Lesson Notes By Weeks and Term v3 - Senior Secondary 1
Particulate Nature of Matter
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Particulate Nature of Matter
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Subject: Chemistry
Class: Senior Secondary 1
Term: 3rd Term
Week: 1
Theme: The Chemical World
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Distinguishbetween physical and chemical changes; distinguishbetween at oms and molecules; describe how the particles are arrangedin the at om; Define:- At omic number- Mass number- Is otopes; Calculate the relative at omicmasses of at oms.
This section provides a detailed explanation of the core concepts related to the particulate nature of matter, building step-by-step towards the learning objectives. 2.
1. Physical and Chemical Changes Matter can undergo various transformations. These changes are broadly categorized into physical and chemical changes.
Physical Change: This is a change in the form or state of matter, but not its chemical composition. No new substance is formed, and the change is generally reversible.
Characteristics: No new substance is produced. The chemical composition remains the same. Often reversible by simple physical means. Usually, small energy changes are involved.
Examples: Melting ice, boiling water, dissolving salt in water, crushing a piece of chalk, cutting wood, grinding maize or pepper. Nigerian Context
Examples: Drying of clothes, melting of ice cream, dissolving granulated sugar in tea, pounding yam, distillation of crude oil components (e.g., kerosene from petroleum).
Chemical Change (Chemical Reaction): This is a change that results in the formation of one or more new substances with entirely different chemical properties from the original substances. This change is typically irreversible.
Characteristics: New substances with new properties are formed. The chemical composition of the substances changes. Usually irreversible or difficult to reverse. Significant energy changes (absorption or release of heat, light, electricity) are involved. Often accompanied by observable signs like gas evolution (bubbles), precipitation (formation of solid), color change, or heat change.
Examples: Burning wood, rusting of iron, cooking an egg, souring of milk, digestion of food. Nigerian Context
Examples: Burning of firewood or kerosene, fermentation of palm wine or cassava for 'garri', rusting of corrugated iron roofing sheets or vehicle bodies, cooking of yam or rice, burning of bush during dry season.
Table 1: Distinction between Physical and Chemical Changes | Feature | Physical Change | Chemical Change | | :------------------- | :---------------------------------------------- | :---------------------------------------------------- | | New Substance | No new substance formed | New substance(s) with different properties formed | | Composition | Chemical composition remains the same | Chemical composition changes | | Reversibility | Generally reversible | Generally irreversible | | Energy Change | Small energy change | Significant energy change (heat, light, sound) | | Nature of Change | Change in state, shape, size | Change in chemical identity | | Examples (Local) | Melting ice block, grinding pepper, dissolving sugar | Burning firewood, rusting of zinc, cooking beans | 2.
2. Atoms and Molecules Atom: The smallest particle of an element that can take part in a chemical reaction. It retains the chemical identity of the element. Historically considered indivisible (Dalton's theory). Composed of subatomic particles (protons, neutrons, and electrons). Atoms are the fundamental building blocks of all matter.
Examples: A single atom of Oxygen (O), Hydrogen (H), Carbon (C).
Molecule: A group of two or more atoms chemically bonded together. It is the smallest unit of a pure substance (element or compound) that can exist independently and still retain the chemical properties of that substance.
Molecules of elements: Composed of two or more identical atoms.
Examples: Oxygen gas (O2), Nitrogen gas (N2), Chlorine gas (Cl2).
Molecules of compounds: Composed of two or more different types of atoms.
Examples: Water (H2O), Carbon dioxide (CO2), Ammonia (NH3).
Table 2: Distinction between Atoms and Molecules | Feature | Atom | Molecule | | :---------------- | :---------------------------------------------- | :------------------------------------------------------------- | | Composition | Smallest unit of an element | Two or more atoms chemically bonded together | | Independence | May or may not exist independently in nature | Can exist independently, retaining properties of substance | | Types | Only one type of particle (elemental) | Can be elemental (e.g., O2) or compound (e.g., H2O) | | Example (Local) | A single carbon atom in charcoal, a single gold atom in a gold ornament | An oxygen molecule (O2) in the air, a water molecule (H2O) in rain | 2.
3. The Structure of the Atom Atoms are not indivisible but consist of even smaller subatomic particles.
Nucleus: Located at the center of the atom, it contains protons and | Can exist independently, retaining properties of substance | | Types | Only one type of particle (elemental) | Can be elemental (e.g., O2) or compound (e.g., H2O) | | Example (Local) | A single carbon atom in charcoal, a single gold atom in a gold ornament | An oxygen molecule (O2) in the air, a water molecule (H2O) in rain | 2.
3. The Structure of the Atom Atoms are not indivisible but consist of even smaller subatomic particles.
Nucleus: Located at the center of the atom, it contains protons and neutrons. The nucleus is positively charged due to the presence of protons and accounts for almost all of the atom's mass. Protons (p+): Positively charged particles (+1 elementary charge). Each proton has a mass approximately equal to 1 atomic mass unit (amu). They define the atomic number and thus the identity of the element.
Neutrons (n0): Electrically neutral particles (no charge). Each neutron has a mass approximately equal to 1 amu, slightly more than a proton. Neutrons contribute to the mass number but not the charge of the nucleus.
Electrons (e-): Negatively charged particles (-1 elementary charge). They orbit the nucleus in specific energy levels or shells. Electrons have a negligible mass (approximately 1/1836th of a proton's mass). In a neutral atom, the number of electrons equals the number of protons.
Diagrammatic Representation of an Atom: (Teacher to draw a simple diagram on the board, illustrating a central nucleus with protons and neutrons, and electrons orbiting in shells.)
Example: A Lithium atom (Li) with 3 protons, 4 neutrons, and 3 electrons.
Nucleus: 3 protons (positive charge), 4 neutrons (no charge). Overall positive charge of +
3. Electron Shells: 2 electrons in the first shell, 1 electron in the second shell. Overall negative charge of -
3. Net charge of the atom: 0 (neutral). 2.
4. Definitions: Atomic Number, Mass Number, Isotopes Atomic Number (Z): The number of protons in the nucleus of an atom. It is a unique identifier for each element. All atoms of a given element have the same atomic number. In a neutral atom, the number of electrons is equal to the number of protons.
Example: For Carbon (C), Z = 6, meaning every carbon atom has 6 protons.
Mass Number (A): The total number of protons and neutrons (also called nucleons) in the nucleus of an atom. Mass Number (A) = Number of Protons (Z) + Number of Neutrons (N)
Example: A carbon atom with 6 protons and 6 neutrons has a mass number of 12 (6+6=12).
Isotopes: Atoms of the same element (i.e., having the same atomic number, Z, and thus the same number of protons) but with different mass numbers (A) due to different numbers of neutrons. Isotopes have identical chemical properties because their electron configurations are the same, but they have different physical properties (e.g., mass, density).
Notation: Isotopes are represented as $\text{_Z^A X}$ or X-
A. Examples: Hydrogen has three common isotopes: Protium ($\text{_1^1 H}$): 1 proton, 0 neutrons (Mass Number = 1) Deuterium ($\text{_1^2 H}$): 1 proton, 1 neutron (Mass Number = 2) Tritium ($\text{_1^3 H}$): 1 proton, 2 neutrons (Mass Number = 3) Carbon-12 ($\text{_6^{12} C}$), Carbon-13 ($\text{_6^{13} C}$), Carbon-14 ($\text{_6^{14} C}$) are isotopes of carbon. All have 6 protons but differ in neutrons (6, 7, and 8 respectively). Carbon-14 is used in radiometric dating for archaeological findings in places like Nok and Igbo-Ukwu. 2.
5. Calculation of Relative Atomic Mass (Ar) The atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom.
Relative Atomic Mass (Ar): Since most elements are mixtures of isotopes, the relative atomic mass is the weighted average mass of all naturally occurring isotopes of an element, relative to 1/12th the mass of a carbon-12 atom. It takes into account both the mass of each isotope and its natural abundance (percentage presence in nature). * Formula: Ar = $\Sigma$ (isotopic mass $\times$ fractional abundance) Where fractional abundance = percentage abundance / 100 Worked Example 1: Chlorine (Cl) has two main isotopes: Chlorine-35 with an mass of a carbon-12 atom.
Relative Atomic Mass (Ar): Since most elements are mixtures of isotopes, the relative atomic mass is the weighted average mass of all naturally occurring isotopes of an element, relative to 1/12th the mass of a carbon-12 atom. It takes into account both the mass of each isotope and its natural abundance (percentage presence in nature).
Formula: Ar = $\Sigma$ (isotopic mass $\times$ fractional abundance) Where fractional abundance = percentage abundance / 100 Worked Example 1: Chlorine (Cl) has two main isotopes: Chlorine-35 with an abundance of 75.77% and Chlorine-37 with an abundance of 24.23%. Calculate the relative atomic mass of chlorine.
Solution:
1. Convert percentage abundances to fractional abundances: Chlorine-35: 75.77% = 0.7577 Chlorine-37: 24.23% = 0.2423
2. Apply the formula: Ar = (Mass of Isotope 1 $\times$ Fractional Abundance 1) + (Mass of Isotope 2 $\times$ Fractional Abundance 2) Ar = (35 $\times$ 0.7577) + (37 $\times$ 0.2423) Ar = 26.5195 + 8.9651 Ar = 35.4846 Therefore, the relative atomic mass of Chlorine is approximately 35.
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8. Worked Example 2: Element Boron (B) has two isotopes: Boron-10 (10B) with 19.9% abundance and Boron-11 (11B) with 80.1% abundance. Calculate its relative atomic mass.
Solution:
1. Fractional abundances: Boron-10: 19.9% = 0.199 Boron-11: 80.1% = 0.801
2. Calculate Ar: Ar = (10 $\times$ 0.199) + (11 $\times$ 0.801) Ar = 1.99 + 8.811 Ar = 10.801 Therefore, the relative atomic mass of Boron is approximately 10.80. 2.
6. Dalton's Atomic Theory (Foundational Concept - for evaluation) John Dalton proposed his atomic theory in the early 19th century, which laid the foundation for modern chemistry. Though some postulates have been refined, its core ideas remain fundamental.
Postulates:
1. All matter is made up of tiny, indivisible particles called atoms.
2. Atoms of a given element are identical in mass and properties; atoms of different elements are different in mass and properties.
3. Atoms cannot be created or destroyed, nor can they be subdivided into smaller particles.
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5. Chemical reactions involve the rearrangement of atoms. 2.
7. Relative Molecular Mass and Chemical Formulae (for evaluation)
Relative Molecular Mass (Mr): The sum of the relative atomic masses of all the atoms present in one molecule of a compound. It is also unitless, expressed relative to 1/12th the mass of a carbon-12 atom.
Formula: Mr = $\Sigma$ (Ar of element $\times$ number of atoms of that element)
Worked Example 3: Calculate the relative molecular mass of water (H2O), given Ar: H=1, O=
1
6. Solution: Mr (H2O) = (2 $\times$ Ar of H) + (1 $\times$ Ar of O) Mr (H2O) = (2 $\times$ 1) + (1 $\times$ 16) Mr (H2O) = 2 + 16 Mr (H2O) = 18 Empirical Formula: The simplest whole-number ratio of atoms in a compound.
Molecular Formula: The actual number of atoms of each element in a molecule of the compound. * Molecular Formula = (Empirical Formula)n, where n = Molecular Mass / Empirical Formula Mass.
Worked Example 4: A compound has an empirical formula of CH2 and a molecular mass of 42 g/mol. Determine its molecular formula. (Given Ar: C=12, H=1).
Solution:
1. Calculate the empirical formula mass (EFM) for CH2: EFM = (1 $\times$ Ar of C) + (2 $\times$ Ar of H) EFM = (1 $\times$ 12) + (2 $\times$ 1) = 12 + 2 = 14 g/mol
2. Calculate 'n': n = Molecular Mass / EFM n = 42 / 14 = 3
3. Determine the molecular formula: Molecular Formula = (CH2)n = (CH2)3 = C3H6 --- 3.
1. Teacher Activities: Introduction (10 minutes): Begin by displaying various common materials (e.g., water, sugar, sand, iron nail, burning candle, cooking oil). Ask students to observe and describe them. Initiate a discussion on what happens when these materials undergo changes (e.g., water boiling, sugar dissolving, nail rusting, candle burning). Introduce the concept of matter being made of particles and the types of changes they undergo. Physical vs.
Chemical Changes (15 minutes): Demonstration: Conduct simple demonstrations: Melting ice (physical). Dissolving sugar/salt in water (physical). Crushing chalk (physical). Burning a small piece of paper/wood/candle (chemical - observe soot, ash, heat). Mixing iron filings with sulfur and heating (chemical - new substance formed, difficult to separate).
Local Relevance: Discuss examples like ripening of fruits (chemical), grinding pepper (physical), cooking yam (chemical), fermentation of 'ogi' (chemical).
Explanation: Clearly define and differentiate between physical and chemical changes using a tabular comparison on the board.
Atoms and Molecules (10 minutes): Visual Aids: Use models (e.g., ball-and-stick models for H2, O2, H2O, CO2) or clear diagrams to illustrate atoms as individual units and molecules as combinations of atoms.
Explanation: Define atoms and molecules, emphasizing their roles as building blocks and stable units, respectively.
Atomic Structure (15 minutes): Diagrams: Draw a clear, labeled diagram of a simple atom (e.g., Helium, Lithium, or Carbon) on the board, showing the nucleus (protons, neutrons) and electron shells/orbits.
Explanation: Describe the properties (charge, mass, location) of protons, neutrons, and electrons. Explain how they are arranged within the atom. Use the analogy of a small central sun (nucleus) with planets (electrons) orbiting it to aid understanding. Atomic Number, Mass Number, Isotopes (15 minutes): Definitions: Provide precise definitions for atomic number (Z), mass number (A), and isotopes.
Examples: Use specific examples like Hydrogen isotopes (Protium, Deuterium, Tritium) and Carbon isotopes (C-12, C-13, C-14) to illustrate these concepts. Explain the notation $\text{_Z^A X}$.
Calculations: Guide students on how to determine the number of protons, neutrons, and electrons given atomic and mass numbers for neutral atoms. Relative Atomic Mass Calculations (15 minutes): Explanation: Explain the concept of relative atomic mass as a weighted average due to isotopes. Worked
Examples: Work through the provided examples (Chlorine, Boron) step-by-step on the board, ensuring students understand each stage of the calculation. Dalton's Atomic Theory & Relative Molecular Mass/Formulae (10 minutes): Explanation: Briefly state Dalton's Atomic Theory as a historical foundation. Worked
Examples: Explain relative molecular mass and walk through calculation for H2O and empirical/molecular formula example for C3H
6. Wrap-up & Q&A (5 minutes): Summarize key concepts and address any student questions. 3.
2. Student Activities: Observation & Recording: Students observe the teacher's demonstrations of physical and chemical changes, record their observations, and identify the type of change.
Discussion: Participate in class discussions on the observed changes and contribute local examples.
Note-taking: Copy definitions, diagrams, and worked examples into their notebooks.
Drawing: Draw and label the structure of a simple atom.
Calculations: Practice calculating the number of protons, neutrons, and electrons for given elements. Work collaboratively on relative atomic mass calculations.
Questioning: Ask clarifying questions during explanations.
Group Work: In small groups, discuss the differences between atoms and molecules, or create additional examples of physical/chemical changes in their environment. ---
Rusting of Iron and Corrosion Control: Concept: Rusting is a chemical change (oxidation) of iron.
Application: Students observe rusting on common items like zinc roofing sheets, vehicle bodies, iron gates, and bridges across Nigeria. Understanding that rusting is a chemical change helps in appreciating methods of prevention like painting (e.g., painting metal fences around homes), galvanizing (e.g., zinc-coated buckets), or oiling (e.g., lubricating tools), which protect the iron from reacting with oxygen and water. This directly impacts the durability of infrastructure and personal assets.
Food Preservation Techniques: Concept: Many food spoilage processes involve chemical changes (e.g., bacterial decomposition, enzymatic browning). Physical changes are also used in preservation.
Application: Drying of 'garri' or 'kulikuli' (physical change - removal of water), salting/smoking of fish or meat (chemical changes altering food chemistry and inhibiting microbial growth), refrigeration (physical change - slowing down chemical reactions). Understanding these concepts helps in making informed decisions about food storage, safety, and reducing food waste in homes and local markets.
Petroleum Refining and Products: Concept: Crude oil is separated into various fractions through physical changes (distillation), and these fractions can be further processed through chemical changes (cracking, reforming) to produce more valuable products.
Application: Students can relate the separation of crude oil into petrol (PMS), kerosene (widely used for cooking), diesel, and lubricants at refineries (e.g., Port Harcourt, Warri) to physical changes. The further conversion of heavier fractions into lighter, more valuable ones (e.g., gasoline) involves chemical changes. This connects to Nigeria's economy, energy supply, and the importance of chemical industries. ---