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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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