Units of Measurement for Physical & Chemical Change

Physical and Chemical Properties (1.1)

Understanding the distinction between physical and chemical properties is fundamental in chemistry. These properties help identify substances and predict their behavior during changes.

• Physical Property: A characteristic of a substance that can be observed or measured without changing its chemical composition. Examples: Viscosity, color, lustre, melting point, boiling point, density.

• Chemical Property: A characteristic that can only be observed during a chemical change, indicating how a substance reacts with other substances. Examples: Reactivity with acid, combustibility, light sensitivity.

• Physical Change: A change that does not alter the chemical composition of a substance. Example: Boiling water (liquid to gas).

• Chemical Change: A change in which the chemical structure of a substance is altered, resulting in the formation of a new substance through a rearrangement of atoms. Example: Rusting of iron (formation of iron(III) oxide).

Additional info: Physical changes are often reversible, while chemical changes typically result in new substances and are not easily reversed.

Examples of Physical and Chemical Changes

Visual representations help distinguish between physical and chemical changes:

• Boiling Water: Water molecules transition from liquid (H2O(l)) to gas (H2O(g)), a physical change.

• Rusting Iron: Iron atoms react with oxygen to form iron(III) oxide (rust), a chemical change.

Energy in Physical & Chemical Change (1.2)

Energy changes often accompany physical and chemical changes. The total energy of an object is the sum of its kinetic and potential energies.

• Kinetic Energy: Energy associated with matter in motion. Example: The movement of molecules in a liquid.

• Potential Energy: Energy associated with position or composition (such as bond energies). Example: Energy stored in chemical bonds.

• Thermal Energy: Energy associated with the temperature of an object, arising from the motion of individual atoms or molecules. Example: Heat absorbed during evaporation.

Additional info: When you sweat, water evaporates from your skin, absorbing heat and making your skin feel cold. This is an example of energy transfer during a physical change.

Law of Conservation of Energy

The law of conservation of energy states that energy cannot be created or destroyed, only converted from one form to another.

• Key Principle: The total energy in a closed system remains constant.

• Example: When a 10 kg object falls, its potential energy is converted to kinetic energy and then to heat upon impact.

Equation:

Additional info: Energy transformations are central to both physical and chemical changes, such as the release of heat during combustion or the absorption of heat during melting.

International System of Units (SI Units)

Measurements in chemistry use the International System of Units (SI), which provides standard units for scientific work.

• Base SI Units:

  Quantity	Unit Name	Symbol
  Length	metre	m
  Mass	kilogram	kg
  Time	second	s
  Temperature	kelvin	K
  Amount of substance	mole	mol
  Electric current	ampere	A
  Luminous intensity	candela	cd

• Derived Units: Formed by combining base units (e.g., volume in m3, speed in m/s, energy in joules).

Additional info: SI units are used globally to ensure consistency in scientific measurements.

Metric Prefixes and Conversions

Metric prefixes are used to express measurements in powers of ten, making it easier to handle very large or small quantities.

• Common Prefixes:

  Prefix	Symbol	Factor
  kilo	k	103
  mega	M	106
  giga	G	109
  centi	c	10-2
  milli	m	10-3
  micro	μ	10-6
  nano	n	10-9

• Conversion Factors: Used to convert between units (e.g., ).

Additional info: Always use conversion factors as equivalent fractions to ensure correct unit cancellation.

Significant Figures (1.4)

Significant figures reflect the precision of a measurement. The rules for determining significant figures are essential for reporting scientific data accurately.

• Rules for Significant Figures:

  • All nonzero digits are significant (e.g., 28.03 has 4 significant figures).

  • Zeros between nonzero digits are significant (e.g., 4087.0301 has 8 significant figures).

  • Leading zeros are not significant (e.g., 0.0032 has 2 significant figures).

  • Trailing zeros after a decimal point are significant (e.g., 45.000 has 5 significant figures).

  • Trailing zeros in a whole number without a decimal point are ambiguous; use scientific notation to clarify.

• Exact Numbers: Numbers from definitions or counting are considered to have infinite significant figures (e.g., 100 people).

Additional info: Scientific notation is used to clearly indicate the number of significant figures in a measurement.

Significant Figures in Calculations

When performing mathematical operations, the number of significant figures in the result must reflect the precision of the input values.

• Multiplication/Division: The result should have the same number of significant figures as the input with the fewest significant figures. Example: (2 significant figures).

• Addition/Subtraction: The result should have the same number of decimal places as the input with the fewest decimal places. Example: (1 decimal place).

Accuracy and Precision

Accuracy and precision are measures of the quality of a measurement.

• Accuracy: How close a measured value is to the true or accepted value.

• Precision: How close repeated measurements are to each other.

• Percent Error Formula:

Additional info: High accuracy and high precision are ideal, but not always achievable due to limitations in equipment and technique.

Solving Chemical Problems (1.5)

Solving chemical problems involves identifying relevant measurements, selecting appropriate equations and conversion factors, and applying dimensional analysis.

• Steps:

  1. Identify the measurements/data given in the question.

  2. Determine what you are solving for.

  3. List equations, conversion factors, and strategies.

  4. Perform calculations, including units at every step.

  5. State the answer with correct units and significant figures.

• Dimensional Analysis (Factor-Label Method): A systematic approach to problem-solving that uses conversion factors to ensure units are correctly converted and canceled. Example: Converting grams to milligrams using .

Additional info: Always check that the final answer has the correct units and reasonable magnitude.

Imperial and SI Conversions

Some problems require conversion between imperial and SI units.

• Length:

• Temperature:

• Speed: Convert km/h to mi/h using appropriate conversion factors.

Additional info: Use provided conversion factors as needed for specific problems.

Scientific Notation

Scientific notation is used to express very large or small numbers in a standard format.

• Format: , where and is an integer.

• Example:

• Example:

Additional info: Scientific notation clarifies significant figures and simplifies calculations.