General Chemistry Fundamentals: Matter, Atomic Theory, Elements, Compounds, and Chemical Transformations

Matter

Definition and Classification

Matter is anything that has mass and occupies space. It is classified based on its composition and properties into ordinary matter and dark matter, and further into pure substances and mixtures.

• Ordinary Matter: Composed of atoms of chemical elements; makes up most observable matter.

• Dark Matter: Not composed of atoms; constitutes about 80% of the universe's matter.

• Chemical Element: Substance composed of one type of atom. Examples: Hydrogen, Helium, Oxygen, Silicon, Aluminum.

• Atomic Number: Number of protons in an atom; determines the element.

Examples of Elemental Abundance:

• Hydrogen: Most abundant element (about 88.6% of all atoms in the universe).

• Helium: Second most abundant (about 11.3%).

• On Earth: Oxygen, Silicon, and Aluminum are most abundant.

Atoms, Molecules, and Compounds

Structure and Properties

Atoms are the fundamental building blocks of matter, composed of a nucleus (protons and neutrons) and electrons. Molecules are groups of atoms bonded together in specific arrangements.

• Atoms: Spherically shaped, chemically indivisible particles.

• Molecules: Two or more atoms bonded together in a fixed arrangement. Example: $\text{O}_2$ (oxygen gas), $\text{CO}_2$ (carbon dioxide).

• Network Solids: Atoms bonded in a continuous network (e.g., diamond, NaCl).

• Elements: Substances made of one type of atom.

• Compounds: Substances composed of atoms of two or more elements chemically bonded.

Pure Substances and Mixtures

Classification and Examples

Pure substances have uniform and definite composition, while mixtures consist of two or more substances physically combined.

• Pure Chemical Substances: Made up of one type of particle (species). Examples: $\text{H}_2\text{O}$ (water), $\text{CO}_2$ (carbon dioxide).

• Mixtures: Physical combinations of two or more substances.

• Homogeneous Mixtures: Uniform composition throughout. Example: clean air, sea water.

• Heterogeneous Mixtures: Non-uniform composition. Example: dusty air, wet sand.

States of Matter

The state of matter depends on the arrangement and interaction of particles.

• Solid: Particles packed closely in fixed positions; definite shape and volume.

• Liquid: Particles close but can move past each other; definite volume, no fixed shape.

• Gas: Particles far apart, move freely; no fixed shape or volume.

Composition of Matter

Matter can be classified based on the types of particles present.

Type	Definition	Examples
Element	One type of atom	Oxygen, Helium
Compound	Two or more types of atoms chemically bonded	Water ($\text{H}_2\text{O}$), Sodium chloride ($\text{NaCl}$)
Homogeneous Mixture	Uniform composition	Air, Sea water
Heterogeneous Mixture	Non-uniform composition	Wet sand, Milk

Physical and Chemical Changes

Definitions and Examples

Physical changes alter the state or appearance of matter without changing its composition. Chemical changes result in the formation of new substances.

• Physical Change: Change in state (solid, liquid, gas) or appearance. Example: melting ice, boiling water.

• Chemical Change: Change in chemical identity. Example: electrolysis of water to produce hydrogen and oxygen gases.

Atomic Theory

Fundamental Laws

Atomic theory explains the nature of matter through several foundational laws.

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction.

  • Example: $\text{Mass of Na} + \text{Mass of Cl}_2 = \text{Mass of NaCl}$

• Law of Definite Proportion: A compound always contains the same proportion of elements by mass.

  • Example: $\frac{16\text{g Oxygen}}{2\text{g Hydrogen}} = 8.0$ in water

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of one element that combine with a fixed mass of the other are small whole numbers.

  • Example: $\frac{2.67}{1.33} = 2/1$ for CO and CO2

• Dalton's Atomic Theory:

  • Elements are composed of tiny, indestructible particles called atoms.

  • Atoms of the same element have the same mass and properties.

  • Atoms combine in simple, whole-number ratios to form compounds.

  • Atoms cannot change into atoms of another element in chemical reactions.

Structure of the Atom

Subatomic Particles

Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons.

• Nucleus: Contains most of the atom's mass and all positive charge.

• Electrons: Negatively charged particles distributed throughout the atom's volume.

• Protons: Positively charged particles in the nucleus.

  • Mass: $1.67262 \times 10^{-27}$ kg

  • Charge: $+1.60218 \times 10^{-19}$ C

• Neutrons: Neutral particles in the nucleus.

  • Mass: $1.67493 \times 10^{-27}$ kg

  • Charge: 0

Atomic Number and Mass Number

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

• Mass Number (A): Sum of protons and neutrons in the nucleus.

• Chemical Symbol: One or two-letter abbreviation for an element.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Mononuclidic Elements: Elements with only one naturally occurring isotope.

• Isotopologues: Chemical substances with the same elemental but different isotopic composition.

Standard Atomic Weight

The standard atomic weight is the weighted average mass of the isotopes of an element.

• Calculated as: $\sum (\text{fraction of isotope } n) \times (\text{mass of isotope } n)$

Ions and Mass Spectrometry

Ions

Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

Mass Spectrometry

Mass spectrometry is a technique used to determine the atomic mass and identify molecules in a sample.

• Separates particles according to their mass.

• Used to determine atomic mass and molecular composition.

The Mole and Molar Mass

Definitions and Calculations

The mole is a fundamental unit in chemistry for counting particles. Molar mass relates the mass of a substance to the number of moles.

• Mole (mol): Amount of material containing $6.02214 \times 10^{23}$ units (Avogadro's number).

• Molar Mass: Mass of 1 mole of atoms in grams; numerically equal to atomic mass in amu.

• Calculation Process:

  1. grams of element $\rightarrow$ (molar mass of element) $\rightarrow$ mol element $\rightarrow$ (Avogadro's #) $\rightarrow$ number of atoms

Example: $12.01 \text{ g C} = 1 \text{ mol C} = 6.022 \times 10^{23}$ atoms

Additional info: Some context and examples have been expanded for clarity and completeness.