CHEMISTRY 1

SEMEPXDOST 2020 Chemistry Presentation Notes

Topics Covered

  • Accuracy and Precision (Page 3)
  • Unit Conversion (Page 8)
  • Laws of Chemical Combination (Page 13)
  • Matter (Page 30)
  • Intermolecular Forces of Attraction (Page 35)
  • Colligative Properties (Page 44)
  • Gas Laws (Page 52)
  • Periodic Trends (Page 59)
  • Solutions and Solubility (Page 66)
  • Chemical Reactions (Page 74)

TOPIC 1: Accuracy and Precision

  • Accuracy: Refers to how closely individual measurements agree with the correct, or "true", value. (Page 4)
  • Precision: A measure of how closely individual measurements agree with one another. (Page 5)
Sample Question
  • Paul conducted an experiment with 5 trials and gathered experimental values of 2.59, 3.03, 3.21, 3.12, and 2.85. If the theoretical value is 6.85, which case describes these values?
    • Answer: A. High precision, low accuracy (Page 7)

TOPIC 2: Unit Conversion

  • Common conversion units and differences between metric and imperial systems must be understood.
  • Various temperature conversion formulas (Page 8).
Common Conversions
  1. Mass & Weight
    • 1 Kg = 1000 g
    • 1 pound = 453.59 g
    • 1 amu = 1.6606 × 10^-24 g
    • 1 ton = 2000 pounds
  2. Volume
    • 1 liter = 1000 cm³ = 1000 mL
    • 1 m³ = 1000 liters
  3. Length
    • 1 inch = 2.54 cm
    • 1 mile = 1.6093 km
  4. Temperature Conversion Formulas:
    • Fahrenheit (°F) = 95\frac{9}{5} (°C) + 32
    • Celsius (°C) = 59\frac{5}{9} (°F - 32)
    • Kelvin (K) = °C + 273.15
    • Rankine (°R) = °F + 459.67 (Page 12)
Metric vs. Imperial System
  • Metric System: Includes meter for length, kilogram for mass, and second for time, known as the Système International (SI). (Page 10)
  • Imperial System: Used in the UK and other Commonwealth countries, including units like inch, mile, and pound. (Page 11)

TOPIC 3: Laws of Chemical Combinations

  • Law of Conservation of Mass: Antoine Lavoisier states that the mass of products in a chemical reaction equals the mass of reactants.

    • Mass can neither be created nor destroyed. (Pages 15-16)
    • Example:
    • $HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)$
    • $H₂ + O₂ → H₂O$ (Page 17)

  • Law of Constant Composition: Joseph Proust states that a chemical compound contains the same elements combined in the same proportion by mass regardless of its source. (Page 19)

    • Example: Pure water always has a mass ratio of hydrogen to oxygen as 1:8. (Page 20)

  • Law of Multiple Proportions: John Dalton explains that when two elements combine to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ratio of small whole numbers. (Page 24)

    • Example: 1.333 g O with CO and 2.666 g O with CO₂ gives a ratio of 1:2. (Page 25)

  • Law of Reciprocal Proportion: Jeremias Richter states that if two different elements combine with the same weight of a third element, their mass ratios are either the same or simple multiples of those ratios. (Page 26)

  • Gay Lussac’s Law of Gaseous Volume: When gases react to form other gases, volumes can be expressed in simple whole numbers under the same conditions of temperature and pressure. (Page 28)

    • Example: 2 H₂(g) + O₂(g) → 2 H₂O. (Page 29)

TOPIC 4: Matter

  • Definition: Matter is anything that occupies space and has mass (Page 31).
  • Classifications of Matter: Based on properties and phases (solid, liquid, gas). (Page 30)
Phases of Matter and Phase Change
  • Gas
  • Liquid
  • Solid
  • Phase changes include:
    • Melting (solid to liquid)
    • Freezing (liquid to solid)
    • Vaporization (liquid to gas)
    • Condensation (gas to liquid)
    • Sublimation (solid to gas)
    • Deposition (gas to solid) (Page 34)

TOPIC 5: Intermolecular Forces of Attraction (IMFA)

  • Types of Forces:
    1. Intramolecular Forces (within molecules)
    2. Intermolecular Forces (between molecules) (Page 35)
Intramolecular Forces
  1. Ionic Bond
  2. Covalent Bond (polar and nonpolar)
  3. Metallic Bond (Page 38)
Intermolecular Forces
  1. Hydrogen Bonds: Forces between a hydrogen atom and electronegative atoms such as F, O, N.
  2. Dipole-Dipole Interactions: Attractive forces between polar molecules.
  3. London Dispersion Forces (LDF): Forces acting between nonpolar substances due to temporary dipoles. (Page 39-41)
Effects of IMFA
  • High IMFA results in:
    • High boiling point
    • High melting point
    • High surface tension
    • High latent heat of fusion, vaporization, sublimation. (Page 42)

TOPIC 6: Colligative Properties

  • Definition: Properties that depend on the quantity of solute particles rather than their identity.
  • Key Colligative Properties Include:
    • Vapor Pressure Lowering: Nonvolatile solute decreases solvent vapor pressure, expressed by Raoult's Law.
    • P<em>soln=(X</em>solvent)(Psolvent)P<em>{soln} = (X</em>{solvent})(P_{solvent}) (Page 48-49)
    • Boiling Point Elevation: Nonvolatile solute raises the boiling point of the solvent.
    • Freezing Point Depression: Solute decreases the freezing point of the solvent. (Page 50-51)

TOPIC 7: Gas Laws

  • Boyle's Law: Volume of a gas is inversely proportional to pressure at constant temperature ( ext{PV = k}). (Page 53)
  • Charles' Law: Volume of a gas is directly proportional to Kelvin temperature at constant pressure ( ext{V/T = k}). (Page 54)
  • Gay-Lussac's Law: Pressure of a gas is directly proportional to Kelvin temperature at constant volume ( ext{P/T = k}). (Page 55)
  • Avogadro's Law: Equal volumes of gases contain equal numbers of molecules at the same conditions. (Page 56)
  • Combined Gas Law: Relates pressure, volume, and temperature of a fixed quantity of gas. (Page 57)
  • Ideal Gas Law: PV=nRTPV = nRT relates pressure, volume, number of moles, and temperature. (Page 58)

TOPIC 8: Periodic Trends

  • Atomic Size and Radius: Decreases left to right due to increased nuclear charge; increases down a group due to electron shielding. (Page 60)
  • Ionization Energy: Energy required to remove an electron. Increases left to right, decreases down a group. (Page 63)
  • Electronegativity: Tendency of an atom to attract electrons. Increases left to right, decreases down a group. (Page 64)
  • Metallic Character: Decreases left to right, increases down a group (Page 62).
Chart of Periodic Trends
  • Ionization Energy: Trends displaying energy needed to remove electrons.
  • Electronegativity: Reflects the tendency of an element to attract electrons in a bond.
  • Metallic Character & Nonmetallic Character: Observed across the periodic table.
  • Atomic Radius: Relationship with chemical properties (Page 65).

TOPIC 9: Solutions and Solubility

  • Types of Solutions: Classified by state of solute and solvent, concentration, conductivity, and saturation degree. (Page 67)
  • Definition of Solubility: The ability of a substance (solute) to dissolve in a solvent (Page 72).
  • Factors Affecting Solubility:
    1. Pressure: Affects gases; solubility increases with partial pressure.
    2. Temperature: Solubility of solids typically increases with temperature; gas solubility decreases with temperature (Page 73).

TOPIC 10: Chemical Reactions

  • Chemical Reaction: Processes through which reactants transform into products. (Page 75)
  • Balancing Chemical Equations:
    • Number of atoms must be equal on both sides.
    • Example: 3H<em>2+N</em>2→2NH33 H<em>2 + N</em>2 → 2 NH_3 (Page 77)
Physical Changes During Reactions
  • Evidence of Reaction: Includes gas evolution, color change, precipitate formation, scent change, temperature change, and volume change (Page 78).
Types of Chemical Reactions
  • Combination: Two reactants form one product.
    Example: 2H<em>2+O</em>2→2H2O2H<em>2 + O</em>2 → 2H_2O (Page 80)
  • Decomposition: A single reactant breaks into multiple products.
    Example: H<em>2O</em>2→H<em>2+O</em>2H<em>2O</em>2 → H<em>2 + O</em>2 (Page 81)
  • Single Replacement: One element replaces another in a compound. (Page 82)
  • Double Replacement (Precipitation): A reaction between two ionic compounds that leads to the formation of an insoluble compound (Page 83).
Energy Changes in Reactions
  • Endothermic Reactions: Absorb heat from surroundings. Feels cold - \Delta H > 0 (Page 84).
  • Exothermic Reactions: Release heat to surroundings. Feels warm - \Delta H < 0 (Page 85).

References

  • Bhatia, L. (2018)
  • Brown, T.L., LeMay, H.E., et al. (2018) Chemistry: The Central Science.
  • Chemistry LibreTexts (2020).
  • Gas Laws. (2018)
  • Helmenstine, A.M. (2018).
  • Khan Academy (2020).
  • LibreTexts (2020).
  • Metric vs. Imperial System (2020).
  • Zumdahl, S. S. (2009).