General Chemistry
Electron configurations describe the arrangement of an atom's electrons within its shells, subshells, and orbitals. Electrons fill orbitals following the Aufbau principle, which states that electrons fill orbitals from lowest to highest energy. Shells can be divided into s, p, d, and f subshells with varying numbers of orbitals. Electrons prefer to stay unpaired as long as possible before doubling up to fill open orbitals within their subshell.
The periodic table is organized in such a way that it can be used to determine electron configurations. Rows correspond to energy levels, and blocks correspond to subshells. The four blocks are the s, p, d, and f blocks. When numbering the rows for electron configurations, the s block starts with row 1, p with row 2, d with row 4, and f with row 6. For larger atoms, noble gas configuration can be used as a form of shorthand. An atom's electron configuration can provide information about its magnetic behavior, with paramagnetic atoms having unpaired valence electrons and diamagnetic atoms having all paired valence electrons. When an atom gains or loses electrons to become an ion, its electron configuration changes, resembling the configuration of a different element.
Lesson Outline
<ul> <li>Introduction to Electron Configurations</li> <ul> <li>Understanding electrons and orbitals</li> <li>Shells, subshells, and orbitals</li> <li>Aufbau principle</li> <li>Using periodic table to determine electron configurations: the "n + l" rule</li> </ul> <li>Electron behavior and magnetism</li> <ul> <li>Paramagnetic atoms</li> <li>Diamagnetic atoms</li> </ul> <li>Electron configurations and the periodic table</li> <ul> <li>Rows and blocks</li> <li>Numbering rows for electron configurations</li> <li>Reading the periodic table</li> </ul> <li>Electron configurations for larger atoms</li> <ul> <li>Noble Gas Configuration</li> <li>Example: Ca -> [Ar]4s2</li> </ul> <li>Electron configurations and ions</li> <ul> <li>Changing electron configurations when forming ions</li> </ul> </ul>
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FAQs
The periodic table is arranged in such a way that it reflects the electron configurations of elements. Each element's position in the periodic table is based on its atomic number, which is equal to the number of protons and electrons in a neutral atom. As you move across each period (horizontal row), electrons are added to the same energy level, filling up the subshells, while moving down a group (vertical column) indicates the start of a new energy level. This arrangement allows for easy determination of an element's electron configuration and its relationship to the other elements around it in the periodic table.
Valence electrons are the outermost electrons in an atom's electron configuration and are situated in the highest energy level. They play a crucial role in determining the chemical properties of an element and its reactivity. Valence electrons are responsible for an atom's tendency to gain, lose, or share electrons during a chemical reaction or bond formation. Elements within the same group on the periodic table tend to have similar valence electron configurations, resulting in similar chemical properties and reactivities.
The Aufbau principle, orbitals, and subshells are key concepts that help understand and predict electron configurations. The Aufbau principle states that electrons fill the available orbitals in order of increasing energy, starting from the lowest energy orbital. Orbitals are regions in an atom where there is a high probability of finding an electron, and they are classified as s, p, d, or f based on their shapes. Subshells are groups of orbitals within the same energy level. Using these concepts, we can build an electron configuration following the Aufbau principle, filling the orbitals in the order of 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, and so on. Moreover, this principle helps to predict the placement of electrons in various orbitals and subshells, which in turn influences the chemical properties of elements.
Noble gas configuration refers to the electron configuration of a noble gas, which has completely filled s and p orbitals in its outermost energy level, resulting in high stability and an unreactive nature. Using noble gas notation, an element's electron configuration can be shorthand represented by enclosing the preceding noble gas in brackets and writing the remaining configuration for the outer electron shells (example: [Xe]6s1). Paramagnetic substances are those that have unpaired electrons in their electron configurations. External magnetic fields partially align these unpaired electrons, making the substance attracted to the magnetic field. Diamagnetic substances, on the other hand, have all electrons paired in their electron configurations. This pairing causes the substance to be repelled rather than attracted to a magnetic field. The difference in these two properties arises due to the number of unpaired electrons present in the electron configuration of the substance.
Ions form when an atom gains or loses one or more electrons, resulting in an unequal number of protons and electrons, and subsequently a net electrical charge. This gain or loss of electrons changes the electron configuration of the atom. Cations are positively-charged ions formed by the loss of one or more valence electrons, typically from metals. This loss of electrons typically results in a stable electron configuration similar to that of the nearest noble gas with a lower atomic number. Anions, on the other hand, are negatively-charged ions formed by the gain of one or more electrons, typically by non-metals. This electron gain typically gives the atom the stable electron configuration of the nearest noble gas with a higher atomic number. Both cations and anions form as a result of the elements' tendency to achieve a stable electron configuration, which is most often achieved by having a completely filled outermost energy level.
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