Atomic Structures and Properties Flashcards
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)
Electron Affinity
the measure of change in energy that occurs when an electron is added to the outer energy level to form an anion
- decreases down (the effect of the nuclear charge will not have as great of an effect on a large atom)
- increases across (effect will be greater on smaller atoms)
What elements can have less than 8 electrons for a stable octet?
Hydrogen
Boron
Beryllium
What elements can have more than 8 electrons for a stable octet?
Sulfer
Phosphorus
Xenon
What does VSEPR stand for?
Valence Shell Electron Pair repulsion theory
Linear
2 total bonds
2 bonds = linear (180)
2 lone pairs = linear (180)
Trigonal Planar
3 total bonds
3 bonds, 0 lone = trigonal planar (120)
2 bonds, 1 lone = bent (<120)
1 bond, 2 lone = 120
Tetrahedral
4 total bonds 4 bonds, 0 lone = Tetrahedral (109.5) 3 bonds, 1 lone = Trigonal pyramidal (107) 2 bonds, 2 lone = bent (104.5) 3 bone, 1 lone = linear (180)
Trigonal Bipyramidal
5 total bonds 5 bonds, 0 lone = trigonal bipyramidal (120/90) 4 bonds, 1 lone = See-saw 3 bonds, 2 lone = T shaped 2 bonds, 3 lone = Linear 1 bond, 4 lone = Linear
Octahedral
6 total bonds (90) 6 bonds, 0 lone = octahedral 5 bonds, 1 lone = square based pyramidal 4 bonds, 2 lone = square planar 3 bonds, 3 lone = t shaped 2 bonds, 4 lone = linear 1 bond, 5 lone = linear
Intramolecular
forced within the molecule, much stronger than inter molecular
Intermolecular
forces between molecules
Properties of Ionic solids
- Brittle
- conduct electricity in liquid form or solution (not solid)
- high melting point
Properties of metallic solids
- electrical conductors
- malleable / ductile
- high melting point
properties of small molecular solids
- low melting point
- not very hard
- non conductors
Properties of covalent network crystals
- extremely hard
- brittle
- very high melting point
- non conductors (except graphite)
- insoluble
