Electron Arrangement and Periodic Trends

Flashcards covering electron arrangement, quantum numbers, electron configurations, and periodic trends based on lecture notes for Chapter 3.

Schrodinger equation

A calculation that suggests electrons behave like waves rather than just particles and is used to organize electrons around the nucleus.

Wave function

A concept developed by Erwin Schrodinger to describe how electrons are organized around the nucleus within the electron cloud.

Orbitals

Regions of space surrounding the nucleus where electrons are found, representing the highest probability density map of an electron existing at a particular point.

Quantum numbers

A series of numbers and letters that characterize orbitals and provide the relative position of an electron in an atom at any given point in time and space.

Principal quantum number (n)

The first quantum number, designated by 'n', which defines the electron shell, its relative size, and energy level. It is a positive integer (1, 2, 3…). Tthe larger the 'n', the larger the shell and higher its energy.

Secondary quantum number (l)

The second quantum number (often a cursive 'l'), which indexes energy differences between orbitals within the same shell and dictates the shape of the orbital. Its values range from 0 to n-1.

Magnetic quantum number (m_l)

The third quantum number, which dictates the number of orbitals available for electrons to be placed within a subshell. Its values range from -l to +l (including 0).

Pauli exclusion principle

A principle stating that no two electrons in an atom may have the same set of four quantum numbers. It indicates that there are two electrons per orbital, which must have opposite spins.

Spin quantum number

The fourth quantum number, which indicates the direction of an electron's spin, represented by values of +1/2 or -1/2.

s orbital

A spherical-shaped orbital, corresponding to a secondary quantum number (l) value of 0.

p orbital

An oblong, figure-eight shaped orbital, corresponding to a secondary quantum number (l) value of 1.

Electron configuration

The distribution of electrons throughout the orbitals within an atom, typically written to show the principal quantum number, orbital type, and number of electrons in that orbital.

Ground state electronic configuration

The lowest energy arrangement of electrons in an atom, which is the most stable state where electrons are 'happy'.

Aufbau principle

A principle used to properly distribute electrons by filling the lowest energy subshells or orbitals first, working upwards in energy.

Noble gas configurations

An abbreviated form of electron configuration where the symbol of the closest previous noble gas is enclosed in brackets, followed by the configuration of the remaining outer electrons.

Valence electrons

Electrons located in the outermost energy level of an atom (the highest principal quantum number 'n'), which are involved in chemical reactions and bonding. Their number correlates to the group number (1A-8A).

Alkali metals

Elements found in Group 1 (1A) of the periodic table.

Alkaline earth metals

Elements found in Group 2 (2A) of the periodic table.

Halogens

Elements found in Group 7A (17) of the periodic table.

Noble gases

Elements found in Group 8A (18) of the periodic table, characterized by full electron shells.

Transition metals

Elements located in the d-block of the periodic table (Groups 3-12).

Main group elements

Elements in the columns labeled with 'A' (Groups 1A, 2A, and 3A-8A, or groups 1, 2, and 13-18) on the periodic table.

Metals

Elements typically found to the left of the metal/nonmetal barrier on the periodic table, characterized by shiny solids, good conductivity (heat and electricity), ductility, and malleability.

Nonmetals

Elements typically found to the right of the metal/nonmetal barrier on the periodic table, characterized by dull appearance, brittleness, and poor conductivity (good insulators).

Metalloids

Elements with material properties that fall in between those of metals and nonmetals, often used as semiconductors (e.g., Silicon).

Periodicity (Periodic trends)

The phenomenon where elements are arranged not only by increasing atomic number but also based on their relative chemical behavior, showing similar properties occurring at regular intervals.

Atomic radius

The relative size of an atom. It increases as you go down a group and decreases as you go across a period from left to right.

Effective nuclear charge

The strength of interaction between the nucleus and an atom's valence electrons. It increases from left to right across a period, leading to smaller atomic sizes.

Ionization energy

The energy required to remove an electron from a neutral atom, forming a positively charged ion (cation). Metals generally have lower ionization energies.

Cation

A positively charged species formed when a neutral atom loses one or more electrons. Cations are typically smaller than their parent neutral atoms.

Electron affinity

The energy change that occurs when an electron is added to a neutral atom to form a negatively charged ion (anion). Nonmetals tend to have large, negative electron affinity values.

Anion

A negatively charged species formed when a neutral atom gains one or more electrons.

Anomalous electron configurations (Chromium & Copper)

Exceptions to typical electron configuration rules (e.g., Chromium and Copper) where a d orbital 'steals' an electron from an s orbital to achieve a more energetically stable half-filled or completely filled d-orbital configuration.