Key concepts in chemistry Revision Notes

Location on the Periodic Table

Metals are predominantly found on the left and in the central blocks (Groups 1, 2, and the transition metals) of the periodic table. Non-metals are located on the right-hand side, typically starting from Group 13 onwards (excluding the metalloids that lie along the border). A 'staircase' line on some periodic tables separates metals from non-metals, with elements along this line sometimes exhibiting properties of both (metalloids).

Atomic Structure and Ion Formation

The classification as metal or non-metal is directly linked to an element's atomic structure, particularly its number of valence electrons. Metals generally have few valence electrons (1, 2, or 3) and readily lose these to achieve a stable electron configuration, forming positive ions. Non-metals typically have more valence electrons (4, 5, 6, 7, or 8) and tend to gain electrons to complete their outer shell, forming negative ions, or they share electrons in covalent bonds. Their tendency to lose or gain electrons dictates their chemical reactivity.

Understanding Electron Shells

Electrons orbit the nucleus in distinct energy levels or shells. Each shell has a maximum number of electrons it can hold: the first shell holds up to 2 electrons, the second up to 8, and the third up to 8 for the first 20 elements. This pattern of filling is crucial for predicting chemical behaviour.

Determining Electronic Configuration

To find the electronic configuration, locate the atomic number of the element, which equals the number of electrons in a neutral atom. Fill the shells sequentially, starting from the innermost shell, until all electrons are placed. For example, an element with 11 electrons will have a configuration of 2.8.1.

Predicting Reactivity

The number of electrons in the outermost shell, known as valence electrons, largely determines an element's chemical reactivity. Atoms tend to gain, lose, or share electrons to achieve a stable outer shell, typically with 8 electrons (octet rule, except for the first shell which is stable with 2), like the noble gases.

Periods and Electron Shells

The periodic table is organised into horizontal rows called periods. The period number an element belongs to directly corresponds to the number of electron shells its atoms contain. For example, elements in Period 2 have two occupied electron shells, while elements in Period 4 have four occupied electron shells. This regularity allows for a quick determination of shell count from an element's position.

Groups and Valence Electrons

Vertical columns in the periodic table are known as groups. The group number (for main group elements, excluding transition metals) indicates the number of electrons in the outermost shell, known as valence electrons. For example, elements in Group 1 possess one valence electron, and elements in Group 16 have six valence electrons. Elements within the same group tend to exhibit similar chemical properties due to having the same number of valence electrons, as these electrons are primarily involved in chemical reactions.

Predicting Configuration from Position

Knowing an element's position in the periodic table allows for a straightforward prediction of its electronic configuration. For example, an element in Period 3, Group 2 will have three electron shells and two valence electrons. This prediction can then be used to understand its reactivity and bonding behaviour because the valence electrons dictate most chemical interactions. The first shell can hold up to 2 electrons, the second up to 8, and subsequent shells can hold more, although 8 is still a key number for stability in the third shell for many elements.

Electron Transfer for Stability

Atoms aim for a stable electron configuration, usually a full outer shell (like noble gases). This is achieved by gaining or losing electrons. Metals tend to lose electrons, while non-metals tend to gain them. This electron transfer leads to the formation of charged particles called ions.

Cations: Group 1 and 2 Metals

Group 1 elements (e.g., sodium) have one electron in their outer shell. They lose this electron to form an ion with a 1+ charge (e.g., Na⁺). Group 2 elements (e.g., magnesium) have two outer electrons. They lose both electrons to form an ion with a 2+ charge (e.g., Mg²⁺). These positively charged ions are called cations.

Anions: Group 6 and 7 Non-metals

Group 7 elements (halogens, e.g., chlorine) have seven electrons in their outer shell. They gain one electron to complete their outer shell, forming an ion with a 1- charge (e.g., Cl⁻). Group 6 elements (e.g., oxygen) have six outer electrons. They gain two electrons to form an ion with a 2- charge (e.g., O²⁻). These negatively charged ions are called anions.

Ionic Bonding

The oppositely charged ions (cations and anions) are attracted to each other by strong electrostatic forces. This attraction forms an ionic bond, holding the ions together in a crystal lattice structure to create an ionic compound.