General, Organic, and Biological Chemistry: Chapter 4 - Atoms and Elements

Elements and Chemical Symbols

• Definition of Elements

  • Elements are considered pure substances from which all other things in the universe are built.

  • They consist of a single type of atom and cannot be broken down into simpler substances by chemical means.

  • Elements are systematically listed on the inside front cover of chemistry textbooks.

  • Examples include:

    • Gold

    • Carbon

    • Aluminum

• Origins of Element Names

  • Element names are derived from various sources, including:

    • Planets: e.g., Mercury, Uranium.

    • Mythological figures: e.g., Titanium (Titans).

    • Minerals: e.g., Boron from borax.

    • Colors: e.g., Chlorine (from the Greek word for greenish-yellow).

    • Geographic locations: e.g., Polonium (Poland), Francium (France).

    • Famous people: e.g., Curium (Marie and Pierre Curie), Einsteinium (Albert Einstein).

• Chemical Symbols

  • A chemical symbol represents the name of an element.

  • It consists of one or two letters.

  • The first letter is always capitalized. If there is a second letter, it must be lowercase.

  • Examples of 1-Letter Symbols:

    • Carbon: $C$

    • Nitrogen: $N$

    • Fluorine: $F$

    • Oxygen: $O$

  • Examples of 2-Letter Symbols:

    • Cobalt: $Co$

    • Calcium: $Ca$

    • Aluminum: $Al$

    • Magnesium: $Mg$

• Symbols Derived from Latin Names

  • Some symbols are based on ancient Latin names rather than modern English:

    • Copper: $Cu$ (from cuprum)

    • Gold: $Au$ (from aurum)

    • Iron: $Fe$ (from ferrum)

    • Silver: $Ag$ (from argentum)

Chemistry Link to Health: The Toxicity of Mercury (Hg)

• Physical Properties

  • Mercury ($Hg$) is a silvery, shiny element.

  • It is unique for being a liquid at room temperature.

• Modes of Entry into the Body

  • Inhalation (breathing in mercury vapor).

  • Contact with the skin.

  • Ingestion of contaminated water or food.

• Biological Impact and Toxicity

  • Mercury is highly toxic to both the central and peripheral nervous systems.

  • Once inside the body, it destroys proteins and disrupts essential cell functions.

• Long-Term Exposure Consequences

  • Permanent damage to the brain and kidneys.

  • Development of mental retardation.

  • Significant decreases in physical development.

• Sources of Contamination

  • Industrial wastes released into the environment.

  • Accumulation in fish and seafood.

  • Improper disposal of batteries.

  • Compact fluorescent bulbs (CFLs).

The Periodic Table: Organization and Groups

• General Organization

  • The periodic table organizes currently known $118$ elements.

  • Elements are placed in order of increasing atomic mass and arranged into groups with similar properties.

• Structure of the Table

  • Groups: Vertical columns of elements that share similar physical and chemical properties.

  • Periods: Horizontal rows of elements. There are $7$ periods in total, counted from the top to the bottom.

• Group Numbering Systems

  • Standard System: Numbers at the top of vertical columns.

  • Letter System:

    • Letter "A" is used for representative elements (Groups $1A$ through $8A$).

    • Letter "B" is used for transition elements (Groups $3B$ through $12B$).

  • Alternative System: Uses simple numbers from $1$ to $18$ from left to right across the table.

• Specific Group Names

  • Alkali Metals (Group 1A/1): Highly reactive metals. Includes Lithium ($Li$), Sodium ($Na$), Potassium ($K$), Rubidium ($Rb$), and Cesium ($Cs$).

  • Alkaline Earth Metals (Group 2A/2): Shiny metals, though less reactive than alkali metals. Includes Beryllium ($Be$), Magnesium ($Mg$), Calcium ($Ca$), Strontium ($Sr$), Barium ($Ba$), and Radium ($Ra$).

  • Halogens (Group 7A/17): Highly reactive nonmetals. Includes Fluorine ($F$), Chlorine ($Cl$), Bromine ($Br$), and Iodine ($I$).

  • Noble Gases (Group 8A/18): Unreactive gases found on the far right of the table.

Classification of Elements: Metals, Nonmetals, and Metalloids

• Physical Separation

  • A heavy zigzag line on the periodic table separates metals from nonmetals.

• Metals

  • Location: To the left of the zigzag line (excluding Hydrogen).

  • Characteristics: Shiny, ductile (can be drawn into wires), and are excellent conductors of heat and electricity.

  • State: Solids at room temperature, except for Mercury ($Hg$), which is a liquid.

• Nonmetals

  • Location: To the right of the zigzag line.

  • Characteristics: Dull, brittle, and poor conductors of heat and electricity (though often good insulators).

  • Properties: Typically have low densities and low melting points.

• Metalloids

  • Location: Situated along the heavy zigzag line.

  • Characteristics: Exhibit properties of both metals and nonmetals.

  • Utility: Better conductors than nonmetals but not as good as metals. They are used extensively as semiconductors and insulators because they can be modified to function as either.

Chemistry Link to Health: Elements Essential to Human Life

• Key Statistics

  • Out of all known elements, $20$ are essential for the survival and well-being of the human body.

  • Four primary elements—Oxygen ($O$), Carbon ($C$), Hydrogen ($H$), and Nitrogen ($N$)—comprise $96\%$ of total body mass.

  • Water ($H_2O$) accounts for $55\%$ to $60\%$ of body mass, which is where most Hydrogen and Oxygen are located.

• Macrominerals

  • These include Calcium ($Ca$), Phosphorus ($P$), Potassium ($K$), Chlorine ($Cl$), Sulfur ($S$), Sodium ($Na$), and Magnesium ($Mg$).

  • Functions:

    • Formation and maintenance of bones and teeth.

    • Maintenance of the heart, blood vessels, muscle contraction, and nerve impulses.

    • Acid-base balance of body fluids.

    • Regulation of cellular metabolism.

The Atom and Atomic Theory

• Definition of the Atom

  • The smallest particle of an element that retains the characteristics and properties of that element.

• Dalton’s Atomic Theory

  • Atoms are tiny particles of matter.

  • Atoms of a specific element are similar to each other but entirely different from atoms of other elements.

  • Atoms of two or more different elements combine in specific ratios to form compounds.

  • Chemical reactions involve the rearrangement of atoms into new combinations; atoms are never created or destroyed during a chemical reaction.

• Subatomic Particles

  • Protons: Particles with a positive ($+$) charge.

  • Electrons: Particles with a negative ($-$) charge.

  • Neutrons: Particles with no charge (neutral).

• Historical Models of the Atom

  • Thomson’s "Plum-Pudding" Model: Proposed that protons and electrons were randomly distributed within a positively charged cloud.

  • Rutherford’s Gold Foil Experiment: Determined that atoms contain a small, dense, positively charged nucleus at the center that deflects positive particles.

• Structure of the Atom

  • Nucleus: Located in the center. Contains protons and neutrons. It represents the vast majority of the atom's mass but occupies a tiny volume.

  • Electron Space: Electrons occupy a large, mostly empty space surrounding the nucleus, accounting for the atom's volume.

• Mass of the Atom

  • Units: Atomic Mass Unit ($amu$) or Dalton ($Da$).

  • $1\,\text{amu} = 1/12$ of the mass of a Carbon-12 ($C-12$) atom.

  • Protons and neutrons each have a mass of approximately $1\,\text{amu}$.

  • Electrons have a mass so small that it is negligible and not included in calculations of an atom's mass.

Atomic Number and Mass Number

• Atomic Number ($Z$)

  • A whole number specific to each element.

  • Equal to the number of protons in an atom.

  • Appears above the element's symbol on the periodic table.

  • Defines the identity of the element (e.g., every Hydrogen atom has $1$ proton; every Copper atom has $29$).

• Neutrality of Atoms

  • For a neutral atom, the net charge is zero.

  • $\text{Number of Protons} = \text{Number of Electrons}$

  • Example: Aluminum ($Al$) has $13$ protons and $13$ electrons ($13+ + 13- = 0$).

• Mass Number ($A$)

  • Represents the total number of particles (protons and neutrons) in the nucleus.

  • Always a whole number.

  • Does not appear on the periodic table.

  • Formula: $\text{Mass Number} = \text{Protons} + \text{Neutrons}$

• Calculating Neutrons

  • $\text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number}$

  • Example Calculation (Lead, Pb):

    • Mass Number = $207$

    • Atomic Number = $82$

    • Neutrons = $207 - 82 = 125$

Isotopes and Atomic Mass

• Isotopes

  • Atoms of the same element that have the same number of protons but different numbers of neutrons.

  • They have different mass numbers.

  • They are distinguished using atomic symbols: ${}^{A}_{Z}X$.

    • $A$: Mass Number (upper left).

    • $Z$: Atomic Number (lower left).

• Atomic Mass

  • The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of that element.

  • Calculation Method:

    1. Multiply the mass of each naturally occurring isotope by its fractional abundance ($\%\text{ abundance} / 100$).

    2. Sum the results for all isotopes.

  • Example: Gallium (Ga):

    • Isotope 1: $60.10\%$ of $^{69}\text{Ga}$ (mass $68.926\,\text{amu}$). Contribution = $68.926 \times 0.6010 = 41.42\,\text{amu}$.

    • Isotope 2: $39.90\%$ of $^{71}\text{Ga}$ (mass $70.925\,\text{amu}$). Contribution = $70.925 \times 0.3990 = 28.30\,\text{amu}$.

    • Atomic Mass = $41.42 + 28.30 = 69.72\,\text{amu}$.

Electron Energy Levels

• Electromagnetic Radiation

  • Energy that moves as waves.

  • Wavelength: The distance between the peaks of the waves.

  • Relationship: High-energy radiation has shorter wavelengths; low-energy radiation has longer wavelengths.

  • Visible range: $700\text{--}400\,nm$.

• Atomic Spectrum

  • When light from a heated element passes through a prism, it separates into distinct lines of color called an atomic spectrum.

  • Each element possesses a unique atomic spectrum.

• Energy Levels (n)

  • Electrons possess specific, quantized energy levels assigned principal quantum numbers ($n = 1, 2, 3, \dots$).

  • As $n$ increases, the energy of the electron increases, and it is located farther from the nucleus.

  • Electrons move to a higher energy level by absorbing energy.

  • Electrons emit a photon of light when they fall back to a lower energy level.

• Sublevels and Orbitals

  • Energy levels contain sublevels: $s, p, d, f$.

  • The number of sublevels in a level is equal to $n$ (e.g., $n=2$ has $2$ sublevels).

  • Energy order: s < p < d < f.

  • Orbital Capacities:

    • $s$ sublevel: $1$ orbital, max $2$ electrons.

    • $p$ sublevel: $3$ orbitals, max $6$ electrons.

    • $d$ sublevel: $5$ orbitals, max $10$ electrons.

    • $f$ sublevel: $7$ orbitals, max $14$ electrons.

• Pauli Exclusion Principle

  • An orbital can hold a maximum of two electrons.

  • If two electrons are in the same orbital, they must have opposite magnetic spins (represented by up and down arrows).

Electron Configurations and Orbital Diagrams

• Orbital Diagrams

  • Boxes represent orbitals; arrows represent electrons and their spins.

  • Orbitals within a sublevel are filled one at a time with single electrons before pairing begins.

• Electron Configuration Notation

  • Shows the placement of electrons in order of increasing energy level and sublevel.

  • Abbreviated Configuration: Uses the symbol of the preceding noble gas in brackets to represent core electrons (e.g., $[He]$ or $[Ne]$).

  • Representative Blocks:

    • s block: Groups $1A$ and $2A$.

    • p block: Groups $3A$ through $8A$.

    • d block: Transition elements (occurs after Calcium, $n-1$ level).

    • f block: Inner transition elements (bottom two rows).

Trends in Periodic Properties

• Valence Electrons

  • Electrons located in the outermost energy level.

  • Responsible for the chemical properties of an element.

  • For representative elements, the Group Number ($1A–8A$) indicates the number of valence electrons.

• Lewis Symbols (Electron-Dot Symbols)

  • Dots represent valence electrons placed around the element's chemical symbol.

  • $1$ to $4$ electrons are single dots; $5$ to $8$ include pairs.

• Atomic Size

  • Determined by the distance between the nucleus and the outermost electrons.

  • Trend: Increases from top to bottom within a group.

  • Trend: Decreases from left to right across a period (due to increased nucleus pull from more protons).

• Ionization Energy

  • The energy required to remove one of the outermost electrons ($\text{Na(g)} + \text{energy} \rightarrow \text{Na}^+(g) + e^-$).

  • Trend: Decreases going down a group (electrons further from nucleus).

  • Trend: Increases going across a period from left to right.

• Metallic Character

  • The tendency of an element to lose valence electrons easily.

  • Trend: Increases going down a group.

  • Trend: Decreases going across a period from left to right.