MCU General Chemistry Lecture 3: Atoms and the Periodic Table Notes

Atomic-Molecular Theory of Matter

• The Atomic-Molecular Theory of Matter states that all matter is composed of small, fast-moving particles called atoms.

• These atoms have the capacity to join together to form molecules.

• This theory is comprised of thousands of individual theories that collectively provide evidence for the existence and behavior of matter at the atomic level.

Historical Progression of Atomic Models

• Democritus (5th Century B.C.)

  • The word "atom" originates from the Greek word "atomos," which means indivisible.

  • The Greek philosopher Democritus first proposed the idea that all matter is constructed of atoms.

• Dalton’s Model

  • Based on experimental results, Dalton developed a theory regarding the structure of matter consisting of four main concepts:

    1. All matter is composed of tiny, indivisible particles called atoms.

    2. Atoms of each specific element are exactly alike and possess the same mass.

    3. An atom of one element cannot be changed into an atom of a different element.

    4. Atoms of different elements can join together to form compounds.

• Thomson’s Model (End of 1800s)

  • Thomson discovered that atoms were not simple, solid spheres.

  • He identified subatomic particles that were very small and negatively charged, which he named electrons.

  • Because atoms were known to be electrically neutral, he reasoned they must contain enough positive charge to balance the negative charge of the electrons.

  • Thomson proposed a model where electrons were embedded into a positively charged sphere, analogous to chocolate chips in cookie dough.

• Rutherford’s Model (1911)

  • By the early 1900s, scientists understood that the positive charge of an atom comes from subatomic particles called protons.

  • Rutherford’s experiments led him to believe that protons are concentrated in a small area at the center of the atom, which he called the nucleus.

• Bohr’s Model (1913)

  • Bohr modified Rutherford’s model by proposing that each electron possesses a specific amount of energy that helps it move around the nucleus.

  • Electrons move around the nucleus in regions called energy levels, which surround the nucleus in rings similar to the layers of an onion.

  • This is often called the planetary model because the energy levels occupied by electrons are like the orbits of planets at different distances from the sun (the nucleus).

• Electron Cloud Model (Modern Model)

  • This is the currently accepted model.

  • It posits that electrons do not move in fixed orbits but dart around within an energy level.

  • The rapid, random motion of electrons creates a "cloud" of negative charge around the nucleus, which gives the atom its size and shape.

• The Modern Atomic Model (1932)

  • In 1932, James Chadwick discovered another particle within the nucleus called the neutron.

  • Neutrons have no electrical charge (they are neutral).

  • The theory states: "At the center of the atom is a tiny, massive nucleus containing protons and neutrons. Surrounding the nucleus is a cloudlike region of moving electrons."

Fundamental Atomic Structure and Subatomic Particles

• Atom Defined: The smallest unit of matter that retains the identity of the substance.

• Structure Regions:

  • Nucleus: Located at the center of the atom; it contains the mass of the atom (protons and neutrons).

  • Electron Cloud: The region surrounding the nucleus that contains most of the physical space in the atom.

• Subatomic Particles:

  • Protons ($p^+$): Positively charged particles located in the nucleus. The number of protons identifies the atom (e.g., 2 protons = Helium, 29 protons = Copper).

  • Neutrons ($n^0$): Neutrally charged particles located in the nucleus. They have no charge and do not have to equal the number of protons or electrons.

  • Electrons ($e^-$): Negatively charged particles located outside the nucleus in the electron cloud. They have relatively no mass compared to protons and neutrons.

Atomic Calculations and Measurements

• Atomic Number ($Z$): Indicates the number of protons in an atom.

  • Examples: Hydrogen ($Z=1$) has 1 proton; Carbon ($Z=6$) has 6 protons.

• Mass Number ($A$): The total number of protons and neutrons in the nucleus.

  • Formula for neutrons: $\text{Number of neutrons} = \text{Mass Number} - \text{Atomic Number}$

  • Example for Lithium: Mass number is $7$, atomic number is $3$. Protons = $3$, Neutrons = $7 - 3 = 4$.

  • Example for Neon: Mass number is $20$, atomic number is $10$. Protons = $10$, Neutrons = $20 - 10 = 10$.

• Neutral Atoms: In a neutral atom, protons equal electrons ($p^+ = e^-$).

  • Example for Helium: Mass number $4$, atomic number $2$. $p^+ = 2$, $n^0 = 2$, $e^- = 2$.

  • Example for Chlorine: Mass number $35$, atomic number $17$. $p^+ = 17$, $n^0 = 18$, $e^- = 17$.

  • Example for Potassium: Mass number $39$, atomic number $19$. $p^+ = 19$, $n^0 = 20$, $e^- = 19$.

• Atomic Weight: The weighted average of all naturally occurring isotopes of an element. This is an experimental decimal number.

  • Example for Nitrogen calculation:

    • $99.63\%$ of $^{14}N$ ($14.00307\,amu$)

    • $0.37\%$ of $^{15}N$ ($15.00011\,amu$)

    • Calculation: $(0.9963 \times 14.00307) + (0.0037 \times 15.00011) = 14.0067 \approx 14.01\,amu$

• Atomic Mass Unit ($amu$): The standard unit for measuring atomic mass.

Electron Quantum Numbers and Shells

• Quantum Numbers:

  1. Principal quantum number ($n$): Represents the distance of an electron from the nucleus (shell).

  2. Second quantum number ($l$): Represents the shape of the electron subshell.

  3. Third quantum number ($m_l$): Represents the number of energy states for each subshell.

  4. Fourth quantum number ($m_s$): Represents the spin moment of an electron.

• Shell Capacities (Bohr Model):

  • $n=1$ (K Shell): 1 Subshell ($s$). Total energy states: 1. Max electrons: $2$.

  • $n=2$ (L Shell): 2 Subshells ($s, p$). Total energy states: 4. Max electrons: $8$.

  • $n=3$ (M Shell): 3 Subshells ($s, p, d$). Total energy states: 9. Max electrons: $18$.

  • $n=4$ (N Shell): 4 Subshells ($s, p, d, f$). Total energy states: 16. Max electrons: $32$.

Electron Configurations

• Electrons have discrete energy states and tend to occupy the lowest available energy state first.

• Stability is achieved when the valence (outer) shell is completely filled.

• Examples of configurations:

  • Hydrogen ($Z=1$): $1s^1$

  • Helium ($Z=2$): $1s^2$ (Stable)

  • Lithium ($Z=3$): $1s^2 2s^1$

  • Beryllium ($Z=4$): $1s^2 2s^2$

  • Boron ($Z=5$): $1s^2 2s^2 2p^1$

  • Carbon ($Z=6$): $1s^2 2s^2 2p^2$

  • Neon ($Z=10$): $1s^2 2s^2 2p^6$ (Stable)

  • Sodium ($Z=11$): $1s^2 2s^2 2p^6 3s^1$

  • Magnesium ($Z=12$): $1s^2 2s^2 2p^6 3s^2$

  • Aluminum ($Z=13$): $1s^2 2s^2 2p^6 3s^2 3p^1$

  • Argon ($Z=18$): $1s^2 2s^2 2p^6 3s^2 3p^6$ (Stable)

  • Krypton ($Z=36$): $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6$ (Stable)

The Periodic Table and Electronegativity

• Electropositive elements: Readily give up electrons to become positive ions (cations).

• Electronegative elements: Readily acquire electrons to become negative ions (anions).

• Electronegativity: Ranges from $0.7$ to $4.0$. Large values indicate a strong tendency to acquire electrons.

• Valence Electrons by Group:

  • Group IA: 1 valence electron ($1e^-$)

  • Group IIA: 2 valence electrons ($2e^-$)

  • Group IIIA: 3 valence electrons ($3e^-$)

  • Group VIIA: 7 valence electrons ($7e^-$)

  • Group VIII (Inert Gases): 8 valence electrons ($8e^-$), except Helium which has 2.

Ions and Ionic Compounds

• Ion: An atom that has gained or lost electrons to become a charged species. The number of protons NEVER changes.

• Anion: A negatively charged ion formed by gaining electrons (e.g., $Cl + e^- \rightarrow Cl^-$).

• Cation: A positively charged ion formed by losing electrons (e.g., $Na \rightarrow Na^+ + e^-$).

  • Example: Sodium ($11p^+, 11e^-$) becomes $Na^+$ ($11p^+, 10e^-$).

  • Example: Chlorine ($17p^+, 17e^-$) becomes $Cl^-$ ($17p^+, 18e^-$).

  • Result: $Na^+$ with $Cl^-$ forms $NaCl$.

Atomic Bonding and Forces

• Bonding Forces: Net force is the sum of attractive and repulsive forces: $F_N = F_A + F_R$.

• Equilibrium State: Occurs when the net force is zero: $F_N = F_A + F_R = 0$.

• Potential Energy ($E$): Related to force by $E = \int F dr$. Net energy is $E_N = E_A + E_R$.

• Types of Primary Bonding:

  1. Ionic Bonding: Occurs between positive and negative ions via electron transfer. Requires a large difference in electronegativity (e.g., $NaCl$). It is nondirectional and common in ceramics.

  2. Covalent Bonding: Involves shared electrons between atoms with comparable electronegativities (e.g., $CH_4$). carbon has 4 valence electrons and needs 4 more, while Hydrogen has 1 and needs 1 more. This bonding is directional.

  3. Metallic Bonding: Arises from a "sea" of donated valence electrons (1, 2, or 3 from each atom). It is the primary bond for metals and their alloys and is nondirectional.

• Secondary Bonding (Van der Waals):

  • Arises from interaction between dipoles.

  • Fluctuating Dipoles: Asymmetric electron clouds (e.g., liquid $H_2$).

  • Permanent Dipoles: Molecule-induced (e.g., liquid $HCl$, polymers).

Bonding Energy Summary

• Ionic: Large bond energy, nondirectional.

• Covalent: Variable bond energy (Large in Diamond, small in Bismuth), directional.

• Metallic: Variable bond energy (Large in Tungsten, small in Mercury), nondirectional.

• Secondary: Smallest bond energy, directional.

Visible Spectrum and Light

• The visible spectrum of light ranges from approximately $400\,nm$ (ultraviolet/violet) to $700\,nm$ (infrared/red).

• Bohr observed line spectra by passing light through a prism to study energy levels.

Questions & Discussion

• What did I learn today? (Self-reflection assessment prompt at the end of the lecture).