# Chap 7

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2017-12-11 20:30

## Section

Question | Answer |
---|---|

define wave | a vibrating disturbance by which energy is transmitted |

characteristics of waves | wavelength(l), amplitude, frequency(n), speed (u) |

define Wavelength (l) | the distance between identical points on successive waves |

Frequency (n) | the number of waves that pass through a particular point in 1 second |

frequency and cycle equation | 1 Hz = 1 cycle/s |

what does the speed of a wave (u) depend on | the type of a wave and the nature of the medium through which the wave is traveling (e.g. air, water, or vacuum) |

equation for the speed of a wave | u = l x n |

units for u = l x n | (u is m/s, l is m, n is Hz) |

who and when proposed that visible light consists of electromagnetic waves | Maxwell 1873 |

components of electromagnetic wave | an electric field component and a magnetic field component |

compare and contrast the 2 components of electromagnetic waves | The two components have the same wavelength, amplitude, frequency, and hence the same speed, but they travel in mutually perpendicular planes |

what is the emission and transmission of energy in the form of electromagnetic waves | electromagnetic radiation |

what is the speed of light (c) in vacuum | 3.00 x 108 m/s |

equation for all electromagnetic radiation | l x n=c (wavelength x frequency = speed of light |

The quantum theory developed by Planck successfully explains | the emission of radiation by heated solids |

Planck’s Quantum Theory | Energy (light) is emitted or absorbed in discrete units (quantum) |

Planck’s constant (h) | = 6.63 x 10-34 J•s |

E=hxn | E=hxn E = h x c /l (n is the frequency of radiation, l is the wavelength of radiation ) |

Energy is emitted in whole-number multiples of | hn |

Photoelectric Effect is | a phenomenon in which electrons are ejected from the surface of metals exposed to light of at least a certain minimum of frequency |

Light has both | wave nature and particle nature |

Photon is | a “particle” of light |

hn = KE + W | KE = hn - W (W is binding energy) (KE is the kinetic energy, W is the work function, a measure of how strongly the electrons are held in the metal) |

Bohr’s Theory of | the Hydrogen Atom |

In bohr's model, light is emitted as e- moves from | one energy level to a lower energy level |

In bohr's model, e- can only have specific | (quantized) energy values |

bohr model equation | Esubn = - RsubH * 1/(n^2) n is (principal quantum number) = 1,2,3,…, RH (Rydberg constant) = 2.18 x 10-18J |

which constant is used in bohr's model | RH (Rydberg constant) = 2.18 x 10^-18J |

Ground state or level | lowest energy state of a system (n = 1) |

Excited state or level | higher in energy than the ground state (n = 2, 3, 4…) |

In the Bohr model, an electron emits a photon when | it drops from a higher-energy state to a lower energy state |

equation for Photon, hn, absorbed or emitted | DE = hn = RsubH * (1/n^2subi) - (1/n^2subf) n (principal quantum number) = 1,2,3,… |

ni > nf: DE is | positive (Absorption) , ni < nf: DE is negative (Emission) |

Calculate the wavelength (in nm) of a photon emitted by a hydrogen atom when its electron drops from the n = 5 state to the n = 3 state | Ephoton = DE = RH (1/n^2subi) - (1/n^2subf) |

LASER | Light Amplification by Stimulated Emission of Radiation |

De Broglie (1924) reasoned that | particles such as e- have wave properties |

De Broglie (1924) reasoned that an e- behaves like a standing wave in the atom, the length of the wave nl equals | the circumference of the orbit 2pr (2pr =nl) |

De Broglie extended Einstein’s wave-particle description of light to | all matter in motion (l = h/(mu) ) u = velocity of e- (in m/s) m = mass of e- (in kg) h in J•s = 6.63 x 10^-34 |

The Heisenberg uncertainty principle states that | it is impossible to know simultaneously both the momentum (p) and the position (x) of a particle with certainty |

The Heisenberg uncertainty principle equation | DxDp >_ h/(4p) Dx= uncertainty in the position Dp = uncertainty in the momentum and p = m x v |

Schrodinger’s equation can only be solved exactly for the _________. We must approximate its solution for _________ | hydrogen atom, multi-electron systems |

In 1926, Schrödinger wrote an equation that described | both the particle (m) and wave (Y) nature of the e- |

Wave function (Y) describes | energy of e- with a given Y and probability of finding e- in a volume of space, Y2 |

an atomic orbital is | a function () that defines the distribution of electron density (2) in space |

Electron density gives the probability that | an electron will be found in a particular region of an atom |

Quantum numbers describe | the distribution of electrons in atoms; they are derived from the Schrödinger equation Y = fn (n, l, ml, ms) |

letters in Y = fn (n, l, ml, ms) | n = principal quantum number, l = angular momentum quantum number, ml =magnetic quantum number, ms = electron spin quantum number |

Principal Quantum number (n) | Main energy level (shell) and distance of e- from the nucleus |

Angular Momentum Quantum number (l) | Shape of the “volume” of space that the e- occupies |

For a given value of n, l = | 0, 1, 2, 3, … n-1 |

Magnetic Quantum number (ml) | Orientation of the orbital in space |

For a given value of l ml = -l, …., 0, …. +l as in | If l = 1 (p orbital), ml= -1, 0, or1 If l = 2 (d orbital), ml= -2, -1, 0, 1, or2 |

Spin Quantum number (ms) | Direction of the electron’s spin |

ms = +½ or -½ | clockwise or not |

how many orbitals does subshells s p d have | one 3 5 |

The total number of orbitals for a given value of n is | n2 |

for s orbital, Its size increases as | the principal quantum number increases |

position and shape of s orbital | spherical, centered in nucleus (only one) |

only way the three p orbitals differ from each other | orientation |

the five d orbitals are identical in | energy |

what is l in s p d f orbitals | 0 1 2 3 |

In a single-electron atom, the energy of the electron is determined solely by | its principal quantum number, n |

In many-electron atoms, the energy of the electron is determined by its principal quantum number, n, AND | the angular momentum quantum number, l |

electrons with the same value of n | shells |

electrons with the same value of n and l | subshells |

electrons with the same value of n and l and ml | orbital |

How many 2p orbitals are there in an atom | If l = 1, then ml = -1, 0, or +1 sp 3 orbitals (3 numbers of ml) |

Pauli exclusion principle | no two electrons in an atom can have the same four quantum numbers |

Paramagnetic atoms are | atoms with one or more unpaired electron spins; they are attracted by a magnet |

Diamagnetic atoms | are atoms in which all electrons are paired; they are slightly repelled by a magnet |

hund's rule | every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin |

Each subshell of quantum number l has (2l+1) orbitals | if l is 1 then 2+1=3 orbitals |

aufbau principle states that | in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels (e.g., 1s before 2s). In this way, the electrons of an atom or ion form the most stable electron configuration possible |

either have incompletely filled d subshells or readily give rise to cations that have incompletely filled d subshells | transition metals |

either have incompletely filled f subshells or readily give rise to cations that have incompletely filled f subshells | Lanthanides or rare earth series |

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