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90 Cards in this Set
- Front
- Back
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Speed of light...
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1. 3.00 X 10^8 m/s
2. 186,000 miles/second |
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How is color determined?
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1. By the different wavelength distances.
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Large wavelength would equal colors closest to...
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1. Red (left side)
2. 750 nm wavelength longest visible light |
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Short wavelength would equal colors closest to...
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1. Violet (right side)
2. 400 nm wavelength shortest visible light |
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Constructive Interference
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1. In Phase: align/overlap
2. Reinforces amplitude (double, etc) |
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Electromagnetic radiation
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1. Light
2. Type of energy embodied in oscillating electric and magnetic fields. |
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Electric field...
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1. region of space where electrically charged particle experiences a force.
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Magnetic field...
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1. region of space where a magnetic particle experiences a force.
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Wavelength...
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1. Distance in space between adjacent crests.
2. Measured in meters, micrometers, etc. |
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Amplitude...
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1. vertical heigh of a crest.
2. Determines how bright a light is... higher amplitude=brighter light. |
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Frequency..
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1. Number of waves per second
2. Cycles per second |
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Electromagnetic spectrum...
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1. Includes all wavelengths of electromagnetic radiation.
2. Short wavelength, high frequency is on the right. 3. Long wavelength, low frequency is on the left. |
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Form of electromagnetic radiation with shortest wavelength...
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1. Gamma ray - produced by sun, stars and unstable nuclei on earth.
2. High energy can damage biological molecules. |
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X-rays...
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1. Next to gamma rays on spectrum.
2. Sufficiently energetic to damage biological tissues over time. |
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Ultra-violet radiation...
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1. Component of sunlight that produces sunburn.
2. Excessive exposure increases risk cancer. |
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Visible light...
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1. Ranging from violet (short high energy) to red (long low energy).
2. Usually can't harm biological molecules unless in high intensity. |
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Order of electromagnetic spectrum from right to left.
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1. Gamma ray
2. X-ray 3. UV ray 4. Visible light 5. Infrared 6. Microwave 7. Radio |
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Infrared radiation...
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1. Heat!!
2. Heat you feel when you place hand near hot object. |
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Microwaves...
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1. Efficiently absorbed by water - causing molecules to bounce and release heat, warming substance.
2. Used for radar and microwaves. |
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Radio Waves...
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1. transmit signals for AM/FM radio, cell phone, and television.
2. Longest wavelength |
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What does change in intensity mean?
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1. Change in brightness
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Diffraction...
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1. Ending of a light wave
2. Bending of a wave around obstacle/slit. |
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Light is a ________.
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1. Wave AND particle
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Two-slit experiment showed what...
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1. If single slit - large blob of light on wall
2. If 2 slit - light/dark/light/dark band of light on wall. 3. Showed that light is a wave! |
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Photoelectric effect...
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1. Observation that many metals emit electrons when light shines upon them.
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What are the angular momentum quantum numbers for s, p, d and f?
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S - 0
P - 1 D - 2 F - 3 |
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What was the Ultraviolet Catastrophe?
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1. Error at short wavelengths in the Rayleigh–Jeans law for the energy emitted by an ideal black-body.
2. Rayleigh-Jeans equation worked well at low frequencies, but failed at high frequencies. 3. Equation said that at high frequencies, high energy emitted. |
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Schrodinger's Cat Experiment...
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1. Radioactive material put in box with cat. At one hour there is a 50/50 shot that the radioactive decayed, killing the cat. The cat may still be alive however, we don't know until we look. At this point the cat is both alive and dead because we don't know.
2. Because of the dual nature of quantum physics (the same thing which makes light both a particle and a wave), at that one-hour mark, the cat is considered both alive AND dead... until you open the box and find out for sure. |
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Blackbody Radiation
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1. Radiation emitted by a black-body
2. Determined by temperature |
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Blackbody
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1. Absorbs all light
2. Absorbs all electromagnetic radiation falling on it. |
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Give an example of a blackbody?
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1. Sun
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One quantum of light is a ________.
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1. Photon
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Thermal Equilibrium
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1. Point at which energy transfer due to temperature difference is zero.
2. Both things have same temperature. |
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Uncertainty Principle
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1. The theory that it is impossible to measure both energy and time (or position and momentum) completely accurately at the same time
2. Werner Heisenberg developed 3. ****Can't know both velocity and position at the same time!!!!!!!!!!! |
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de Broglie Equation
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1. Wavelength = planks constant/mass * velocity
2. Electron wavelength is related to its kinetic energy. 3. (Faster electron moving, higher kinetic energy, shorter wavelength.) |
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Willhelm Wien (Wien's Law)
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1. First to attempt to explain blackbody radiation
2. Worked poorly at low frequencies!! |
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Atomic Spectroscopy
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1. Study of electromagnetic radiation absorbed and emitted by atoms.
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Bohr's Model of Atom
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1. Electrons traveled around nucleus in circular orbits.
2. Orbits could only exist at specific, fixed distances from nucleus. 3. Energy of each orbit was fixed (quantized) 4. Said that only when electron jumped from one stationary state to another, it gave off or absorbed radiation. |
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In Bohr's model, transitions between states that were closer together produced light of ________ energy.
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1. Lower
2. Longer wavelengths |
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Balmer and Rydberg and the Rydberg constant.
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1. Developed equation that predicted wavelengths of the hydrogen emission spectrum.
2. Balmer developed six series of spectrum in the visible light and UV areas. 3. Rydberg Constant: 1.097 X 10 ^7 |
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What is Planck's constant?
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1. 6.626 X 10 ^ -34 J*s
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Radiant energy (photon energy) is ________________ to its frequency.
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1. Directly proportional
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Radiant energy (photon energy) is ________________ to its wavelength.
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1. Inversely proportional
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Energy of a photon equation
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1. E=nhv
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Max Planck Accomplishments - known for
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1. Originator of Quantum Theory
2. Nobel Prize in Physics 1918 3. Did work on blackbody radiation as well. |
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Binding energy
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1. Energy required to knock an electron off metal surface.
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Einstein said that light comes in _______________ called ________.
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1. Packets
2. Quanta |
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Schrondinger's Equation
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1. Predicts that wavefunctions can form standing waves, called stationary states/orbitals.
2. Hv=Ev 3. Allows for partial solving to determine orbital shape. |
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Hund's Rule
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1. Greater total spin state usually makes the resulting atom more stable.
2. When electrons fill degenerate (same) orbitals, they first fill them singly with parallel spins. |
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Wavefunction?
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1. mathematical function that describes the wavelike nature of an electron.
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Pauli Exclusion Principle
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1. No, two electrons can have the same 4 quantum numbers.
2. No two fermions (one with same spin) can share the same quantum level. |
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JJ Thomson Plum Pudding Model
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1. Electrons enveloped in a positive charged cloud/pudding. They are just mushed together
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Rutherford Atom Model
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1. Central mass location called nucleus
2. Electrons surround nucleus but most mass is in the very center, not spread out like pudding model. |
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Excitation and Radiation
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1. When atom absorbs energy, electron in lower energy level is EXCITED and promoted to higher energy level.
2. But this makes atom unstable and so it quickly falls back to lower energy level. 3. By doing this, light is released. |
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Blackbody radiation is __________, where as hydrogen emission is ___________.
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1. Blackbody - continuous
2. Hydrogen - discontinous |
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What are quanta?
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1. Little packets of light!
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Partial solving of Schrodinger's Equation leads to_______________.
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1. The shape of the orbitals that electrons exist in.
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Can you knock electrons off metal by increasing the intensity of a light?
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1. NO! Only an increase in frequency (or change in color basically) can knock off electrons.
2. Changing intensity of a light won't matter! |
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What is ground state?
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1. Lowest energy state an electron can have
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In Bohr's electron model, can electrons exist on the lines in between levels?
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1. NO! Electrons have to be in one level or the other, they can't be partially in one or the other.
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What does the l stand for?
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1. Angular momentum quantum number
2. determines the shape of the orbital |
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What does the ml stand for?
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1. Magnetic Quantum number
2. Specifies the orientation of the orbital |
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What does the n stand for?
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1. Principle quantum number
2. Determines overall size and energy of an orbital |
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Einstein describes light as a________ in his concept of the__________________.
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1. Particle
2. Photoelectric effect |
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Describe the idea of indeterministic probability versus deterministic?
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1. Deterministic is like the trajectory of a baseball - hits same spot when taking same force each time
2. Indeterministic is like the trajectory of particle, sometimes hits here and sometimes there, but overall follows somewhat same path. |
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The periodic table is valuable because it is___________.
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1. Predictive
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Mendeleev claimed that the periodic properties of elements were a function of their _________________. However, a better interpretation should state that periodic properties are a function of _________________.
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1. Atomic masses
2. Atomic number |
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Aufbau Principle...
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1. The pattern of orbital filling in an atom.
2. Aufbau: build-up 3. Fill in electrons beginning with lowest energy orbitals first. |
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Core Electrons...
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1. Electrons in a complete principle energy level and those in d and f sublevels.
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Electron configurations can be abbreviated by using a ________________ fore core electrons.
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1. Nobel Element
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Transition Metals
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1. d block
2. Atomic radii as move right in transition elements stays about the same. |
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Inner Transition Metals
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1. f block
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Sublevel Splitting
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1. Splitting of sublevels is caused by electron-electron repulsions.
2. 3d wants to have 5 or 10 because its stable, will pull from 4s!!!!!!!!!!!!! |
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Elements whose existence were predicted by Mendeleev's Table...
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1. Gallium
2. Scandium 3. Samarium 4. Holium 5. Thulium 6. Gadolinium 7. Neodymium 8. Praesodymium 9. Dysprosium 10. Germanium |
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Full shells and the "full shell look" are ________ configurations of electrons in elements and ions.
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1. Stable
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Effective Nuclear Charge (Zeff)
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1. The actual nuclear charge experienced by an electron.
2. Charge of the nucleus MINUS the charge of the shielding electrons. 3. NET NUCLEAR CHARGE 4. Protons - non-valence electrons |
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Ionization Energy
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1. Energy required to remove an electron from the atom or ion in a gaseous state.
2. 1st Ionization Energy: energy required to remove the 1st electron. 3. NE |
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Trends in Ionization Energy in columns
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1. Ionization energy decreases as you move down a column/family because electrons in the outermost level become farther away from the nucleus and are less tightly held.
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Trends in Ionization Energy in rows
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1. Ionization increases as you move to the right across a row/period because electrons in outermost level experience greater effective nuclear charge.
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Trends in Metallic Character in columns
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1. As you move down a column/family, metallic character increases.
2. SW |
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Trends in Metallic Character in rows
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1. As you move right across a row/period, metallic character decreases.
2. SW |
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Shielding...
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1. Core electrons shield electrons in the outermost energy levels from nuclear charge.
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Trends for Atomic Radii in columns
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1. Move down a column, larger atomic radii
2. Due to principle quantum number increase, which means larger orbitals. 3. SW |
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Trends in Atomic Radii in rows
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1. Move right, smaller atomic radii
2. Due to effective nuclear charge increase resulting in a stronger attraction between electrons and nucleus. 3. SW |
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Ionic Radii Trends with cations and atoms/anions and atoms.
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1. Cations smaller than atoms
2. Anions larger than atoms |
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Paramagnetic versus diamagnetic
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1. Attracted to external magnetic field - paramagnetic
2. Not attracted to external magnetic field - diamagnetic. |
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VanderWall Radius
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1. 1/2 distance between center of each nucleus of two atoms side by side.
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Covalent Radius
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1. Bonding radius
2. 1/2 distance from one nucleus to another. |
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Penetration
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1. Electrons spend some time overlapping in other orbitals.
2. Causes outer shell electrons to experience greater nuclear charge. (4s versus 3d) |
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Electron Affinity
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1. More negative as you go right
2. E 3. Energy change associated with gaining of an electron by an atom in gaseous state. |
