Models of the atom & quantum theory

Niels Bohr (1913) Previous research had concluded that light has a dual nature. It could act as both particles and waves Bohr proposed a model that explained why only certain frequencies of light were emitted from hydrogen.

Bohr’s model of the atom Suggested that the electron in a hydrogen atom moves around the nucleus in a circular path i.e. Like planets around a sun Atoms can have multiple “orbits” that have a definite, fixed amount of energy The lowest allowable energy state of an atom is the ground state This is closest to the nucleus The smaller the orbit, the lower the energy and vise versa Orbits

Excited electrons In the ground state, atoms do not radiate energy When energy is add from an outside source, electrons move to a higher-energy orbit This movement raises the electron to an “excited” state Once excited, the electron can drop from the higher- energy orbit to a lower-energy orbit Results in the atom emitting a photon

Ladder analogy Like the rungs of the ladder on the right, the energy levels in an atom are not equally spaced The higher the energy level occupied by an electron, the less energy it takes to move from the energy level to the next highest energy level

Problem with bohr’s model Bohr’s model had some limits It could not explain the spectra of light emitted for any other element It did not fully account for the chemical behavior of atoms

Electrons as waves French scientist De Broglie pointed out the electron orbits in Bohr’s model were similar to the behavior of waves Electrons, like waves, could be: Diffracted (bent) Bending a wave as it passes by the edge of an object Interfere with each other When waves overlap. Results in a reduction of energy in some areas and an increase of energy in others.

Uncertainty principle Werner Heisenberg – 1927 Showed that it is impossible to take any measurement of an object without disturbing the object! Detection of electrons Detected by their interaction with photons Have about the same mass as an electron Found by “bumping” a photon into an electron Doing so changes both wavelength of photon and the position and velocity of the electron Heisenberg’s Uncertainty Principle: It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

Schrödinger's wave equations Remember the problem with Bohr’s model? Only relevant to hydrogen… Erwin Schrodinger – 1926 Developed an equation that treated electrons in atoms as waves His new model not only could be applied to hydrogen, but all the other elements as well!!! YAY This new model, in which electrons are treated as waves is called: The Quantum Mechanical Model of the Atom Other names: Wave Mechanical Model or Charge Cloud Model

Schrödinger sound familiar?

Location of Electrons Electrons described as waves only have a certain probability of being found within a particular volume of space around the nucleus Schrödinger’s wave function predicts a 3D region around the nucleus that described the electron’s probable location Atomic Orbital: Description of the 3D region around the nucleus (looks like a fuzzy cloud) Clouds show the region of probable electron locations Size and shape of cloud depends on the energy of the electrons that occupy them

Activity Probability of finding an electron around the nucleus activity

Electrons around the nucleus Orbitals can differ based on size and shape There are four types of orbitals S, P, D and F Each orbital can hold two electrons An S orbital is shaped like a sphere A 1s orbital is smaller than a 2s orbital 1s 2s

Orbital Shape

Electrons in the orbitals Orbital type Number of Types Electrons that fit into this “shell” s 1 2(1) = 2 p 3 2(3) = 6 d 5 2(5) = 10 f 7 2(7) = 14 The right column tells us the total electrons that fit into all types of that one orbital

Rules of arrangement Three rules, or principles define how electrons can be arranged in an atom’s orbitals Aufbau Principle Pauli Exclusion Principle Hund’s Rule

Aufbau Principle States: each electron occupies the lowest energy orbital available First step in determining the ground-state electron configuration is to learn the sequence of atomic orbitals from lowest to highest energy Each box in the Aufbau diagram represents an orbital

Electron configuration Based on Aufbau’s Principle, the order of the orbitals should be: 1s – 2s – 2p – 3s – 3p – 4s – 3d – 4p – 5s – 4d – 5p – 6s… All elements follow this rule with few exceptions

Pauli exclusion principle Electrons in orbitals can be represented by arrows in boxes Each electron has an associated spin Arrow pointing up represents electron spinning in one direction Arrow point down represents electron spinning in opposite direction Empty box represents an unoccupied orbital Pauli Exclusion Principle: A maximum of two electrons can occupy a single atomic orbital, but only if the electrons have opposite spins = 0 e- = 1 e- = 2 e-

Hund’s Rule Remember that negatively charged electrons repel each other Hund’s Rule: Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital

Electron configuration The electron configuration on an atom is a shorthand method of writing the location of electrons by sublevel The first number describes the principle energy level The letter represents the sublevel The sublevel is followed by a superscript with the number of electrons in the sublevel If the 2p sublevel contains 2 electrons, it is written 2p2

Writing electron configurations First, determine how many electrons are in the atom Iron has 26 electrons Arrange the energy sublevels to increasing energy 1s 2s 2p 3s 3p 4s 3d… Fill each sublevel with electrons until you have used all the electrons in the atom Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6 The sum of the superscripts equals the atomic number of iron (26)

Example Phosphorus, an element used in matches, has an atomic number of 15. Write the electron configuration of a phosphorus atom. Answer: 1s2 2s2 2p6 3s2 3p3

Practice electron configurations Write the complete electron configurations for: Be F S Ca Things to double check: All superscripts add up to total electrons Orbital order follows the Aufbau Principle

answers Be: 1s2 2s2 F: 1s2 2s2 2p5 S: 1s2 2s2 2p6 3s2 3p3 Ca: 1s2 2s2 2p6 3s2 3p6 4s2