Atomic Electronic Structure

• The Bohr theory successfully introduced energy levels but lacked details of atomic structure.
• Atoms interact through their electrons; therefore understanding electron behavior is crucial.

Electronic Structure of Atoms

• An atom's electronic structure is its electron arrangement.
• Electronic structure depends on the number of electrons, their distribution, and their energies.
• Quantum theory led to the knowledge of electronic structure.

Light and Electronic Structure

• Understanding atomic electronic structure requires analyzing emitted or absorbed light.
• To understand models of electronic structure, understanding light is needed.

Development of Quantum Theory

• Quantum theory was developed in two stages.
• The electron was initially considered a particle to explain atomic spectra and electronic configurations.
• The electron was then treated as a wave which assisted in stereochemistry and calculating light molecule properties.

Electromagnetic Radiation and Atomic Interactions

• Electromagnetic radiation and atoms can interact by:
  • The emission of light from hot objects in blackbody radiation.
  • The emission of electrons from metal surfaces due to light, known as the photoelectric effect.

Quantized Energy and Photons

• In 1900, Max Planck (1858–1947) theorized that atoms release or absorb energy only in discrete amounts, called quanta.
• Planck was awarded the Nobel Prize in Physics in 1918 for work on the quantum theory.

Photoelectric Effect

• Einstein's theory of light quantization explained the photoelectric effect.
• Alkali metals such as Cesium (Cs), Lithium (Li), Sodium (Na), Potassium (K), and Rubidium (Rb) release electrons when exposed to light.
• Energy from photons striking the metal transfers to electrons.
• Electrons must overcome attractive forces to be released from the metal.
• Threshold frequency is required.
• Each metal has a particular minimum threshold frequency for the photoelectric effect to occur:
• Electrons will not escape if photons have less energy than this threshold, regardless of light intensity.
• Light possesses both wave and particle-like qualities.
• This dual nature is a characteristic of matter.
• Automatic door openers use the photoelectric effect.

Spectrum

• Rainbows are examples of continuous spectrums, produced by sunlight dispersal via raindrops or mist.
• Monochromatic radiation consists of a single wavelength, such as radiant energy from a laser.
• Most radiation sources produce radiation containing many different wavelengths.
• A spectrum is produced when radiation from sources separates into its different wavelengths.
• A prism can disperse white light from a light bulb into a spectrum of colors.

Bohr's Model of the Hydrogen Atom

• Rutherford discovered the nature of the atomic nucleus to have electrons orbiting the nucleus.
• In explaining the line spectrum of hydrogen, Bohr started with Rutherford's atomic model.
• Electrons were assumed to move in circular orbits around the nucleus.
• According to classical physics, electrically charged particles moving in a circular path should continuously lose energy through electromagnetic radiation.
• Bohr's model includes the following assumptions:
  • The prevailing laws of physics alone cannot adequately describe atoms.
  • Energies are quantized (as Planck suggested).
  • Only orbits of certain radii are permitted, corresponding to definite energies.
  • Electrons in allowed energy states do not radiate energy, preventing them from spiraling into the nucleus.

Limitations of the Bohr Model

• The Bohr model cannot be applied to atoms with multiple electrons such as lithium and helium.
• Only atoms with one electron were applicable for this model.
• The Bohr model explained only spherical orbits, but there was no explanation for elliptical orbits.
• The model could not explain Heisenberg's uncertainty principle.

Wave Behavior of Matter

• The concept that radiant energy had a dual nature became commonly used in the years following Bohr's hydrogen atom model.
• Radiation was determined to have wavelike or particle-like behavior.
• Louis de Broglie (1892-1987) suggested wave-particle quality should apply to matter.
• Electrons can be seen as having properties of a wave.
• Every particle exhibits wavelike characteristics.
• A wavelength is related to its mass and speed (v).
• The de Broglie relation which is: λ = h/mv
• The quantity mv for any object equals momentum.
• Heavy particles traveling at low speeds have small wavelengths, while small particles traveling at low speed have a large wavelength.
• De Broglie equation can calculate the wavelength of a moving mass.

The Uncertainty Principle

• Classical physics states that when considering a ball rolling down a ramp, one can calculate its exact position, direction, and speed at any time.
• Discovery of wave properties of matter raised questions about classical physics.
• Waves extend in space, and their location is not precisely defined.
• It is impossible to know both the exact momentum of the electron and its exact location in space at the same time.
• Thus, it is inappropriate to imagine electrons as moving in defined circular orbits about the nucleus.
• De Broglie's hypothesis and Heisenberg's principle led to the formation of the atomic structure theory.
• Precise definition of the instantaneous location and the momentum of the electrons is abandoned with the wavelike description of matter.
• Instead of defining exact position and momentum, the probability of finding the electron in a given volume of space is used.

Quantum Mechanics and Atomic Orbitals

• Erwin Schrödinger (1887-1961) proposed Schrödinger's wave equation in 1926.
• Schrödinger’s equation incorporates both the wavelike and the particle-like behavior of the electron.
• Schrödinger’s equation can be solved exactly only for a species containing a hydrogen-like system.
• Schrödinger's work created a quantum mechanics or wave mechanics approach to subatomic particles.
• Schrödinger equation describes electrons in atoms.
• Electrons are treated as having wavelike properties.
• In contrast to the planner orbits of the Bohr atom, electrons move in the three-dimensional space surrounding the nucleus.
• Electrons are quantized; atoms and molecules can only exist in certain energy states.
• Atoms or molecules change energies when emitting or absorbing radiation.
• Quantum numbers describe the allowed energy states of atoms and molecules.
• Neither the exact position nor the exact momentum of an electron can be determined simultaneously.
• The spatial region around the nucleus where an electron is most likely located can be predicted for each electron in an atom.

Atomic Orbitals

• An atomic orbital will most probably locate a specific energy electron
• Psi squared (ψ²) represents the probability density or electron density.
• The “location” of an electron requires four quantum numbers, which describe energy levels and the shapes of electron distributions.
• Principal quantum number (n), has whole number values, and describes the main energy level that an electron occupies:
  • As n increases, electron is further from the nucleus.
  • Roughly equivalent to the n of the energy levels in the Bohr atom.
• Angular momentum quantum number, l, defines the shape of the orbital:
  • Sublevels or sub-shells are possible Within a shell.
  • l, may take integral values from 0 up to and including (n-1) for each value of n.
  • The value of "l" is designated by the letters s, p, d, and f.
• The magnetic quantum number: m(l) -Designates the specific orbital within a subshell. -Orbital within a given subshell differ in their orientations in space, but have identical energies. -Integers ranging from -l to +l.
• Depend on the value of l.
  • When l = 1, i.e. p subshell, there are 3 permissible values of ml.
  • ml = (-l), ......0.........(+l).
  • -1, 0, +1. i.e. there are three atomic orbital associated with a p sub shell.
  • These orbitals are px, py and pz.
• The spin quantum number: ms
  • Refers to the orientation of the magnetic field and spin of an electron.
  • For every set of n, l, and ml values, ms can take the value of +1/2 (spin up) or -1/2 (spin down)

Quantum Number Examples

• Principal quantum number(n) is 2: -Allowed values of l = 0 and 1, i.e. s and p subshells present.
• ml has values of -1;0;+1 -The corresponding value of l = 1;
• l has values 0,1, 2 and 3 -The corresponding value of n is >3

Representing Orbitals

• Possible types of atomic orbitals for n=4 are: -s, p, d, f
• n = 4 and l = 2 describes the d atomic orbital.
• Three quantum numbers that describe a 2s atomic orbital:
  • n = 2, l = 0, and ml = 0
• The principal quantum number distinguishes the 2s and 4s atomic orbitals

Orbitals of Many Electrons

• Atom's electron configuration distributes electrons among the subshells. The diagrammatical representation of the filling of electrons is as follows

Filling Orbitals

• Filling variations may occur after 3p.
• Orbitals are filled in order of energy, lowest energy orbitals are filled first, only then entering higher energy orbitals. -Order as follows; 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p< 6s < 5d 4f < 6p < 7s < 6d 5f
• Hund’s rule states that electrons may not be spin-paired in degenerate orbitals until each orbital of the set contains one electron.
• Degenerate orbitals attain lowest energy when electrons occupy separate orbitals with unpaired spins.
• Electrons occupying orbitals singly in a degenerate set have parallel spins, having identical ms values.
• Pauli exclusion principle says that no two electrons in the same atom can have the same set of n, l, ml and ms quantum numbers.

Applying the Rules

• Construct an orbital diagram for the electron configuration of oxygen, with atomic number 8.
  • How many unpaired electrons does an oxygen atom have?
• Write the electron configuration for phosphorus, element 15.
  • How many unpaired electrons does a phosphorus atom possess?

Condensed Electron Configuration

• The nearest element to the noble gas element having a lower atomic number would be listed by its chemical symbol in brackets.

Anomalous Electron Configurations

• Chromium and Copper exhibit anomalous behavior due to the closeness of the 3d and 4s orbital energies.
• Copper, [Ar] 3d¹⁰4s¹, forms two ions. - In the Cu⁺ ion the electronic structure is [Ar] 3d¹⁰. - The more common Cu²⁺ ion has the structure [Ar] 3d⁹.
• A few similar cases among the heavier transition metals:
  • Those with partially filled 4d or 5d orbitals and among the f-block metals.

Atoms and Counting Atoms

• The structure of atoms is as follows;

Atoms, lons and Practice

• Protons have a positive (+) charge, and electrons have a negative (-) charge.
• In a neutral atom, the number of protons equals the number of electrons, resulting in a net charge of zero (0).
  • For instance, helium (atomic number 2) in a stable state consists of 2 protons and 2 electrons.
• In neutral atom, the atomic number = no. of protons = no. of electrons
• Gain or loss of electrons forms an atom:
  • When an atom GAINS electrons it becomes NEGATIVE
  • When an atom LOSES electrons, it becomes POSITIVE

Aluminum Practice

• Fill in the missing details for Aluminum (Al), without a periodic table:
  • Protons = 13
  • Electrons =?
  • Neutrons = 14
  • Atomic Number=?
  • Atomic Mass =?

Nuclear Symbol Notation Practice

• Using Nuclear Symbol notation (no periodic table):
  • Complete the details for: -¹⁰⁸₄₇Ag
    • Protons =
    • Electrons =
    • Neutrons =
    • Atomic Number =
    • Atomic Mass =

Isotopes Practice

• Use a periodic table to complete the details for Uranium-235:
  • Protons = -Electrons = -Neutrons = -Atomic Number = -Atomic Mass =

Practice with lons

• Use the periodic table to add data for ³⁹₁₉K¹⁺
  • Charge =
  • Protons =
  • Electrons =
  • Neutrons =
  • Atomic Number =
  • Atomic Mass =

Periodic Tables.

• The periodic table arranges the chemical elements in a tabular form and is organized by their atomic number.
• Periodic Tables uses properties of the element in its electron configurations, and recurring chemical properties.
• The table typically has 18 columns, and 7 rows, with a double row of elements below.
• The rows and columns are called periods and groups.
• Other groups have names, such as halogens or noble gases, alkali, etc
• Periodic tables incorporate recurring trends.
• Use can derive relationships between the properties of the elements.
• Can predict properties of elements not yet discovered or synthesized.

Mendeleev's Periodic Table

• Dmitri Mendeleev is credited for creating the first most widely recognized periodic table, published in 1869.
• Mendeleev developed his periodic table to illustrate periodic trends using known elements.
• Mendeleev observed similar chemical and physical properties to neighboring elements.
• Mendeleev organized elements by increasing mass.
• Mendeleev predicted also predicted properties of then-unknown elements.

Modern Periodic Tables

• The table can be arranged into four rectangular blocks; s-block (left), p-block (right), d-block (middle), and f-block (below).
• Representative elements or main group elements are for which all inner subshells are fully occupied and the outer s and p-subshells are filling.
• s- block elements are on the left and p- block elements are on the right.
• Groups IA, IIA, VIIA and VIIIA are known as alkali metals, the alkaline-earth metals, the halogens and noble gas
• Transition elements are elements in which electrons enter d or f subshells as atomic number increases
  • The elements of the d-block are also known as the d- transition elements
• Inner transition elements have elements of the f block and are split into two main groups: First f-transition series (lanthanides)- 58Ce through 71Lu & the Second f-transition series (actinides)- 90Th through 103Lu.
• Characteristic valence electron configuration of the group 7A (halogens) elements?
• Using the periodic table, locate the elements and define with the following configurations:
• Representative
• Transitional
• Metal
• Non Metal

Electronic Configuration Questions

• (a)[Xe] 5d¹⁰ 6s² 6p² (b)[Kr] 4d⁷ 5s¹
• (c) 1s² 2s² 2p⁶ 3s²
• (e) tellurium -(d) [Kr] 4d¹⁰ 5s² 5p⁶