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Gr. 12 Structure and Properties of Matter Cheat Sheet by

The History of the Atom

Democritus (300 B.C.)
First person to conceive the idea of tiny, indivi­sible partic­les called atoms
John Dalton (1805) - The billiard ball model
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)
J.J. Thomson (1897) - The raisin bun model
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
Ernest Ruther­ford (1903) - The beehive model
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
Niels Bohr (1913) - The planetary model
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 ≤ ln - 1
Magnetic quantum number
ml
Direct­ion­/or­ien­tation
-l ≤ -mll
Spin quantum number
ms
Spin
±1/2

Shape of Electron Orbitals (l and ml)

Value of l
Symbol
Shape
# of suborb­itals (ml)
0
s (sharp)
1 (ml = 0)
1
p (princ­ipal)
3 (ml = -1, 0, 1)
2
d (diffuse)
5 (ml = -2, -1, 0, 1, 2)
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
AX2
Linear
3
Trigonal planar
AX3
Trigonal planar
  
AX2E
Bent
4
Tetrah­edral
AX4
Tetrah­edral
  
AX3E
Trigonal pyramidal
  
AX2E2
Bent
5
Trigonal bipyra­midal
AX5
Trigonal bipyra­midal
  
AX4E
See-saw
  
AX3E2
T-shape
  
AX2E3
Linear
6
Octahedral
AX6
Octahedral
  
AX5E
Square pyramidal
  
AX4E2
Square planar
See a visual table here.

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), K3PO4 (potassium phosph­ate), CuSO4 (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: I2(s) (iodine), At2(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
                               

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