Chapter 1: Structure of the Atom

Important Atom Discoveries

• John Dalton’s Atomic Theory
  • All matter is made up of tiny, indivisible particles called atoms.
  • Atoms of the same element are identical in size, mass, and other properties.
  • Atoms of different elements have different properties.
  • Atoms combine in simple, whole-number ratios to form compounds.
  • Chemical reactions involve the rearrangement of atoms, but the atoms themselves are not created or destroyed.
• John Dalton’s Law of Multiple Proportions
  • John Dalton's law of multiple proportions states that when two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element can be expressed in small whole numbers.
• Discoveries of charges and masses of the neutron, proton, and electron
  • Neutron: Charge = 0, Mass = 1.0087 atomic mass units (amu)
  • Proton: Charge = +1, Mass = 1.0073 amu
  • Electron: Charge = -1, Mass = 0.00055 amu

Atomic Structure

Light and the Atom

• The ground state of an atom is its lowest energy state, where all of its electrons are in their respective lowest energy levels or orbitals. In this state, the atom is considered to be in its most stable configuration.
• When an atom absorbs energy, its electrons can move to higher energy levels, creating an excited state. However, these excited states are temporary, and the electrons will eventually return to their ground state by releasing the excess energy in the form of light or heat.

Wavelength, Frequency, and Energy of Light

• Electromagnetic radiation is defined by wavelength and frequency
  • wavelength is the distance between to repeating points (minima to minima or maxima to maxima)
  • Frequency is the number of waves per second
• (wavelength)(frequency) = speed of light
  • speed of light = 3.0 x 10^8

Bohr Model of the Atom

• The Bohr Model of an atom is a simplified representation of an atom that shows electrons orbiting around the nucleus in specific energy levels or shells. It was proposed by Niels Bohr in 1913 and is still used today as a basic model for understanding atomic structure.

The Size of the Atom

• Momentum is correlated to the size of the electron’s orbit
• The Bohr radius is the distance between the nucleus and the electron in a hydrogen atom in its ground state.

The Wave-Mechanical Model of the Atom

• Wave equations describe the behavior of waves in different mediums.
  • The most common wave equation is the one-dimensional wave equation, which describes waves traveling in a single direction.
  • The wave equation is a partial differential equation, which means it involves derivatives of a function with respect to multiple variables.
  • The wave equation can be used to model a wide range of wave phenomena, including sound waves, electromagnetic waves, and water waves.
  • The wave equation is a linear equation, which means that the principle of superposition applies. This allows us to add together different waves to create more complex wave patterns.
• Heisenberg Uncertainty Principle
  • The Heisenberg Uncertainty Principle states that it is impossible to simultaneously determine the exact position and momentum of a particle. This is due to the fact that the act of measuring one property of a particle will inevitably disturb the other property. In other words, the more precisely we know the position of a particle, the less precisely we can know its momentum, and vice versa.

Structure of the Atom

Principal Energy Levels (Shells)

• Principal energy levels refer to the major energy levels of an atom, designated by the principal quantum number (n). These levels are arranged in increasing order of energy, with the lowest energy level being closest to the nucleus. Each energy level can hold a certain number of electrons, with the maximum number of electrons in a level being given by the formula 2n^2. The first energy level (n=1) can hold a maximum of 2 electrons, the second energy level (n=2) can hold a maximum of 8 electrons, and so on.

Sublevels (Subshells)

• Electrons in an atom occupy subshells, which are designated by letters (s, p, d, f).
• Each subshell can hold a maximum number of electrons, which is determined by the formula 2n^2, where n is the principal quantum number.
• The s subshell can hold a maximum of 2 electrons, the p subshell can hold a maximum of 6 electrons, the d subshell can hold a maximum of 10 electrons, and the f subshell can hold a maximum of 14 electrons.
• The subshells are arranged in order of increasing energy, with the s subshell being the lowest in energy and the f subshell being the highest in energy.

Orbitals

• Electron orbitals are regions of space around the nucleus of an atom where electrons are likely to be found.
• Each orbital can hold a maximum of two electrons with opposite spins.
• The number and arrangement of orbitals in an atom is determined by its electron configuration.
• There are four types of orbitals: s, p, d, and f, each with a different shape and energy level.
• The s orbital is spherical in shape and has the lowest energy level, while the f orbital is more complex and has the highest energy level.
• The p and d orbitals have different shapes and energy levels depending on their orientation in space.
• The arrangement of electrons in orbitals determines the chemical properties of an element and how it interacts with other elements.

Electronic Configurations

Complete Electronic Configurations

• Electron configuration refers to the arrangement of electrons in an atom or molecule. Also known as Aufbau Ordering
• It is represented by a series of numbers and letters that indicate the energy level, orbital shape, and number of electrons in each orbital.
• The first number represents the energy level or shell, while the letter represents the orbital shape.
• The second number represents the number of electrons in that orbital.
• The complete electron configuration of carbon is 1s²2s²2p².

Important Exceptions to Aufbau Ordering

• A complete d subshell is filled but the s subshell is not, in order to give the atom more stability

• The complete electron configuration of copper is:

  1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^10

Abbreviated Electronic Configurations

• Abbreviated electronic configurations are a shorthand method of representing the electron configuration of an atom.
• The noble gas configuration of the previous noble gas is used as a starting point.
• The symbol of the noble gas is written in brackets to represent its full electron configuration.
• The remaining electrons are written after the noble gas symbol, using the appropriate sublevel notation.
• For example, the abbreviated electronic configuration of chlorine (Cl) is [Ne] 3s2 3p5, where [Ne] represents the electron configuration of neon (Ne) and 3s2 3p5 represents the remaining electrons in the chlorine atom.

Valence Electrons

• Valence electrons are the outermost electrons in an atom.
• They are involved in chemical reactions and bonding with other atoms.
• The number of valence electrons determines the chemical properties of an element.
• Elements in the same group of the periodic table have the same number of valence electrons.
• Valence electrons are represented by dots in Lewis dot structures.

Hund’s Rule

• Hund's rule states that electrons will occupy separate orbitals within a subshell before pairing up.
• Each orbital can hold a maximum of two electrons with opposite spins.
• Electrons will fill the lowest energy orbitals first before moving to higher energy orbitals.
• This rule applies to atoms in their ground state and helps to explain the magnetic properties of elements.

Pauli Exclusion Principle

• The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers.
• In simpler terms, this means that electrons in an atom must occupy different energy levels and have different spins.

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