Unit 2: The Atom Flashcards
Democritus
all matter is made of indestructible atoms
Dalton
all elements are made of identical atoms and they can combine to make compounds
Thomson
● plum-pudding model
● discovered electron
Rutherford
positive charged nucleus surrounded by eletrons
Bohr
● discovered energy levels
● 2n^2
Chadwick
discovered neutron
Proton
● Positively Charged
● 1.673*10^-27 kg
● 1 amu (1.0073 amu (atomic mass units))
● In the nucleus
Neutron
● Not charged
● 1.675*10^-27 kg
● 1 amu (1.0087 amu)
● In the nucleus
Electron
● Negatively charged
● 9.109*10^-31 kg
● 0 amu (0.00055 amu)
● Orbiting the nucleus in specific regions
Atomic Structure
● Nucleus
➜ Central region
➜ Very dense
➜ Held together by a “strong force”
● The rest is empty space
● Bee in a cathedral analogy - Nucleus = bee, rest of atom = cathedral
Periodic table element (in order going down)
● Atomic # - # of protons
● Element Symbol - 1-2 letters, capitalize first
● Element name
● Average atomic mass
Isotopes
● Different atoms of the same element that have different #s of neutrons
● Identified by mass #
● Mass # = # of protons + # of neutrons for an isotope
● DIFFERENT than average atomic mass on periodic table
Isotope example: Carbon -13
Mass # - 13 C - Symbol
Atomic # - 6
Carbon - 14
Symbol: C - periodic table
Atomic #: 6 - periodic table
Mass #: 14 - given/protons + neutrons
# of protons: 6 - atomic #
# of neutrons: 8 - protons + neutrons = mass #
# of electrons: 6 - electrons = protons in a neutral atom
Average atomic mass
● Given on periodic table
● Not a whole # - NOT mass #
● Takes into account the amount of each isotope naturally present
Average Atomic Mass formula
● Mass # x (% abundance x 100) = ___ amu
● Add values together from each isotope
Which isotope has the greatest abundance: Kr-83 or Kr-84?
Kr-84; the molar mass of Kr (83.796) is closer to 84 than 83.
Energy levels
● 1st/most basic level of electron organization
● rings from Bohr model
● numbered 1, 2, 3, etc.
● bigger number = higher energy/farther from nucleus
● “floors”
Sublevels
● divisions in energy levels
● s, p, d, f
● E.L. 1 - s
● E.L. 2 - s, p
● E.L. 3 - s, p, d
● E.L. 4 - s, p, d, f
Orbitals
● S - 1 orbital
● P - 3 orbitals
● D - 5 orbitals
● F - 7 orbitals
● “rooms”
Electron Configuration: Helium
1s^2
● 1 - energy level
● s - sublevel
● 2 - # of electrons in sublevel
Pauli exclusion principal
● Only 2 electrons max in each orbital
● Electrons must spin in opposite directions
Aufbau Principle
● Fill lowest energy level first
● bau = below
Hund’s Rule
● Electrons must fill every orbital in a sublevel before doubling up
● Bus rule
Electron configuration exceptions
● Chromium - 1s2 2s2 2p6 3s2 3p6 3d5 4s1
- 4s and 3d are very close in energy, half full is more stable
● Copper - 1s2 2s2 2p6 3s2 3p6 3d10 4s1
- Full 3d is more stable than full 4s
Kernel Structure
● Shorthand version of electron configuration
Light
● behaves as a particle and a wave
- wavelength
- frequency
● Wavelength and frequency are inversely proportional
● Energy and wavelength are directly proportional
Heisenberg Uncertainty Principle
● orbital “blur” is the closest to finding electrons
● it is impossible to know the exact location and velocity of a particle at the same time
● observe electrons with light
● electrons absorb light (energy)
● electron’s velocity changes
Match the colors with the wavelength, frequency, and energy
violet - short wavelength, high frequency, high energy
red - long wavelength, low frequency, low energy
Absorption
● electrons absorb energy from the electromagnetic spectrum
● allows them to jump to a higher energy level
● “excited” state
● farther from the nucleus
● black bars
Emission
● electrons release the energy as light
● “ground” state
● closer to the nucleus
● colored bars
Why are colors produced?
● When electrons return to their ground state, they emit energy equivalent to the energy of the colors they emit
● High - purple, low - red
Mendeleev
● organized periodic table
● noticed patterns
Metals
● Usually solids at room temp
● shiny
● ductile/malleable
● good conductors
Non-metals
● solid, liquid, or gas at room temp
● brittle
● poor conductors
● hydrogen
Metalloids
● boron, silicon, germanium, arsenic, antimony, tellurium, polonium, astatine
● touches zigzag
● properties of both
● semiconductors with other elements
Periodic table sections
● Main group elements - s and p block
● Transition metals - d block
● Inner transition metals - f block
● Alkali metals - column 1
● Alkaline Earth Metals - column 2
● Halogens - column 17 (-tennessine)
● Noble gases - last column
Coulombic attraction
force of attraction between positive and negative charges
Effective nuclear charge
● A measure of how much positive pull of the nucleus an electron feels
● Increases across the rows of the periodic table
- more protons = more pull
Electron shielding
● Electrons in higher energy levels feel less of a pull from the nucleus
- farther away
- repelled by inner electrons
- “shielded” by inner electrons
● Increases down the columns of the periodic table as atoms have more electrons/energy levels
Atomic radius
● Radius (nucleus to outer edge of last energy level)
● Increases down a column (more energy levels, less pull)
● Decreases across a row (increase in protons attracts electrons closer to the nucleus)
First Ionization Energy
● Energy needed to remove 1 electron
● Decreases down a column (electron shielding = less attraction = less energy to go against)
● Increases across a row (# of protons increases, harder to remove electrons)
● Matches with electronegativity
Electrogativity
● Measure of how “greedy” an atom is for electrons in a chemical bond
● Decreases down a column (more electron shielding = nucleus has less force)
● Increases across a row (effective nuclear charge = more attraction to electrons)
