Atomic Structures and Properties Flashcards
(39 cards)
Democritus’ theory
All matter is made up of indivisible particles called “atoms” and void, which is the empty spaces between atoms
John Dalton’s Theory
- Atoms are indivisible and can’t be broken down
- all matter is made up of tiny particles
- all atoms of one element are are identical
- atoms create compounds by combining atoms of different elements
JJ Thomson’s Theory
- Cathode rays of negative charge, that were deflected by a magnetic field & how much energy they carry
- atom is really a positive field with negative charges embedded within it’s matrix (raisin bun)
The Rutherford model
conducted the alpha particle experiment (gold foil)
- most particles went through
- a small fraction had a large deflection
- a minute fraction rebounded
Rutherford’s nuclear atom model
- atom is mostly empty space
- All positive charge is concentrated in a small volume called the nucleus
- electrons revolve around the nucleus like planets in the solar system
Plank / Einstein - Quantum
- plank determined that energy is absorbed by atoms in certain fixed amounts known as quanta
- Einstein extended the theory by determining that radiant energy is also quantized
- discrete energy packets = photons
Einsteins theory
electromagnetic radiation has characteristics of both a wave and a stream of particles
The Bohr model
- Electrons revolve around the nucleus in certain allowed orbits; each orbit corresponds to a specific amt of EN
- As long as the electron remains in the same orbit, it neither emits nor absorbs EN
- As the electron jumps from one orbit to another, EN is absorbed or emitted
- EN difference between orbits corresponds to specific wavelengths
- Bohr’s calculations only worked for hydrogen
DeBroglie (wave mechanical model)
-understanding that any small particle, such as an electron in motion, has associated wave behaviour
Schrodinger (wave mechanical model)
- considered the behaviour of the inside of an atom
- The positive nucleus is surrounded by a cloud of electrons waves, electrons can only have quantized energy levels because the requirement for whole # of wavelengths for electron waves
Heisenberg uncertainty principle (wave mechanical model)
-impossible to know both the velocity and location of an electron at the same time
S orbital
- Spherical in shape
- size increases as “n” increases
- There is only 1 S orbital in a sublevel
P orbital
- Dumb-bell shaped
- Aligned along x, y, z axis
- only 3 p orbitals in a sublevel
- size increases as “n” increases
D orbital
- has 4 lobes per orbital
- aligned according to x, y, z axis
- only 5 d orbitals in a sublevel
F orbital
- don’t need to know the shape
- only 7 f orbitals in a sublevel
Principal quantum number (n)
-ENERGY LEVEL
refers to the major (or principal) energy levels in an atom
-the higher n is the farther away the electrons are from the nucleus
Angular momentum number (l)
-SUBLEVEL (S, P, D, F) energy sublevel -shape of orbital -l = n - 1 0 = s 1 = p 2 = d 3 = f
Magnetic quantum number (ml)
-ORBITAL ORIENTATION
Orientation of the orbital
-specifies the exact orbital within each sublevel
-l to +l example: l = 2 ml = -2, -1, 0, 1, 2
Spin quantum number (ms)
-SPIN OF ELECTRON
Electron spin +1/2 -1/2
an orbital can hold 2 electrons that spin in opposite directions
Pauli exclusion principal
- no two electrons in an atom can have the same 4 quantum numbers
- each electron has a unique address
- Change of spin on electron
Aufbau principal
each electron goes into the lowest available energy state, once that is full the next lowest starts filling
Hund rule
bus seat principal
-orbitals with the same energy levels (three 2p) electrons will occupy all empty orbitals first before a second electron goes into the orbitals
Atomic size
- down the periodic table the size increases (more energy levels, and inner electrons shielding the nucleus)
- across the periodic table the size decreases (more protons inside the nucleus, electrons pulled inwards)
Ionization energy
- The amount of energy required to pull off an electron
- decrease as you go down (easier to pull an electron off a bigger atom)
- increase as you go across (more difficult to remove an electron off of a smaller atom)
