Grade 10 Science Notes: Chemistry, Biology, Ecology, and Optics

Posted on Jun 26, 2024 in Chemistry

Grade 10 Chemistry

1. Atomic Structure

Atoms and Elements

• Atoms: Smallest unit of matter, consists of protons, neutrons, and electrons.
• Protons: Positively charged, found in the nucleus.
• Neutrons: Neutral, found in the nucleus.
• Electrons: Negatively charged, orbit the nucleus.
• Atomic Number (Z): Number of protons in an atom.
• Mass Number (A): Total number of protons and neutrons.
• Isotopes: Atoms of the same element with different numbers of neutrons.

Electron Configuration

• Electron Shells: Energy levels where electrons are found.
• Configuration: Distribution of electrons across different shells (e.g., 2-8-1 for sodium).
• Valence Electrons: Electrons in the outermost shell, determine chemical reactivity and bonding.

2. Periodic Table

Structure

• Groups: Vertical columns (1-18), elements in a group have similar properties.
• Periods: Horizontal rows, properties change progressively across a period.

Categories

• Metals: Good conductors of heat and electricity, malleable, ductile.
• Non-metals: Poor conductors, brittle when solid.
• Metalloids: Properties of both metals and non-metals.

Important Groups

• Group 1: Alkali metals (e.g., Lithium, Sodium), 1 valence electron.
• Group 2: Alkaline earth metals (e.g., Magnesium, Calcium), 2 valence electrons.
• Group 13: Boron group (e.g., Boron, Aluminum), 3 valence electrons.
• Group 14: Carbon group (e.g., Carbon, Silicon), 4 valence electrons.
• Group 15: Nitrogen group (e.g., Nitrogen, Phosphorus), 5 valence electrons.
• Group 16: Oxygen group (e.g., Oxygen, Sulfur), 6 valence electrons.
• Group 17: Halogens (e.g., Fluorine, Chlorine), 7 valence electrons.
• Group 18: Noble gases (e.g., Helium, Neon), 8 valence electrons (except Helium, which has 2).

3. Chemical Bonding

Ionic Bonds

• Formation: Transfer of electrons from a metal to a non-metal.
• Ions: Charged atoms (cations are positive, anions are negative).
• Properties: High melting and boiling points, conduct electricity when molten or dissolved.

Covalent Bonds

• Formation: Sharing of electron pairs between non-metals.
• Molecules: Groups of atoms bonded covalently (e.g., H₂O, CO₂).
• Properties: Low melting and boiling points, poor electrical conductivity.

Metallic Bonds

• Formation: Sea of shared free electrons among metal atoms.
• Properties: Good electrical and thermal conductivity, malleable, ductile.

4. Chemical Reactions

Types of Reactions

• Synthesis: A + B → AB
• Decomposition: AB → A + B
• Single Displacement: A + BC → AC + B
• Double Displacement: AB + CD → AD + CB
• Combustion: Hydrocarbon + O₂ → CO₂ + H₂O

Balancing Equations

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
• Steps:
  1. Write the unbalanced equation.
  2. Count atoms of each element.
  3. Use coefficients to balance atoms.
  4. Check the balance.

Acids and Bases

• Acids: Produce H⁺ ions in solution (e.g., HCl).
• Bases: Produce OH⁻ ions in solution (e.g., NaOH).
• pH Scale: Measures acidity/alkalinity (0-14 scale; <7 acidic, >7 basic, 7 neutral).

Neutralization

• Reaction: Acid + Base → Salt + Water
• Example: HCl + NaOH → NaCl + H₂O

5. Naming Compounds

Ionic Compounds

• Binary Ionic Compounds: Consist of two elements.
• Naming Rule: Metal name + non-metal name with “-ide” suffix.
• Example: NaCl = Sodium Chloride
• Transition Metals: Include the charge in Roman numerals.
• Example: FeCl₂ = Iron(II) Chloride

Covalent (Molecular) Compounds

• Binary Covalent Compounds: Consist of two non-metals.
• Naming Rule: Use prefixes to indicate the number of each element.
• Prefixes: 1 = mono-, 2 = di-, 3 = tri-, 4 = tetra-, 5 = penta-, 6 = hexa-, 7 = hepta-, 8 = octa-, 9 = nona-, 10 = deca-.
• Example: CO₂ = Carbon Dioxide, P₂O₅ = Diphosphorus Pentoxide

Acids

• Binary Acids: Contain hydrogen and one other non-metal.
• Naming Rule: “Hydro-” prefix + non-metal root + “-ic” suffix + “acid.”
• Example: HCl = Hydrochloric Acid
• Oxyacids: Contain hydrogen, oxygen, and another element.
• Naming Rule: Root of the polyatomic ion + “-ic” (if ion ends in “-ate”) or “-ous” (if ion ends in “-ite”) + “acid.”
• Example: H₂SO₄ (sulfate) = Sulfuric Acid, H₂SO₃ (sulfite) = Sulfurous Acid

Polyatomic Ions

• Common Polyatomic Ions:
  Ammonium: NH₄⁺
  Hydroxide: OH⁻
  Nitrate: NO₃⁻
  Sulfate: SO₄²⁻
  Carbonate: CO₃²⁻
  Phosphate: PO₄³⁻
• Naming Rule: Name of the metal + name of the polyatomic ion.
• Example: NaNO₃ = Sodium Nitrate, CaCO₃ = Calcium Carbonate.

Steps to Balance a Chemical Equation

1. Write the Unbalanced Equation
   Start with the reactants on the left side and the products on the right side.
   Example: C₃H₈ + O₂ → CO₂ + H₂O
2. Count the Number of Atoms of Each Element
   List the number of atoms of each element present in the reactants and products.
   Example:
   Reactants: C₃H₈ (3 C, 8 H) and O₂ (2 O).
   Products: CO₂ (1 C, 2 O) and H₂O (2 H, 1 O).
3. Balance One Element at a Time
   Start with an element that appears in only one reactant and one product. Adjust coefficients (numbers in front of compounds) to balance each element.
   Example:
   Balance Carbon (C): C₃H₈ + O₂ → 3CO₂ + H₂O (Now 3 C atoms on both sides)
   Balance Hydrogen (H): C₃H₈ + O₂ → 3CO₂ + 4H₂O (Now 8 H atoms on both sides)
   Balance Oxygen (O): C₃H₈ + 5O₂ → 3CO₂ + 4H₂O (Now 10 O atoms on both sides: 5 O₂ → 10 O, 3 CO₂ → 6 O, 4 H₂O → 4 O)

Grade 10 Ecology

1. Ecosystems

• Components: Biotic (living) and Abiotic (non-living)
• Types: Terrestrial (forests, deserts) and Aquatic (freshwater, marine)
• Energy Flow: Sun → Producers (plants) → Consumers (herbivores, carnivores) → Decomposers (fungi, bacteria)

2. Food Chains and Webs

• Food Chain: Linear sequence of organisms where nutrients and energy pass.
• Food Web: Complex network of interconnected food chains.

3. Trophic Levels

• Primary Producers: Plants, algae
• Primary Consumers: Herbivores
• Secondary Consumers: Carnivores that eat herbivores
• Tertiary Consumers: Carnivores that eat other carnivores

4. Biogeochemical Cycles

• Water Cycle: Evaporation, condensation, precipitation
• Carbon Cycle: Photosynthesis, respiration, decomposition, fossil fuels
• Nitrogen Cycle: Nitrogen fixation, nitrification, assimilation, ammonification, denitrification

5. Population Dynamics

• Carrying Capacity: Maximum population size an environment can sustain
• Limiting Factors: Resources, predation, disease, climate
• Growth Models: Exponential and logistic growth

6. Biodiversity

• Importance: Ecosystem stability, genetic variation, resource availability
• Threats: Habitat destruction, pollution, overexploitation, invasive species, climate change

Grade 10 Biology

1. Cell Biology

• Cell Theory: All living things are made of cells, the cell is the basic unit of life, all cells come from pre-existing cells
• Cell Structures:
  • Prokaryotic Cells: No nucleus, simple structure (bacteria)
  • Eukaryotic Cells: Nucleus, complex organelles (animals, plants)
• Organelles: Nucleus, mitochondria, ribosomes, endoplasmic reticulum, Golgi apparatus, lysosomes, chloroplasts (in plants), cell membrane, cell wall (in plants)

2. Genetics

• DNA Structure: Double helix, nucleotides (adenine, thymine, cytosine, guanine)
• Replication: DNA copying process
• Transcription and Translation: From DNA to mRNA to protein
• Mendelian Genetics: Dominant and recessive traits, Punnett squares, genotype, phenotype
• Genetic Variation: Mutation, recombination, sexual reproduction

3. Evolution

• Natural Selection: Survival of the fittest, adaptation, speciation
• Evidence of Evolution: Fossils, comparative anatomy, molecular biology, biogeography
• Mechanisms: Mutation, gene flow, genetic drift, selection

4. Human Biology

• Body Systems:
  • Circulatory: Heart, blood, blood vessels
  • Respiratory: Lungs, trachea, bronchi
  • Digestive: Stomach, intestines, liver, pancreas
  • Nervous: Brain, spinal cord, nerves
  • Skeletal: Bones, joints
  • Muscular: Muscles, tendons
  • Reproductive: Male and female reproductive organs

Grade 10 Optics

1. Properties of Light

• Nature: Electromagnetic wave, particle (photon)
• Speed: 3 x 10⁸ m/s in vacuum
• Behavior: Reflection, refraction, diffraction, interference

2. Reflection

• Laws of Reflection: Angle of incidence equals angle of reflection
• Types of Mirrors:
  • Plane Mirror: Upright, virtual image
  • Concave Mirror: Converging, real or virtual image
  • Convex Mirror: Diverging, virtual image

3. Refraction

• Snell’s Law: n₁ sin θ₁ = n₂ sin θ₂
• Index of Refraction: n = c/v (c: speed of light in vacuum, v: speed of light in medium)
• Lenses:
  • Convex (Converging) Lens: Focuses light, real or virtual image
  • Concave (Diverging) Lens: Spreads light, virtual image

4. Optical Instruments

• Microscope: Uses lenses to magnify small objects
• Telescope: Uses lenses or mirrors to observe distant objects
• Camera: Uses lens to focus light on a photosensitive surface

5. Human Eye

• Structure: Cornea, lens, retina, optic nerve
• Function: Light enters through cornea and lens, focused on retina, converted to neural signals, transmitted to brain

6. Wave-Particle Duality

• Concept: Light exhibits both wave-like and particle-like properties
• Experiments: Double-slit experiment (wave), photoelectric effect (particle)

7. Color

• Visible Spectrum: Red, orange, yellow, green, blue, indigo, violet (ROYGBIV)
• Color Addition: Primary colors (red, green, blue) combine to form white light
• Color Subtraction: Primary pigments (cyan, magenta, yellow) combine to absorb light and produce color