Chapter 2

Dalton's Atomic Theory and the Fundamentals of Atoms

• Elements are composed of extremely small particles referred to as atoms. Atoms are the fundamental building blocks recognized in the periodic table and chemistry.
• All atoms of a given element are identical across the board. This identity entails having the same size, mass, and chemical properties. For example, any particular atom of oxygen is identical to any other atom of oxygen.
• Atoms of one element are different from the atoms of all other elements. Comparing oxygen to hydrogen reveals that these two different atoms are completely distinct in their size, mass, and chemical properties.
• Compounds are composed of atoms of more than one element. To form a compound like $H_2$, more than one element must come together.
• In any compound, the ratio of the numbers of atoms of any two of the elements present are either an integer or a simple fraction. When atoms form a compound, they do so in a specific, fixed ratio. For instance, in carbon monoxide ($CO$), the ratio of carbon to oxygen is fixed at $1:1$. Changing this ratio to $1:2$ results in a completely different compound, carbon dioxide ($CO_2$).
• A chemical reaction involves only the separation, combination, or rearrangement of atoms. It does not result in the creation or destruction of atoms. This principle implies that bonds in reactants are broken and rearranged to form new products without losing or creating matter.

Law of Multiple Proportions and Conservation of Mass

• The Law of Multiple Proportions states that different compounds made up of the same elements differ in the number of atoms of each kind that combine, and they do so in simple whole-number ratios.
  • Example: Carbon Monoxide ($CO$) vs. Carbon Dioxide ($CO_2$). Both are made of carbon and oxygen. In $CO$, the ratio of oxygen to carbon is $1$. In $CO_2$, the ratio of oxygen to carbon is $2$. The number of carbon atoms remains the same, but the variation in oxygen atoms creates different substances.
• The Law of Conservation of Mass dictates that during a chemical reaction, we should not create or destroy atoms. The number of reactant atoms on the left-hand side of an equation must equal the number of product atoms on the right-hand side.
• Mathematical Representation of Conservation:
  • If responding $16$ atoms of $X$ with $8$ atoms of $Y$ to form a compound $X_2Y$:
  • Reactants: $16X + 8Y$
  • Products: $8$ moles of $X_2Y$
  • Verification: $8 \times 2 = 16$ atoms of $X$; $8 \times 1 = 8$ atoms of $Y$.
  • This ensures the number of moles is equivalent across both sides, maintaining mass balance.

The Structure of the Atom and Subatomic Particles

• The atom is structured like a sphere with a central nucleus.
• The nucleus contains the neutrons and protons.
• Electrons orbit the nucleus.
• There are three primary subatomic particles: protons, neutrons, and electrons. Understanding them involves identifying the discovering scientists, their charge nature, and their relative sizes.

Discovery of the Electron: JJ Thomson and Millikan

• JJ Thomson utilized the Cathode Ray Tube to investigate the electron. He was able to measure the charge-to-mass ratio ($e/m$) of the electron.
• The specific value measured for this ratio is $-1.76 \times 10^8\,C/g$ (coulombs per gram). For this discovery, Thomson was awarded a Nobel Prize.
• Bar Magnet Experiment: Thomson used a bar magnet against the cathode ray tube.
  • Without the magnet, there was no deflection.
  • Using the positive end showed no deflection, but using the negative end caused a deflection.
  • This indicated the negatively charged nature of the electron based on the principle that like charges repel.
• Millikan's Oil Droplet Experiment: Millikan measured the mass of the electron using findings from Thomson's ratio.
• The mass of an electron was determined to be approximately $9.1 \times 10^{-28}\,g$. This extremely small exponent indicates that the electron's mass is very tiny.
• Plum Pudding Model: Proposed by JJ Thomson, this model suggests the atom is a uniform sphere of positive charge with electrons embedded within it, similar to raisins in a pudding. This postulation was based on the idea that a positive region must exist to hold negatively charged electrons in place (opposite charges attract).

Radioactivity and the Rutherford Experiment

• Rutherford identified three types of radioactive particles by passing them through an electric field:
  • Beta ($\beta$) particles: Attracted to the positive plate, indicating they are negatively charged.
  • Alpha ($\alpha$) particles: Deflected toward the negative plate, indicating they are positively charged.
  • Gamma ($\gamma$) rays: Passed through without deflection, indicating they are neutral.
• Gold Foil Experiment: Rutherford fired alpha particles at a thin gold foil.
  • Observations: Most particles passed through, but some were deflected back significantly.
  • Conclusion 1: An atom's positive charge is concentrated in a tiny central region called the nucleus.
  • Conclusion 2: Protons have an opposite charge to electrons.
  • Conclusion 3: The mass of a proton is approximately $1840$ times the mass of an electron.
• Size Metaphor: If a nucleus were the size of a marble, the atom would be the size of a $50$-yard or $500$-yard stadium, illustrating that the nucleus is very small compared to the total volume of the atom.

Discovery of the Neutron: James Chadwick

• James Chadwick discovered the third subatomic particle: the neutron.
• Experiment: He bombarded a beryllium nucleus with alpha particles.
• Products: The reaction produced Neutrons, Carbon-$12$, and "NHg" (energy/particles).
• Nature of the Neutron:
  • It is neutral (charge of $0$).
  • Mass is approximately equal to the mass of a proton ($1.67 \times 10^{-24}\,g$).

Summary of Subatomic Particles

• Electron: Charge is $-1$. Mass is approximately $10^{-28}\,g$.
• Proton: Charge is $+1$. Mass is approximately $10^{-24}\,g$.
• Neutron: Charge is $0$. Mass is approximately $10^{-24}\,g$.
• Comparative Size: The mass of the proton is approximately equal to the mass of the neutron ($m_p \approx m_n$). Together, they are significantly larger ($\approx 1840$ times) than the electron.

Atomic Number, Mass Number, and Isotopes

• Atomic Number ($Z$): the number of protons in the nucleus. This determines the identity of the element.
• Mass Number ($A$): the total number of protons plus the number of neutrons. $A = \text{protons} + \text{neutrons}$.
• Neutral Atoms: In a neutral atom, $\text{Atomic Number (Z)} = \text{number of protons} = \text{number of electrons}$.
• Chemical Representation: ${^A_Z}X$, where $X$ is the element symbol, $A$ is the superscript (mass number), and $Z$ is the subscript (atomic number).
• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons in their nuclei.
• Examples of Isotopes:
  • Hydrogen Isotopes:
    • Hydrogen-$1$ (${^1_1}H$): $1$ proton, $0$ neutrons.
    • Deuterium (${^2_1}H$ or $D$): $1$ proton, $1$ neutron.
    • Tritium (${^3_1}H$): $1$ proton, $2$ neutrons.
  • Uranium Isotopes:
    • Uranium-$235$ (${^{235}_{92}}U$): $143$ neutrons ($235 - 92$).
    • Uranium-$238$ (${^{238}_{92}}U$): $146$ neutrons ($238 - 92$).

The Periodic Table Structure

• Groups: Vertical columns. Elements in the same group often share properties.
• Periods: Horizontal rows.
• Notable Group Names:
  • Group 1: Alkali metals.
  • Group 2: Alkaline earth metals.
  • Group 7: Halogens.
  • Group 8: Noble gases.
• Classifications:
  • Metals: Highlighted in green; usually on the left/center. Form cations ($+$).
  • Non-metals: Highlighted in blue; usually on the right. Form anions ($-$).
  • Metalloids: Highlighted in gray; share properties of both metals and non-metals.
• Transition Metals: Located in the center; tendency to form more than one type of cation (e.g., $Fe^{2+}$ and $Fe^{3+}$).
• Note on Periodic Table Layout: There is a "flip" in representation compared to standard notation. On the table, the atomic number ($Z$) is usually above the symbol, and the mass ($A$) is below it.

Chemistry in Action: Elemental Abundance

• Natural Abundance in Earth's Crust: Oxygen is the most abundant ($45.5\%$), followed by Silicon.
• Natural Abundance in the Human Body: Oxygen is the most abundant, followed by Carbon.

Molecules and Ions

• Molecule: An aggregate of two or more atoms in a definite arrangement held together by chemical forces.
  • Diatomic Molecules: Contain only two atoms.
    • Homonuclear (elements): $H_2, N_2, O_2, F_2, Cl_2, Br_2, I_2$.
    • Heteronuclear (compounds): $HCl, CO$.
  • Polyatomic Molecules: Contain more than two atoms. Examples: $O_3$ (ozone), $H_2O$, $NH_3$ (ammonia), $CH_4$ (methane).
• Ions: An atom or group of atoms with a net positive or negative charge.
  • Cation: Ion with a positive charge. Formed when a neutral atom loses one or more electrons (e.g., $Na \rightarrow Na^+ + 1e^-$, where proton count $11$ stays the same but electrons decrease to $10$).
  • Anion: Ion with a negative charge. Formed when a neutral atom gains one or more electrons (e.g., $Cl + 1e^- \rightarrow Cl^-$, where proton count $17$ remains same but electrons increase to $18$).
  • Monoatomic Ions: Contain only one atom (e.g., $Na^+, Cl^-, Ca^{2+}, O^{2-}, Al^{3+}, N^{3-}$).
  • Polyatomic Ions: Multiple atoms carrying a charge (e.g., $OH^-, CN^-, NH_4^+, NO_3^-$).

Predicting Ion Charges

• Group 1: Forms $+1$ charge.
• Group 2: Forms $+2$ charge.
• Group 3: Forms $+3$ charge.
• Group 5: Forms $-3$ charge (needs $3$ electrons to reach $8$).
• Group 6: Forms $-2$ charge (needs $2$ electrons to reach $8$).
• Group 7: Forms $-1$ charge (needs $1$ electron to reach $8$).
• Group 8: Noble gases; generally do not gain or lose electrons.

Formulas and Models

• Molecular Formula: Shows the exact number of atoms of each element in the smallest unit of a substance.
• Structural Formula: Shows how atoms are connected, often using dash lines to represent bonds.
• Ball-and-Stick Model: Shows connectivity; utilizes specific colors for atoms.
  • Oxygen: Red.
  • Nitrogen: Blue.
  • Carbon: Black.
  • Hydrogen: Gray.
• Space-Filling Model: Provides a more realistic picture by considering the relative sizes of atoms (e.g., showing Oxygen is larger than Hydrogen).
• Empirical Formula: The simplest whole-number ratio of atoms in a substance.
  • Example 1: Acetylene ($C_2H_2$) molecular formula reduces to empirical formula $CH$.
  • Example 2: Glucose ($C_6H_{12}O_6$) reduces to $CH_2O$.
  • Example 3: Nitrous Oxide ($N_2O$) cannot be reduced; thus, the molecular and empirical formulas are the same ($N_2O$).