Alkali metals-react w water to form strong bases
Group 2
Alkaline earth metals-also react with water but not as reactive as group 1.
Group 7
Halogens-salt forming. React w metals to form salt
Group 8
Noble gases-do not react well. Not reactive.
Groups in D region
Transition metals
Groups in F region
Inner transition metals
Periods (series)
Horizontal rows, left to right
Period 1
H and He-2 elements
Period 2 and 3
8 elements each
Period 4 and 5
18 elements each
Period 6
Contains Lanthanides series (inner transition) 32 elements
Period 7
Contains actinide (inner transition) 32 elements
Metals-properties
85%. Left side of zig zag line. Conductive, malleable, dense, solid except Hg, ductile (wire), luster
Non metals-properties
Not conductive, gases, low m.p., insulating, dull, brittle
Metalloids
Border the zig zag line, show properties of metals and nonmetals. Al is considered metal. Properties-semi conductive
Isoelectronic
The name given to ions that have the same electron configuration as Atoms of noble gases
Ions
Charged particles that gain electrons (-charge) Or lose electrons (+charge) To obtain the same electron configuration as a noble gas to become stable
Metals usually lose electrons to become on isoelectronic with the noble gas in the previous energy level
Nonmetals usually gain electrons to become isoelectronic with the nearest noble gas.
On the periodic table from top to bottom within a group
Atomic radius gets bigger because they’re more energy levels and the valence electrons are further from the nucleus
From left to right within a Period
Atomic radius get smaller- protons and electrons added as you go from left to right (More attraction to nucleus) But no new (bigger) energy levels are added
Ionization energy
Energy needed to remove an electron from a neutral atom(Gas phase)
Low ionization energy
Characteristic of metals because It gets them closer to eight which is more stable they want to give them away. LOSE
High ionization energy
Characteristic of nonmetals because they need electrons to get eight to make them more stable so they want to take them away. GAIN
Within a group of representative elements the ionization energy generally
Decreases with increasing atomic size because they’re further from the nucleus and lose electrons easier
Within a period Ionization energy generally
Increases since elements become less willing to give up electrons
Ionization energy gets incredibly larger if the Adam has already
Lost the valence electrons necessary to achieve an octet
