Lesson Notes By Weeks and Term v5 - Grade 10

Lesson Video

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Performance objectives

• Define chemical bonds and the octet rule.
• Differentiate between ionic, covalent, and metallic bonds.
• Draw Lewis dot diagrams for simple molecules.
• Relate bond type to the properties of a substance.
• Use electronegativity to predict bond type.

Lesson summary

This week, we delve into the fascinating world of chemical bonding and how it dictates the properties of substances. Understanding chemical bonding is crucial for comprehending the structure and behaviour of matter, from the air we breathe to the materials used to build our homes and infrastructure. Think about the strength of the steel used in bridges or the properties of the plastics used to make everyday items. These properties are all directly related to the type of chemical bonds holding the atoms together.

Lesson notes

2.1 Chemical Bonds: The Glue Holding Matter Together A chemical bond is a force of attraction that holds atoms together to form molecules or ionic compounds. Atoms bond together to achieve a more stable electron configuration, usually resembling that of a noble gas (8 valence electrons, except for Helium which needs 2). This tendency to achieve 8 valence electrons is known as the octet rule. 2.2 Types of Chemical Bonds There are three main types of chemical bonds: Ionic Bonds: Formed through the transfer of electrons from one atom to another, resulting in the formation of ions (charged particles). These bonds typically occur between a metal (which loses electrons to form a positive ion or cation) and a nonmetal (which gains electrons to form a negative ion or anion). The electrostatic attraction between the oppositely charged ions holds them together in a crystal lattice.

Example: Sodium chloride (NaCl), common table salt. Sodium (Na) has one valence electron and chlorine (Cl) has seven. Sodium loses its valence electron to chlorine, forming Na + and Cl - ions. The electrostatic attraction between these ions creates the ionic bond.

Properties:* High melting and boiling points, hard and brittle, conduct electricity when dissolved in water (electrolyte).

Covalent Bonds: Formed through the sharing of electrons between two atoms. Covalent bonds typically occur between two nonmetals. The shared electrons are attracted to the nuclei of both atoms, holding them together.

Example:* Water (H 2 O). Oxygen (O) has six valence electrons and needs two more to complete its octet. Each hydrogen (H) atom shares one electron with the oxygen atom, forming two covalent bonds.

Properties:* Lower melting and boiling points compared to ionic compounds, can be solids, liquids, or gases at room temperature, generally do not conduct electricity. Covalent compounds can be polar or nonpolar, influencing their solubility.

Lewis Dot Diagrams: A visual representation of covalent bonding, showing the valence electrons as dots around the element symbols. Shared electron pairs are represented by lines.

Example: Water (H 2 O)* H : O : H (Each dot represents a valence electron. The lines can be used to represent each shared pair H-O-H).

Metallic Bonds: Found in metals. In metallic bonding, valence electrons are delocalized and free to move throughout the metal structure, forming a "sea of electrons". The positively charged metal ions are held together by their attraction to the sea of electrons.

Example:* Copper (Cu). Copper atoms readily lose their valence electrons, and these electrons are free to move throughout the copper structure.

Properties:* Excellent conductors of electricity and heat, malleable (can be hammered into sheets), ductile (can be drawn into wires), lustrous (shiny). 2.3 Electronegativity and Bond Type Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The electronegativity difference between two bonding atoms can be used to predict the type of bond that will form: Large electronegativity difference (greater than 1.7): Ionic bond. Intermediate electronegativity difference (between 0.4 and 1.7): Polar covalent bond. The electrons are shared unequally, creating a partial positive charge (δ+) on the less electronegative atom and a partial negative charge (δ-) on the more electronegative atom. Small electronegativity difference (less than 0.4): Nonpolar covalent bond. The electrons are shared equally. 2.4 Worked

Examples: Predict the type of bond that will form between potassium (K) and oxygen (O). Electronegativity of K = 0.82 Electronegativity of O = 3.44 Electronegativity difference = 3.44 - 0.82 = 2.62 Since the electronegativity difference is greater than 1.7, an ionic bond will form. Potassium will lose electrons to oxygen to form potassium oxide (K 2 O). Draw the Lewis dot diagram for carbon dioxide (CO 2 ). Carbon (C) has 4 valence electrons. Oxygen (O) has 6 valence electrons. Carbon is the central atom. Connect each oxygen atom to the carbon atom with a single bond (C-O). This uses 2 valence electrons from carbon and 1 from each oxygen. Carbon still needs 4 more electrons to complete its octet, and each oxygen needs

6. Form double bonds between the carbon and each oxygen atom (O=C=O). Each double bond contributes 2 electrons towards the octet of the central atom. Place the remaining valence electrons as lone pairs on the oxygen atoms so that each oxygen atom also has 8 electrons.

The Lewis structure is: :O=C=O: Explain why metals are good conductors of electricity. Metals have a "sea of delocalized electrons." This means that the valence electrons are not bound to individual atoms but are free to move throughout the metallic structure. When a voltage is applied across a metal, these electrons can easily move, carrying an electric current. This free movement of electrons is the reason metals are excellent conductors of electricity.