Chem video 1

Classroom Policies

• Computers are allowed for note-taking or viewing the PowerPoint, but refrain from activities that may distract others, such as video games or online shopping. If needed, sit in the back.
• Office hours are available immediately after class in the Bush building; lab materials will be available for purchase.
• Not everything on the slides needs to be written down, especially if it's common sense.

Chapter 1: Introduction to Science

• Science begins with observation, leading to curiosity.
• Science focuses on the physical world, examining phenomena through the five senses.
• Scientific thought evolves as new information emerges. Rigidity hinders progress; adaptability is essential.
• Avoid definitive statements like "I proved it"; maintain flexibility in acknowledging potential errors based on new evidence.
• Key terms:
  • Hypothesis: A tentative interpretation or explanation. It's subject to change and aims to explain, not merely describe.
  • Law: A description, not an explanation. Example: the law of gravity describes the phenomenon of objects falling but doesn't explain why they fall.
  • Theory: A well-supported explanation, less tentative than a hypothesis. It's as close to the ultimate truth and has withstood repeated testing. While hypotheses may change with new evidence, theories are more robust.
• Historical Example
  • Lavoisier's experiment challenged the phlogiston theory, which posited a fire-like element released during burning.
  • Lavoisier conducted experiments in closed containers, measuring mass before and after burning. He observed no mass change, contradicting the phlogiston theory.
  • This observation aligns with the law of conservation of mass: in a chemical reaction, matter is neither created nor destroyed.
  • The law of conservation of mass is a description, not an explanation.
• The law of conservation of mass example
  • When combining 32 grams of sulfur and 56 grams of iron, the product's mass should equal the total mass of the reactants (88 grams). Gas (if given off) must also be accounted for.
• Theory VS Laws
  • A hypothesis, when repeatedly supported, can evolve into a scientific theory, providing explanations for observed laws.
  • Darwin's theory of natural selection, which arose independently in two locations, exemplifies a well-tested explanation that started as a hypothesis.
  • Theories explain the causes behind observations and laws, enabling predictions in different scenarios.
• John Dalton's Atomic Theory
  • Matter comprises small, indestructible particles called atoms. This explains why matter is conserved during chemical reactions.
  • Dalton revived Democritus's earlier concept but supported it with scientific data.
  • Theories are not mere speculations but are supported by substantial evidence.
  • Modern technology allows direct visualization of atoms, such as iron atoms on a copper surface.
• Experiments
  • Unsupported hypotheses require revision, emphasizing the importance of well-designed and controlled experiments.
  • A controlled experiment includes a control group (normal conditions) for comparison. Human studies are challenging due to variability among subjects.
  • Ideally, an experiment should manipulate only one independent variable to isolate its effect on the dependent variable.
  • The independent variable is graphed on the x-axis, while the dependent variable is on the y-axis.

Chemistry and Matter

• Chemistry studies matter, biology studies life, and physics studies forces and motion. Integration among these sciences is pervasive.
• Chemicals are ubiquitous, present in everything around us. Soda contains sugar molecules, carbon dioxide, and water. Pencil "lead" is actually graphite (carbon). The human body comprises proteins, DNA, and carbohydrates.
• Basic Definitions
  • Atom: The simplest unit of chemistry, composed of protons, neutrons, and electrons.
    • Proton: P^+; positive charge.
    • Neutron: N^0; neutral charge.
    • Electron: e^-; negative charge.
  • Element: A specific type of atom (e.g., helium, fluorine, oxygen), distinguished by its number of protons.
  • Molecules or Formula Units: Atoms linked by chemical bonds.
    • Molecules: Non-metal atoms bonded together (e.g., carbon dioxide CO2, nitrogen gas N2). They are discrete units.
    • Formula Units: Metal and non-metal atoms combined (e.g., iron oxide Fe2O3). It represents a ratio within a crystal lattice.
  • Compound: A substance formed when two or more different atoms are chemically bonded. Can be a molecule or formula unit, different atoms combined.
• Soda Composition
  • A single drop contains approximately 1 \times 10^{21} atoms.
  • Soda is a mixture of molecules: carbon dioxide (gas), water (liquid), and sugar (sweetness).
  • Carbon structure
    • Carbon dioxide is linear, with the formula O=C=O. Each carbon-oxygen bond is polar because oxygen attracts electrons more strongly than carbon.
    • Despite the polar bonds, carbon dioxide is nonpolar because the two dipoles cancel each other out.
    • The lack of net polarity results in weak intermolecular forces, causing carbon dioxide to exist as a gas at room temperature.
  • Water structure
    • Water is bent, not linear.
    • The bent shape creates a dipole moment, with oxygen having a partial negative charge and hydrogen a partial positive charge. \delta^+ and \delta^-.
    • The polarity of water makes it a good solvent, allowing water molecules to stick together.
• Carbon Dioxide in soda
  • Carbon dioxide is forced into the solution under pressure. When the pressure is released (opening the can), carbon dioxide escapes, forming bubbles.

Safety Considerations

• Do not mix household chemicals:
  • Bleach and vinegar produce chlorine gas.
  • Ammonia and bleach can create a potentially explosive mixture.
  • Rubbing alcohol and bleach can form chloroform.
  • Hydrogen peroxide and vinegar can create peracetic acid, an irritant.
• Avoid mixing different types of batteries due to varying discharge rates, which can cause corrosion and damage equipment.
• Avoid mixing Red Bull and milk.
• Be cautious with grapefruit juice and certain medications as they it can increase absorption of prescribed medicine make drugs more lethal.
• Acetaminophen and alcohol can cause liver damage, and ibuprofen/aspirin with alcohol is harmful to the stomach lining.

Environmental concerns

• Chemistry is a doubled edge-sword.
  • CFCs caused the Ozone depletion.
  • Global warming is a great problem.
  • Plastic wastes.

Chapter 2: Measurements in Science

• Types of Measurements
  • Qualitative: Descriptive, subjective, and opinion-based.
  • Quantitative: Involves numbers (e.g., mass, length, volume, temperature).
• Uncertainty
  • All measurements have uncertainty, which needs to be accounted for. The question is how to take the measurement.
• Key Terms
  • Precision: Refers to grouping. It describes how close a set of measurements are to one another, regardless of whether they are correct.
  • Accuracy: Refers to the bullseye/true value. It describes how close a measurement is to the accepted or true value.
• Targets
  • Neither accurate nor precise: scattered and far from the bullseye.
  • Precise but not accurate: grouped together but not near the bullseye.
  • Accurate but not precise: scattered around the bullseye, but the average position is the center.
  • Both accurate and precise: grouped tightly around the bullseye.
• Equation for % Error
  • The difference between the measured (experimental) value and the true (accepted) value divided by the true value, expressed as a percentage.
  • \text{Percent Error} = \frac{|\text{Experimental Value} - \text{Accepted Value}|}{\text{Accepted Value}} \times 100\%
• Examples
  • Mass measurement example
    • Experimental value: 17.7 grams
    • Accepted value: 21.2 grams
      \text{Percent Error} = \frac{|17.7 - 21.2|}{21.2} \times 100\% = 16.5 \%
  • Volume measurement example
    • Experimental value: 4.26 mL
    • Accepted value: 4.15 mL
      \text{Percent Error} = \frac{|4.26 - 4.15|}{4.15} \times 100\% = 2.65 \%

Scientific Notation

• Why use it?
  • It simplifies writing very large or very small numbers.
• Parts of Scientific Notation
  • Coefficient: The numerical part (e.g., 6.022 in 6.022 \times 10^{23}).
  • Base: Always 10 in scientific notation.
  • Exponent: Indicates the power of 10 (e.g., 23 in 6.022 \times 10^{23}).
• Positive Exponents
  • Numbers larger than one; the decimal moves to get a large number.
• Negative Exponents
  • Numbers smaller than one; the decimal moves to get a small number.
• Coefficient Format
  • One non-zero number to the left of the decimal point.
• Changing incorrect Scientific Notation
  • Move decimal to the right spot
  • The expontent will move depending on the movement needed for the coefficient to be properly expressed.
  • Example
    • 352.9\times 10^{-4}

Converting from Scientific Notation

• 9.678 \times 10^4 = 96780
• 1.98 \times 10^{-3}=0.00198
• 4.2 \times 10^7 = 42000000
• 1.22 \times 10^{-5}=0.0000122