The History of the Atom
Democritus (300 B.C.) |
First person to conceive the idea of tiny, indivisible particles called atoms |
|
Pure substances are made up of atoms |
Atoms of the same element are exactly alike |
Atoms cannot be created, destroyed, or divided into smaller particles |
Compounds are formed by joining 2 or more elements |
William Crookes (1875) - Discovery of the electron |
Created an electric discharge tube (a cathode ray tube) with a screen and magnet |
Discovered the bar magnet could deflect/move the cathode rays (they have a charge) |
If he added a paddle wheel inside the tube, it moved (the rays had mass) |
|
Using Crooke's cathode ray tube, determined rays were made up of negatively charged particles called electrons |
Electrons were 2000x lighter than hydrogen, the lightest known element |
|
Conducted the gold foil experiment; if the atom was like Thomson proposed, any alpha particles sent through it would pass straight through |
Most of the particles went through, but some were scattered |
Determined that atoms were mostly empty space, with a small, dense, positively charged nucleus in the centre with e¯ scattered around it |
1932 - determined with James Chadwick that the mass of the nucleus did not equal the mass of the protons only, i.e. electrically neutral neutrons |
|
Proposed that electrons are not allowed to orbit anywhere, but rather they occupy certain defined (fixed) orbits |
Based off experiments with hydrogen atoms and spectroscopes |
Electrons can jump to higher orbits when they are given energy in quantized amounts (no partial amounts), usually in the form of photons (light particles) |
Quantum Mechanical Model of the Atom
Louis de Broglie (1924) proposed that if light waves properties of particles, then particles can have properties of waves |
Erwin Schrodinger (1933) realized that a wave theory and mathematical equations were needed to explain atoms with more than 2 e_ |
Schrodinger's Wave Function |
Contains 3 variables called quantum numbers ( n, ln
, ml
) to help determine a region in space where the electron spends 90% of its time (the atomic orbital) |
A fourth number ( ms
) was added so that all characteristics of atoms could be explained |
Heisenberg's Uncertainty Principle: it is impossible to know both the exact location and speed of an e_ at a given time |
Quantum Theory and Chemical Bonding
Valence Bond Theory: atomic orbitals of one atom can overlap with atomic orbitals of another atom to share a common region of space |
Molecular Orbital Theory: when orbitals overlap, they combine to form new orbitals called molecular orbitals (hybridization); the greater the overlap, the more stable the bond |
Double/Triple Bonds: Sigma (σ) bonds (end-to-end overlap of orbitals) and pi (π) bonds ("sideways" orbitals—usually p orbitals—overlap above and below the plane of the bond) |
Single bond = 1 σ bond; Double bond = 1 σ bond + 1 π bond; Triple bond = 1 σ bond + 2 π bonds |
|
|
Quantum Numbers
Quantum number |
Symbol |
Meaning |
Possibilities |
Principal quantum number |
n |
Energy level |
n Є ℕ (any whole number > 0) |
Secondary quantum number |
l |
Shape of orbital |
0 ≤ l ≤ n - 1 |
Magnetic quantum number |
|
Direction/orientation |
|
Spin quantum number |
|
Spin |
|
Shape of Electron Orbitals (l and ml)
Value of l |
Symbol |
Shape |
|
0 |
s (sharp) |
|
|
1 |
p (principal) |
|
|
2 |
d (diffuse) |
|
|
3 |
f (fundamental) |
|
7 ( ml
= -3, -2, -1, 0, 1, 2, 3) |
VSEPR Theory
VSEPR: Valence Shell Electron Pair Repulsion |
Helps determine the structure around an atom by minimizing the repulsive force between e¯ pairs |
Bonded and lone pair e¯ position themselves as far away as possible from each other |
Lone pairs of e¯ on a central atom repels a little more than bonding pairs; they push the bonding pairs closer together |
VSEPR Molecule Shapes
# of e¯ groups |
e¯ configuration |
AXE formula |
Molecular shape |
2 |
Linear |
|
Linear |
3 |
Trigonal planar |
|
Trigonal planar |
| |
|
|
Bent |
4 |
Tetrahedral |
|
Tetrahedral |
| |
|
|
Trigonal pyramidal |
| |
|
|
Bent |
5 |
Trigonal bipyramidal |
|
Trigonal bipyramidal |
| |
|
|
See-saw |
| |
|
|
T-shape |
| |
|
|
Linear |
6 |
Octahedral |
|
Octahedral |
| |
|
|
Square pyramidal |
| |
|
|
Square planar |
Bond vs Molecular Polarity
Bond Polarity: the even/uneven distribution of e¯ across one bond (can be single/double/triple); determined by ΔEN (difference in electronegativity) |
Molecular Polarity: the even/uneven distribution of e¯ across an entire molecule; determine many properties of the substance |
3 important factors to molecular polarity: presence/absence of polar bonds, shape of the molecule, and presence/absence of lone e¯ pairs |
It is possible to have a non-polar molecule with polar bonds within, if the shape cancels out any vectors created by the bonds.
|
|
Ionic Crystals
Solids in which positive and negative ions are arranged in a crystal lattice |
Boiling/melting point |
High |
Malleability |
Brittle |
Conductivity |
Poor as solid, high as solution |
Solubility in water |
Very soluble |
Hardness |
Very hard (very scratch-resistant) |
Types of forces acting on molecule |
Ionic bonds |
Examples: NaCl (table salt), K 3
PO 4
(potassium phosphate), CuSO 4
(copper (II) sulfate)
Metallic Crystals
Solids composed of individual molecules held together by intermolecular forces (IMFs); "neutral" molecules that form complex crystal lattice in solid state |
Boiling/melting point |
Vary widely |
Malleability |
Ductile (very flexible) |
Conductivity |
High as a solid |
Solubility in water |
Slightly soluble |
Hardness |
Varied |
Types of forces acting on molecule |
Metallic bonds |
Examples: Au (gold), Ag (silver), Ni (nickel), Fe (iron), Co (cobalt), Cu (copper), Zn (zinc), Cr (chromium)
Ionic vs Metallic Bonds
Ionic Bond: Highly electropositive ion (cation) gives up extra e¯ and gives them to highly electronegative ion (anion), then bond through very strong electrostatic attraction between the two ions, creating an ionic crystal structure |
Metallic Bond: Many metal atoms shed a "sea" of e¯ that engulf the metal ions (e¯ are delocalized); pulled from all directions, the metal ions can barely move and pack tightly together in crystalline structures |
Both ionic and metallic crystals take an immense amount of energy to break the bonds between ions; however, since the metal ions are inside the "sea" of e¯, metallic crystals are much more malleable than normal ionic crystals (the e¯ mitigate the effect of shifting and sudden repulsion between the ions).
Molecular Crystals
Solids composed of individual molecules held together by intermolecular forces (IMFs); "neutral" molecules that form complex crystal lattice in solid state |
Boiling/melting point |
Low |
Malleability |
N/A |
Conductivity |
Poor as solids |
Solubility in water |
Varied |
Hardness |
Soft (easy to scratch) |
Types of forces acting on molecule |
IMFs - weaker than ionic/metallic bonds |
Examples: I 2(s)
(iodine), At 2(s)
(astatine)
Covalent Network Crystals
Solids in which the atoms form covalent bonds in an interwoven network; most contain C or Si atoms |
Boiling/melting point |
Very high |
Malleability |
N/A |
Conductivity |
Poor as solids |
Solubility in water |
Varied |
Hardness |
Extreme hardness or softness |
Types of forces acting on molecule |
Covalent bonds (strength increases with more bonds); sometimes IMFs (usually LDF) |
Examples: Diamond, graphite, silicone (not Si (silicon)), semiconductors, buckyballs, nanotubes
|
Created By
Metadata
Favourited By
Comments
No comments yet. Add yours below!
Add a Comment
Related Cheat Sheets
More Cheat Sheets by nescafeabusive32