chap 2

Chapter 2: Atomic Structure and the Periodic Table

Chapter Objectives (1 of 2)

2.1 Locate elements in the periodic table on the basis of their group and period designations, which are based on shared properties and electron configurations. (Section 2.1)

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2.2 Describe the charge, relative mass, and location of the three subatomic particles that constitute an atom: electrons, protons, and neutrons. (Section 2.2)

2.3 Write the atomic symbol for a given set of subatomic particles, including the mass number and atomic number. (Section 2.3)

2.4 Define the terms isotope and atomic weight, emphasizing the significance of isotopes in both chemical behavior and atomic mass calculations. (Section 2.4)

2.5 Write balanced equations for radioactive decay and other nuclear processes, outlining the conservation of mass and charge in these transformations. (Section 2.5)

Chapter Objectives (2 of 2)

2.6 Solve problems using the half-life concept to determine the remaining quantity of a radioactive substance after a given period. (Section 2.5)

2.7 Describe the applications of radiation in health and medicine, including diagnostic imaging and cancer treatments. (Section 2.5)

2.8 Write correct electron configurations for each element through atomic number 56, which reflects the distribution of electrons in atomic orbitals and relates to chemical properties. (Section 2.6)

2.9 Describe periodic trends in atomic radius, first ionization energy, and electronegativity, helping to explain the behavior of elements in reactions and bonding. (Section 2.7)

Organization of the Periodic Table

Periodic Table Development

The periodic table was first created by Dmitri Mendeleev, who organized the elements based on their atomic mass and chemical properties, predicting the existence and properties of undiscovered elements.

Key Terms

• Elemental Symbol: A one- or two-letter designation for an element, often derived from its English or Latin name (e.g., H for Hydrogen, He for Helium).

• Group: A vertical column in the periodic table that contains elements with similar chemical properties and the same number of valence electrons.

• Period: A horizontal row in the periodic table, indicating the number of electron shells in the element's atoms.

Categories of the Periodic Table

Categories of the Periodic Table (1 of 3)

• Metals: Located in the left two-thirds of the periodic table, metals are characterized by their high thermal and electrical conductivity, malleability (ability to be shaped), ductility (ability to be stretched), and shiny appearance or luster.

• Nonmetals: Found in the right one-third of the periodic table, nonmetals can be gases, liquids, or brittle solids and exhibit properties that are generally opposite to those of metals, such as poor electrical conductivity.

• Metalloids: Positioned between metals and nonmetals on the periodic table, metalloids possess intermediate properties and are often semiconductors, which makes them useful in electronics.

Categories of the Periodic Table (2 of 3)

• Representative Elements: Groups 1-2 and 13-18, these elements are characterized by having a wide range of properties and are involved in key chemical reactions.

• Transition Metals: Comprising Groups 3-12, these metals usually have multiple oxidation states and form colored compounds.

• Inner-transition Metals: This category includes the Lanthanides (elements 57-71) and the Actinides (elements 89-103), known for their complex electron configurations and radioactive properties.

Specific Families of the Periodic Table

Elements can be classified into specific families based on similar properties that influence chemical relationships. Examples include:

• Neon (Ne): Noble gas in a group marked by inertness (Group 18).

• Cesium (Cs): An alkali metal known for its reactivity and found in Group 1.

• Gold (Au): A transition metal in Group 11, valued for its rarity and use in electronics and jewelry.

• Germanium (Ge): A metalloid located in Group 14, important for semiconductor technology.

Subatomic Particles (Section 2.2)

Characteristics of Subatomic Particles

• Electron

  • Symbol: e−

  • Charge: -1

  • Mass: 9.07 × 10^-28 g

  • Location: Found in atomic orbitals surrounding the nucleus, electrons are responsible for the chemical properties of elements.

• Proton

  • Symbol: p

  • Charge: +1

  • Mass: 1.67 × 10^-24 g

  • Location: Located in the atomic nucleus, protons define the identity of an element (atomic number).

• Neutron

  • Symbol: n

  • Charge: 0 (neutral)

  • Mass: 1.67 × 10^-24 g

  • Location: Also found in the nucleus, neutrons contribute to the mass of an atom and influence nuclear stability.

Historical Development of Atomic Theory

• John Dalton (1766-1844): Proposed that atoms are indivisible, cannot be created or destroyed, and that atoms of different elements vary in mass and properties.

• J.J. Thomson: Developed the plum pudding model, illustrating that atoms contain negatively charged electrons embedded within a positively charged matrix, leading to the discovery of the electron.

• Ernest Rutherford: Conducted the gold foil experiment, demonstrating that atoms are mostly empty space and proposing the existence of a dense nucleus containing protons.

Atomic Symbols (Section 2.3)

Definitions

• Atomic Symbol: Represents both the name and the number of subatomic particles (protons and neutrons) in an element, essential for identifying elements in chemical reactions.

• Atomic Number (Z): The number of protons in an atom's nucleus, defining the element and its position in the periodic table.

• Mass Number (A): The total number of protons and neutrons in the nucleus; important for identifying isotopes.

Isotopes and Atomic Weight (Section 2.4)

Isotopes

• Isotopes: Atoms with the same atomic number (same element) but different mass numbers due to differing numbers of neutrons, affecting their nuclear stability and radioactive properties.

  • Example: Hydrogen has three isotopes: Protium (one proton, no neutrons), Deuterium (one proton, one neutron), and Tritium (one proton, two neutrons).

Atomic Weight

• Atomic Weight: Calculated as the weighted average of the mass numbers of all naturally occurring isotopes of an element, accounting for their relative abundances.

• Atomic Mass Unit (amu): A unit for expressing atomic and molecular weights, where 1 amu is defined as one twelfth of the mass of a carbon-12 atom (approximately 1.661 × 10^-24 g).

Radioactive Nuclei (Section 2.5)

Types of Nuclear Radiation

• Alpha Radiation: Consists of helium nuclei, each containing two protons and two neutrons, ejected from unstable nuclei during radioactive decay.

• Beta Radiation: Involves the emission of electrons from the nucleus created during the transformation of a neutron into a proton, resulting in an increase in atomic number.

• Gamma Radiation: High-energy electromagnetic radiation emitted from a nucleus, often accompanying alpha and beta decay, with no mass or charge.

• Positron Emission: A form of beta decay where a positron (positive electron) is emitted, decreasing the proton count in the nucleus.

Nuclear Reactions

• Nuclear Reaction: A process in which two nuclei collide, leading to the formation of new elements or the transformation of existing elements and the release of energy.

• Reactants and Products: Reactants are the substances that begin a reaction while products are formed as a result of the reaction, essential for writing balanced nuclear equations.

• Daughter Nuclei: The stable end products resulting from the decay of radioactive parent nuclei, which can further decay into other elements or isotopes.

Radiation in Health and Medicine

Applications

• Radiation is instrumental in modern diagnostics (e.g., X-rays, CT scans) and therapeutics (e.g., targeted radiation therapy for cancer).

• Tracer Studies: Radioactive tracers are used in medical diagnostics to observe physiological processes, allowing for non-invasive imaging.

• Dosage Calculations: Essential for ensuring safety and efficacy in radiation therapy, taking into account the type of radiation and individual patient factors.

Trends Within the Periodic Table (Section 2.7)

Key Trends

• Atomic Radius: The distance from the nucleus to the outermost electrons; increases down a group (due to added shells) and decreases across a period (due to increased nuclear charge pulling electrons closer).

• First Ionization Energy: The energy required to remove the outermost electron; increases up a group (due to decreasing atomic radius) and across a period (due to increased nuclear charge).

• Electronegativity: A measure of the tendency of an atom to attract electrons; increases with atomic number up a group and across a period due to effective nuclear charge.

Summary of Periodic Trends

Understanding these periodic trends helps predict how elements will interact in chemical reactions, guiding the formation of compounds and biological molecules.

Additional Topics

Electron Configurations

The arrangement of electrons in the atomic orbitals reveals how they occupy energy levels and subshells, which is crucial for chemical bonding and reactivity.

Summary of Notable Properties

The periodic table provides a systematic framework for examining similarities and differences among groups of elements, facilitating predictions about their behaviors.

Types of Radioactive Isotopes

Methods of using radioactive isotopes in medical applications, including targeting cancerous tissues and studying biological processes, are critical for advancing healthcare and understanding fundamental principles of nuclear chemistry.