Atomic Structure & Properties

• This module focuses on expanding the knowledge of atomic structure and its relation to chemical properties.
• Emphasis is on electron arrangement and how it influences atomic behavior.

Introduction to Module 2

• The chemical properties of atoms are determined by their structure, especially the arrangement of electrons.
• Prior knowledge of protons, neutrons, and electrons is assumed
• Models like the Bohr Model are insufficient to fully describe observed atomic properties.
• Will further examine atomic structure and relate it to chemical properties.

Lesson 1: The Electromagnetic Spectrum

• Light is generated by the movement of electrons.
• Substances emit light when heated, revealing information about their nature.

Characteristics of Waves

• Light travels in waves, so understanding wave characteristics is important.
• Wavelength (λ) is the distance from a point on one crest to the same point on the next crest, measured in meters (m).
• Frequency (f) is the number of wavelengths passing a point per unit time, measured in Hertz (Hz) or s⁻¹.

Frequency, Wavelength, and Energy

• In 1900, Max Planck discovered the frequency of light is directly related to the energy released by glowing objects when heated.
• Wavelength is inversely related to frequency; as wavelength increases, frequency decreases, and vice versa.
• Wavelength is inversely related to the energy of light; greater wavelength means lower energy.
• Radio waves have a wavelength of about 2 m and a frequency of about 1.5 x 10^8 Hz.
• X-rays have a wavelength of about 1.25 x 10^-10 m and a frequency of about 2.4 x 10^18 Hz.
• Radio waves are less harmful than X-rays due to the lower energy.

The Electromagnetic Spectrum

• Sunlight through a prism creates a spectrum, which is a rainbow of colours.
• Each colour in the rainbow represents light of a different frequency or wavelength which is called a continuous spectrum
• The visible spectrum is a small portion of the electromagnetic spectrum, around 10^-6 m wavelength.
• James Clerk Maxwell determined that light consists of electrical and magnetic waves moving perpendicularly, making it part of the electromagnetic spectrum.

Line Spectra

• Passing electric current through hydrogen gas in a tube makes it glow
• Passing the light of the glowing gas through a slit and prism produces a spectrum with distinct lines.
• This type of spectrum is an emission spectrum (or line spectrum), representing separate wavelengths of light emitted by the gas.
• Unlike continuous spectra, colours in a line spectrum do not blend.
• Each coloured line corresponds to a specific wavelength or frequency of light.
• Spectroscopy techniques determine a substance's emission spectrum.
• The line spectrum of each element is unique, like a fingerprint.
• Burning elements causes them to emit a distinctive color of light.
• Salts containing metals produce characteristic colours when placed in a flame, known as a flame test.

Applications of Line Spectra

• The aurora borealis (northern lights) is a result of excited atmospheric gases releasing light.
• Line spectra are used in manufacturing fireworks.

Lesson 2: The Quantum Mechanical Model

• Lesson focuses on outlining of the quantum mechanical model of the atom.

Neils Bohr Explains Line Spectra

• In 1922, Neils Bohr developed a model that explained hydrogen's line spectra.
• Electrons in a hydrogen atom are arranged in stable orbits around the nucleus depending on their energy
• Orbits are like unequally spaced rungs on a ladder
• Electron may gain energy and jump from its normal ground state to a higher, excited state when radiation is absorbed by an atom.

The Quantum Mechanical Model of the Atom

• Prior knowledge relies on the the Bohr model
• Niels Bohr made significant contributions to understanding atoms, but lacked a complete description of electronic behaviour in the atom
• Werner Heisenberg developed the Heisenberg Uncertainty Principle
• According to Heisenberg, it is impossible to know both the momentum and the position of an electron simultaneously with certainty.
• In 1926, Erwin Schrödinger developed an equation describing the energies and behaviour of small particles like electrons which led to the new field of quantum mechanics.
• Quantum mechanics suggests that energy comes in indivisible bundles called 'quanta'.

Principal Energy Levels

• According to Schrödinger's equation, Bohr's orbits become principal quantum numbers (n), also called principal energy levels
• Within each energy level, there is a set of orbitals or sublevels.
• Schrödinger expanded on Bohr's model of electrons orbiting at specified distances, determining that electrons move more or less randomly around the nucleus
• Bohr’s orbits are boundaries in which electrons are most likely found.
• The lowest energy level is n = 1; the highest is n = 7.

Electron Orbitals

• Electrons fill the space around the nucleus of an atom
• Electron orbitals are three-dimensional regions where there is a 90% probability of finding electrons
• Orbitals are described in terms of size, shape, and orientation in space.
• There are five types of orbitals, named with the letters s, p, d, f, and g
• Regions of probability result in fuzzy electron clouds.
• The s-orbitals are spherical in shape
• Each energy level has one s-orbital
• The s-orbital in the first energy level is given the designation 1s, the s-orbital in the second energy level has the designation 2s, and so on
• The position of the nucleus is assumed to be the point where the three axes intersect.
• The p-orbitals have two regions, or lobes, of high probability
• Electrons spend equal time in both lobes of the p-orbital
• The p-orbitals are only present in the second energy level and higher
• There are three kinds of p-orbitals, depending on their position in space (along the x-axis, along the y-axis, or along the z-axis).
• The d-orbitals are only present in the third energy level and up
• They have more varied shapes than p-orbitals.
• There are five types of d-orbitals.
• The f-orbitals are only present in the fourth energy level and up
• There are seven different types of f-orbitals and their shapes are considerably more complex than the d-orbitals (some containing as many as eight lobes).
• The g-orbitals occur only in energy levels 5 and up
• Only the the s, p and d orbitals up to energy level 4 are covered in this course

Lesson 3: Electron Configuration

• Using the Quantum Mechanical Model, scientists can predict the location of electrons according to principal energy levels and sublevels
• The characteristics of elements are due to the arrangement of their electrons.

Electron Orbitals: A Review

• Orbitals are regions of space around the nucleus of the atom where there is a high probability of finding an electron
• Each orbital is given the letter designation of s, p, d, or f
• Each of these orbitals has a different, characteristic shape.
• There is a different number of each type of orbital possible in each energy level, n
• The total number of orbitals in each energy level can be calculated by squaring the principal number (n²)
• Principal energy level 1 has n² = 1 and orbital type is 1 s-orbital.
• Principal energy level 2 has n² = 4 and orbital types of 1 s-orbital + 3 p-orbitals.
• Principal energy level 3 has n² = 9 and orbital types of 1 s-orbital + 3 p-orbitals + 5 d-orbitals.
• Principal energy level 4 has n² = 16 and orbital types of 1 s-orbital + 3 p-orbitals + 5 d-orbitals + 7 f-orbitals.

The Pauli Exclusion Principle

• Electrons spin on their own axis and this spinning produces a magnetic field
• According to Wolfgang Pauli, two identical electrons cannot occupy the same orbital, however said that two electrons can only occupy the same orbital if they have opposite spins.
• The number of electrons per energy level will be 2 x n² (or 2n²)
• Principal energy level 1 can contain the number of electrons (2n²) = 2.
• Principal energy level 2 can contain the number of electrons (2n²) = 8.
• Principal energy level 3 can contain the number of electrons (2n²) = 18.
• Principal energy level 4 can contain the number of electrons (2n²) = 32.

Electron Configuration

• Electron configuration is in orbitals around a nucleus
• Aufbau Principle: electrons must be added to the lowest energy state before a higher one

Writing Electron Configuration

• Assign a number for each principal energy level, as well as a number to each electron.
• The number in the coefficient position refers to the principal energy level of the electron
• The small-case letter refers to the name of the orbital
• The superscript number refers to the number of electron(s) in that orbital.

Noble Gas Configuration

• Noble gas configuration can be written for any element
• Kernel is also called the highest noble gas structure within the configuration

Exceptions to the Rules

• There are certain cases where the orbitals do not fill in the order you have just learned
• Chromium and copper, and their families undergo what is called electron promotion.

Lesson 4: Valence Electron Configuration

• Relates the electron configuration of an element to its valence electron(s) and its position on the periodic table.
• Orbitals relates to the Periodic Table in a very direct way
• Period number is the same as the principal energy level.
• The alkali metals and alkali earth metals are the two families that make up the "s" block.
• The six families from group 13 (the Boron family) to group 18 (the Noble Gases) make up the "p" block.
• The transition metals from group 3 to group 12 make up the "d" block.
• The actinides and the lanthanides make up the "f" block.

Valence Electrons

• The electronic configuration of an atom is given by listing its subshells with the number of electrons in each subshell.
• The electrons on the highest numbered subshells are the valence electrons
• It is the valence electrons that are lost or gained in chemical reactions, and these electrons will always be from the "s" and "p" orbitals.
• The electron configuration of an element is related both to its valence electrons and to its position on the periodic table
• The alkali metals have 1 valence electron (s¹)
• The alkali earth metals have 2 valence electrons (s²)
• The boron family has 3 valence electrons (s²p¹)
• The carbon family has 4 valence electrons (s²p²)
• The nitrogen family has 5 valence electrons (s²p³)
• The chalcogen family has 6 valence electrons (s²p4)
• The halogen family has 7 valence electrons (s²p5)
• The noble gases have 8 valence electrons (s²pº)

Electron Configurations of Ions

• An ion is the result of either losing the valence electrons (leaving a positive ion) or gaining enough electrons to fill that valence subshell.

Lesson 5: Periodic Trends

• Identifies and accounts for periodic trends among the properties of elements, and relates the properties to electron configuration.
• This includes atomic radii, ionic radii, ionization energy, and electronegativity
• In 1870, Dmitri Mendeleev arranged the 65 elements known at that time into a periodic table which created the periodic law (when the elements are arranged according to their mass, they have a regular periodic repeating of similar properties)

Force on an Electron

• The positively charged nucleus of the atom applies a force on each negatively charged electron, holding the electrons around the atom.
• The force on an electron is dependent upon three factors: nuclear charge, distance of the electron from the nucleus and the shielding effect.

Atomic Radii

• Atomic radius is the distance between two nuclei when two like atoms are bonded together
• Scientists measure atoms by the process of x-ray crystallography

Trends in Atomic Radii

• As you move from left to right across the periodic table, the atomic radii decrease.
• As you move down a group, the atomic radii increase.

Ionic Radii

• For negative ions, one or more electrons are added to the atom. The more electrons are added to an atom, the lower the force of the nucleus on each electron
• For positive ions, one or more electrons are removed from the atom. As a result, there is more than one proton for every electron in the ion
• As an ion becomes more negative, the ionic radius increases
• As an ion becomes more positive, the ionic radius decreases

Electronegativity

• Electronegativity is the ability of an atom in an element to attract electrons when the atom is in a compound.
• As you move from left to right across the periodic table, the electronegativity values increase.
• As you move down a group, the electronegativity values decrease.

Ionization Energy

• Ionization energy (IE) is defined as the amount of energy needed to remove an electron from a gaseous atom.
• There are two major trends related to ionization energy and the periodic table
• As you move down a group, the ionization energy decreases
• As you move across a period from left to right, the ionization energy increases.