Electron configurations reveal how electrons are arranged in atoms, following key principles like the Aufbau principle and Hund's rule. This knowledge is crucial for understanding atomic structure and predicting chemical behavior, forming the foundation for exploring electron properties in atoms.
The periodic table's organization reflects electron configurations, making it a powerful tool for predicting atomic properties. By understanding core and valence electrons, we can explain chemical reactivity and physical characteristics, connecting atomic structure to observable phenomena.
Write the electron configuration for an atom or ion using the Aufbau principle, Hund's rule, and the Pauli exclusion principle
Principles for writing electron configurations
- Aufbau principle electrons fill orbitals starting with the lowest available energy levels before filling higher levels
- Electrons occupy orbitals in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
- Hund's rule electrons fill orbitals of equal energy with one electron each before pairing them in the same orbital
- For example, in the 2p subshell, each of the three p orbitals (2pₓ, 2p_y, 2p_z) will have one electron before any orbital has two electrons
- Pauli exclusion principle no two electrons in an atom can have the same four quantum numbers, meaning each orbital can hold a maximum of two electrons with opposite spins
- The two electrons in an orbital must have opposite spin quantum numbers (m_s = +1/2 and -1/2)
Writing electron configurations
- Electron configurations are written with the subshell label (1s, 2p) followed by a superscript indicating the number of electrons in that subshell
- For example, the electron configuration of carbon (6 electrons) is 1s² 2s² 2p²
- For atoms with partially filled d or f subshells, the subshell label is preceded by the element symbol in brackets
- For example, the electron configuration of manganese (25 electrons) is [Ar] 4s² 3d⁵, where [Ar] represents the electron configuration of argon (1s² 2s² 2p⁶ 3s² 3p⁶)
- Electron configurations for ions are written by adding or removing electrons from the neutral atom configuration, starting with the highest energy electrons
- For example, the electron configuration of the Fe²⁺ ion is [Ar] 3d⁶, as the two 4s electrons are removed first
Predict the electron configuration of an atom using the periodic table
Relationship between the periodic table and electron configurations
- The periodic table is arranged in order of increasing atomic number, which corresponds to the number of electrons in a neutral atom
- Elements in the same group (vertical column) have similar electron configurations in their outermost energy levels
- For example, all Group 1 elements (alkali metals) have one valence electron in an s orbital
- The period (horizontal row) indicates the highest principal quantum number (energy level) occupied by electrons in an atom
- For example, elements in Period 3 have electrons in the n=3 energy level
- The block (s, p, d, or f) in which an element is located corresponds to the subshell being filled with electrons
- For example, elements in the p-block are filling their p subshells
Determining electron configurations using the periodic table
- The electron configuration of an atom can be determined by following the order of increasing atomic number and applying the Aufbau principle, Hund's rule, and the Pauli exclusion principle
- Start with the element's period and block to determine the highest energy level and subshell being filled
- Add electrons to subshells according to the Aufbau principle until the total number of electrons equals the element's atomic number
- Apply Hund's rule and the Pauli exclusion principle when filling subshells with multiple orbitals (p, d, and f subshells)
Identify and differentiate between core electrons and valence electrons
Core electrons
- Core electrons are the electrons in the inner, completely filled energy levels of an atom
- Core electrons are not involved in bonding and are less important in determining an element's chemical properties
Valence electrons
- Valence electrons are the electrons in the outermost energy level (highest principal quantum number) of an atom
- The number of valence electrons determines an element's chemical properties, such as its reactivity and the types of bonds it can form
- In the electron configuration notation, valence electrons are represented by the subshells with the highest principal quantum number
- For example, in the electron configuration of carbon (1s² 2s² 2p²), the 2s and 2p electrons are the valence electrons
- For main group elements, the number of valence electrons corresponds to the group number
- Group 1 elements (alkali metals) have 1 valence electron
- Group 2 elements (alkaline earth metals) have 2 valence electrons
Determine the relationship between electron configuration and the physical and chemical properties of elements
Chemical properties
- Elements with similar electron configurations in their outermost energy levels (valence electrons) exhibit similar chemical properties
- Elements in the same group have similar reactivities and form similar compounds
- The stability of an atom is related to its electron configuration, with atoms having completely filled or half-filled subshells being more stable
- Noble gases (Group 18) have completely filled outer subshells and are chemically inert
- The reactivity of an element is determined by its ability to gain, lose, or share electrons to achieve a stable electron configuration (usually that of a noble gas)
- Metals tend to lose electrons to form cations, while nonmetals tend to gain electrons to form anions
Physical properties
- The magnetic properties of an element are influenced by the presence of unpaired electrons in its electron configuration
- Atoms with unpaired electrons are paramagnetic and attracted to magnetic fields, while atoms with paired electrons are diamagnetic and repelled by magnetic fields
- The electron configuration of an element can also affect its physical properties, such as melting point, boiling point, and electrical conductivity, due to the interactions between electrons and the atomic nucleus
- Elements with strong metallic bonding (many delocalized electrons) tend to have high melting points and good electrical conductivity (metals)
- Elements with covalent bonding (shared electrons) tend to have lower melting points and poor electrical conductivity (nonmetals)