Chemistry and Physics Notes

Law of Conservation of Mass

• Definition: Matter cannot be created nor destroyed, only changed from one form to another.
• Equation: mass of reactants = mass of products
• Application:
  • Gases in open systems: When a gas is a reactant, the system's mass increases. When a gas is a product, the system's mass decreases as it escapes.
  • The law holds true when considering the total mass, including any gases that enter or leave the system.
• Challenges in Demonstrating in Open Systems:
  • Difficulty in precisely measuring or capturing gaseous reactants or products.
  • Variations in room conditions (temperature, pressure) can lead to inaccuracies.

Open vs. Closed Systems

• Closed System:
  • No substances can enter or escape (e.g., precipitation reaction).
  • No atoms are lost or gained.
  • Mass remains constant before, during, and after the reaction.
• Open System:
  • Substances can enter or leave the vessel (e.g., gaseous products escaping).
  • When a gas escapes, include the mass of the escaped gas in the calculation.
  • The mass of the reaction vessel decreases as gas escapes or increases as gaseous reactants are added.

Glossary of Terms

• Physical vs. Chemical Changes
• Reactants vs. Products
• Chemical Equation vs. Chemical Formula
• Molecule: Chemical made up of 2 or more atoms
• Atom: The smallest unit of an element that maintains the properties of the element.
• Compound: Two or more different types of atoms bonded together.

Evidence of Chemical Change

• Gas produced
• Odour produced
• Change in colour
• Change in temperature
• New substance produced
• A precipitate is formed
• Production of light or sound

Endothermic vs. Exothermic Reactions

Collision Theory

• For a reaction to occur:
  • Collision: Reactants must come into contact.
  • Orientation: Reactants must collide in the correct orientation.
  • Activation Energy: Reactants must collide with enough energy (minimum energy for a reaction to proceed).
    • More energy means particles can transfer more energy.
    • Higher frequency of collisions increases likelihood of successful collisions.

Conditions Impacting Reaction Speed

• Temperature:
  • Higher temperature → particles gain more energy → faster collisions → more frequent collisions.
• Concentration:
  • More particles per unit volume (or pressure for gas) → more frequent collisions.
• Surface Area:
  • Higher surface area → more chances of collision → higher collision frequency.
• Catalyst:
  • Speeds up reaction without being used up.
  • Lowers activation energy by providing alternative reaction pathways.
  • High activation energy slows down reaction.

Acids and Bases

• Common Acids:
  • Hydrochloric Acid (HCl)
  • Nitric Acid (HNO3)
  • Sulfuric Acid (H2SO4)
• Common Bases:
  • Sodium hydroxide (NaOH)
  • Calcium carbonate (CaCO3)
• pH Scale:
  • 1-7 is acidic
  • 7-14 is basic
  • Numbers further from 7 indicate stronger/more corrosive substances.
• Measuring pH:
  • Indicator: Chemical dyes (e.g., universal indicator).
  • pH probe: Indicates a numerical pH value.
• Acid Definition:
  • Any substance that forms an aqueous solution with a pH less than 7.
  • Hydrogen ions (H+) make it acidic, measured in molars (amount of hydrogen ions).
• Base Definition:
  • Any substance with a pH greater than 7.
  • Alkalis: Bases that dissolve in water to form a solution with a pH greater than 7.

Solubility

• Measure of a substance's ability to dissolve.
• Solute: The substance being dissolved.
• Solvent: The substance doing the dissolving.
• Aqueous Solution: A soluble salt dissolved in a solvent, like water.

Solubility Rules

• Soluble Compounds:
  • Compounds of sodium, potassium, and ammonium.
  • All nitrates.
  • All chlorides, except silver and lead (II).
  • All sulfates, except barium, calcium, and lead (II).
• Insoluble Compounds:
  • Sodium, potassium, and ammonium carbonates.
  • All other carbonates.
  • Sodium, potassium, and calcium hydroxides.
  • All other hydroxides.

Electrolysis of Water

• Decomposition of Water: $2H<em>2O → 2H</em>2 + O_2$
• Chemical reaction breaking down the compound.
• Collecting oxygen gas on one side and hydrogen gas on the other.
• Tests:
  • Hydrogen Test: 'Pop' test.
  • Oxygen Test: Glowing splint is relit when placed in oxygen.

Types of Reactions with Acids

1. Acid + Metal → Salt + Hydrogen (Single displacement)
   • Example: Hydrochloric acid (HCl) + Magnesium (Mg) → Magnesium chloride (MgCl2) + Hydrogen (H2)
2. Acid + Carbonate → Salt + Carbon Dioxide + Water
   • Example: Nitric acid (HNO3) + Calcium carbonate (CaCO3) → Calcium nitrate (Ca(NO3)2) + Carbon dioxide (CO2) + Water (H2O)
3. Acid + Base → Salt + Water (Neutralisation)
   • Example: Hydrochloric acid (HCl) + Sodium hydroxide (NaOH) → Sodium chloride (NaCl) + Water (H2O)

Precipitation Reactions

• Double displacement.
• Precipitate: The solid product of a chemical reaction when two solutions containing dissolved ions mix together.

Combustion

• Burning of fuel: A reaction in which a substance (usually a hydrocarbon) reacts with oxygen, releasing energy (light and heat).
  • One of the reactants must be oxygen.
  • Products are carbon dioxide and water vapour: fuel + O2 → CO2 + H2O
• Example: Candle burning
  • Heat melts wax (physical change).
  • Liquid wax moves up the wick and vaporizes into a gas.
  • Gas combusts with oxygen.
  • Products: Carbon dioxide, water, heat, light.
• Incomplete combustion (not enough oxygen): Carbon or carbon monoxide is formed.

Corrosion

• Rusting: The process by which metals slowly break down by reacting with substances in their environment.
• Iron (Fe) + Oxygen (O2) + Water (H2O) → Hydrated Iron Oxide (Rust)
• Conditions required:
  • Both metal and oxygen must be present.
  • Only the surface of the metal corrodes.
  • Sometimes a protective layer forms.
• How to prevent rust:
  • Barrier Methods: Prevent oxygen from touching the metal.
    • Painting
    • Oiling or greasing
    • Electroplating
  • Sacrificial Methods: Add a more reactive metal.

Displacement Reactions

• A + BC → AC + B
• More reactive metals displace less reactive ones.

Synthesis Reactions

• Multiple reactants form a single product.
• A + B → C
• Example: Hydrogen gas (H2) + Oxygen gas (O2) → Water (H2O)

Decomposition Reactions

• A single reactant breaks into two or more products.
• AB → A + B
• Example: Calcium carbonate (CaCO3) → Calcium oxide (CaO) + Carbon dioxide (CO2)

Perpetual Motion

Newton's Laws of Motion

1. Law of Inertia: The property of something to resist a change in motion.
2. $F = ma$ (Force = mass x acceleration)
3. For every action, there is an equal but opposite reaction.

Net Force

• The resulting force when all forces acting on an object are considered.
• If an object is still, the net force is zero.

Speed, Distance, and Time

• $Speed = \frac{Distance}{Time}$

Scalar vs. Vector

• Scalar: Only numerical (quantitative) e.g speed.
• Vector: Numerical and directional (qualitative and quantitative) e.g velocity.

Distance and Displacement

Instantaneous vs. Average Speed

• Instantaneous Speed: The speed of an object at any instant.
• Average Speed: The mean rate of motion of an object.

Speed and Velocity

Acceleration

• The rate that velocity changes.
• $Acceleration = \frac{(Final Speed - Initial Speed)}{Time}$
  • If an object is changing speed or velocity, it must be accelerating.
  • A change in direction or speed is a result of unbalanced forces.

Motion Graphs

• Constant Speed
• Accelerating or Decelerating
• Displacement-Time Graph: Can go below the x-axis because direction is assigned (positive as north, negative as south).

Balanced and Unbalanced Forces

Free Body Diagrams