Density and Measurement in Science

Density of an Irregularly Shaped Object

• Mass of the object: 103 g
• Initial water volume: 22.4 mL
• Water level after adding object: 93.8 mL
• Volume of the object: 93.8 mL - 22.4 mL = 71.4 mL
• Density of the object: 103 g / 71.4 mL = 1.44 g/mL
• Accepted density: 1.77 g/mL
• Percent error: |1.44 g/mL - 1.77 g/mL| / 1.77 g/mL * 100% = 18.6%

Graduated Cylinder vs. Beaker

• Graduated cylinder provides more precision in measurements due to more graduations, resulting in more significant figures.
• Significant figures express the uncertainty of a measurement.
• Important practices when using a graduated cylinder:
  • Read at eye level
  • Read from the bottom of the meniscus
  • Ensure the cylinder is on a flat surface

Scientific Experiment Terminology

• Hypothesis: A prediction of the outcome of an experiment. (e.g., If the size of the parachute increases, then the time taken for the parachute to fall will increase, because a larger parachute will catch more air)
• Independent Variable: The changed factor in an experiment. (e.g., Size of the parachute)
• Dependent Variable: The measured factor in response to the independent variable. (e.g., Time taken for the parachute to fall )
• Control: A baseline for comparison; condition where the independent variable is not changed. (e.g., No control in a parachute experiment, varying the parachute is the independent variable)
• Constant: Items kept the same throughout the experiment, (e.g., Parachute material, shape in the parachute experiment)

Erlenmeyer Flask

• Shape: Sloped sides
• Use: Holding liquids, suitable for swirling

Ruler Measurements

• Ruler A measurement: 4.8 cm (2 significant figures)
• Ruler B measurement: 4.84 cm (3 significant figures)

Measurement Units and Tools

• Mass of a pencil: grams (g), electronic balance
• Length of a desk: centimeters (cm), ruler or meter stick
• Temperature of coffee: degrees Celsius (°C), thermometer or temperature probe

Melting Point Precision and Accuracy

• Student A is more precise (measurements close together) and accurate (close to accepted value). Student B's measurements were more dispersed.

Scientific Notation and Standard Form

• 0.00443 in scientific notation: 4.43 x 10^-3
• 35889 in scientific notation: 3.5889 x 10⁴
• 3.71 x 10^-3 in standard form: 0.00371
• 8.825 x 10⁵ in standard form: 882500

Atomic Structure Discoveries

• Bohr: Electrons orbit the nucleus in specific energy levels.
• Rutherford: Atom is mostly empty space with a dense, positive nucleus.
• Thomson: Discovered the electron using a cathode ray experiment

Periodic Table

• Mendeleev: Created a periodic table arranged by atomic mass.
• Modern Periodic Table: Organized by atomic number (number of protons)

Periodic Table Locations of Elements

• Metals: Left of the staircase
• Nonmetals: Right of the staircase
• Metalloids: Along the staircase
• Transition Metals: Middle section (groups 3-12)
• Halogens: Group 17
• Alkali Metals: Group 1
• Noble Gases: Group 18
• Alkaline Earth Metals: Group 2

Properties of Elements in the Same Group/Period

• Same Group: Same number of valence electrons, similar chemical properties.
• Same Period: Valence electrons on the same shell/principal energy level.

Subatomic Particles

• Proton: Positive charge, ~1 amu, located in the nucleus.
• Neutron: No charge, ~1 amu, located in the nucleus.
• Electron: Negative charge, ~0 amu (much smaller than proton), located in orbitals around the nucleus.

Nuclear Notation Examples and Isotopes

• Sodium-24: ²⁴Na₁₁ (11 protons, 13 neutrons, 11 electrons)
• Sulfur-36: ³⁶S₁₆ (16 protons, 20 neutrons, 16 electrons)

Carbon Isotopes and Radioactivity

• Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.
• Carbon-14 is radioactive because its neutron number makes the nucleus unstable, causing it to decay for stability. Carbon-12 and Carbon-13 are stable.

Valence Electrons by Group

• Group 1-2: 1 to 2 valence electrons.
• Group 13-18: 3 to 8 valence electrons.

Atom Stability

• Atoms become stable (lower potential energy) with a full valence electron shell (octet). This usually occurs through forming chemical bonds. Noble gases are already stable with full valence.

Ion Formation

• K: K⁺
• Be: Be²⁺
• Br: Br^-
• Al: Al³⁺

Potassium and Iodine Compound

• Formula: KI (Potassium Iodide)
• Potassium loses electrons because it is a metal, transferring to iodine

Barium and Sulfur Compound

• Bond Type: Ionic bond (metal and nonmetal)
• Formula: BaS (Barium Sulfide)

Compound Classification

• a. Low melting point, doesn't conduct electricity in solution: Covalent compound
• b. Ductile, good conductor of electricity: Metal
• c. High melting point, dissolves well in water: Ionic compound

Intermolecular Bonding and Melting Point

• Stronger intermolecular bonds lead to higher melting points.

Intermolecular Bond Strength Order

• London dispersion forces < dipole-dipole forces < hydrogen bonding < ionic bonds

Physical vs. Chemical Change

• Physical Change: Alters physical properties but not the composition (e.g., melting ice)
• Chemical Change: Alters the chemical composition (e.g., rusting)
• Indicators of Chemical Change: Color change, smell, light, temperature change, precipitate formed, gas formed

Lewis Structures for NH3 and CF4

• NH₃: Nitrogen forms single bonds with three hydrogens. Satisfies nitrogen's need for 3 more electrons for octet configuration. Each hydrogen completes its need for one more electron.
• CF₄: Carbon forms single bonds with four fluorines. This results in each fluorine atom completing its octet, and carbon fulfilling its need of 4 electrons.

Nuclear vs. Chemical Reactions

• Nuclear Reaction: Changes to the atom's nucleus
• Chemical Reaction: Changes in the arrangement of atoms to create new substances. The three nuclear types are: nuclear decay, nuclear fission (large nucleus splits), and nuclear fusion (small nuclei combine).

Balanced Chemical Reactions and Reaction Types

• a. C₃H₈ + 5O₂ → 3CO₂ + 4H₂O (Combustion)
• b. 2NH₃ → N₂ + 3H₂ (Decomposition)
• c. Al(NO₃)₃ + 3NaOH → Al(OH)₃ + 3NaNO₃ (Double Replacement)
• d. MnO₂ + 2Mg → 2MgO + Mn (Single Replacement)
• e. 2Al + 3Cl₂ → 2AlCl₃ (Synthesis)

Energy Changes in Reactions

• Negative ΔH_rxn (enthalpy of reaction): Exothermic (energy released)
• Positive ΔH_rxn: Endothermic (energy absorbed)

Chemical Equation Calculations

• a. Balanced equation: CH₄ + 2H₂O → CO₂ + 4H₂
• b. ΔH_rxn = 162kJ (Calculation shown in details in given text)
• c. Endothermic (ΔH_rxn is positive)
• d. Carbon-oxygen bond in CO₂: Polar covalent, Double bond( sharing 4 electrons)
• e. Moles of H₂ produced from 2 moles of CH₄: 8 moles
• f. Hydrogen-hydrogen bond in H₂: nonpolar covalent single bond (Sharing 2 electrons). Electrons shared equally.

Mole Calculations for CH₂F₂

• a. Molar mass of CH₂F₂: 52.03 g/mol
• b. Mass of 4.66 mol CH₂F₂: 242.46 g
• c. Number of molecules in 17.6 g CH₂F₂: 2.04 x 10²³ molecules

Intermolecular vs. Intramolecular Bonds

• Intermolecular: Forces between molecules
• Intramolecular: Forces within a molecule
• When boiling water, intermolecular bonds are being broken /disrupted.

Endothermic/Exothermic Processes and Energy Transformations

• Photosynthesis: Endothermic; light energy converted to chemical potential energy.
• Cell respiration: Exothermic; chemical potential energy into heat, kinetic, and potential energy.
• Combustion of fossil fuels: Exothermic; chemical potential energy converted to heat and kinetic energy.

Coal and Nuclear Power Plant Comparison

• Coal: Chemical reaction (combustion), releases CO₂ and other pollutants, major source of air pollution and global warming.
• Nuclear: Nuclear reaction (fission), produces radioactive waste, requires large amounts of water.

Strontium-90 Half-Life

• Half-life of Strontium-90: 25 years (Based on Graph provided).

This quiz covers the concepts of density calculations, precision measurement techniques using graduated cylinders versus beakers, and key scientific terminology. Test your understanding of how to accurately determine density and apply significant figures in measurements.