Lesson Notes By Weeks and Term v5 - Grade 10

Lesson Video

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Performance objectives

• Define and differentiate ionic, covalent, and metallic bonds.
• Explain the role of electronegativity in bond formation.
• Draw Lewis dot diagrams for simple molecules.
• Relate bond type to the properties of a substance.
• Predict the type of bonding in a given substance.

Lesson summary

This week, we delve into the fascinating world of chemical bonding and the particles of substances. Understanding how atoms combine to form molecules and larger structures is fundamental to comprehending the properties of everything around us – from the air we breathe to the materials that build our homes and the food we eat. This knowledge is especially important in South Africa, where industries like mining, agriculture, and manufacturing rely heavily on understanding the chemical properties of various substances. For example, understanding chemical bonding helps us to develop corrosion-resistant materials for mining equipment or to create more effective fertilizers for agriculture.

Lesson notes

2.1 Introduction to Chemical Bonds: Chemical bonds are the attractive forces that hold atoms together to form molecules, ions, and extended networks. The drive for atoms to form bonds arises from their desire to achieve a stable electron configuration, usually resembling the noble gases with a full outer electron shell (octet rule, with exceptions like Hydrogen aiming for two electrons). There are three primary types of chemical bonds: ionic, covalent, and metallic. 2.2 Ionic Bonds: Definition: Ionic bonds form through the electrostatic attraction between oppositely charged ions (cations and anions). This occurs when one atom readily loses electrons (typically a metal) and another atom readily gains electrons (typically a nonmetal). This transfer of electrons results in the formation of ions.

Formation: The atom that loses electrons becomes a positively charged ion (cation), while the atom that gains electrons becomes a negatively charged ion (anion). These oppositely charged ions are then attracted to each other, forming an ionic bond.

Electronegativity: Ionic bonds typically form when there is a large electronegativity difference (generally greater than 1.7) between the two atoms involved. Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond.

Properties: Ionic compounds generally have high melting and boiling points, are hard and brittle, and conduct electricity when dissolved in water or melted (but not in the solid state). This is because strong electrostatic forces hold the ions in a fixed lattice structure. The mobility of the ions when dissolved or melted allows for charge transport.

Examples: Sodium chloride (NaCl - common table salt), magnesium oxide (MgO), and calcium fluoride (CaF2).

Example 1: Formation of Sodium Chloride (NaCl) Sodium (Na) has one valence electron, while chlorine (Cl) has seven. Sodium readily loses its valence electron to achieve a stable electron configuration. Na → Na + + e - Chlorine readily gains an electron to achieve a stable electron configuration. Cl + e - → Cl - The resulting ions, Na + and Cl - , are oppositely charged and attract each other electrostatically, forming NaCl. NaCl forms a giant ionic lattice structure, where each Na + ion is surrounded by Cl - ions, and each Cl - ion is surrounded by Na + ions. 2.3 Covalent Bonds: Definition: Covalent bonds form when atoms share electrons to achieve a stable electron configuration. This typically occurs between two nonmetal atoms.

Types: Covalent bonds can be single (sharing one pair of electrons), double (sharing two pairs of electrons), or triple (sharing three pairs of electrons). The more shared electrons, the stronger and shorter the bond.

Electronegativity: Covalent bonds typically form when there is a small electronegativity difference (generally less than 1.7) between the two atoms involved.

Polarity: If the electronegativity difference is zero, the bond is considered nonpolar covalent (e.g., H-H). If there is a significant electronegativity difference, the bond is polar covalent. In a polar covalent bond, the electrons are shared unequally, resulting in a partial positive charge (δ+) on the less electronegative atom and a partial negative charge (δ-) on the more electronegative atom.

Properties: Covalent compounds can exist as gases, liquids, or solids at room temperature. They generally have lower melting and boiling points than ionic compounds. They typically do not conduct electricity because there are no mobile ions or electrons.

Examples: Water (H2O), methane (CH4), carbon dioxide (CO2), and diamond (C).

Example 2: Formation of Water (H2O) Oxygen (O) has six valence electrons, and hydrogen (H) has one. Oxygen needs two more electrons to achieve a stable electron configuration. Each hydrogen atom shares one electron with the oxygen atom. Oxygen now has eight electrons in its valence shell (two from each hydrogen atom), and each hydrogen atom has two electrons (shared with oxygen). The O-H bonds are polar covalent because oxygen is more electronegative than hydrogen, resulting in a partial negative charge on the oxygen atom and partial positive charges on the hydrogen atoms. 2.4 Metallic Bonds: Definition: Metallic bonds form between metal atoms. Metal atoms readily lose their valence electrons, forming positive ions (cations) that are surrounded by a "sea" of delocalized electrons.

Formation: The delocalized electrons are not associated with any particular atom but are free to move throughout the metal structure. This creates a strong attractive force between the positive ions and the electron sea.

Properties: Metals generally have high melting and boiling points (although mercury is an exception), are malleable (can be hammered into sheets), ductile (can be drawn into wires), and excellent conductors of heat and electricity. The delocalized electrons are responsible for the high electrical and thermal conductivity.