Physical and Chemical Properties, SI Units, and Measurement Uncertainty

Comparison of Physical and Chemical Properties

• Physical Properties: These are properties that can be measured or observed without changing the identity or composition of the substance.

  • Measured Physical Properties: Includes density, melting point, boiling point, and electrical conductivity.
  • Observed Physical Properties: Includes color and hardness.
  • Hardness Measurement (Vickers Method): Hardness is often measured using an indenter. In the Vickers method, a diamond or tungsten indenter with a specific diamond shape is dropped onto a sample using a weight. A microscope is then used to measure the depth of the indentation; the deeper the indentation, the lower the hardness of the material.

• Chemical Properties: These describe the way a substance may change or react to form other substances. Measuring these properties typically involves consuming or destroying the sample.

  • Flammability: The ability of a substance to burn in the presence of oxygen.
  • Acidity: Measured using a $pH$ meter.
  • Toxicity: Considered a chemical property because it involves the substance interacting chemically with a biological system.
  • Reactivity/Oxidation: The susceptibility of a substance to chemical processes, such as iron rusting. Rust formation involves the conversion of iron (a shiny metal) into red iron oxide ($Fe_2O_3$). While the physical appearance changes, the underlying process is the breaking and forming of chemical bonds.

Understanding Physical and Chemical Changes

• Physical Changes: Changes in the physical state or appearance of a material without a change in its identity.

  • Phase Changes: Boiling, freezing, or melting. For example, boiling water converts it from a liquid to a vapor/gas state, but the water molecules ($H_2O$) remain intact.
  • Vapor vs. Gas: In chemistry, a "vapor" refers specifically to the gaseous state of a substance that is normally a liquid at room temperature (e.g., water vapor). A "gas" refers to a substance that exists in the gaseous state at room temperature and pressure (e.g., oxygen gas).

• Chemical Changes: These involve the breaking and forming of chemical bonds to create new substances.

  • Example: Burning wood is a chemical change where sunlight, water, and carbon dioxide stored in the wood are reacted. While technically reversible in nature (trees growing), it is not easily reversed in a laboratory setting.

• Distinguishing Between the Two:

  • Molecular Interaction: Physical changes involve disrupting interactions between molecules (intermolecular forces), while chemical changes involve breaking the bonds within molecules (intramolecular forces).
  • The Reversibility Test: A key indicator of a physical change is how easily it can be reversed. Freezing and melting ice or boiling and condensing water are easily reversible. For instance, water vapor in the air hitting a cold car windshield at night slows down and sticks together, forming droplets via condensation; when the sun comes out, they return to the vapor state.

Extensive vs. Intensive Properties

• Intensive Properties: These are independent of the amount of substance present.

  • Examples: Density, temperature, and color. The density of iron is approximately $7.87\,g/cm^3$ regardless of whether you have a small nail or a massive steel girder. Water boils at $100^{\circ}C$ whether it is a small cup or a giant pot.

• Extensive Properties: These depend on the amount of substance present.

  • Examples: Mass, volume, and heat. While a small cup and a large pot of water may be at the same temperature ($100^{\circ}C$), the large pot contains significantly more total heat energy.

Periodicity and the Periodic Table

• Structure: The table consists of horizontal rows called periods and vertical columns called groups.
• Chemical Similarity: Elements within the same group share similar chemical properties. For example, group 16 elements (Oxygen, Sulfur, Selenium, Tellurium, Polonium) tend to react in similar ratios (e.g., $H_2O$ and $H_2S$ exhibit a 2:1 ratio of hydrogen to the group element).
• Physical Diversity: Elements within a group do not necessarily share physical properties. Oxygen is a gas at room temperature, while sulfur is a solid. Within Group 16, the elements transition from non-metals (oxygen) to metalloids and finally metals (polonium).
• Metals vs. Non-metals: Metals are typically shiny and conduct heat/electricity well. Non-metals are usually dull and act as insulators.
• Halogens (Group 17): Includes Fluorine, Chlorine, Bromine, and Iodine. While they all react with sodium in a 1:1 ratio (chemical similarity), their physical properties vary wildly (e.g., Fluorine's melting point is $-220^{\circ}C$ while Iodine's is $114^{\circ}C$).

Scientific Measurements and the SI System

• Components of a Measurement: Number, unit, and uncertainty.

• The SI System: The International System of Units (Le Systme International d'Units) was adopted in 1964. It is distinct from the standard metric system and is based on seven fundamental base units, including:

  • Meter ($m$): Length.
  • Kilogram ($kg$): Mass.
  • Second ($s$): Time.
  • Kelvin ($K$): Temperature.
  • Mole ($mol$): Amount of substance.

• Derived Units: Units that are combinations of the seven base units.

  • Force (Newton): $1\,N = 1\,kg \cdot m/s^2$.
  • Volume: Derived from length ($Volume = length^3$). Common units include the Liter ($L$), where $1\,L = 1\,dm^3$, and the milliliter ($mL$), where $1\,mL = 1\,cm^3$ (also abbreviated as CC in medical contexts).
  • Density: Typically expressed as $g/cm^3$ or $g/mL$ for solids and liquids, and $g/L$ for gases.

Modern Definitions of Measurement Units

• Prototypes: Historically, units were defined by physical objects. The meter was a platinum-iridium bar kept in France. The kilogram was a metal cylinder stored under multiple bell jars in argon to prevent contamination or mass loss. However, these objects were found to change over time (e.g., losing mass as atoms "jumped off" the surface).
• Constants: Recently (notably in 2019), all SI base units were redefined based on fundamental physical constants.
  • The Meter: Now defined by the speed of light ($c$). It is the distance light travels in a vacuum in $\frac{1}{299,792,458}$ of a second.
  • The Kilogram: Now defined using Planck’s constant ($h$).
  • The Second: Based on a radiation event in Cesium ($Cs$) atoms, which is why official timekeepers are called "atomic clocks."

Temperature and Absolute Zero

• Temperature Scales:
  • Celsius ($^{\circ}C$): Based on water freezing at $0^{\circ}C$ and boiling at $100^{\circ}C$.
  • Kelvin ($K$): The absolute temperature scale. The increment size is identical to Celsius, but the zero point is different.
• Absolute Zero ($0\,K$): The theoretical temperature where all molecular motion ceases. It is exactly $-273.15^{\circ}C$.
• Conversion Formula: $T(K) = T(^{\circ}C) + 273.15$
• Energy Relationship: Temperature in Kelvin is directly proportional to the average kinetic energy ($KE$) of the system: $\text{KE} \propto K$. Doubling the temperature in Kelvin doubles the energy; doubling ${^{\circ}C}$ does not.

Uncertainty and Significant Figures

• Precision vs. Accuracy:

  • Accuracy: How close a measurement is to the true or accepted value.
  • Precision: How close multiple measurements are to each other (reproducibility).
  • Examples: High precision with low accuracy often indicates an instrument is out of calibration.

• Significant Figures (Sig Figs):

  • The Cost of Precision: Increasing decimals requires more expensive equipment. A balance measuring to $1.0\,g$ may cost $\$50$, while one measuring to $1.0158\,g$ may cost $\$1500$. It is vital not to overstate or understate precision.
  • Rules for Counting:
    1. Non-zero digits are always significant.
    2. Captive zeros (between non-zeros) are significant.
    3. Leading zeros are never significant.
    4. Trailing zeros are significant only if a decimal point is present.
  • Scientific Notation: Used to express significant figures without ambiguity (e.g., $1.00 \times 10^{3}$ clearly has three sig figs).

• Calculations with Sig Figs:

  • Addition and Subtraction: Result is rounded to the same number of decimal places as the value with the fewest decimal places.
  • Multiplication and Division: Result is rounded to the same number of total sig figs as the value with the fewest sig figs.
  • Rounding the Five: Advanced rounding (banker's rounding) suggests rounding to the nearest even number when the digit to be dropped is exactly 5 to avoid statistical bias (e.g., both $2.45$ and $2.55$ would round to $2.5$).

Dimensional Analysis

• Concept: A method of problem-solving that uses the units of the measurements to help guide through the necessary steps. Units cancel out just like numbers in fractions.
• Conversion Factors: Ratios used to express the same quantity in different units. For example, $\frac{2.54\,cm}{1\,in}$ or $\frac{1\,in}{2.54\,cm}$.
• Example Calculation: To convert $34\,inches$ to centimeters:     $34\,in \times \frac{2.54\,cm}{1\,in} = 86.36\,cm$     Applying sig figs (starting with two), the result is rounded to $86\,cm$.