Atomic Theories Flashcards
How did bohr expnd on rutherfords theory of the atom?
Quantized the shells
Electron ins ground state vs excited state
Ground has lower enegery
When an electron falls from higher engery level to lower enegry levels, how is energy released?
as a photon
Difference between previous model of atom vs modern quantum mechanical model?
2D vs 3D
Further from the nucleus the ____ energy an electron has.
More
region of space in which high probablity of finding an electron
Orbital
term used to label energy levels of electrons
n=
How are s orbitals different from p orbitals
different shape s =O p= 8
How does atom valence elctron configuration determine place on periodic table?
Valence determines block or group number
What two elements are exceptions to the way we normally write electron configurations? Write
the expected and the actual confi uration of each. What rules are followed? What ru es are
violated?
Cu= [Ar] 4s1 3d10
Cr= [Ar]4s0 3d5
full subshells are stable
Transition metals
titanium, chronium, mecury
Discovered the nucleaus
Ruther
Cathode ray tubes
Ruther
Discovered proton
Ruther
Most of atom empty space
Rutherford
Partial Positive
if something is a partial positive it’s electronegativity will be less than the other elements electronegativity
•if the molecule has more than two elements you look at the individual bonds
Partial Negative
if something is a partial negative it’s electronegativity will be higher than the other elements
Polar Bonds
to know if a molecule has polar or non-polar bonds you look at the difference in electronegativity
•if the difference is above 0.3 its polar bonds
•if the difference is below 0.3 its non-polar bonds
Polar Molecules
to know if a molecule is polar or non-polar, you can look at its symmetry
•if its symmetrical the molecule is non-polar
•if its asymmetrical the molecule is polar
London Forces
if the molecule is non-polar you will only have london forces
Dipole Dipole
if the molecule is polar you will have dipole dipole forces (only if you don’t have hydrogen bonds)
Hydrogen Bonds
if you only have a H with a 𝒮+and a N, O, or F with a 𝒮- you will have a hydrogen bond
Density Formula
Mass/ Volume
Ion
an atom that loses or gains electrons (to form a full valence shell) and results in a charge. It is a charged particle.
Cations
lose electrons
become more positive
Anions
gains electrons
becomes more negative
Orbitals
3D space around a nucleus where there is a probability of finding electrons.
S subshell
Principal Energy Level: +1
Shape: spherical
# of orbitals: 1
Max # of Electrons: 2
P subshell
Principal Energy Level: +2
Shape: Dumbell like
# of orbitals: 3
Max # of Electrons: 6
D subshells
Principal Energy Level: +3
Shape: vary
# of orbitals: 5
Max # of Electrons: 10
F subshells
Principal Energy Level: +4
Shape: Vary
# of orbitals: 7
Max # of Electrons: 14
Aufbau Principal
each electron is added to the lowest energy level available.
Hund’s Rule
for orbitals at the same energy level, one electron occupies each sub orbital before the electrons pair up.
Pauli Exclusion Principal
only two electrons (with opposite spins) can occupy each orbital
Lewis Bonding Theory
atoms/ions are stable if they have an octet of electrons •electrons are most stable in pairs
•atoms form chemical bonds (ionic or covalent) to become stable
VSEPR Theory
VSEPR (Valence Shell Electron Pair Repulsion)
•valence electrons stay as far away as possible to minimize repulsion
•when looking at a molecule we look specifically at the central atom (the one that has the most bonding electrons) to determine the 3-D geometry
