Atomic Fundamentals Flashcards
atom
a unit of matter
proton
neutron
electron
subatomic particles with positive, neutral, and negative charge
atomic #
the # of protons; defines an element
atomic mass
the # of nucleons (protons+neutrons), aka the mass #; measured in amu
ion
a charged atom; opposite is a “neutral atom”
isotopes
atoms of the same element but different mass #s (still stable though)
average atomic mass
(mass of isotope A x percent abundance of isotope A) + (mass of isotope B x percent abundance of isotope B)
1. multiply each mass & % to decimal forms
2. add products and apply rules (SF dp rules of the GIVEN MASS)
percent abundance
a percentage showing how common an isotope is in the given sample
Dalton’s Atomic Theory
1804
1. all matter is made of tiny particles, atoms, which can’t be subdivided, created, or destroyed (!)
2. atoms from the same element are identical in composition and properties, while dif. atoms have dif. II (!)
3. atoms of different elements combine in simple, whole # rations to form compounds
4. in chemical reactions, atoms combine, separate, or rearrange but are never created, destroyed, or changed
J.J. Thomson’s Theory of Electrons
1897:
Thomson discovered electrons by using a cathode ray tube, showing particles of matter with negative charge, subatomic, and the same across all elements.
The Thomson Model
The plum and pudding model
Rutherford’s Gold Foil Experiment
1908:
The discovery of the nucleus when [—] bounced back, later discovered protons
Rutherford’s Criticisms
- mass # doesn’t equal protons and electrons (James Chadwick discovers neutrons later)
- how do protons stick together? “strong nuclear force”
Speed
always 3.00 x 10^8 m/s
wavelength
“lambda”: measured in nm (10^-9m)
energy
“E”: measured in Joules; low wavelength = high E, high wavelength = low E
emission spectrum
each element’s unique variety of wavelengths
The Bohr Model
- electrons absorb light (energy)
- electrons jump to excited state (up to outer shells)
- electron relaxes back to ground state
- energy is released as light
relationship between energy & wavelengths
small E = less photon emission = long wavelengths
big E = more photon emission = short wavelengths
why do different elements emit different wavelengths?
different energies jump shells at different distances
wave-particle duality
treating electrons as resulting particles of inconsistent behavior, as “loops” vibrating with whole # cycles
subshells
aka sublevels, next level inside shells; s, p, d, and f are different electrons wave shapes
orbitals
next level in subshells; are pockets that hold 2 e-
electron configurations
describes how electrons are arranged around the nucleus of any given atom
orbital filling rules
- Aufbau principle: electrons fill lower energy orbitals before occupying higher energy orbitals
- Pauli Exclusion Principle: at most 2 electrons per orbital, each w dif. spin
- Hund’s Rule: each orbital in a sublevel takes 1 electron before 2 pair up
know how-to:
- aufbau diagram
- extended/full e- config.
- noble gas shorthand
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