PDFDocument_Chapter_2

Atoms and Elements

Understanding Atoms

• Atoms are the fundamental building blocks of matter, forming the basis of everything in the universe. They are incredibly small, typically measuring around 0.1 nanometers in diameter.

• Cutting graphite results in carbon atoms, which are indivisible in nature according to classical physics. However, they can be split into subatomic particles through nuclear reactions.

• The term "atom" is derived from the Greek word "atomos," meaning "indivisible," which reflects the historical belief that atoms could not be further divided.

Imaging Atoms

Scanning Tunneling Microscopy (STM)

• A powerful technique that allows scientists to image and manipulate individual atoms with high precision. It utilizes quantum tunneling of electrons between a sharp tip and a conducting surface.

• Can visualize individual iron atoms (colored red) on a copper surface (colored blue), helping to illustrate the connection between the macroscopic properties of materials and their microscopic structures.

Characteristics of Atoms

• There are approximately 91 naturally occurring elements in nature, with over 20 synthetic elements created in laboratories. Each element has unique properties and behaviors.

Historical Perspectives on Atomic Theory

Early Ideas

• Philosophers such as Leucippus and Democritus were pioneers in proposing that matter consisted of tiny, indestructible particles called atoms. They speculated that these atoms varied in size, shape, and movement.

Advancements by John Dalton (1766-1844)

• Dalton's work provided substantial evidence supporting atomic theories. He analyzed the nature of gases and proposed that chemical reactions involve the rearrangement of atoms.

• His work laid the foundation for modern atomic theory, which transformed chemistry into a more systematic science.

Modern Atomic Theory

Key Laws Leading to Atomic Theory

1. Law of Conservation of Mass: This fundamental law states that mass is neither created nor destroyed in chemical reactions, establishing that the total mass of reactants equals the total mass of products.

2. Law of Definite Proportions: This law indicates that a compound always contains the same proportion of elements by mass, regardless of its source (e.g. water always has a mass ratio of 2:16 for hydrogen to oxygen).

3. Law of Multiple Proportions: This law describes how different compounds made of the same elements will have mass ratios that are simple whole-number ratios (e.g., CO vs. CO₂ has mass ratios of 1:1.33 and 1:2.67 respectively).

Detailed Look at Proportions

Law of Definite Proportions

• For example, the decomposition of water (H₂O) always yields the same acceptable mass ratio of hydrogen (2g) to oxygen (16g), simplifying to a ratio of 1:8.

Law of Multiple Proportions

• Carbon monoxide (CO) has a different mass ratio of oxygen (1.33:1), compared to carbon dioxide (CO₂) which shows a different mass ratio (2.67:1), demonstrating the distinct proportions of elements in different compounds.

Dalton's Atomic Theory

Key Principles Include:

• Matter is composed of atoms, which are the smallest units of matter that retain the properties of an element.

• Atoms of a single element are identical in size, mass, and properties, differing from atoms of other elements.

• Atoms combine in simple whole-number ratios to form compounds, illustrating the foundation of chemical reactions.

• Atoms of one element cannot transform into atoms of another element during chemical reactions, reinforcing the concept of conservation of matter.

Discovery of the Electron

J.J. Thomson's Experiment

• Thomson utilized cathode rays to discover electrons, which are negatively charged subatomic particles, thus contributing to the understanding of atomic structure.

• He calculated the charge-to-mass ratio of electrons, which is approximately -1.76 x 10³ C/g, showcasing their significant interaction in chemical reactions.

Millikan's Experiment

• Millikan's oil drop experiment accurately determined the charge of a single electron, further confirming the existence and properties of this subatomic particle.

Plum-Pudding Model

• Thomson proposed the Plum-Pudding Model, where he suggested that atoms are composed of a positively charged ‘soup’ with negatively charged electrons embedded, giving a uniform structure.

Rutherford's Gold Foil Experiment

Findings

• Rutherford's experiments revealed that atoms consist mostly of empty space, containing a dense central nucleus that holds most of the mass and positive charge.

• Proposed the nuclear model of the atom, where electrons orbit around the positively charged nucleus, leading to a new understanding of atomic structure.

• Stated that atoms are electrically neutral due to the equal numbers of protons (positively charged) and electrons (negatively charged).

Subatomic Particles

Types of Subatomic Particles

• Protons: Positively charged particles found in the nucleus (mass = 1.007 amu).

• Neutrons: Neutral particles, also found in the nucleus (mass = 1.008 amu).

• Electrons: Negatively charged particles orbiting the nucleus (mass = 0.00055 amu).

• Together, protons and neutrons determine the mass of the atom, with electrons contributing little to the mass while being crucial for chemical interactions.

Defining Elements

Atomic Number (Z)

• The atomic number, or Z, corresponds to the number of protons in an atom and is critical in defining an element.

• Each element is identified by its unique chemical symbol on the periodic table (e.g., He symbolizes helium, C symbolizes carbon).

Periodic Table Overview

Elements and Isotopes

• Each element has distinct characteristics, while isotopes of the same element vary in neutron count.

• Example: Neon has isotopes Ne-20, Ne-21, and Ne-22, which differ by their neutron count but share the same number of protons.

Mass Number (A)

• The mass number is the sum of protons and neutrons in an atom’s nucleus, denoted as A (mass number) and Z (atomic number).

• Isotopic notation examples: Carbon-12 can be expressed as C-12 or simply 12C.

Mendeleev's Contributions

Periodic Law

• Mendeleev recognized the periodic patterns of elements based on their atomic masses, leading to the organized structure of the periodic table we use today.

• He predicted the existence of elements such as germanium by identifying gaps in the properties of elements, emphasizing the predictive power of the periodic law.

Classifying Elements

Types

• Metals: Exhibit high conductivity of electricity and heat, are malleable (can be shaped) and ductile (can be stretched).

• Nonmetals: Poor conductors of heat and electricity, typically not malleable; often gain electrons during chemical reactions.

• Metalloids: Exhibit properties of both metals and nonmetals, often functioning as semiconductors in electronic applications.

Groups of Elements

• Noble Gases (Group 8A): Unreactive due to their full valence shell; includes helium, neon, and argon.

• Alkali Metals (Group 1A): Highly reactive metals, characterized by low density and a tendency to lose one electron to form cations; includes lithium, sodium, and potassium.

• Alkaline Earth Metals (Group 2A): Reactive but less so than alkali metals; includes calcium, which is essential for biological functions.

Chemical Ions

Formation of Ions

• In ionic bonding, metals tend to lose electrons, forming cations (positively charged ions), while nonmetals typically gain electrons, forming anions (negatively charged ions). This transfer of electrons facilitates the formation of various compounds.

Atomic Mass and Molar Mass

Definitions

• Atomic mass: The average mass of an element's isotopes, which is displayed on the periodic table and is expressed in atomic mass units (amu).

• Molar mass: The mass of one mole of an element, numerically equal to its atomic mass in grams per mole (g/mol), allowing for the conversion between atomic scale and macroscopic quantities.

The Mole Concept

Understanding Moles

• A mole is defined as exactly 6.022 x 10²³ particles (Avogadro's number), facilitating the understanding of quantities in chemistry.

• The mole concept is critical for converting between masses and number of atoms or molecules in reactions and calculations.

Conversion Between Mass and Moles

Calculating Atom Counts

• To determine how many atoms are present in a sample, convert the mass of the sample to moles using its molar mass, and then convert moles to atoms using Avogadro's number.

Conclusion

• A comprehensive understanding of atomic structure, atomic theory, periodic table organization, and the concept of moles is essential in chemistry, enabling chemists to quantify and comprehend the properties and reactions of materials. This knowledge serves as the backbone of chemical sciences, influencing both theoretical studies and practical applications.