2.1

Properties of Atoms

• Nature's materials consist of fundamental substances called elements combined in various ways.

• Historical progression of element definition:

  • From the seventeenth century through the end of the nineteenth century, elements were defined as pure substances that could not be broken down further by the methods of chemistry.

  • It was later recognized that each element contains only one type of atom, which serves as the basic unit of matter.

  • By $1850$, approximately $60$ elements had been identified, including common elements such as oxygen, copper, gold, and sodium.

  • Currently, $118$ elements are known:

  • $94$ occur naturally.

  • $24$ have been created artificially in laboratory settings.

• Chemical Symbols: Elements are identified by a chemical symbol consisting of a one- or two-letter abbreviation:

  • Carbon is represented by $\text{C}$.

  • Hydrogen is represented by $\text{H}$.

  • Helium is represented by $\text{He}$.

Subatomic Particles and Atomic Structure

• Atoms are composed of three primary types of subatomic particles:

  • Nucleus: A dense central core composed of protons and neutrons.

  • Protons: Positively charged particles located within the nucleus, each assigned a mass of $1$.

  • Neutrons: Electrically neutral particles located within the nucleus, each assigned a mass of $1$.

  • Electrons: Negatively charged particles that move around the nucleus at a distance and possess negligible mass.

Atomic Parameters and Variations

• Atomic Number: The total number of protons in an atom's nucleus. The atomic number specifies an atom as a particular element:

  • Hydrogen has an atomic number of $1$ ($1$ proton).

  • Carbon has an atomic number of $6$ ($6$ protons).

• Atomic Mass: The combined mass of an atom's protons and neutrons.

• Isotopes: Atoms of the same element that contain equal numbers of protons but differing numbers of neutrons, changing the atomic mass.

  • Isotopes of Carbon:

  • Carbon-12 (${}^{12}\text{C}$): Comprises approximately $99\%$ of carbon atoms. Contains $6$ protons and $6$ neutrons (atomic mass = $12$).

  • Carbon-13 (${}^{13}\text{C}$): Comprises approximately $1\%$ of carbon atoms. Contains $6$ protons and $7$ neutrons (atomic mass = $13$).

  • Carbon-14 (${}^{14}\text{C}$): Found in a very small fraction of carbon atoms. Contains $6$ protons and $8$ neutrons (atomic mass = $14$).

  • Chemical notation represents atomic mass as a superscript to the left of the element symbol (e.g., ${}^{12}\text{C}$).

Electrical Charge and Ions

• Neutral Atoms: Under typical conditions, an atom possesses an equal number of protons and electrons. For example, a carbon atom has $6$ protons and $6$ electrons, so the positive and negative charges cancel each other out, making the atom electrically neutral.

• Ions: Electrically charged atoms resulting from chemical processes where an atom gains or loses electrons:

  • Positively Charged Ion: Formed when an atom loses an electron.

  • Negatively Charged Ion: Formed when an atom gains an electron.

  • Notation: Charge is specified as a superscript to the right of the chemical symbol. For example, $\text{H}^+$ indicates a hydrogen ion that has lost an electron and carries a positive charge.

Orbital Structure and Electron Energy Levels

• Orbital: A region of space where an electron is present most of the time. The exact path or location of an electron at any instant cannot be known or predicted.

  • Hydrogen Orbital: Consists of a single spherical orbital occupied by one electron.

  • Orbital Visualization: Visualized as a cloud of points that is denser where the electron is more likely to be found.

  • Orbital Capacity: The maximum number of electrons in any single orbital is $2$

Energy Levels and Distance

• Orbitals differ in size and shape depending on distance from the nucleus.

• Energy Relationship: Electrons in orbitals close to the nucleus possess less energy than electrons in orbitals farther away.

  • Bicycle Elevation Analogy: Biking up a hill expends energy that becomes stored in the increased elevation of the bike and rider. Some of that stored energy is released when coasting down the hill. Elevation corresponds to an electron's distance from the nucleus. An electron gives up energy when it moves to a closer orbital, which explains why electrons fill up orbitals close to the nucleus before occupying those farther away.

Electron Shells and Configurations

• Shells (Energy Levels): Groups of orbitals existing at a given energy level.

  • First Shell: Consists of a single spherical orbital holding up to $2$ electrons (lowest energy level).

  • Second Shell: Consists of four orbitals holding up to $8$ total electrons:

  • One spherical orbital (larger in diameter than the first shell's orbital, possessing slightly less energy than the dumbbell orbitals at this level).

  • Three dumbbell-shaped orbitals.

• Carbon Electron Configuration (total $6$ electrons):

  • First Energy Level (1st shell): $2$ electrons occupy the small spherical orbital.

  • Second Energy Level (2nd shell): $4$ electrons are distributed among four possible orbitals:

  • $2$ electrons in the outermost spherical orbital.

  • $1$ electron in each of two dumbbell-shaped orbitals.

  • $1$ dumbbell-shaped orbital is empty.

  • Because a full set at this energy level contains $8$ electrons, carbon would require a total of $4$ additional electrons to completely fill all orbitals in its second shell.

Organization of the Periodic Table of Elements

• Periodic Table: A tabular arrangement of chemical elements organized in order of increasing atomic number (number of protons). Its development is credited to the nineteenth-century Russian chemist Dmitri Mendeleev to organize elements by chemical properties.

Rows (Periods)

• Across any horizontal row, each element has one more proton and one more electron than the element to its left.

• For the first three horizontal rows, elements in the same row share the same number of shells, meaning they have the same number and types of orbitals available to be filled by electrons.

• Electrons fill the shell across a row until a full complement of electrons is reached at the right-hand side.

• Row 2 Progression: Begins with Lithium ($\text{Li}$, $3$ protons) and ends with Neon ($\text{Ne}$, $10$ protons), achieving a complete complement of $8$ electrons in its outer shell.

Columns (Groups or Families)

• Vertical columns are called groups or families.

• Members of a group all have the same number of electrons in their outermost shell.

  • Example: Carbon ($\text{C}$) and Lead ($\text{Pb}$) both have $4$ electrons in their outermost shell.

• The number of outer shell electrons determines in large part how elements interact with other elements to form diverse molecules.

Molecules and Chemical Bonds

• Molecule: Groups of two or more atoms attached together that act as a single unit.

• Chemical Bond: A form of attraction between atoms that holds them together when they combine.

Valence Electrons and Covalent Bonds

• Valence Electrons: Electrons located in the outermost orbitals farthest from the nucleus, existing at the highest energy level of the atom.

• Covalent Bond: Formed when two atoms share valence electrons.

  • Mechanism: When outermost orbitals of two atoms come into proximity, two atomic orbitals (each containing $1$ electron) merge into a single molecular orbital containing a full complement of $2$ electrons.

  • Chemical Formula: Written as the letter abbreviation for each element followed by a subscript giving the number of that type of atom in the molecule (e.g., $\text{H}_2$ for hydrogen gas).

  • Structural Formula: A covalent bond between atoms is denoted by a single line connecting the two chemical symbols (e.g., $\text{H}-\text{H}$).

• Double Bond: Formed when two adjacent atoms share two pairs of electrons, denoted by a double line connecting chemical symbols. In this process, four orbitals (each occupied by a single electron) merge to form two molecular orbitals.

The Octet Rule and Molecular Stability

• Molecules tend to be most stable when the two atoms forming a bond share enough electrons to fill the outermost shell.

  • Outer shell electron capacity:

  • Hydrogen outer shell: Holds $2$ electrons.

  • Carbon, Nitrogen, and Oxygen outer shell: Holds $8$ electrons.

• Octet Rule: The tendency of elements to prefer eight electrons in their outermost shell to achieve stability.

• Bonding Examples:

  • Carbon ($4$ valence electrons) combines with $4$ hydrogen atoms ($1$ valence electron each) to form methane ($\text{CH}_4$).

  • Nitrogen ($5$ valence electrons) combines with $3$ hydrogen atoms to fulfill shell requirements.

Self-Assessment Questions & Core Concepts Review

• Question 1: In the early 1900s, Ernest Rutherford produced a beam of very small positive particles and directed it at a thin piece of gold foil just a few atoms thick. Most of the particles passed through the foil without changing their path; very rarely, a particle was deflected. What conclusions can you draw from this experiment about the structure of an atom?

  • Conclusion: The fact that most positive particles pass through unimpeded demonstrates that an atom is composed mostly of empty space. The rare deflection of positive particles indicates that the positive charge and the vast majority of the atom's mass are concentrated in a extremely small, dense central area called the nucleus.

• Question 2: What are atoms made up of? Describe each component.

  • Components:

  • Protons: Positively charged subatomic particles located in the central nucleus with an atomic mass defined as $1$.

  • Neutrons: Electrically neutral subatomic particles located in the central nucleus with an atomic mass defined as $1$.

  • Electrons: Negatively charged subatomic