Fundamental Concepts in General, Organic, and Biological Chemistry

Study Guide - Smart Notes

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1. States of Matter

Types of States

Matter exists in different physical forms known as states of matter. Each state has distinct properties based on particle arrangement and energy.

Solid: Definite shape and volume. Particles are closely packed and vibrate in place. Example: Ice.
Liquid: Definite volume but no definite shape. Particles are less tightly packed and can move past each other. Example: Water.
Gas: No definite shape or volume. Particles move freely and are far apart. Example: Oxygen.
Plasma: Ionized gas with free electrons and ions. Example: Lightning, neon signs.
Bose-Einstein Condensate: State at temperatures near absolute zero where particles behave as a single quantum entity. Example: Laboratory-created condensates.

2. Changes of State

Phase Transitions

Changes between states of matter involve energy transfer and are called phase transitions.

Fusion (Melting): Solid to liquid.
Solidification (Freezing): Liquid to solid.
Evaporation: Liquid to gas.
Condensation: Gas to liquid.
Sublimation: Solid to gas.
Deposition: Gas to solid.

3. Methods for Separating Mixtures

Separation Techniques

Mixtures can be separated into their components using physical methods:

Filtration: Separates solids from liquids using a porous barrier.
Distillation: Separates substances based on differences in boiling points.
Chromatography: Separates components based on movement through a medium.
Evaporation: Removes a liquid from a solution to leave a solid residue.
Decantation: Separates liquid from solids by pouring off the liquid.

4. Classification of Matter

Types of Substances

Matter is classified based on composition:

Pure Substances: Elements and compounds with fixed composition.
Mixtures: Physical combinations of two or more substances. Can be homogeneous (uniform) or heterogeneous (non-uniform).

5. Physical and Chemical Properties

Types of Properties

Physical Properties: Observed without changing the substance's identity. Examples: Color, density, melting point.
Chemical Properties: Describe a substance's ability to undergo chemical changes. Examples: Reactivity, flammability.

Physical and Chemical Changes

Physical Change: No new substance is formed. Example: Melting ice.
Chemical Change: New substances are formed. Example: Rusting iron.

6. Subatomic Particles and Atomic Structure

Subatomic Particles

Proton (p): Positive charge, located in the nucleus.
Neutron (n): No charge, located in the nucleus.
Electron (e): Negative charge, located in electron cloud around nucleus.

7. Atomic Models

Development of Atomic Theory

Dalton's Model: Atoms are indivisible particles.
Thomson's Model: "Plum pudding" model; electrons in a positive sphere.
Rutherford's Model: Small, dense, positive nucleus; electrons orbit nucleus.
Bohr's Model: Electrons in fixed orbits with quantized energy levels.
Quantum Mechanical Model: Electrons in orbitals; probability-based locations.

8. Atomic Number and Mass Number

Definitions

Atomic Number (Z): Number of protons in the nucleus; defines the element.
Mass Number (A): Total number of protons and neutrons in the nucleus.

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

9. Identifying Isotopes and Atomic Structure

Isotope Notation

Isotopes are written as , where X is the element symbol, A is mass number, Z is atomic number.
Example: is carbon-14.

10. Organization of the Periodic Table

Structure and Groups

7 periods (rows), 18 groups (columns).
Group Classifications: Alkali metals, alkaline earth metals, halogens, noble gases, transition metals, lanthanides, actinides.
Metals: Left and center; Nonmetals: Right; Metalloids: Stair-step region.

11. Periodic Trends

Trends Across Periods and Groups

Atomic Radius: Decreases across a period, increases down a group.
Ionization Energy: Increases across a period, decreases down a group.
Electronegativity: Increases across a period, decreases down a group.

12. Electron Configuration Rules

Filling Order and Principles

Aufbau Principle: Electrons fill orbitals of lowest energy first.
Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.
Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Example: The electron configuration for sodium (Na, atomic number 11) is:

13. Electron Configurations and Valence Electrons

Identifying Valence Electrons

Valence electrons are those in the outermost shell; they determine chemical reactivity.
For representative elements, the group number indicates the number of valence electrons.

Identifying an Atom from Electron Configuration

Count the total number of electrons in the configuration.
This number equals the atomic number and identifies the element.
Example: has 11 electrons, which is sodium (Na).