Ch2

Chapter Outline

  • Section 2.1: Chemical Symbols

  • Section 2.2: The Laws of Chemical Combination

  • Section 2.3: The History of the Atom

  • Section 2.4: Subatomic Particles, Isotopes, and Ions

  • Section 2.5: Atomic Masses

  • Section 2.6: The Periodic Table

Section 2.1: Chemical Symbols

  • Chemical Symbols: Recognize and understand the symbols for commonly used elements.

Table 2.1: Elements Whose English Names and Symbols Begin with Different Letters

English Name

Symbol

Basis for the Symbol

Antimony

Sb

Stibium

Gold

Au

Aurum

Iron

Fe

Ferrum

Lead

Pb

Plumbum

Mercury

Hg

Hydrargyrum

Potassium

K

Kalium

Silver

Ag

Argentum

Sodium

Na

Natrium

Tin

Sn

Stannum

Tungsten

W

Wolfram

  • Chemical Formulas: Chemists write chemical symbols together in formulas to identify compounds.

    • Example: CO signifies carbon monoxide; Co represents cobalt.

    • Chemical formulas indicate the relative number of each element in a compound: H2O indicates 2 atoms of H and 1 atom of O.

    • Parentheses in formulas imply multiplication of atoms within by the subscript: Ca(NO3)2 contains 1 Ca atom, 2 N atoms, and 6 O atoms.

Section 2.2: The Laws of Chemical Combination

  • Classical Laws of Chemical Composition: Used to calculate quantities involved in combinations of elements.

Law of Definite Proportions

  • States that any given compound is composed of definite proportions by mass of its elements.

    • Example: In pure water, the mass ratio of oxygen to hydrogen is 8:1.

    • Ratio: Proportion of an element in a compound is the ratio of the mass of the element to the total mass of the compound.

    • Percent by Mass: The percent by mass of an element is the proportion of the element multiplied by 100%.

Law of Multiple Proportions

  • For any two (or more) compounds composed of the same elements, the ratio of the masses of any one element is a small, whole-number ratio.

    • Example: For every 1 g of C in CO, there are 1.33 g of O, and for every 1 g of C in CO2, there are 2.66 g of O.

Section 2.3: The History of the Atom

  • Dalton’s Atomic Theory: Explains classical laws of chemical combination:

    • Atoms combine to form compounds.

    • Atoms do not break down but exchange partners to create new compounds.

    • Mass is conserved during a chemical reaction.

Dalton’s Theory and the Laws

  • Law of Conservation of Mass: Atoms combine in small, whole-number ratios, resulting in conserved mass.

  • Law of Definite Proportions: All samples of a pure compound have identical mass proportions due to specific ratios of atoms.

  • Law of Multiple Proportions: Atoms combine in varying ratios to form different compounds, with the ratio of the masses being a whole number.

Key Experiments and Discoveries

  • Rutherford’s Nuclear Model: Experiment series that lead to the understanding of the atom's nuclear structure.

Section 2.4: Subatomic Particles, Isotopes, and Ions

  • Mass Number: Identify mass number using the count of protons, electrons, and neutrons.

  • Table 2.2: Properties of Subatomic Particles

    Particle

    Charge (e)*

    Mass (u)†

    Location in Atom

    Proton (p)

    +1

    1.0073

    In the nucleus

    Neutron (n)

    0

    1.0087

    In the nucleus

    Electron (e)

    -1

    0.000549

    Outside the nucleus

  • Relative Sizes and Masses: If an electron were the mass of a marble, a proton or neutron would be the mass of a bowling ball.

    • If the nucleus were the size of a marble, the entire atom would be the size of a stadium.

Atomic Number, Z

  • Atoms maintain electroneutrality (number of protons = number of electrons).

    • Atomic number, Z, indicates the number of protons in the nucleus: Z = p.

    • Each element has a unique atomic number on the periodic table (e.g., Z = 1 for hydrogen, Z = 8 for oxygen).

Isotopes

  • Atoms of the same element can have differing numbers of neutrons.

  • Isotopes: Same atomic number (number of protons) but different mass numbers due to different neutron counts.

  • Each isotope is identified by its mass number, A: A = p + n.

  • Example: Hydrogen-1 (^1H) with 1 proton and 0 neutrons, Hydrogen-2 (^2H) with 1 proton and 1 neutron.

Calculation of Neutron Count

  • Neutrons can be calculated as:
    n = A - Z. In the context of atomic structure: - **Z** represents the **atomic number**, which indicates the number of protons in an atom's nucleus (Z = p). Each element has a unique atomic number. - **A** represents the **mass number**, which is the total count of protons and neutrons in an atom's nucleus (A = p + n). The mass number identifies a specific isotope of an element.

Section 2.5: Atomic Masses

  • Atomic Mass Calculation: Calculate atomic mass by considering the isotopes' masses and their natural abundances.

  • Atomic Mass Scale: Atomic masses measured on a relative scale; initially, oxygen isotopes were used, but C-12 is the modern standard.

    • C-12 assigned a mass of exactly 12 u.

Atomic Mass Computation

  • Atomic mass of an element calculated as the weighted average of its naturally occurring isotopes.

Section 2.6: The Periodic Table

  • Contributions of Mendeleev and Meyer: Their work laid the foundation for the periodic table.

  • Classification of Elements: Elements organized based on their location in the periodic table.

  • Periods and Groups: Identify periods (horizontal rows) and groups (vertical columns) by name and number.

Figures

  • Figure 2.15: Rothschild's Cathode Ray Tube.

  • Figure 2.18: Modern Periodic Table and various group names, including:

    • 1A: Alkali metals

    • 2A: Alkaline earth metals

    • 7A: Halogens

    • 8A: Noble gases.