Lesson Notes By Weeks and Term v3 - Senior Secondary 1
Chemical Combination
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Chemical Combination
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Subject: Chemistry
Class: Senior Secondary 1
Term: 3rd Term
Week: 1
Theme: The Chemical World
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2.1.1 The First Twenty Elements The periodic table is an organized arrangement of elements. For chemical combination, understanding the first twenty elements is fundamental, as their electronic structures dictate their reactivity. | Atomic Number (Z) | Symbol | Name | Electronic Configuration | | :---------------- | :----- | :---------- | :----------------------- | | 1 | H | Hydrogen | 1 | | 2 | He | Helium | 2 | | 3 | Li | Lithium | 2, 1 | | 4 | Be | Beryllium | 2, 2 | | 5 | B | Boron | 2, 3 | | 6 | C | Carbon | 2, 4 | | 7 | N | Nitrogen | 2, 5 | | 8 | O | Oxygen | 2, 6 | | 9 | F | Fluorine | 2, 7 | | 10 | Ne | Neon | 2, 8 | | 11 | Na | Sodium | 2, 8, 1 | | 12 | Mg | Magnesium | 2, 8, 2 |...
| Dative (Coordinate) Bonding | | :----------------- | :-------------------------------------------------- | :--------------------------------------------------- | :------------------------------------------------------------- | | Electron involvement | Complete transfer of electrons | Mutual sharing of electrons | Sharing of electrons, but both come from one atom | | Atoms involved | Metal + Non-metal | Non-metal + Non-metal | Between an atom with a lone pair and an electron-deficient atom | | Resulting particles | Ions (cations and anions) | Molecules | Ions or molecules with specific sites for donation/acceptance | | Forces | Strong electrostatic attraction between ions | Weaker intermolecular forces (often) between molecules | Strong intramolecular force, similar to covalent | | Solubility (in water) | Generally soluble | Varies (polar soluble, non-polar insoluble) | Varies (often soluble if ionic component is present) | | Electrical Conductivity | Good conductors when molten or in solution | Poor conductors (except for specific cases like graphite) | Poor conductors (unless overall charged species) | | Physical State | Mostly solids | Solids, liquids, or gases | Solids, liquids, or gases | | Melting/Boiling Points | High | Generally low | Varies | Atoms combine to achieve a stable electronic configuration, typically an octet (8 valence electrons) or a duplet (2 valence electrons for very small atoms like H, He, Li, Be) in their outermost shell, similar to the noble gases. This process involves the transfer or sharing of valence electrons. 2.2.1 Types of Chemical Bonding Electrovalent (Ionic)
Bonding: Formation: Occurs between a metal and a non-metal. It involves the complete transfer of one or more valence electrons from a metal atom (which forms a positive ion/cation) to a non-metal atom (which forms a negative ion/anion). The resulting oppositely charged ions are held together by strong electrostatic forces of attraction.
Characteristics: Usually results in solid compounds with high melting and boiling points, typically soluble in water, and conductors of electricity when molten or in aqueous solution.
Example: Formation of Sodium Chloride (NaCl) Na (2, 8, 1) loses 1 electron to become Na+ (2, 8) [stable octet]. Cl (2, 8, 7) gains 1 electron to become Cl− (2, 8, 8) [stable octet]. Na+ and Cl− ions are attracted to form NaCl.
Example: Formation of Magnesium Chloride (MgCl2) Mg (2, 8, 2) loses 2 electrons to become Mg2+ (2, 8). Each Cl (2, 8, 7) gains 1 electron to become Cl− (2, 8, 8). Two Cl atoms are needed for one Mg atom. Mg2+ and 2Cl− ions are attracted to form MgCl
2. Covalent Bonding: Formation: Occurs typically between two non-metal atoms. It involves the sharing of one or more pairs of valence electrons between atoms. Each shared pair of electrons constitutes a single covalent bond.
Characteristics: Can form solids, liquids, or gases. Generally have lower melting and boiling points than ionic compounds, often insoluble in water (or form solutions that do not conduct electricity).
Types of Covalent Bonds: Single bond: One shared pair of electrons (e.g., H-H in H2, Cl-Cl in Cl2, H-Cl in HCl).
Double bond: Two shared pairs of electrons (e.g., O=O in O2).
Triple bond: Three shared pairs of electrons (e.g., N≡N in N2).
Example: Formation of Water (H2O) Oxygen (2, 6) needs 2 electrons. Each Hydrogen (1) needs 1 electron. Oxygen shares one electron with each of the two hydrogen atoms, forming two single covalent bonds. H : O : H (Lewis structure showing shared pairs).
Example: Formation of Methane (CH4) Carbon (2, 4) needs 4 electrons. Each Hydrogen (1) needs 1 electron. Carbon shares one electron with each of the four hydrogen atoms, forming four single covalent bonds.
Dative (Coordinate)
Covalent Bonding: Formation: A special type of covalent bond where both shared electrons in the bond come from only one of the participating atoms. The atom donating the electron pair is called the donor, and the atom accepting the pair is called the acceptor.
Characteristics: Once formed, a dative bond is indistinguishable from a normal covalent bond. It is represented by an arrow (→) pointing from the donor atom to the acceptor atom.
Example: Formation of Ammonium Ion (NH4+) Ammonia (NH3) has a lone pair of electrons on the nitrogen atom. A hydrogen ion (H+), which has no electrons, can accept this lone pair.
H3N: → H+ forms NH4+.
Example: Formation of Hydronium Ion (H3O+) Water (H2O) has two lone pairs on the oxygen atom. A hydrogen ion (H+) can accept one of these lone pairs.
H2O: → H+ forms H3O+. 2.2.2 Differentiating Bond Types: | Feature | Electrovalent (Ionic) Bonding | Covalent Bonding | Dative (Coordinate) Bonding | | :----------------- | :-------------------------------------------------- | :--------------------------------------------------- | :------------------------------------------------------------- | | Electron involvement | Complete transfer of electrons | Mutual sharing of electrons | Sharing of electrons, but both come from one atom | | Atoms involved | Metal + Non-metal | Non-metal + Non-metal | Between an atom with a lone pair and an electron-deficient atom | | Resulting particles | Ions (cations and anions) | Molecules | Ions or molecules with specific sites for donation/acceptance | | Forces | Strong electrostatic attraction between ions | Weaker Naming compounds allows for clear communication among chemists. 2.3.1 Conventional (Common) Names These are historical names that are widely used but do not follow systematic rules. Students should be familiar with common ones.
H2O: Water NH3: Ammonia CH4: Methane NaCl: Table salt NaOH: Caustic soda (Sodium hydroxide)
CaCO3: Limestone/Marble/Chalk (Calcium carbonate)
H2SO4: Oil of vitriol (Sulphuric acid) 2.3.2 IUPAC (International Union of Pure and Applied Chemistry) Names IUPAC nomenclature provides a systematic way to name compounds based on their chemical composition. For Binary Ionic Compounds (Metal + Non-metal): Name the metal cation first (using its elemental name). Name the non-metal anion second, changing its ending to "-ide". If the metal forms multiple ions, indicate its charge (oxidation state) with a Roman numeral in parentheses after the metal's name (e.g., Iron(II) chloride for FeCl2, Iron(III) chloride for FeCl3).
Examples: NaCl: Sodium chloride MgCl2: Magnesium chloride CaO: Calcium oxide Al2O3: Aluminium oxide FeS: Iron(II) sulphide CuO: Copper(II) oxide For Binary Covalent Compounds (Non-metal + Non-metal): The first element in the formula is named first (unless it's oxygen in some cases). The second element is named with its suffix changed to "-ide". Use prefixes to indicate the number of atoms of each element. 1: mono- (often omitted for the first element) 2: di- 3: tri- 4: tetra- 5: penta- 6: hexa-
Examples: CO: Carbon monoxide CO2: Carbon dioxide SO2: Sulfur dioxide NO2: Nitrogen dioxide CCl4: Carbon tetrachloride N2O4: Dinitrogen tetroxide PCl3: Phosphorus trichloride For Compounds with Polyatomic Ions: Name the cation first, then the polyatomic anion (e.g., SO42− is sulfate, NO3− is nitrate, CO32− is carbonate, OH− is hydroxide, PO43− is phosphate, NH4+ is ammonium).
Examples: CaCO3: Calcium carbonate Na2SO4: Sodium sulfate NH4Cl: Ammonium chloride NaOH: Sodium hydroxide Matter exists predominantly in three states: solid, liquid, and gas. These states are distinguishable by their physical properties, which are explained by the kinetic theory of matter. 2.4.1 Distinguishing States of Matter: | Property | Solid | Liquid | Gas | | :---------------- | :------------------------------------------ | :----------------------------------------- | :------------------------------------------ | | Shape | Definite | Indefinite (takes shape of container) | Indefinite (takes shape of container) | | Volume | Definite | Definite | Indefinite (fills entire container) | | Compressibility | Very low | Very low | High | | Fluidity | No (rigid) | Yes (flows) | Yes (flows) | | Particle Arrangement | Tightly packed in fixed positions, ordered | Loosely packed, can slide past each other | Very far apart, randomly arranged | | Intermolecular Forces | Very strong | Moderate | Very weak / Negligible | | Particle Movement | Vibrate about fixed positions | Random, translational, rotational, vibrational | Rapid, random, translational (high kinetic energy) | | Density | High | Moderate (usually less than solid) | Very low | 2.4.2 Kinetic Theory of Matter The kinetic theory of matter (or kinetic molecular theory) explains the macroscopic properties of gases, liquids, and solids in terms of the motion of their constituent particles (atoms, molecules, or ions).
Postulates of the Kinetic Theory: Matter is composed of tiny particles: All matter is made up of very small particles (atoms, molecules, or ions) that are in constant, random motion.
Particles possess kinetic energy: These particles possess kinetic energy, and their motion is directly related to temperature. Higher temperature means higher average kinetic energy and faster particle movement.
Spaces between particles: There are spaces (intermolecular spaces) between these particles. The size of these spaces varies between states of matter.
Forces of attraction: There are attractive forces (intermolecular forces) between particles. The strength of these forces varies and influences the state of matter.
Collisions are elastic: Particles collide with each other and with the walls of their container. These collisions are perfectly elastic, meaning no net loss of kinetic energy during collisions. Volume of particles is negligible (for ideal gases): For gases, the actual volume occupied by the particles themselves is negligible compared to the total volume of the container. (This postulate is primarily for ideal gases, but the concept of particle size relative to space holds for all states). 2.4.3 Applications of Kinetic Theory to States of Matter: Solids: Particles are tightly packed in fixed positions due to very strong intermolecular forces. They only vibrate about their mean positions, explaining their definite shape and volume, high density, and inability to flow or be easily compressed.
Liquids: Particles are less tightly packed than in solids, with weaker (but still significant) intermolecular forces. They can slide past each other, allowing liquids to flow and take the shape of their container, yet maintain a definite volume. They are largely incompressible due to relatively small spaces between particles.
Gases: Particles are very far apart, with negligible intermolecular forces. They move rapidly and randomly, colliding frequently. This explains why gases expand to fill their container (indefinite shape and volume), are highly compressible (large spaces between particles), and have very low density.
Phase Changes: The kinetic theory also explains phase changes. Increasing temperature increases kinetic energy.
Melting: When a solid gains enough kinetic energy, its particles overcome the strong forces holding them in fixed positions and start to slide past each other, forming a liquid.
Boiling/Evaporation: When a liquid gains enough kinetic energy, its particles overcome the intermolecular forces entirely and escape into the gaseous state.
Condensation/Freezing: The reverse processes involve particles losing kinetic energy, leading to stronger intermolecular forces dominating and transitioning back to liquid or solid states.
Water Treatment and Purification in Nigeria: The process of water purification involves chemical combinations. For instance, coagulation uses chemicals like aluminium sulfate (Al2(SO4)3), an ionic compound, which combines with impurities to form larger particles that settle out. Chlorination involves elemental chlorine (Cl2, a covalent molecule) or chlorine compounds reacting with water and microbes. Understanding ionic and covalent bonds helps comprehend why these chemicals react the way they do to ensure clean drinking water for communities, especially relevant with water scarcity and sanitation challenges in many parts of Nigeria. Cement and Concrete Production for Infrastructure Development: Nigeria has significant infrastructure projects (roads, bridges, buildings). Cement is a key component of concrete, which forms through a complex series of chemical combinations. Limestone (calcium carbonate, CaCO3 - ionic compound) and clay are heated, leading to new ionic and covalent compounds like tricalcium silicate, dicalcium silicate, etc. These compounds react with water (a covalent molecule) in specific ways to form a hardened matrix. Understanding chemical bonding helps in optimizing cement formulations for stronger and more durable infrastructure, crucial for Nigeria's development goals.
Domestic Cooking Gas (LPG) and Kerosene: LPG (Liquefied Petroleum Gas), commonly used for cooking in Nigerian homes, consists mainly of propane (C3H8) and butane (C4H10). Kerosene is a mixture of longer-chain hydrocarbons. These are all covalent compounds. Their gaseous or liquid states at room temperature and ease of combustion (chemical combination with oxygen) are explained by their molecular structure, weak intermolecular forces, and the kinetic theory of matter. This knowledge is vital for understanding fuel efficiency, safe handling, and the chemistry behind combustion, impacting everyday household safety and energy usage.