Atoms and Elements Study Notes

Elements and Chemical Symbols

• Definition of Elements: Elements are pure substances from which everything else is built.
• Chemical Symbols:
  • Chemical symbols are one or two-letter abbreviations used to represent an element’s name.
  • Capitalization Rule: When a chemical symbol contains two letters, the second letter is never capitalized.
    • Example: "Co" is the chemical symbol for cobalt.
    • Example Variation: "CO" specifies the presence of two distinct elements, carbon (C) and oxygen (O).
• Required Memorization: Students are required to memorize the names and symbols of the first twenty elements plus Bromine (Br) and Iodine (I).
• Reference List of Elements and Symbols (Examples):
  • Aluminum: $Al$
  • Barium: $Ba$
  • Carbon: $C$
  • Calcium: $Ca$
  • Chlorine: $Cl$
  • Hydrogen: $H$
  • Iodine: $I$
  • Nitrogen: $N$
  • Oxygen: $O$
  • Phosphorus: $P$
  • Silver: $Ag$
  • Gold: $Au$
  • Sulfur: $S$
  • Sodium: $Na$
  • Magnesium: $Mg$
  • Potassium: $K$
  • Fluorine: $F$

Dalton's Atomic Theory and the Nature of Atoms

• The Atom: An atom is the smallest particle of an element. All elements listed on the periodic table are composed of atoms. Atoms are the building blocks of everything and are too small to see with the naked eye.
• Historical Context: The existence of atoms did not become a formal scientific theory until 1808, when John Dalton (1766–1844) developed an atomic theory proposing that atoms were responsible for the combination of elements into compounds.
• Postulates of Dalton's Atomic Theory:
  1. All matter is made up of tiny particles called atoms.
  2. All atoms of a given element are identical to one another and different from atoms of other elements.
  3. Atoms of two or more different elements combine to form compounds. A particular compound is always composed of the same kinds of atoms and the same number of each kind of atom.
  4. A chemical reaction involves the rearrangement, separation, or combination of atoms. Atoms are never created or destroyed in a chemical reaction.

Laws of Conservation and Percent Composition

• Law of Conservation of Mass: Mass is neither created nor destroyed during a chemical reaction. The total mass of the reactants must equal the total mass of the products.
• Percent Composition Formula:
  • $\frac{\text{Mass of One}}{\text{Mass of Total}} \times 100 = \text{\% Composition of One}$
• Example Calculation #3:
  • Mass of Oxygen = $44.4\,g$
  • Mass of Hydrogen = $5.6\,g$
  • Combined Total Mass = $50.0\,g$
  • The mass is conserved between the reactants and products.

Subatomic Particles and Atomic Models

• Discovery Period: By the end of the 1880s, experiments with electricity demonstrated that atoms were composed of subatomic particles: protons, neutrons, and electrons.
• Electrical Charges:
  • Electrical charges can be positive or negative.
  • Repulsion: Two positive charges repel each other; two negative charges repel each other.
  • Attraction: Unlike charges (positive and negative) attract each other.
• J. J. Thomson (1897):
  • Discovered cathode rays were streams of small negatively charged particles called electrons.
  • Proposed the "plum pudding" model, which hypothesized that protons and electrons were scattered/distributed throughout the atom.
• Ernest Rutherford (1910s):
  • Hypothesized a model with protons concentrated in the center and electrons circling them.
  • Proposed that there is significant empty space around the central protons.
• Gold Foil Experiment (1911):
  • Thomson and Rutherford's experiments led to a new structural conclusion.
  • Results showed no scattering; most particles went straight through the atom, confirming the empty space model.
  • Conclusion: Atoms contain a small, positively charged region in the center called the nucleus, and a region of space around the nucleus occupied by electrons.
• James Chadwick (1932):
  • Discovered the nucleus also contained neutral particles called neutrons.

Properties of Protons, Neutrons, and Electrons

• Nucleus: Contains protons (positive charge) and neutrons (neutral charge). Protons and neutrons account for almost all the mass of the atom and are packed into the tiny volume of the nucleus.
• Electron Cloud: Rapidly moving electrons (negative charge) surround the nucleus and account for the large volume of the atom.
• Atomic Mass Unit (amu):
  • Standard unit used by chemists to describe atomic mass.
  • Defined as the mass of a hydrogen atom with 1 proton.
  • The mass of all elements is compared to the mass of the Carbon-12 ($^{12}C$) atom.
  • General values: Protons and neutrons both have a mass of approximately $1\,amu$. Electrons have a much smaller weight and generally do not affect the total atomic mass.

Equations and Atomic Fundamentals

• Equation 1: Atomic Number:
  • $\text{Atomic Number (Z)} = \text{Number of Protons}$
  • The atomic number is specific for each element and is the same for all atoms of that element.
  • Examples: Hydrogen ($H$) = 1; Carbon ($C$) = 6; Copper ($Cu$) = 29; Gold ($Au$) = 79.
• Equation 2: Atomic Charge:
  • $\text{Charge} = \text{Number of Protons} - \text{Number of Electrons}$
  • An atom is electrically neutral (net charge of zero) if it has an equal number of protons and electrons.
• Equation 3: Mass Number:
  • $\text{Mass Number (A)} = \text{Number of Protons} + \text{Number of Neutrons}$
  • The mass number represents the total particles in the nucleus. It does not appear on the periodic table because it refers to a single atom.
• Atomic Mass vs. Mass Number:
  • Atomic mass is listed below the symbol on the periodic table.
  • It is the weighted average of all naturally occurring isotopes of an element, compared to $^{12}C$.

Calculations and Atomic Symbols

• Atomic Symbol Format: \text{^{mass\ number}_{atomic\ number}Symbol^{charge}}
  • Example: \text{^{24}_{12}Mg^{+2}}
• Example Calculations for Charge and Electrons:
  • For \text{^{24}_{12}Mg^{+2}}: $+2 = 12 - X \rightarrow X = 10\,electrons$.
  • Oxygen with charge $-2$ and atomic number 8: $-2 = 8 - X \rightarrow X = 10\,electrons$.
  • Silicon (atomic number 14) with 17 electrons: $X = 14 - 17 \rightarrow Charge = -3$.
  • Silver (atomic number 47) with 48 electrons: $X = 47 - 48 \rightarrow Charge = -1$.
• Calculating Neutrons:
  • Potassium ($A = 39$, $Z = 19$): $39 - 19 = 20\,neutrons$.
  • Zinc ($A = 65$, $Z = 30$): $65 - 30 = 35\,neutrons$.

Isotopes

• Definition: Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.
• Carbon Isotopes:
  • $^{12}C$: 6 protons, 6 neutrons, 6 electrons.
  • $^{13}C$: 6 protons, 7 neutrons, 6 electrons (used for radioisotopic dating).
  • $^{14}C$: 6 protons, 8 neutrons, 6 electrons.
• Other Isotope Examples:
  • Phosphorus isotope: \text{^{31}_{15}P} (15 protons, 16 neutrons).
  • Oxygen isotope: \text{^{16}_{8}O} (8 protons, 8 neutrons).
  • Zinc isotope: \text{^{65}_{30}Zn} (30 protons, 35 neutrons).

The Periodic Table

• Dmitri Mendeleev (1872):
  • Created the periodic table by arranging elements by increasing atomic mass.
  • Arranged elements into groups with similar properties.
  • Discovered properties occur periodically in regular cycles.
  • Predicted gaps for elements that were unknown at the time based on reactivity patterns with oxygen and water.
• Organization of the Modern Periodic Table:
  • Groups: Vertical columns containing elements with similar properties.
  • Periods: Horizontal rows of elements.
    • Period 1: Contains 2 elements ($H, He$).
    • Periods 2 and 3: Contain 8 elements each.
• Group Numbering:
  • Letter A: Used for representative elements ($1A$ to $8A$).
  • Letter B: Used for transition elements.
  • Numbers 1–18: Alternative system used for columns from left to right.
• Named Families/Groups:
  • Group 1A (1): Alkali metals ($Li, Na, K, Rb, Cs, Fr$).
  • Group 2A (2): Alkaline earth metals ($Be, Mg, Ca, Sr, Ba, Ra$).
  • Group 6A (16): Chalcogens ($O, S, Se, Te, Po$).
  • Group 7A (17): Halogens ($F, Cl, Br, I, At$).
  • Group 8A (18): Noble gases ($He, Ne, Ar, Kr, Xe, Rn$).

Classification of Elements

• Metals:
  • Located to the left of the heavy zigzag line.
  • Properties: Shiny solids (except mercury, which is liquid), ductile (shaped into wires), malleable (hammered into sheets), good conductors of electricity, and high melting points.
• Nonmetals:
  • Located to the right of the zigzag line.
  • Examples: $H, C, N, O, Cl, S$.
  • Properties: Dull appearance, brittle (not malleable/ductile), poor conductors of heat/electricity, low melting points, and low densities.
• Metalloids:
  • Border the zigzag line (except Aluminum, which is a metal).
  • List: $B, Si, Ge, As, Sb, Te, Po, At$.
  • Properties: Exhibit mixed properties of both metals and nonmetals.
  • Semiconductors: Better conductors than nonmetals but not as good as metals; can be modified into insulators or conductors.

Questions & Discussion

• Question: Select the correct symbol for magnesium.
  • Response: (1) Mg.
• Question: Select the correct name for K.
  • Response: (2) potassium.
• Question: Is the mass of an electron greater than the mass of a proton?
  • Response: False.
• Question: Does the nucleus contain only protons and neutrons?
  • Response: True.
• Question: Match metals in Group 4A (14).
  • Response: Tin (Sn) and Lead (Pb).
• Question: Match metalloids in Group 4A (14).
  • Response: Silicon (Si) and Germanium (Ge).