Chemistry: Chapter 7: Atomic Structure & Periodicity Essay

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electromagnetic radiation
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one of the ways that energy travels through space Gamma, X-Rays, UV, Visible, IR, Micro,Radio (FM, short, AM)
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wavelength
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-distance between two consecutive peaks or troughs in a wave -measured in meters
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frequency
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-number of waves that pass a given point per second -measured in hertz (sec⁻¹)
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speed
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measured in meter / sec 3.0E8 m/s
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inverse relation
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(wavelength)(frequency) = speed of light
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Quantum Theory
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-Energy is gained or lost in whole number multiples of the quantity hv (plank’s Constant & an integer) /E= nhv
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quanta
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energy is transferred to matter in packets of energy
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Planck’s constant
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h = 6.626E-34 (J)(s)
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Gamma
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Wavelength =10⁻¹² m
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X-Rays
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Wavelength = 10⁻¹⁰ m
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UV
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Wavelength = 10⁻⁸ m
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Visible
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4 to 7 x 10⁻⁷ m blue green yellow orange red wavelength increases frequency decreases energy decreases
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IR
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Wavelength = 10⁻⁴ m
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Micro
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Wavelength = 10⁻² m
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Radio
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FM –> wavelength = 1 short –> wavelength = 10² AM –> wavelength = 10⁴
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Particle Nature of Matter
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1. EM radiation is a stream of particles – photons 2. Energy and mass are inter-related Einstein
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Dual Nature of Light
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1. light travel through space as a wave 2. light transmits energy as a particle 3. particle’s have wavelength, exhibited by diffraction patterns -large particles have very short wavelengths -all matter exhibits both particle and wave properties de Broglie
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Photoelectric Effect
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-when light strikes a metal surface, electrons are ejected -if the threshold frequency has been reach, increasing the intensity only increases the number of electrons ejected -if the frequency is increased, the ejected electrons will travel faster
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Continuous spectra
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contains all wavelengths of light
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Emission spectra
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Bright Line Spectra 1. excited electrons in an atom return to lower energy states 2. Energy is emitted in the form of a photon of definite wavelength 3. Definite change in energy corresponds to: -definite frequency -definite wavelength 4. Only certain energies are possible within any atom
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Absorption Spectra
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-A spectrum of electromagnetic radiation transmitted through a substance, showing dark lines or bands due to absorption of specific wavelengths -when electrons get excited they absorb light
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Bohr Model
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Quantum model 1. The electron moves around the nucleus only in certain allowed circular orbits 2. Bright line spectra confirms that only certain energies exist in the atom, and atom emits photons with definite wavelengths when the electron returns to a lower energy state 3. Energy levels available to the electron in the hydrogen atom n = an integer Z = nuclear change J = energy
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Calculating the energy of the emitted photon
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1. calculate electron energy in outer level 2. calculate electron energy in inner level 3. calculate the change in energy (/E) /E = energy of final state – energy of initial state 4. use the equation to calculate the wavelength of the emitted photon
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Energy Change in Hydrogen Atoms
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1. Calculate energy change between any two energy levels
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Shortcomings of the Bohr model
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1. Bohr’s model does not work for atoms other than hydrogen 2. Electron’s do not move in circulate orbits
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The Quantum Mechanical Model of the Atom
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The electrons as a standing wave 1. Standing waves do not propagate through space 2. Standing waves are fixed at both ends 3. Only certain size orbits can contain whole numbers of half wavelength (fits the observation of fixed energy quantities “quanta”)
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The Shrodinger Equation
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1. for the motion of one particle, along the x axis in space: 2. Solution of the equation has demonstrated that E (energy must occur in integer multiples 3. General equation: H = a set of mathematical functions “operator” = a wave function
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orbitals
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specific wave functions 1. Are not circular orbits for electron 2. are areas of probability for locating electrons
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Heisenberg Uncertainty Principle
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1. There is a fundamental limitation on how precisely we can know both the position and momentum of a particle at a given time 2. The more accurately we know the position of any particle, the less accurately we know its momentum, and vice-versa
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Physical Meaning of a Wave Function
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1.Square of the absolute value of the wave function gives a probability distribution 2. Electron density map indicates the most probable distance from the nucleus 3. Do not describe -how an electron arrived at its location -where the electron will go next -when the electron will be in a particular location
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principal quantum number (n)
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-integral values: 1, 2, 3…. -indicates probable distance from the nucleus higher numbers = greater distance = less tightly bound = higher energy
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angular momentum quantum number (l)
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orbital quantum number -integral values from 0 to n – 1 for each principal quantum number n -indicate the shape of the atomic orbitals
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magnetic quantum number (ml)
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-integral values from l to -l, including zero -related to the orientation of the orbital in space relative to the other orbitals
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spin quantum number (ms)
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-an orbital can hold only two electrons, and they must have opposite spins -spin can have two values, +1/2 & -1/2
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nodes
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internal regions of zero probability
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electron configuration
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1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p→6s→ 4f→5d→6p s = 1 / p = 3 / d = 5 / f = 7
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size of orbitals
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-defined as the surface that contains 90% of the total electron probability -orbitals of the same shape grow larger as n increases
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s Orbitals
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-spherical shape -nodes (s orbitals of n=2 or greater)
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p Orbitals
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-two lobes each -occur in levels n=2 and greater Each Orbital lies along an axis (2px, 2py, 2pz)
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d Orbitals
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-occur in levels n=3 and greater -two fundamental shapes 1. four orbitals with four lobes each, centered in the plane indicated in the orbital label (dxz, dyz, dxy, dx²-y²) 2. fifth orbital is uniquely shaped – two lobes along the z-axis and a belt centered in the xy plane (dz²)
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f orbitals
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-occur in levels n=4 and greater -highly complex shapes -not involved in bonding in most compounds
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energy level shifting
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-for many electron atoms, electron repulsion’s increase energy, causing orbitals to “overlap” -d orbitals shift 1 level -f orbitals shift 2 levels
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degenerate orbitals
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all orbitals with the same value of n have the same energy (hydrogen only)
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ground state
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lowest energy state
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excited state
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when the atom absorbs energy, electrons may move to higher energy orbitals
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Pauli Exclusion Principle
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in a given atom no two electrons have have the same set of four quantum numbers
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Polyelectronic Atoms
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Internal Atomic energies 1. kinetic energy of moving electrons 2. potential energy of attraction between nucleus and electrons 3. potential energy of repulsion between electrons Electronic Correlation Problem 1. Electron pathways are not known, so electron repulsive forces cannot be calculated exactly 2. we approximate the average repulsion’s of all other electrons
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screening or shielding
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1. electrons are attracted to the nucleus 2. electrons are repulsed by other electrons 3. electrons would be bound more tightly if other electrons weren’t present
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variations in energy within the same quantum level
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1. atoms other than hydrogen have variations in energy for orbitals having the same principle quantum number 2. electrons fill orbitals of the same n value in preferential order -Ens < Enp < End < EnF 3. electron density profiles show that s electrons penetrate to the nucleus more than other orbital types -closer proximity to the nucleus = lower energy
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Aufbau Principle
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as protons are added one by one to the nucleus to build up elements, electrons are similarly added to these hydrogen-like orbitals
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Hund’s Rule
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The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals
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valence electrons
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-electrons in the outermost principle quantum level of an atom -elements in the same group (vertical column) have the same valence electron configuration
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transition metals
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“d” block
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lanthanide and actinde series
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-the sets of 14 elements following lanthanum and actinium -“f” block
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representative elements
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-main-group -groups 1A through 8A -Configurations are consistent
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Metalloids
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-semi-metals -found along the border between metals and nonmetals (stair-case) -exhibit properties of metals and nonmetals
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Ionization energy
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the energy required to remove and electron from an atom -increases for successive electrons -increases across a period WHY? electrons in the same quantum level do no shield as effectively as electrons in inner levels -irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove -decreases with increasing atomic number within a group WHY? electrons farther from the nucleus are easier to remove Multiple (I₂,I₃…) and “jumps”: taking away more than one electron -more energy to take second electron than first // core electron take massive amounts of energy
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Electron Affinity
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the energy change associated with the addition of an electron -increases across a period -decreases down in a period WHY? electrons farther from the nucleus experience less nuclear attraction -some irregularities due to repulsive forces in the relatively small p Orbitals
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Atomic Radius
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1. determination of radius -half of the distance between radii in a covalently bonded diatomic molecule “covalent atomic radii” 2. Periodic Trends -decreases across a period WHY? increased effective nuclear charge due to decreased shielding -increase down a group WHY? addition of principal quantum levels
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alkali metals
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1. easily lose valence electrons (reducing agents) -react with halogens to form salts -react violently with water (Lithium is not the more reaction because the heat of reaction is insufficient to melt lithium and expose all of its surface area 2. Large hydration energy -positive ionic charge makes ions attract to polar water molecules 3. Radius and Ionization energy follow expected trend
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Ion size compared to Atom size
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metals -decreases WHY? less electrons but same number of protons pulling non-metals -increase WHY? proton:electron ratio smaller
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metallic properties
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ability to loose electrons (shiny, condubility, malleability) -across a period –> decreases -down a family –> increases WHY? valence electrons far from nucleus
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non-metallic properties
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ability to gain electrons -across a period –> increases -down a family –> decreases
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Electronegativity
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ability to attract a shared electron in a bond Fluorine – 4.0 Fr +Cs ~~ .7 (non-metal)
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Atomic structure
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Protons, Neutrons, and Electrons.
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Periodicity
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the quality of recurring at intervals
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Electromagnetic
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pertaining to or exhibiting magnetism produced by electric charge in motion
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Radiation
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(medicine) the treatment of disease (especially cancer) by exposure to radiation from a radioactive substance
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wavelengths
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physical energy
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Frequency
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the number of observations in a given statistical category
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Energy
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an imaginative lively style (especially style of writing)
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Speed is constant
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2.99792458 x 10^8 m/s
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Nature
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the essential qualities or characteristics by which something is recognized
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Matter
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that which has mass and occupies space
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Max Planck
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German physicist whose explanation of blackbody radiation in the context of quantized energy emissions initiated quantum theory (1858-1947)
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Quantum Theory
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(physics) a physical theory that certain properties occur only in discrete amounts (quanta)
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Planck’s constant
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h= 6.626 x 10^-34 J S
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photons
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light quanta
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atomic mass
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(chemistry) the mass (in atomic mass units) of an isotope of an element
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Density
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mass/volume
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Volume
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physical objects consisting of a number of pages bound together
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Speed
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distance travelled per unit time
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Frequency
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the number of observations in a given statistical category
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Energy
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(physics) the capacity of a physical system to do work
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Orbital
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of or relating to an orbit
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Shapes
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used to enhance a publication and convey meaning
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Atom
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(physics and chemistry) the smallest component of an element having the chemical properties of the element
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Atomic radius
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one-half the distance between the nuclei of identical atoms that are bonded together
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Hydrogen
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a nonmetallic univalent element that is normally a colorless and odorless highly flammable diatomic gas

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