Chapter 4:Atoms and Elements

Introductory Chemistry, Seventh Edition

Chapter 4: Atoms and Elements

Experiencing Atoms in the Sea and Mountains (1 of 2)

  • Atoms are the foundation of our sensations.

  • Composition of Seaside Rocks:

    • Typical seaside rocks are made up of silicates, which are compounds composed of silicon and oxygen atoms.

  • Composition of Seaside Air:

    • Contains nitrogen (N) and oxygen (O) molecules.

    • May also contain amines, which are organic compounds characterized by the presence of nitrogen.

  • Focus Chemical:

    • Triethylamine, an amine emitted by decaying fish, is responsible for the fishy smell often associated with seaside environments.

Experiencing Atoms in the Sea and Mountains (2 of 2)

  • Small Size and Large Number of Atoms in a Pebble:

    • Conceptual Analogy: If each atom within a small pebble were scaled to the size of the pebble itself, that pebble would exceed the height of Mount Everest, illustrating the minuscule size of atoms compared to macroscopic objects.

Atoms and Elements (1 of 2)

  • Fundamental Concepts:

    • Atoms compose matter, and their properties directly determine the properties of matter.

    • Definition of Atom:

    • An atom is defined as the smallest identifiable unit of an element.

    • Definition of Element:

    • An element is a pure substance that cannot be broken down into simpler substances through chemical means.

    • Natural Elements:

    • Approximately 91 different elements exist naturally, translating to about 91 distinct types of atoms.

Atoms and Elements (2 of 2)

  • Synthetic Elements:

    • Scientists have successfully created around 20 synthetic elements that do not occur naturally in the environment.

    • Defining Controversy:

    • The exact count of naturally occurring elements is debated, as some elements once thought to be synthetic might occasionally occur in trace amounts in nature.

Historical Perspective on Matter
Democritus and Leucippus: Matter is Made of Particles

  • Timeframe:

    • Active around 460–370 B.C.E., Democritus, alongside his mentor Leucippus, initiated the concept that matter consists of tiny, indestructible particles called atomos (meaning "indivisible").

  • Core Theory:

    • If matter is divided successively, it will eventually lead to the smallest, indivisible particles.

Atomic Theory: John Dalton

  • Year of Formalization:

    • Dalton developed a widely accepted atomic theory in 1808, over 2000 years after Democritus.

  • Main Tenets of Dalton’s Atomic Theory:

    1. Each element is composed of tiny, indestructible particles known as atoms.

    2. All atoms of a specific element possess identical mass and other distinguishing properties when compared to atoms of other elements.

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

Modern Evidence for Atomic Theory
Writing with Atoms

  • Technological Advances:

    • Utilizing a scanning tunneling microscope (STM), scientists at IBM have manipulated individual atoms to create words and images, producing the miniature film: "A Boy and His Atom."

Discovery of Electrons: J. J. Thomson

  • Key Discoveries:

    • J. J. Thomson (1856–1940) identified electrons, the first subatomic particles, with the following characteristics:

    • Electrons carry a negative charge.

    • Electrons are significantly smaller and lighter than atoms.

    • Electrons are uniformly present across various substances.

    • Electron Charge Balancing:

    • Thomson hypothesized that atoms must include a positive charge to counterbalance the negative charge of electrons.

Thomson’s Plum-Pudding Model

  • Conceptual Model:

    • In Thomson's plum-pudding model, electrons (depicted in yellow) exist within a positively charged 'pudding' (depicted in red). This model represents an early theoretical structure of atoms.

Rutherford’s Gold Foil Experiment (1 of 3)

  • Experimental Setup:

    • Conducted in 1909, Rutherford directed alpha particles at a thin sheet of gold foil.

  • Results:

    • While most particles traveled straight through the foil, a small fraction was deflected at sharp angles.

Rutherford’s Gold Foil Experiment (2 of 3)

  • Expected vs. Actual Results:

    • Expected Result (Plum-Pudding Model): If Thomson's model were correct, alpha particles would pass through gold foil with minimal deflection.

    • Actual Result:

    • Predominantly, alpha particles penetrated the foil smoothly,

    • However, a handful were deflected or bounced back, indicating unexpected atomic structure.

Rutherford’s Gold Foil Experiment (3 of 3)

  • Visual Representation:

    • Comparison illustrates the predicted path of alpha particles under the plum-pudding model and their encountered paths in a nuclear model, yielding new insights regarding atomic structure.

Rutherford: Nuclear Theory of the Atom (1 of 4)

  • Paradigm Shift:

    • Most of an atom's mass and all its positive charge reside in a tiny core termed the nucleus.

    • The volume of the atom largely consists of empty space, occupied by negatively charged electrons.

    • Balance of Charge:

    • The number of electrons surrounding the nucleus equals that of positively charged particles (protons) in the nucleus, resulting in electrical neutrality overall.

Rutherford: Nuclear Theory of the Atom (2 of 4)

  • Distribution of Mass:

    • Concentration of mass resides in the nucleus, with electrons dispersed throughout a broader region.

  • Visuals give a metaphorical representation of the nucleus and electrons within an atom.

Rutherford: Nuclear Theory of the Atom (3 of 4)

  • Density Insights:

    • The dense nucleus constitutes over 99.9% of an atom's mass while occupying only a minor fraction of its volume.

    • Electrons, although occupying a greater volume, contribute minimally to the mass.

    • Conclusion: Matter exhibits less uniformity than appearances suggest.

Rutherford: Nuclear Theory of the Atom (4 of 4)

  • Incredible Density Conversion:

    • For instance, a singular grain of sand made of solid atomic nuclei would have a mass of approximately 5 million kilograms.

    • Related Phenomenon: Astronomers regard black holes and neutron stars as exemplars of this extreme density characteristic.

Protons, Neutrons, and Electrons: Relative Mass

  • Comparative Masses:

    • Scenario: If a proton's mass equals that of a baseball, an electron’s mass compares to that of a rice grain.

    • Proton Mass: Nearly 2000 times greater than that of an electron.

    • Neutrons possess a mass similar to protons while electrons remain nearly negligible in mass.

Protons, Neutrons, and Electrons: Electrical Charge (1 of 2)

  • Charge Properties:

    • Electrical charge serves as a fundamental property innate to protons and electrons.

    • Attraction and Repulsion:

    • Opposites attract: positive and negative charges.

    • Like charges repel: positive-positive or negative-negative combinations.

    • Neutralization:

    • A pair of one proton and one electron will have a full charge cancellation (0).

Protons, Neutrons, and Electrons: Electrical Charge (2 of 2)

  • Models of Charge Dynamics:

    • Visual representations depict the interaction between positive (red) and negative (yellow) charges, demonstrating attraction and repulsion phenomena.

Protons, Neutrons, and Electrons: Summary
Table 4.1: Subatomic Particles

Subatomic Particle

Mass (kilogram)

Mass (amu)

Charge

Proton

1.67262imes10−271.67262 imes 10^{-27}

1.0073

1+

Neutron

1.67493imes10−271.67493 imes 10^{-27}

1.0087

0

Electron

0.00091imes10−270.00091 imes 10^{-27}

0.00055

1−

Electrical Storm Provides Evidence of Charge in Matter

  • Neutrality in Matter:

    • Typically, matter exhibits charge neutrality with balanced counts of positive and negative charges.

  • Lightning Phenomenon:

    • During electrical storms, negative charge accumulates on clouds, whereas the ground becomes positively charged.

    • This uneven charge distribution leads to significant and dramatic rebalancing, often witnessed as lightning.

Elements: Defined By Their Number of Protons (1 of 2)

  • Unique Elemental Identity:

    • Elements are defined by the quantity of protons they possess.

    • The number of protons in an atomic nucleus is what categorizes the atom as a specific element.

    • Altering the proton count would lead to the identification of a different element.

    • Symbolic Representation:

    • The number of protons is referred to as atomic number (Z).

Elements: Defined By Their Number of Protons (2 of 2)

  • Comparisons:

    • Helium nucleus contains 2 protons represented as He.

    • Aluminum nucleus counts 13 protons represented as Al.

Periodic Table of Elements: Organized by Atomic Number

  • Structure:

    • The periodic table exemplifies the arrangement of all known elements by increasing atomic numbers.

Periodic Table: Names and Symbols of the Elements (1 of 2)

  • Symbolization:

    • Chemical symbols predominantly derive from modern English names, e.g., C for carbon, Br for bromine.

    • Exceptions involve symbols based on Latin, such as:

    • Potassium (K) from kalium.

    • Sodium (Na) from natrium.

Periodic Table: Names and Symbols of the Elements (2 of 2)

  • Additional Element Examples:

    • Lead: Pb (plumbum)

    • Mercury: Hg (hydrargyrum)

    • Iron: Fe (ferrum)

    • Silver: Ag (argentum)

    • Tin: Sn (stannum)

    • Copper: Cu (cuprum)

Periodic Table: Origins of the Names of the Elements (1 of 3)

  • Naming Origins:

    • Early chemists often named new elements based on their properties, for instance:

    • Argon: named from Greek argos, meaning “inactive.”

    • Bromine: derived from bromos, meaning “stench.”

Periodic Table: Origins of the Names of the Elements (2 of 3)

  • Geographical Naming:

    • Some elements were named after countries:

    • Polonium after Poland

    • Francium after France

    • Americium after the USA

Periodic Table: Origins of the Names of the Elements (3 of 3)

  • Honorary Naming:

    • Elements named after notable scientists:

    • Curium is named in honor of Marie Curie, recognized for her contributions to radioactivity and discovery of new elements.

Periodic Table of Elements

  • Comprehensive Listing:

    • The periodic table encompasses each element's name, symbol, and atomic number.

    • Additional alphabetical listings are found in various sections of the text.

Looking for Patterns: Dmitri Mendeleev

  • Contributions:

    • Dmitri Mendeleev (1834 -1907) organized elements according to increasing relative mass and discovered recurring properties, leading to the formulation of the periodic law.

Looking for Patterns: Recurring Properties

  • Visual Representations:

    • Coloring each element denotes property classifications.

    • Structure alignment allows for elements with similar properties to be positioned vertically (similar to Mendeleev’s original table).

Looking for Patterns: Periodic Law

  • Implementation:

    • Mendeleev’s periodic law captures a multitude of observations but does not provide all underlying explanations; such insights would come from further theories.

Locating Metals, Nonmetals, Metalloids on the Periodic Table

  • Classification Framework:

    • Broad grouping of elements into metals, nonmetals, and metalloids based on characteristics.

Properties of Metals

  • Location and Traits:

    • Metals occupy the left side of the periodic table and share similar characteristics:

    • Good conductors of heat and electricity.

    • Malleability (can be hammered into sheets).

    • Ductility (can be drawn into wires).

    • Lustrous (shiny appearance).

    • Tend to lose electrons during chemical reactions.

  • Universal Examples:

    • Iron (Fe), Magnesium (Mg), Chromium (Cr), Sodium (Na).

Properties of Nonmetals (1 of 2)

  • Location and Traits:

    • Nonmetals are found on the upper right side of the periodic table, separated from metals by a zigzag line descending from boron to astatine.

    • Varied Properties:

    • Some exist as solids at room temperature, others as gases, and one (bromine) is liquid.

    • Tend to be poor conductors of heat and electricity.

Properties of Nonmetals (2 of 2)

  • Chemical Behavior:

    • In reactions, nonmetals typically gain electrons.

  • Common Examples:

    • Oxygen (O), Nitrogen (N), Chlorine (Cl), Bromine (Br), Iodine (I).

Properties of Metalloids (1 of 2)

  • Location and Traits:

    • Metalloids, or semimetals, are positioned along the zigzag separation of metals and nonmetals, exhibiting mixed properties.

  • Semiconductor Characteristics:

    • Their intermediate electrical conductivity is variable and controllable, making them crucial in electronics.

Properties of Metalloids (2 of 2)

  • Examples in Industries:

    • Silicon (Si), Arsenic (As), Germanium (Ge).

    • Noteworthy Application: Silicon is pivotal within computer technology and electronic manufacturing.

Main Group Elements and the Transition Elements

  • Group Characteristics:

    • The periodic table classifies elements into:

    • Main group elements with predictable properties based on their table position; and

    • Transition elements with less predictable properties based on position alone.

Groups in the Periodic Table

  • Column Classification:

    • Columns are identified as groups or families, where elements within a group generally share similar properties and may have group names.

Groups in the Periodic Table: Alkali Metals

  • Properties:

    • The alkali metals group is characterized by high reactivity, including:

    • Lithium, Sodium (exhibiting violent reactions with water), Potassium, Rubidium, and Cesium.

    • Notably, hydrogen is not classified as an alkali metal, despite its position in Group 1A.

Groups in the Periodic Table: Alkaline Earth Metals

  • Reactivity:

    • Alkaline earth metals exhibit reactive tendencies but are not as reactive as alkali metals, comprising:

    • Beryllium, Magnesium (reacting with water), Calcium, Strontium, and Barium.

Groups in the Periodic Table: Halogens

  • Reactivity:

    • Halogens are very reactive nonmetals, including:

    • Chlorine (noted for its greenish-yellow gas state and strong odor), Fluorine, Bromine, Iodine (solid at room temperature), and Astatine.

Groups in the Periodic Table: Noble Gases

  • Chemical Stability:

    • Noble gases are chemically inert, exhibiting stability and non-reactivity, encapsulating:

    • Helium, Neon, Argon, Krypton, and Xenon.

Ions: Gaining and Losing Electrons (1 of 4)

  • Formation of Ions:

    • In a variety of chemical reactions, atoms may gain or lose electrons, resulting in charged particles termed ions.

    • Definitions:

    • Positive ions are designated as cations.

    • Negative ions are known as anions.

  • Charge Notation:

    • The ion charge is denoted in the upper right corner of the symbol, formatted with numerical charge magnitude preceding the sign.

Ions: Gaining and Losing Electrons (2 of 4)

  • Lithium Ion Example:

    • In chemical reactions, lithium atoms lose one electron:

    • Lithium Ion = Symbol: Li+Li^+, contains 3 protons and 2 electrons.

Ions: Gaining and Losing Electrons (3 of 4)

  • Fluoride Ion Example:

    • Fluorine atoms gain one electron, forming:

    • Fluoride Ion = Symbol: F−F^-, contains 9 protons and 10 electrons.

Ions: Gaining and Losing Electrons (4 of 4)

  • Naming Ions:

    • The lithium ion is named lithium cation, while the fluoride ion is named fluoride anion, adopting the stem of the element's name (fluor) plus the suffix “–ide.”

Ions and the Periodic Table (1 of 2)

  • Group Number Association:

    • The number indicated above letter A for each main-group column (1-8) reveals the count of valence electrons for elements within that column.

  • Charge Prediction:

    • Charge acquired by an element is correlated to its position relative to the noble gases in the periodic table.

Ions and the Periodic Table (2 of 2)

  • Charge Trends:

    • Ions may be predicted based on their group number, with main-group elements forming ions that align their valence electrons with those of the nearest noble gas.

Isotopes

  • Definition:

    • Atoms within a single element share the same proton count, yet they might differ in neutron count.

    • Unique atoms with identical proton numbers yet varying neutron counts are termed isotopes.

  • Natural Abundance:

    • Each element has its unique isotopic distribution, signifying unique percent abundances in nature.

Isotopes: Natural Abundance of Isotopes in Neon

  • Neon Isotopes Explained:

    • Naturally occurring neon comprises three isotopes:

    • Ne-20 (10 neutrons);

    • Ne-21 (11 neutrons);

    • Ne-22 (12 neutrons). Each of these isotopes features 10 protons.

Isotopes: Mass Number

  • Definition of Mass Number (A):

    • It signifies the total count of protons (Z) and neutrons within an atom’s nucleus.

    • Neutron count can be calculated as the difference between the atomic number and the mass number.

Isotopes: Symbol Notation

  • Isotope Symbols Representation:

    • Commonly illustrated in the format: (ZAX)(^{A}_{Z}X), where Z represents the atomic number and A denotes the mass number.

    • Example: For Ne-20, the symbol format is (1020Ne)(^{20}_{10}Ne).

Isotopes: Isotope Symbols

  • Alternative Notation for Isotopes:

    • An alternative approach includes the chemical symbol (or name) followed by a hyphen and the mass number; e.g., for neon isotopes:

    • Ne-20, Ne-21, Ne-22.

Calculating Atomic Mass as the Weighted Average (1 of 2)

  • Atomic Mass Definition:

    • Each element's atomic mass listed in the periodic table denotes the average mass of its constituent atoms.

    • Equation for calculation:
      extAtomicMass=ext(fractionabundance)imesext(massofisotope)ext{Atomic Mass} = ext{(fraction abundance)} imes ext{(mass of isotope)}

Calculating Atomic Mass as the Weighted Average (2 of 2)

  • Fraction Abundance Explanation:

    • Defined as the percentage natural abundance divided by 100, facilitating weighted average calculations for isotopes of an element.

  • Example Calculations:

    • Naturally occurring chlorine comprises:

    • 75.77% Cl-35 (mass 34.97 amu)

    • 24.23% Cl-37 (mass 36.97 amu).

Calculating Atomic Mass as the Weighted Average: Sample Problem

  • Gallium Isotope Example:

    • Gallium features two naturally occurring isotopes:

    • Ga-69 (mass 68.9256 amu, natural abundance 60.11%)

    • Ga-71 (mass 70.9247 amu, natural abundance 39.89%).

    • Calculate atomic mass based on these figures.

Calculating Atomic Mass as the Weighted Average: Sample Solution

  • Methodology:

    • Utilize fractional abundances alongside isotopic masses to compute the atomic mass according to previously established definitions.

Radioactive Isotopes

  • Instability and Emission:

    • Certain isotopes possess unstable nuclei, which emit energetic subatomic particles transforming into isotopes of different elements.

    • Radioactive emissions are defined as nuclear radiation.

    • Types of radioactive isotopes exhibit varying lifetimes of radioactivity.

Radioactive Isotopes in the Environment

  • Potential Harms:

    • Nuclear radiation poses risks to living organisms due to interactions that damage biological molecules.

    • Variances in half-lives of radioactive isotopes exist; for instance:

    • Pb-185 exhibits short-lived radiation.

    • Pu-239 remains radioactive over extended periods (thousands to billions of years).

Radioactive Isotopes in Medicine

  • Beneficial Uses:

    • Certain radioactive isotopes offer medical advantages:

    • Technetium-99 (Tc-99) is commonly used in medical diagnostics, assisting doctors in imaging internal organs and identifying diseases.

Review (1 of 5)

  • Atomic Theory History:

    • Ancient Greeks posited that matter comprises small, indestructible particles.

    • Dalton formulated a contemporary atomic theory:

    • Matter consists of atoms.

    • Unique properties distinguish atoms of different elements.

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

Review (2 of 5)

  • Rutherford's Model:

    • Nuclear structure of the atom posits a small, dense nucleus containing all positive charge and most atomic mass, complemented by electrons occupying the majority of atomic volume.

Review (3 of 5)

  • Subatomic Particles:

    • Mass comparisons highlight similarities between protons and neutrons, with each at about 1 amu and insignificant mass attributed to electrons.

    • Charge characteristics:

    • Proton: 1+

    • Electron: 1−

    • Neutron: neutral.

Review (4 of 5)

  • Periodic Table Insights:

    • Organizes elements by ascending atomic number.

    • Elements within columns possess similar properties classified as groups or families, differentiating metals (electron losers) from nonmetals (electron gainers) and metalloids in between.

Review (5 of 5)

  • Ions and Isotopes Overview:

    • Formation of ions occurs via electron gain or loss leading to cations (positive) and anions (negative).

    • Isotopes represent variations of elements differing in neutron count, characterized by their mass number, A, serving as a weighted average reflection of isotopic variances.

Learning Objectives for Chapter 4 (1 of 2)

  • LO: Recognize that all matter comprises atoms.

  • LO: Explain how Thomson and Rutherford's experiments contributed to the nuclear theory of the atom.

  • LO: Describe properties and charges of electrons, neutrons, and protons.

  • LO: Utilize the periodic table to ascertain an element's atomic symbol and atomic number.

  • LO: Classify elements by group utilizing the periodic table.

Learning Objectives for Chapter 4 (2 of 2)

  • LO: Determine an element's ion charge based on proton and electron counts.

  • LO: Identify the number of protons and electrons in a given ion.

  • LO: Establish atomic numbers, mass numbers, and isotope symbols for isotopes.

  • LO: Calculate atomic mass utilizing percent natural abundances alongside isotopic masses.

Atoms are the foundation of our sensations and the building blocks of matter that compose all substances around us. The study of atoms and elements is crucial in understanding the composition and behaviors of different materials.

Composition of Seaside Rocks

  • Typical seaside rocks are predominantly made up of silicates, which are compounds that consist of silicon (Si) and oxygen (O) atoms. Silicates form a significant portion of the Earth's crust and are integral in forming various geological formations.

Composition of Seaside Air

  • The air at seaside locations contains a rich mixture of nitrogen (N) and oxygen (O) molecules, which are essential for life. It may also include organic compounds known as amines, characterized by their containing nitrogen atoms. These compounds can contribute to the unique odors experienced near the coast.

Focus Chemical: Triethylamine

  • Triethylamine, a specific type of amine, is a volatile organic compound that is emitted by decaying organic matter, such as fish. This compound is responsible for the characteristic fishy smell often associated with seaside environments, illustrating how chemical compounds can influence our sensory experiences in specific locations.

Small Size and Large Number of Atoms in a Pebble

  • Conceptual Analogy: If each atom within an ordinary small pebble were scaled to the size of the pebble itself, the height of this enlarged pebble would significantly exceed that of Mount Everest. This analogy helps illustrate the minuscule size of atoms compared to macroscopic objects, reinforcing the idea that matter is made up of an immense number of tiny, indivisible particles.

Atoms and Elements

  • Fundamental Concepts: Atoms are the smallest identifiable units of elements, and their properties directly determine the characteristics of the matter they compose.

  • Definition of Atom: An atom is defined as the smallest identifiable unit of an element, consisting of a nucleus (composed of protons and neutrons) surrounded by electrons.

  • Definition of Element: An element is a pure substance that cannot be broken down into simpler substances through chemical means. Elements are fundamental to chemistry as they are the basic building blocks for all matter.

  • Natural Elements: Approximately 91 different elements exist naturally on Earth, correlating to around 91 distinct types of atoms recognized in the periodic table. Each element has unique characteristics and behaviors that are critical in chemical reactions.

Synthetic Elements

  • Scientists have succeeded in creating around 20 synthetic elements, which do not occur naturally in the environment. The creation of these elements often involves complex nuclear reactions and play crucial roles in advanced scientific research and applications.

  • Defining Controversy: The exact number of naturally occurring elements remains a subject of debate among scientists, as some elements previously thought to be synthetic may occur sporadically in trace amounts in natural environments, further complicating elemental classification.

Historical Perspective on Matter
Democritus and Leucippus: Matter is Made of Particles

  • Timeframe: Active around 460–370 B.C.E., Democritus, alongside his mentor Leucippus, introduced the revolutionary idea that matter consists of tiny, indivisible particles called atomos, meaning "indivisible." This foundational concept laid the groundwork for future atomic theory.

  • Core Theory: They theorized that if matter is successively divided, it would eventually lead to the smallest, indivisible entities, marking a pivotal shift in scientific thought toward understanding matter's composition.

Atomic Theory: John Dalton

  • Year of Formalization: John Dalton developed a widely accepted atomic theory in 1808, more than 2000 years after Democritus, greatly enhancing our understanding of atomic structure and behavior.

  • Main Tenets of Dalton’s Atomic Theory:

    1. Each element is composed of tiny, indestructible particles known as atoms.

    2. All atoms of a specific element have identical mass and other defining properties, distinguishing them from atoms of other elements.

    3. Atoms combine in simple, whole-number ratios to form chemical compounds, foundational to understanding chemical reactions.