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.

• 

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