Atoms and Elements: Structure, Laws, and the Periodic Table

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Atoms and Elements

Introduction

This study guide covers the foundational concepts of atomic structure, the laws that led to modern atomic theory, and the organization and classification of elements in the periodic table. Understanding these topics is essential for connecting macroscopic chemical phenomena to the microscopic world of atoms.

Atomic Structure

Atoms: The Fundamental Units

Atom: The smallest identifiable unit of an element, retaining its chemical properties.
There are about 91 naturally occurring elements and over 20 synthetic elements.
Elements are pure substances made of only one type of atom (e.g., copper), while compounds (e.g., salt) consist of atoms of different elements.

The Structure of the Atom

Each atom consists of two main parts:
- Nucleus: The dense center containing protons and neutrons.
- Electron Cloud: Surrounds the nucleus; contains electrons.
The nucleus contains over 99% of the atom's mass but is extremely small compared to the atom's overall size.

Subatomic Particles

Proton (p+): Positively charged, located in the nucleus, mass ≈ 1 amu.
Neutron (n0): Neutral (no charge), located in the nucleus, mass ≈ 1 amu.
Electron (e-): Negatively charged, located outside the nucleus, mass ≈ 0.00055 amu.

Particle	Mass (kg)	Mass (amu)	Charge (relative)	Charge (C)
Proton	1.67262 × 10-27	1.00727	+1	+1.60218 × 10-19
Neutron	1.67493 × 10-27	1.00866	0	0
Electron	0.00091 × 10-27	0.00055	-1	-1.60218 × 10-19

Modern Atomic Theory and Laws

Key Laws Leading to Atomic Theory

Law of Conservation of Mass (Antoine Lavoisier): In a chemical reaction, matter is neither created nor destroyed. The total mass of substances remains unchanged.
Law of Definite Proportions (Joseph Proust): All samples of a given compound have the same proportions of their constituent elements. Also called the law of constant composition.
Law of Multiple Proportions (John Dalton): When two elements form different compounds, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios.

Examples and Applications

Law of Conservation of Mass Example: When a log burns, the mass of the ash is less than the original log, but the matter is conserved as gases released into the air.
Law of Definite Proportions Example: Decomposition of 18.0 g of water yields 16.0 g oxygen and 2.0 g hydrogen. Mass ratio:
Law of Multiple Proportions Example: Carbon monoxide (CO) and carbon dioxide (CO2) have oxygen-to-carbon mass ratios of 1.33:1 and 2.67:1, respectively. The ratio of ratios is

Dalton's Atomic Theory

Each element is composed of tiny, indestructible particles called atoms.
All atoms of a given element have the same mass and properties.
Atoms combine in simple, whole-number ratios to form compounds.
Atoms of one element cannot change into atoms of another element in chemical reactions; they only change how they are bound together.

Elements and the Periodic Table

Defining Elements

The number of protons in the nucleus defines the element and is called the atomic number (Z).
The mass number (A) is the sum of protons and neutrons.
Elements are represented by a chemical symbol (e.g., C for carbon, Au for gold).

Periodic Table Organization

Elements are arranged by increasing atomic number.
Each element has a unique name and symbol (1 or 2 letters; first letter capitalized, second letter lowercase).
Chemical symbols may derive from English, Latin, or Greek names (e.g., Au from "Aurum" for gold).

Classification of Elements

Metals: Good conductors, malleable, ductile, shiny, tend to lose electrons.
Nonmetals: Poor conductors, not malleable or ductile, tend to gain electrons.
Metalloids: Exhibit mixed properties, often semiconductors.

Groups and Periods

Groups (vertical columns): Elements with similar properties; numbered 1-18.
Periods (horizontal rows): Numbered 1-7.
Main-group elements: Groups 1, 2, and 13-18.
Transition elements: Groups 3-12.

Special Groups

Noble Gases (Group 18/8A): Chemically inert, e.g., helium, neon, argon.
Alkali Metals (Group 1/1A): Highly reactive metals, e.g., sodium, potassium.
Alkaline Earth Metals (Group 2/2A): Fairly reactive, e.g., calcium, magnesium.
Halogens (Group 17/7A): Very reactive nonmetals, e.g., fluorine, chlorine.

Isotopes

Definition and Properties

Atoms of the same element with different numbers of neutrons are called isotopes.
Isotopes have the same number of protons but different mass numbers.
The natural abundance of isotopes is the relative percentage of each isotope in a natural sample.

Symbol	Number of Protons	Number of Neutrons	Mass Number (A)	Natural Abundance (%)
Ne-20	10	10	20	90.48
Ne-21	10	11	21	0.27
Ne-22	10	12	22	9.25

Calculating Neutrons

Number of neutrons = Mass number (A) - Atomic number (Z)
Example: Argon isotope with A = 40, Z = 18: Neutrons = 40 - 18 = 22

Atomic Mass

Average Atomic Mass

The atomic mass (atomic weight) of an element is the weighted average of the masses of its isotopes, based on their natural abundances.
Formula:

Example: Chlorine has two isotopes, Cl-35 (75.77%, 34.97 amu) and Cl-37 (24.23%, 36.97 amu):

Summary Table: Key Atomic Concepts

Concept	Definition	Example
Atom	Smallest unit of an element	Copper atom
Element	Substance made of one type of atom	Gold (Au)
Isotope	Atoms of same element, different neutrons	Ne-20, Ne-21, Ne-22
Atomic Number (Z)	Number of protons	Carbon: Z = 6
Mass Number (A)	Protons + Neutrons	Carbon-12: A = 12
Atomic Mass	Weighted average of isotopic masses	Cl: 35.45 amu

Additional info:

Understanding atomic structure and the periodic table is foundational for all further study in chemistry, including chemical reactions, bonding, and properties of matter.