Chemistry: Atoms, Elements, and Ions

• Chemistry is the study of matter, its properties, and how it changes.
• Matter is anything that possesses mass and occupies space.

Structure of an Atom

• Atoms constitute the fundamental building blocks of matter.
• Atoms consist of protons, neutrons, and electrons.
• Protons carry a positive charge, neutrons are neutral, and electrons have a negative charge.
• The number of protons determines the identity of the element.
• Protons and neutrons are located in the nucleus.
• Electrons reside in orbitals surrounding the nucleus.
• Atomic number (Z) signifies the quantity of protons within an atom's nucleus.
• Mass number (A) represents the sum of protons and neutrons in an atom's nucleus.
• Isotopes are variants of an element with differing neutron counts.
• Ions are atoms or molecules with a net electric charge due to gaining or losing electrons.
• Cations are positively charged ions formed by electron loss.
• Anions are negatively charged ions formed by electron gain.

Classification of Matter

• Matter is classified as either a pure substance or a mixture.
• Pure substances possess a fixed composition and distinct properties.
• Elements and compounds are types of pure substances.
• An element is a substance that cannot be broken down by chemical means (e.g., gold, oxygen).
• A compound consists of two or more elements chemically combined in a fixed ratio (e.g., water, salt).
• A mixture involves the physical combination of two or more substances.
• Mixtures can be homogeneous or heterogeneous.
• Homogeneous mixtures exhibit uniform composition (e.g., saltwater, air).
• Heterogeneous mixtures exhibit non-uniform composition (e.g., sand and water, salad).
• Solutions are homogeneous mixtures where a solute dissolves in a solvent.

States of Matter

• Solids, liquids, and gases are the three common states of matter.
• Solids maintain a fixed shape and volume.
• Liquids maintain a fixed volume but adopt the shape of their container.
• Gases lack a fixed shape or volume, expanding to fill the available space.
• Changes of state are physical changes, including melting (solid to liquid), freezing (liquid to solid), boiling/vaporization (liquid to gas), condensation (gas to liquid), sublimation (solid to gas), and deposition (gas to solid).

Properties of Matter

• Properties of matter can be physical or chemical.
• Physical properties are observable or measurable without altering the substance's composition (e.g., color, density, melting point, boiling point).
• Chemical properties describe how a substance interacts with others or changes composition (e.g., flammability, reactivity with acid).
• Intensive properties remain independent of the amount of substance present (e.g., temperature, density).
• Extensive properties are dependent on the amount of substance present (e.g., mass, volume).

Lab Safety

• Always use appropriate personal protective equipment (PPE), including safety goggles, gloves, and lab coat.
• Be aware of the location and proper use of safety equipment like fire extinguishers, safety showers, and eyewash stations.
• Comprehend all experimental procedures before starting.
• Handle chemicals carefully to avoid direct contact.
• Eating, drinking, or smoking is prohibited in the laboratory.
• Discard chemical waste as instructed.
• Immediately report accidents or spills to the instructor.
• Identify potential laboratory hazards like flammable materials, corrosive substances, and sharp objects.
• Maintain a clean and organized workspace.
• Do not conduct unauthorized experiments.
• Behave responsibly and avoid horseplay.

Measurement and Units

• The International System of Units (SI units) serves as the standard scientific unit system.
• SI base units include:
  • Meter (m) for length.
  • Kilogram (kg) for mass.
  • Second (s) for time.
  • Kelvin (K) for temperature.
  • Mole (mol) for the amount of substance.
• Metric prefixes modify SI units to express multiples or fractions (e.g., kilo-, centi-, milli-).
• Volume is commonly measured in liters (L) or milliliters (mL), where 1 L = 1000 mL.
• Temperature is measured in Celsius (°C) or Kelvin (K), where K = °C + 273.15.
• Density is the mass-to-volume ratio (d = m/v), typically in g/mL or g/cm³.

Reading a Metric Ruler

• Millimeters (mm), centimeters (cm), and meters (m) are the length units on a metric ruler.
• Align the object with the ruler's zero mark.
• Measure at the object's edge.
• Estimate to the nearest tenth of a millimeter for greater precision.
• 1 cm = 10 mm
• 1 m = 100 cm = 1000 mm

Significant Figures

• Significant figures are digits known with certainty plus one estimated digit.
• Rules for determining significant figures:
  • Non-zero digits are always significant.
  • Zeros between non-zero digits are significant.
  • Leading zeros are not significant.
  • Trailing zeros to the right of the decimal point are significant.
  • Trailing zeros in a whole number with no decimal point are not significant.
  • Trailing zeros in a whole number with the decimal point shown are significant.
• Rules for calculations:
  • Multiplication/division: the result reflects the fewest significant figures among factors.
  • Addition/subtraction: the result reflects the fewest decimal places.

Scientific Notation

• Scientific notation expresses large or small numbers as A × 10^B, where 1 ≤ A < 10 and B is an integer.
• To convert to scientific notation, move the decimal point to have one non-zero digit to its left.
• A leftward decimal shift yields a positive exponent.
• A rightward decimal shift yields a negative exponent.

Error Analysis

• Accuracy indicates closeness to the true value.
• Precision indicates the closeness of repeated measurements.
• Random errors have equal probabilities of being too high or too low.
• Systematic errors are consistently too high or too low.
• Percent error is calculated as [(experimental value - accepted value) / accepted value] × 100%.
• Standard deviation quantifies the spread of data points.
• Uncertainty estimates the deviation of a measured value from the true value.