Electronegativity and Molecular Polarity

This page delves deeper into the concepts of electronegativity and molecular polarity, which are crucial for understanding chemical bonding in A Level Chemistry.

Definition: Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond.

The page explains trends in electronegativity across the periodic table:

1. Across a period: Electronegativity increases from left to right.

   • Reason: Increasing nuclear charge (proton number) leads to greater attraction of bonding electrons to the nucleus.

2. Down a group: Electronegativity decreases.

   • Reason: Increased shielding effect and distance from the nucleus reduce the attraction between bonding electrons and the nucleus.

Highlight: Fluorine is the most electronegative element in the periodic table.

The document introduces the concept of dipole moments in polar covalent bonds. When two atoms of different electronegativities form a covalent bond, the more electronegative atom attracts the bonding electrons more strongly, creating a slight negative charge (δ-) on that atom and a slight positive charge (δ+) on the less electronegative atom.

Vocabulary: Dipole moment - a measure of the separation of positive and negative electrical charges in a system of atoms or molecules.

The page emphasizes that the overall polarity of a molecule depends not only on the presence of polar bonds but also on the molecular geometry:

Example: Carbon tetrachloride (CCl₄) has polar C-Cl bonds, but the tetrahedral arrangement of these bonds results in a non-polar molecule as the dipole moments cancel out.

This concept is crucial for understanding molecular interactions and properties in OCR A Level Chemistry Structure and bonding questions.

The document concludes by reinforcing the importance of understanding electronegativity and molecular polarity in predicting and explaining chemical behavior, which is essential for answering A level Chemistry electron configuration questions and understanding intermolecular forces in A Level Chemistry OCR.

Electron Orbitals and Molecular Shapes

This page introduces the concept of electron orbitals and their role in determining molecular shapes. It covers the basic principles of orbital filling and explores various molecular geometries.

Definition: An orbital is a region of high probability for finding an electron in an atom.

Highlight: Each orbital can hold a maximum of two electrons.

The page discusses the order of orbital filling, which generally follows increasing energy levels (1s, 2s, 2p, 3s, 3p, 3d, 4s, etc.). However, there are some important exceptions to remember:

Example: Copper (Cu) has an electron configuration that ends with 3d¹⁰ 4s¹ instead of the expected 3d⁹ 4s².

Example: Chromium (Cr) has a configuration of 3d⁵ 4s¹ rather than 3d⁴ 4s².

These exceptions occur due to the stability of half-filled or fully-filled d-orbitals.

The page then transitions to molecular shapes, introducing linear and non-linear molecules. It presents various molecular geometries, including:

1. Linear (e.g., CO₂)
2. Trigonal planar (e.g., BF₃)
3. Tetrahedral (e.g., CH₄)
4. Pyramidal (e.g., NH₃)

Vocabulary: Bond angle - the angle formed between two adjacent bonds in a molecule.

The document explains that molecular shapes are determined by the number of bonding pairs and lone pairs of electrons. It also introduces the concept of molecular polarity, noting that the shape of a molecule can affect whether it is polar or non-polar.

Highlight: Lone pairs of electrons repel more strongly than bonding pairs, which can affect bond angles and overall molecular shape.