The History of the AtomDemocritus (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 AtomLouis 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 BondingValence 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 NumbersQuantum 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 | ml | Direction/orientation | -l ≤ -ml ≤ l | 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 TheoryVSEPR: 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 |
Bond vs Molecular PolarityBond 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 CrystalsSolids 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 CrystalsSolids 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 BondsIonic 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 CrystalsSolids 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 CrystalsSolids 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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