Light, Energy, and Electrons Chapter Test

  • Excited/Ground State

    • Ground State
    • Lowest energy state
    • This means that e- are found in shells closer to the nucleus
    • n =1
    • Excited State
    • Higher potential energy of an atom
    • n = 2 or higher
    • A form of heat, light, electrical, or mechanical energy is needed to go from the ground to an excited state
    • As electrons increase in energy, they move away from the nucleus and into outer shells

  • Absorption/Emission

    • Absorption (take in)
    • Energy moves electrons from a ground state to a higher energy state
    • Heat, light, electrical, chemical mechanical energy
    • Emission (give off)
    • Lets electrons fall back down to a lower energy state
    • Usually light
    • Energy must be absorbed for an electron to move to a higher state (one with a higher n value)
    • Energy is emitted when the electron moves to an orbit of lower energy (one with a lower n value)
    • The overall change in energy associated with "orbit jumping" is the difference in energy levels between the ending (final) and initial orbits

  • Wavelength/Frequency/Energy (ROY G BIV) (Both equations)

    • The wavelength (λ) of light is defined as the distance between the crests or troughs of a wave motion.
    • Wavelengths found in the electromagnetic spectrum (range of light) can be measured in units as large as 103 meters (radio waves) to 10-11 meters (gamma waves).
    • For the wavelengths of visible light (the light we see in color) the most common units used are nanometers (10-9 meters) and Angstroms (10-10 meters).
    • Frequency (ν) is the number of occurrences of a repeating event per unit time.
    • In the case of light, frequency refers to the number of times a wavelength is repeated per second. The unit used most often to describe frequency is Hz which means "per second" or /s.
    • The relationship between wavelength and frequency is related through the speed of light.
    • c = λν
      • c = 3.00 x 10^8 m/s
      • c is the speed of light
      • v is frequency
      • λ is wavelength
    • E=hv
      • h = 6.63x10^-34 J.s
      • E stands for energy (in Joules)
      • v stands for frequency [in reciprocal seconds – written s^-1 or Hertz (Hz)- 1Hz = 1 s^-1)
    • h is Planck’s constant.
    • If the frequency is known, it can easily be converted to wavelength using the speed of light and vice versa.
    • The wavelengths and frequencies of the light emitted by an atom (its emission spectrum) is determined by its electronic structure.
    • As each electron moves from a higher energy level (orbit) to a lower one, a different color is emitted.
    • Each shade of color has a unique wavelength based on the unique distance and energy.
    • As a wavelength increases in size, its frequency and energy (E) decrease.
    • As the frequency increases, the wavelength gets shorter.
    • As the frequency decreases, the wavelength gets longer.
    • Electromagnetic Spectrum

  • Quantum Numbers (names and their meaning only)

    • Principle Quantum Number (n)
    • Indicates the main energy level (shell) occupied by the e- (distance from the nucleus)
    • Shell number (1st shell is closest to nucleus, 2nd is further, and so on)
    • Come from the Bohr Model
    • Values of n can only e positive integers (1, 2, 3, etc.)
    • As n increases, the orbital becomes larger; the electron has a higher energy and is farther away from the nucleus
    • Angular Momentum Quantum Number (l)
    • Indicates the general type of shapes of the orbitals
    • Nickname is subshell of n
    • Designated s, p, d, f
    • Values of l are zero and all positive integers less than equal to n-1
    • Magnetic Quantum Number (ml)
    • Indicates which exact orbital the electron is in
      • Describes the orientation of the orbital
    • Because an s orbital is spherical, it only has one orientation (ml = 0)
    • p orbitals can have three different orientations, one along the x-axis, one along the y-axis, and one along the z-axis
    • Spin Quantum Number (ms)
    • Indicates the two spin states of an e- in an orbital
    • Only 2 e- fit in each orbital, and they spin in opposite directions (up and down)
    • Possible m, values are -1/2, + 1/2
    • Spin is represented by dashes inside circles
      • Orbital notation

  • Shells

    • Distance from the nucleus (principle quantum number)
    • Represent ranges in energy

  • Subshells

    • Represent shapes (s, p, d, f)
    • One or more orbitals with the same set of n and l values
    • Each shell is divided into the number of subshells equal to the principal quantum number, n, for that shell.
    • The first shell consists of only the 1s subshell; the second shell consists of two subshells, 2s and 2p; the third of three subshell, 3s, 3p and 3d, and so forth.
    • Each subshell is divided into orbitals. Each s subshell consists of one orbital; each p subshell of three orbitals, each d subshell of five, and each f subshell of seven orbitals.
    • Angular momentum quantum number

  • Number of subshells in a shell

    • The number of subshells in a shell is equal to the shell number
    • 1st shell - 1 subshell
    • 2nd shell - 2 subshells
    • 3rd shell - 3 subshells

  • Electron Filling Order: 1s 2s 2p…

    • Electron filling tree
    • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
    • Also known as the Aufbau principle

  • Orbitals

    • Three dimensional space that electrons most probably occupy
    • Defined by n, l, and ml
    • The math equation treats electrons like waves
    • You can solve the equation to get the shape in space in which electrons are
    • Shapes look like “clouds” of probability

  • Number of orbitals per subshell: S P D F/ Number of electrons per orbital and per subshell

    • S subshell
    • Spherical shaped
    • 1 orbital, 2 e-
    • P subshell
    • Peanut shaped
    • 3 orbitals, 6 e-
    • D subshell
    • Double peanut shaped
    • 5 orbitals, 10 e-
    • F subshell
    • Flower shaped
    • 7 orbitals, 14 e-
    • Each subshell’s name comes from the old spectroscopic description of the lines corresponding to these orbitals
    • 1st subshell in a shell = s subshell → sharp
    • 2nd subshell in a shell = p subshell → principal
    • 3rd subshell in a shell = d subshell → diffuse
    • 4th subshell in a shell = f subshell → fundamental

  • Aufbau Principle

    • “Building up”
    • An electron occupies that lowest energy possible
    • The levels follow a pattern of increasing energy
    • Fill starting at nucleus (Bohr Models)
    • P subshell → 3 orbitals
    • Fill left to right

  • Pauli Exclusion Principle

    • No 2 electrons have the same spin if they are in the same orbital

  • Hund’s Rule

    • Electrons do not pair up until there are no more empty orbitals in that subshell

  • Orbital Notation

    • Representation of electron configuration in which orbital is represented by a circle and dashes
    • Each dash represents the number of electrons in each subshell

  • Electron Configuration

    • The correct order electrons are filled in
    • The most stable, or ground, electron configuration of an atom is that in which the electrons are in the lowest possible energy level
    • All subshells contain a certain number of orbitals
    • May be occupied by a single e- or by 2e- having opposite spins
    • Like cups
    • Shells don’t always get filled from 1 to 2 to 3 etc. because some subshells overlap

  • Valence/Core electrons

    • Valence = outermost
    • Valence electrons are electrons in the outer shells
    • Core electrons are electrons in the inner shells
    • Count the total electrons in the highest shell number
    • Do not count electrons in d subshells
    • Do count s and p

  • The Periodic Table and ordering of electrons

    • Rows (periods)
    • All of the elements in the row have the same number of orbitals
    • Columns (groups)
    • All of the elements in the column have the same number of (valence) electrons
    • Share similar chemical and physical properties because they possess the same # of valence electrons