Chemistry Unit 5 Test Study Guuide

I.  Wave Nature of Light


  • electrons spread through space as an energy wave

  • light travels through space in the form of radiant energy; it travels through waves or as fast-moving particles

  • visible light: a type of electromagnetic radiation

  • electromagnetic radiation: a form of energy that exhibits wave-like behavior as it travels through space

  • amplitude: how much energy a wave carries; the more energy, the higher the amplitude; the vertical distance from the origin to the crest 

  • wavelength: the distance between a point on a wave and the nearest point just like it (measured in m, nm; 1 m = 109 nm)

  • frequency: the number of wavelengths that pass a fixed point each second; as frequency increases, energy increases and wavelength decreases (measured in Hz or s-1)

  • crests: the highest points of a wave

  • troughs: the lowest points of a wave

  • hertz: 1 wavelength passing by a fixed point per second; measures frequency; equal to 1 s-1

  • electromagnetic spectrum (from lowest to highest frequency) - 

  1. radio waves

  2. cell phone waves

  3. microwaves

  4. radar waves

  5. infrared waves 

  6. visible light waves 

  7. ultraviolet light

  8. X-rays

  9. gamma rays


  • continuous spectra: contains all colors (R O Y G B I V)

  • line spectra: sharp loans observed in a spectrum of light emitted or absorbed by an element 

  • flame tests: many elements give off characteristic light which can be used to help identify them (atoms are excited by heat or electricity)

  • c = λν ; speed of light = wavelength x frequency

  • speed of light: 3.00 x 108 m/s

  • the speed for all electromagnetic waves is the speed of light in a vacuum, 3.00 x 108 m/s

  • E = hν ; energy = Planck’s constant x frequency

  • Planck’s constant: 6.626 x 10-34 J x s

  • if you need to find E, energy, but the frequency is not given, use the formula - E = h x (c/λ) 


II.  Particle Nature of Light


  • quantum: the minimum amount of energy that can be gained or lost by an atom; matter gains or loses energy only in small, specific amounts, quanta


III.  Quantum Mechanical Model of the Atom


  • ground state: the state at which an atom has its lowest energy level

  • excited state: state of an atom caused by the excitation of the electron by absorbing energy

  • when atoms absorb energy -> electrons move onto higher atomic levels; these electrons then lose energy by emitting light when they return to lower energy levels

  • principal energy levels: the regions of space in which electrons can move about the nucleus

  • Heisenberg uncertainty principle: states that is is fundamentally impossible to know precisely both the velocity and position of a particle at the same time; you can determine where the electron is (position/orientation), but not where it is going (speed/momentum), and vise versa

  • de Broglie equation: predicts that all moving particles have wave characteristics; λ = h/mv (wavelength = Planck’s constant/ mass of the particle x velocity)

  • atomic orbital: a region of space in which there is a high probability of finding an electron; each orbital contains a maximum of 2 electrons

  • quantum numbers: specify the address of each electron in an atom and describe the properties and the electrons in orbitals - 

  1. n, principal quantum number: electron’s energy depends principally on this (tells which level the electron is in, ie n=1, n=2, . . . ); are positive numbers beginning with 1 (closest to the nucleus and moving out); as n increases, so does the energy of the electron

  2. l, angular quantum number: tells the shape of the orbital; for orbitals of the same principal energy level, l distinguishes different shapes 

  3. m1, magnetic quantum number: indicates the orientation of an orbital around the nucleus; s orbital is sphere shaped, p orbital is shaped dumbbell shaped, d orbital is like 2 dumbbells, and f orbital is flower shaped)

  4. ms spin quantum number: has 2 possible values (clockwise ½ or counterclockwise -½), which indicate the 2 fundamental spin states of an electron in an orbital; identifies the 2 possible spin orientations of an electron in an orbital (clockwise ½ or counterclockwise -½)


  • 2n2: formula for the number of electrons that can fit in a shell/principal energy level (ie n=3, 2 x 32 = 18; 18 electrons can fit and be contained in n=3)

  • quantum theory: describes mathematically the wave properties of electrons and other very small particles

  • sublevels/energy sublevels: are contained within the principal energy levels (rooms inside the main floors, the principal energy levels); include s, p, d, and f -

  1. s: contains 1 orbital, 2 electrons; spherical shaped, therefore 1 possible orientation

  2. p: contains 3 orbitals, 6 electrons; dumbbell shaped

  3. d: contains 5 orbitals, 10 electrons; 2 dumbbell shaped

  4. f: contains 7 orbitals, 14 electrons; flower shaped


IV.  Electron Configuration and Orbital Notation


  • electron configuration: tells us in which orbital the electrons for an element are located, following 3 rules - 

  1. Aufbau principle: states that an electron occupies the lowest energy orbital in order of increasing energy

  2. Pauli exclusion principle: states that a maximum of 2 electrons can occupy a single orbital, but only if the electrons have opposite spins; no 2 electrons can fill 1 orbital with the same spin; also, no 2 electrons in the same atom can have the same set of 4 quantum numbers

  3. Hund’s rule: states that orbitals of equal energy are each occupied by 1 electron before any orbital is occupied by a second electron, and all electrons in sumpy occupied orbitals must have the same spin