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Lesson 1, Lesson 5, Lesson 2, Lesson 4, Lesson 3 - Coggle Diagram
Lesson 1
Recent History:
- Lavoisier: Conservation of mass
Nothing is lost or made, all is transformed
- Proust: Definite properties
Same properties across all samples of a given compound
- Dalton: Multiple properties
Atoms of same element are indivisible and have the same mass
Atoms of dif. elements combine in fixed ratios
Electron:
Discovery of electron charge: Millikan
- Oil drop experiment
Atomized oil falling between two charged plates, charge required to suspend the atoms used, in conjunction with mass to determine charge
Discovery of the electron: Thomson
- Cathode ray tube
Tube with very little air, light rays going from cathode to anode, makes few gas molecules fluoresce, the rays were deflected by magnetic/electric fields, therefore they carried a charge
Atomic models:
Plum pudding: Thomson
- Positive batter
- Negative "berries" throughout
Nuclear model: Rutherford
- Discovery of the proton: Gold foil experiment
Positively charged alpha radiation shot at thin gold foil, some particles deflected at large angles, this would require a dense positive charge (positive nucleus)
The neutron: Chadwick
- Mass of electrons + protons ≠ mass of most elements
Presence of neutral particles
Ancient History:
- Democritus:
Universe made of small, indivisible "atoms"
- Aristotle:
Universe comprised of 4 elements
Lesson 5
Atomic radius:
- Distance from core to last orbital
Biggest: Alkali metals
Smallest: Noble gasses
Trends:
Periodic table:
- Group (column):
Down: + orbitals, + electrons, bigger radii
- Period (row):
Right to left: Noble to alkali = bigger radii
Alkali = less valence = less stable = less attraction = bigger
Attraction on e = nb. protons - nb. core e
Cations and anions:
- Cations (+): lose e, energy shared less, more attraction, so smaller
- Anions (-): gain e, energy shared more, less attraction so bigger
Ionization energy:
- Down a group: + orbitals, - IE, easier to ionize
- Exceptions:
Elements prioritize symmetry
B resists less than Be
O resists less than N
Electron affinity:
- Increases: Left to right, bottom to top
- Radius gets smaller, attraction from nucleus gets bigger
Alkali metals:
- Low melting point
- Gas and water produced when contact w/ water
- Oxidizes in air
-
Lesson 2
Classical physics:
- Matter = particules
- Light energy = waves of electromagnetic energy
- Poses 3 problems
Problem 1: Blackbody radiation
- All objects above 0 K emit radiation
This would mean humans (at 310 K) would emit X-Rays
Answer 1: Planck
- Atoms emit/absorb radiation in discrete quantities (quanta)
- Quanta: Smallest amount of absorbable/emittable energy (electromagnetic radiation)
- E of a single quantum is proportional to the freq. of the radiation
E = hv
- Since E is always emitted in multiples of hv, it acts like a particle
Problem 2: Photoelectric effect
- Metals emit electrons/current when hit by light
- The speed of the emission should vary by wavelength and intensity
- Would produce lag when turning smt on
Experiments proved otherwise (min. freq. needed regardless of current, no lag time)
Answer 2: Einstein
- Light should be thought of as a stream of photon particles rather than a wave
- Using Planck's quantum theory: electrons must absorb one photon of enough energy to jump (min. freq. needed)
- Electron makes jump as soon as it absorbs photon
Problem 3: Atomic Spectroscopy
- When atoms absorb energy, they emit light
- Spectroscopy: single gas emits light when given electrical energy to rid the gained energy
- Different gas = different absorption
More electrons means more light
Answer 3: Rydberg
- Finds a general formula for the Hydrogen line spectrum
- Certain wavelengths required for electrons to jump orbits
Answer 3.5: Bohr's model
- 1) Single electron
- 2) In orbit, it does not emit energy
- 3) If it changes orbit, it's energy must change by 1 photon
Each jump represents a freq.
- Initial state: fundamental level (resting
- Excited: moving up/down
- Ionized: free
- Level determines light emission
Light's wave nature:
- Wavelength:
Distance between peaks
Spd. of Light = Wve. Lngth. X Freq.
- Frequency:
Nb of complete cycles passing a given point
- Amplitude:
Distance from peak to midline
- Velocity:
Speed of wave
Lesson 4
Energy splitting and multielectron atoms:
- In H (w/ only 1 electron), the energy of the orbital depends solely on the value of n
- For multielectron atoms, it depends also on l
Coulomb's law:
- The potential energy between two same charges decreases as they get further away (repulsion)
- The potential energy between two different charges is negative and becomes more negative as they get closer
Shielding:
- Positive charge in nucleus (attracts electrons)
- Core electrons shield this attraction
- Easier to remove electrons
Penetration:
- Some electrons penetrate closer to nucleus
- This lowers it's energy, adding stability
Schrodinger:
- His equation can be modified to account for electron repulsion
- It's only approximate
Lesson 3
Dual nature of light: wave-particle duality
- Both a particle and a wave
-
- m: mass, v: velocity
- Electrons will act like waves
Uncertainty principle: Heisenberg
- Problem with wave-like particles: impossible to know both momentum and position simultaneously
- The electron does not orbit nucleus predictably
Schrodinger: Orbital model
- We can predict where the electron should be
Quantum mechanics:
- Probability that the electron will be in a certain place
Orbitals:
- Wave functions that describe a specific distribution of electron density
Quantum number:
- Describes part of the orbital
n
- Values: 1 to ∞
- Describes a shell
- As n increases, orbital gets larger, electrons a further from nucleus and less stable
l
- Values: 0 to n-1
- Describes shape of orbital
- Defines a subshell
mL
- Values: -l to l
- Describes orbital orientation
ms
- Values: 1/2 or -1/2
- Defines electron spin
Hydrogen orbitals:
1s
2s & 3s
- Sphere
- Larger in size
- Electron density is 0 at nodes
Nb of nodes: n-1
2p
- Tear drop lobes
- Differ in orientation (x,y,z)