CHEM101: General Chemistry I

Read this text, which gives some of the history of the development of atomic theory. In the section, "Atoms Become Real", pay close attention to the law of conservation of mass-energy and the law of definite proportions. These laws define how we describe chemical reactions. Also, pay close attention to Dalton's Atomic Theory. John Dalton was the first person to propose a cohesive theory for how atoms make up matter.

As stated above, we define the elements by their number of protons. We define the atomic number, Z, as the number of protons in an atom. For a neutral atom (not a charged ion), the number of electrons must equal the number of protons. However, the number of neutrons can vary within atoms of a given element.

Atoms of the same element with different numbers of neutrons are called isotopes. Most elements have multiple isotopes. For a given isotope, we define the mass number, A, as the atomic number plus the number of protons. We can write this as A = Z + N where A is mass number, Z is atomic number, and N is number of neutrons.

Read this section, which explains how we write atomic symbols with atomic numbers and mass numbers.

When we see the mass of an element on the periodic table, we are seeing the weighted average of the masses of all isotopes of that element. Then read the section "Isotopic Mixtures and Abundances" near the bottom of the page. This section describes how to determine the average atomic mass of an element if we know the isotope masses and their relative abundance. Try the practice problems to do these types of calculations yourself.

Read this text, which defines Avogadro's Number and moles. It also shows examples of conversions between particles and moles, and conversions between moles and mass.

Watch these two videos, which provide step-by-step tutorials on how to convert between the number of atoms and moles using Avogadro's Number as a conversion factor, and how to convert between mass and moles using molar mass as a conversion factor. After you watch one example, you may want to pause the video and try an example on your own. We will be using these types of calculations later in this course.

Read this text, which outlines many of the experiments done in the early part of the twentieth century that led to quantum theory. Pay attention to the equation E = hv, which relates frequency (v) to energy (E) of a wave by Plank's Constant (h).

Read this text. Section 1 describes Young's Double Slit Experiment, which showed that atom-sized particles exhibit wave-like properties by producing wave-like interference patterns. The illustrations in this section show the differences between macroscopic particles and atom-sized particles.

Section 2 provides an overview of the properties of waves, including amplitude, frequency, wavelength, and energy. Pay attention to the examples where energy, frequency, and wavelength are calculated for a given wave. Also, note the relative energies in the electromagnetic spectrum.

Section 3 shows that different atom types emit different frequencies of energy, or spectra, which are characteristic of that element. We can use these atomic spectra to identify elements in a sample.

Finally, section 4 describes how we can calculate the wavelength of a moving atom-sized particle using the deBroglie relation. It also outlines the Heisenberg Uncertainty Principle, which states that it is impossible to know both the position and momentum of a moving atom-sized particle with any precision.

The first model of the atom to use quantum mechanics was the Bohr Model of the hydrogen atom. You have no doubt seen this model, as it is frequently used in logos. The Bohr model, often called the planetary model, consists of the nucleus surrounded by electrons in fixed orbits circling the nucleus, much like planets orbit the sun. This model properly accounted for the hydrogen atomic spectrum and introduced the idea of quantum numbers – numbers that define the energy of an electron. However, as we learn in our next resource, this model is incorrect.

The current view of atomic structure is that electrons exist in a cloud surrounding the nucleus, rather than in fixed orbits. The electrons exist in orbitals, areas of high probability of finding the electron. Different orbitals have different shapes and sizes, which correspond to different energy levels.

Read this text, which reviews the Bohr model and then shows how more advanced quantum mechanics gave rise to the orbital model of the atom.

This video reviews the different orbital shapes and names. The electrons in an atom are arranged in shells that surround the nucleus, with each successive shell being farther from the nucleus. Electron shells consist of one or more subshells, and subshells consist of one or more atomic orbitals. Electrons in the same subshell have the same energy, while electrons in different shells or subshells have different energies.

This video reviews the types of orbitals and shows how we can count electrons and place them in orbitals using the periodic table.

Electron configurations describe where electrons are located around the nucleus of an atom. For example, the electron configuration of lithium, 1s²2s¹, tells us that lithium has two electrons in the 1s subshell and one electron in the 2s subshell.

Our third video shows how to use the Aufbau Principle when filling electrons in orbitals for larger elements. This principle tells us that we fill electrons from the lowest energy to the highest energy.

The Aufbau principle states that electrons fill lower-energy atomic orbitals before filling higher-energy ones (Aufbau is German for 'building-up'). By following this rule, we can predict the electron configurations for atoms or ions. The Aufbau principle is most useful for the first 20 elements: from Sc on, the Aufbau principle does not accurately predict the order of electron filling in atoms.

This text describes building the periodic table, one electron at a time.

Our final video in this sequence introduces the idea of an outer shell, or valence electrons, in an electron configuration. Because valence electrons are in orbitals in the outermost shell, they are the electrons that can be involved in chemical reactions.

Valence electrons are the electrons in the outermost shell, or energy level, of an atom. For example, oxygen has six valence electrons, two in the 2s subshell and four in the 2p subshell. We can write the configuration of oxygen's valence electrons as 2s²2p⁴.

Because the elements in the periodic table are listed in a way that relates to their valence electrons, we can infer many properties of elements based on their relative position on the periodic table. Many of these properties explain the reactivity of the elements.

Read this text, which defines properties including atomic radius (size), ionization energy (energy required to remove an electron from an atom), electron affinity (tendency of an atom to gain electrons), and electronegativity (tendency of an atom to pull electrons toward itself in a bond).