Test 3 Flashcards
4th quantum spin(s)
Electron is a spinning charge, can either spin up or down. Represented by either +1/2 or -1/2
Auf bau principal
lowest energy occupied first
Polyexclusion principal
each electron has a unique set of quantum numbers
Hund’s principal
parallel spins in different boxes before pairing up to minimize C-e repulsion
Magnetic properties
-related to unpaired electrons
Diamagnetic
Unaffected by magnetic field, all electrons paired
Paramagnetic
Attracted by magnetic field, has unpaired electrons
Electron configuration of Ions
Cations loose e
Highest n always looses first
If more than one highest n highest l lost first
Electron configuration of Anions
gain e to get to noble gas configuration
add e as usual
Valence electrons
part of s and p orbitals and non filled f and d orbitals
non valence electrons
filled d and f orbitals are not valence electrons
Electron configuration of Cr
[Ar}4s^13d^5
Electron configuration of Cu
[Ar] 4s^13d^10
Effective nuclear charge
“charge” felt by other electrons. Ability to pull in electrons
How to calculate Zeff
Z(# of protons) - S(screening electrons closer to the nucleus)
Periodic table ->(Zeff)
Zeff increases as #of protons increases, but not the # of screening electrons
Periodic Table goes down(Zeff)
Zeff is constant
Periodic trends in atomic size
n and Zeff
As n increases, radius increases (PT goes down)
As Zeff increases, radius decreases (PT ->)
Size of transition elements
their sizes do not change much
Size of atoms vs cations
Cations(+) < atoms
Because Zeff increases
Size of atoms vs anions
Anions(-) > atoms
Because Zeff decreases
Trends in first ionization energies
closer is harder to ionize and vice versa. As Zeff increases leads to harder removal. As n increases, becomes easier to remove.
2nd and higher ionization energie
I1 < I2 < I3 < …
Ionic bonding
A cation and an Anion
electrons are given
Covalent bond
Generally non-metal + non-metal
electron sharing
Metallic bonding
electrons are mobile in what is called an “electron sea”
conductors
Why do covalent bonds share electrons
To get lower in energy
Lattice energy
the energy it takes to separate one mol of ionic into infinite separation
Coulomb’s law equation
E = (kQ1Q2)/R
Electronegativity
The force that an atom attracts electrons in a covalent bond.
As # of bonds increase
strength of the bond does too
Formal charge
Charge on an atom in a molecule by assigning one e- to each end of a single bond. Can be calculated by using original # of electrons - current #
Sum of the formal charges
charge on species. represents the charge of the molecule. Wants each formal charge to be as small as possible
Electron poor
When an atom is lacking an electron to achieve octet rule, but maintains 0 formal charge
electron rich
an atom has too many valence electrons, but still maintains 0 formal charge
odd electron
an atom has an unpaired valence electron, but still maintains 0 formal charge
Bond energy
it takes energy to break a bond and releases energy when a bond is formed. The average value can be found on the table at the back of the data sheet
molecular geometry
shapes of molecules and bond angles. It takes at least 3 atoms to make an angle
steric number (sn)
number of electron groups around an atom. Lone pairs + atoms attached to central atom
example of electron poor molecule
BF3
example of odd electron
NO
example of electron rich
PF5
Angle and name of steric number 2
180 and linear
angle and name of sn 3
120 degrees and trigonal pyramid
name and angle of sn 4
109 degrees and tetrahedral
name and angle of sn 5
6 90 degrees, 2 120 degrees and 1 180 degrees, trigonal bypyramid
angle and name of sn 6
90 and 180 degrees, octahedral/square pyramidal
dipole moment and polarity
if a force of attraction from a polar bond creates a net vector that is not 0, the molecule is polar
Anion
negatively charged ion
cation
positively charge ion
Orbital hybridization
When orbitals intersect during bonding
Types of hybridization
SN 2: sp, p+p SN 3: sp2, p Sn 4: sp3 SN5: sp3d SN6: sp3d2
Single bond type of hybridization
Sigma
Double bond type of hybridization
pi
Triple bond type of hybridization
sigma and pi
Intermolecular forces
forces between molecules
Dipole dipole interaction
Polar molecules, the more polar it is, the higher the dipole-dipole force
London dispersion force
Depends on the number of electrons
the higher the molar mass, the higher the force
Hydrogen bonding
If H is connected to a small electro negative element(F,O,N), it is attracted to a lone pair on another electro negative element
Strength of each inter molecular forces
Hydrogen bonding > dipole - dipole > london dispersion force
boiling point related to inter molecular forces
The higher the inter molecular forces, the higher the boiling point
Isomer
Two molecules with the same molecular formula but different structures(Ie Ch3OCH3 and CH3CH2OH)
vapor pressure in relation to molar mass
the smaller the molar mass the higher the vapor pressure
vapor pressure compared to intermolecular forces
higher the forces, lower the vapor pressure
Vapor pressure in relation to temperature
vp increases as t increases
