Chapter 1: Matter and Measurement

Classifying Matter: Pure Substance or Mixture

Matter is anything that takes up space and has mass. It can be classified based on its composition and properties.

• Pure Substances: These have a fixed composition and distinct properties. They are further classified as elements or compounds.

• Mixtures: Combinations of two or more substances. Mixtures are classified as homogeneous (uniform composition) or heterogeneous (non-uniform composition).

Element Symbols

Each element is represented by a unique symbol, usually consisting of one or two letters. The first letter is always capitalized.

• 1-Letter Symbols: H (hydrogen), B (boron)

• 2-Letter Symbols: He (helium), Ca (calcium)

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

The Periodic Table of Elements

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties into columns.

• Groups: Vertical columns, elements in the same group have similar properties.

• Periods: Horizontal rows, elements in the same period have the same number of electron shells.

Physical Change

A physical change alters the form of matter but not its chemical identity. Examples include changes in state (solid, liquid, gas).

• Key Point: The composition of the substance remains unchanged during a physical change.

• Example: Melting ice to form liquid water.

Chemical Reaction

A chemical change results in a change in the chemical composition of a substance. When a substance undergoes a chemical change, it is involved in a chemical reaction.

• Key Point: Chemical reactions produce new substances with different properties.

• Example: Burning charcoal to produce carbon dioxide and ash.

Chemical Equations

Chemical equations represent the reactants and products in a chemical reaction, including their physical states and stoichiometric coefficients.

• Example Equation:

• Physical States: (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous

Scientific Notation

Scientific notation is used to express very large or very small numbers. It consists of a coefficient and a power of ten.

• Format:

• Example:

Significant Figures

Significant figures indicate the precision of a measured value. All nonzero digits are significant; zeros may or may not be significant depending on their position.

• Rules:

  • All nonzero digits are significant.

  • Leading zeros are not significant.

  • Captive zeros (between nonzero digits) are significant.

  • Trailing zeros are significant only if there is a decimal point.

• Example: 2.005 (4 significant figures)

Significant Figures in Calculations

• Addition/Subtraction: The answer is rounded to the least significant decimal place.

• Multiplication/Division: The answer is rounded to the number of significant figures in the least precise measurement.

• Example: (rounded to one decimal place)

SI Units and Metric Prefixes

The International System of Units (SI) is the modern metric system. Standard units include kilogram (kg) for mass, liter (L) for volume, and meter (m) for length.

• Metric Prefixes:

  Prefix	Abbreviation	Multiplier
  Tera	T
  Giga	G
  Milli	m
  Micro	μ
  Nano	n

Conversion Factors

Conversion factors are used to convert between different units of measurement.

• Example:

• Application:

Mass, Volume, and Density

Mass is the amount of material in an object, measured in grams (g). Volume is the space occupied, measured in milliliters (mL) or cubic centimeters (cm3).

• Density: Density is the ratio of mass to volume.

• Example: Water has a density of .

Temperature

Temperature measures the average kinetic energy of particles. Common units are Celsius (°C), Fahrenheit (°F), and Kelvin (K).

• Conversion:

Energy and Heat

Energy is the capacity to do work or supply heat. It exists as potential energy (stored) or kinetic energy (motion).

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

• Unit: Joule (J)

Calorie and Specific Heat

The calorie is the amount of energy required to raise the temperature of one gram of water by one degree Celsius.

• Conversion:

• Specific Heat (SH): The amount of heat needed to raise the temperature of 1 gram of a substance by 1°C.

• Example: Water has a high specific heat compared to metals.

States of Matter

Matter exists in three states: solid, liquid, and gas. Each state has distinct properties.

Chapter 2: Radioactivity

Atomic Particles

Atoms consist of subatomic particles: protons, neutrons, and electrons.

• Proton: Charge +1, located in nucleus, mass ≈ 1 amu

• Neutron: Charge 0, located in nucleus, mass ≈ 1 amu

• Electron: Charge -1, located outside nucleus, mass ≈ 1/2000 amu

Atomic Number vs. Mass Number

• Atomic Number (Z): Number of protons in an atom; defines the element.

• Mass Number (A): Total number of protons and neutrons in an atom.

• Isotopes: Atoms of the same element with different mass numbers due to varying numbers of neutrons.

Types of Radiation

Radioactive decay involves the emission of particles or energy from unstable nuclei.

• Alpha (α) Particle: Positively charged, 2 protons and 2 neutrons ()

• Beta (β) Particle: Negatively charged, high-energy electron ()

• Gamma (γ) Ray: Neutral, high-energy electromagnetic radiation

• Positron: Positive charge, same mass as electron ()

• Neutron: No charge

Biological Effects of Radiation

Ionizing radiation can damage living cells by ejecting electrons, making atoms more reactive and potentially harmful to biological tissues.

• Penetration: Higher energy radiation penetrates deeper into tissues.

Radioactive Decay

• Alpha Decay: Loss of an alpha particle.

• Beta Decay: Loss of a beta particle.

• Gamma Emission: Loss of a gamma ray, often accompanies other decay processes.

Half-Life

The half-life of a radioactive isotope is the time required for half of the atoms in a sample to decay.

• Short half-lives: Used in medical applications for rapid decay and elimination.

• Long half-lives: Used for dating archaeological samples.

• Measurement: Geiger counter

• Example: Carbon-14 has a half-life of 5730 years.

Chapter 3: How Elements Combine

Electron Arrangement

Electrons occupy energy levels around the nucleus, with each level holding a maximum number of electrons.

• Formula: Maximum electrons per level = , where n is the energy level.

• Example: n=1 can hold 2 electrons; n=2 can hold 8 electrons.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

The Octet Rule

Atoms tend to react to achieve eight electrons in their valence shell, attaining a stable noble gas configuration.

• Noble Gases: Group 8A elements; naturally unreactive due to full valence shell.

• Other Atoms: May lose, gain, or share electrons to achieve an octet.

Ion Formation and Naming

Ions are formed when atoms gain or lose electrons, resulting in a net charge.

• Cation: Positively charged ion (loss of electrons)

• Anion: Negatively charged ion (gain of electrons)

• Naming: Add 'ion' to the element name (e.g., sodium ion). For metals with multiple charges, indicate the charge in parentheses (e.g., iron(II) ion).

Naming Nonmetals and Polyatomic Ions

• Nonmetals: Suffix '-ide' replaces the last part of the element name (e.g., fluoride).

• Polyatomic Ions: Most end in '-ate'; '-ite' indicates one fewer oxygen atom. Examples: hydroxide (OH-), cyanide (CN-), ammonium (NH4+), chlorate, acetate, sulfate, carbonate, phosphate.

Additional info: Some content was inferred and expanded for clarity and completeness, including standard definitions, formulas, and examples relevant to GOB Chemistry.