Chemistry Fundamentals: Properties, Measurements, and Significant Figures
Conservation of Mass and Chemical Reactions
- Mass is conserved in all chemical processes; the amount of hydrogen and the amount of oxygen that combine to form water equals the mass of the resulting water. This is illustrated by the statement that the mass of reactants equals the mass of products in chemical reactions. The formation of water from hydrogen and oxygen can be written as a chemical equation:
2\,\mathrm{H2} + \mathrm{O2} \rightarrow 2\,\mathrm{H_2O} - The same mass balance applies in reverse in processes like fuel cells, where hydrogen and oxygen form water while generating electricity.
- In everyday devices (e.g., phones), a multitude of chemical processes occur to make devices and materials (coatings, glass, electronics), all governed by the same conservation principles.
- Temperature and phase changes can be discussed without a thermometer in simple terms: e.g., water freezes below 0 °C under one atmosphere; this is a qualitative indicator of temperature relative to zero.
Atoms, Molecules, and Properties
- Atoms can bond to form molecules; molecules are groups of atoms bonded together.
- The microscopic domain contains atoms and molecules, and their arrangement determines properties and behavior of materials.
- Differentiation between physical and chemical properties:
- Chemical properties are related to how atoms bond to other atoms; changes in bonding indicate chemical processes.
- Physical properties do not necessarily involve changes in bonding.
- Examples of physical vs chemical processes:
- Water condensing from gas to liquid is a physical change (no change in water bonding).
- Iron rusting is a chemical change (iron bonds with oxygen to form iron oxide).
- Chromium plating on iron blocks oxidation, illustrating a chemical difference due to bonding/bonding environment.
- For fluorine as an example:
- Fluorine has physical properties (e.g., a yellow gas).
- It also has chemical properties (it reacts with most substances).
- The melting point of fluorine is a physical property: melts at
-220\,^{\circ}\mathrm{C} - Finely divided metals burning in fluorine is a chemical property.
- A specific chemical statement: 19\ \text{g F}2 \text{ reacts with } 1\ \text{g H}2 (emphasizing chemical reactions and stoichiometry).
- Summary: distinguishing physical vs chemical properties relies on whether bonding changes occur during observation or reaction.
Physical vs Chemical Properties: Overview and Examples
- Physical property examples: color/state (gas), melting point, boiling point, density, odor, etc.
- Fluorine as a yellow gas is a physical property.
- The term "gas" is a physical property descriptor.
- Chemical property examples: reactivity with substances, combustion, corrosion, and other reactions that involve forming new bonds.
- Fluorine reacts with most substances (chemical property).
- Finely divided metals burning in fluorine is a chemical property.
- Burning and reacting are indicators of chemical processes.
- Note: When describing properties, the context (chemical bonding environment) determines whether a property is physical or chemical.
Extensive vs Intensive Properties
- Extensive properties depend on the amount of substance:
- Mass is extensive: the mass of 1 cup of water differs from the mass of 1 gallon of water.
- Volume is extensive (depends on amount).
- Intensive properties do not depend on the amount of substance:
- Density is intensive: the density of water is the same whether you have a cup or a gallon.
- Temperature is intensive: equal at the same conditions regardless of the amount.
- Boiling point is intensive: a cup or a gallon of water boils at the same temperature under the same pressure.
- Practical takeaway: use extensive properties to quantify amount, and intensive properties to characterize the substance independent of amount.
Measurement, SI Units, and Basic Conversions
- SI base units discussed:
- Length: the meter (m).
- Mass: the kilogram (kg); 1\,kg = 1000\,g.
- Temperature scales: Celsius (°C), Fahrenheit (°F), and Kelvin (K).
- Common prefixes (applied to meters, grams, liters, etc.):
- kilo (k) = 1000, centi (c) = 10^{-2}, milli (m) = 10^{-3}.
- Examples: 1\,km = 1000\,m; 1\,cm = 0.01\,m; 1\,mm = 0.001\,m.
- Volume: defined by the liter (L):
- 1\,L = 1000\,mL = 1000\,cm^3.
- 1\,mL = 1\,cm^3.
- Temperature scales:
- Celsius to Kelvin: K = °C + 273.15
- Water freezing/melting at 0 °C and 1 atm, boiling at 100 °C under 1 atm.
- Fahrenheit scale: water freezes at 32 °F and boils at 212 °F under standard conditions.
- Additional units and concepts:
- Speed and density units (e.g., meters per second, kilograms per cubic meter, g/cm^3, g/L).
- Area units: square meters, etc.
- A note on dimensional analysis: density ρ = m / V.
- Reading measurement devices and the idea of precision:
- Graduated cylinders and the concept of the meniscus must be read at the bottom of the curve.
- Measurements are estimates beyond the smallest scale of the instrument; record with appropriate significant figures.
- Definition: significant figures convey the precision of a measurement.
- Rules for which digits are significant:
- All nonzero digits are significant.
- Zeros between nonzero digits are significant (captured zeros).
- Leading zeros are not significant (they indicate the decimal place).
- Trailing zeros after the decimal point are significant.
- Trailing zeros before the decimal point are ambiguous in this class; they are treated as not significant unless a decimal point is shown.
- Examples (as discussed in the transcript):
- 3090: signification depends on decimal indication; in this class trailing zeros before a decimal point are not necessarily significant (ambiguous without context).
- 70.607: five significant figures (7, 0, 6, 0, 7).
- 0.00832407: six significant figures (8, 3, 2, 4, 0, 7).
- Expressing numbers unambiguously using scientific notation:
- 1300 with two significant figures can be written as 1.3\times 10^{3}.
- To express four significant figures you can write 1.300\times 10^{3}, etc.
- Practical takeaway: use scientific notation to communicate the intended precision clearly.
- Addition and subtraction (alignment by decimal places):
- The result should be reported to the least precise decimal place among the operands.
- Example from transcript: 43.25 + 12.3 = 55.55, but reported to the tenths place due to the least precise decimal place (one decimal place) → 55.5.
- If the numbers have different decimal places, truncate/round to the smallest number of decimal places present.
- Subtraction example note: when subtracting numbers with different significant figures, the final result’s precision is governed by decimal places, not the number of significant figures.
- Multiplication and division (sig figs rule):
- The final answer should have the same number of significant figures as the factor with the fewest significant figures.
- Example concept: If a three-sig-figure number is multiplied by a four-sig-figure number, the product should have three significant figures.
- Rounding rules (how to round when discarding digits):
- If the dropped digit is less than 5, round down.
- If the dropped digit is greater than 5, round up.
- If the dropped digit is exactly 5, some conventions:
- The book’s convention is to round 5 to even (banker's rounding).
- Some people argue for always rounding up (odd rounding).
- Examples from the transcript (illustrative, with notes about rounding conventions):
- 31.57 rounded to two significant figures → 32.
- 8.1649 rounded to three significant figures → 8.16.
- 0.051065 rounded to four significant figures → 0.05107.
- 0.028675 rounded according to the drop digit’s value (7) → 0.0287.
- 18.384 rounded to three significant figures (example given as 18.3 in the transcript; note this appears inconsistent with standard rounding rules; typically 18.38 rounds to 18.4 for three sig figs).
- 6.8752 rounded to four significant figures → 6.88.
- Emphasis: the number of significant figures kept depends on the operation (add/subtract vs multiply/divide) and on the rounding convention chosen.
Worked Examples and Practice Problems (Added Context)
- Reading a graduated cylinder: estimation and significant figures in volume measurements
- A measurement like 21.6 mL is a practical estimate between marked lines; the last digit is an estimate and contributes to the significant figures.
- Discussion of reporting measurements with appropriate significant figures (e.g., 21.6 mL has three significant figures).
- A practical addition example from the transcript:
- Sum: 1.0023\ \text{g} + 4.383\ \text{g} = 5.3853\ \text{g}
- Because addition/subtraction is governed by decimal places, the result should be reported with the least number of decimal places among the operands (in this case, three decimal places), giving 5.385\ \text{g}.
- Mass-energy and chemical balance (concept): the mass of reactants equals the mass of products in closed systems.
- Chemical equation example (water formation): 2\,\mathrm{H2} + \mathrm{O2} \rightarrow 2\,\mathrm{H_2O}
- Temperature conversion to Kelvin: K = °C + 273.15
- Density: \rho = \frac{m}{V}
- Volume relationships: 1\ \text{L} = 1000\ \text{mL} = 1000\ \text{cm}^3, 1\ \text{mL} = 1\ \text{cm}^3
- Length and mass prefixes: 1\,km = 1000\,m, 1\,g = 1\,\text{g}, 1\,kg = 1000\,\text{g}.
- Length: base unit is the meter (m); distance/time units for speed include m/s or km/s depending on context.
- Standard states on measurement: record measurements with appropriate significant figures and acknowledge that some digits are estimates based on device precision.