Fundamentals of General, Organic, and Biological Chemistry - Chapter 2: Atoms and the Periodic Table

Chapter 2: Atoms and the Periodic Table

2.1 Atomic Theory and the Structure of Atoms

• Atom: The smallest particle an element can be divided into while retaining its identity. The term comes from the Greek atomos, meaning "indivisible."

• Learning Objective: Explain the major assumptions of atomic theory, and name and identify the properties of the subatomic particles that make up an atom.

• Atomic Theory (Four Assumptions): Chemistry is founded on these four fundamental assumptions:

  • All matter is composed of atoms.

  • The atoms of a given element differ from the atoms of all other elements.

  • Chemical compounds consist of atoms combined in specific ratios. Only whole atoms can combine.

  • Chemical reactions change only the way atoms are combined in compounds.

• Subatomic Particles: Atoms are composed of tiny subatomic particles.

  • Protons: Carry a positive electrical charge.

  • Neutrons: Have a mass similar to that of a proton but are electrically neutral.

  • Electrons: Have a mass that is only $\frac{1}{1836}$ that of a proton and carry a negative electrical charge.

• Relative Mass Scale: The masses of atoms and subatomic particles are expressed on a relative mass scale.

  • Atomic Mass Unit (amu): The unit for describing the mass of an atom, based on the mass of a carbon-12 atom.

• **Atomic Structure: ** * Nucleus: Protons and neutrons are packed closely together in a dense core called the nucleus. This is the central part of the atom.

  • Electron Cloud: Surrounding the nucleus, electrons move rapidly through a large volume of space.

  • Analogy for Size: The relative size of a nucleus in an atom is comparable to that of a pea in the middle of a stadium.

• Electrical Forces in an Atom:

  • Attraction: Opposite electrical charges attract each other. Negatively charged electrons are held near the positively charged nucleus.

  • Repulsion: Like charges repel each other. Electrons try to get as far away from one another as possible.

2.2 Elements and Atomic Number

• Learning Objective: Identify atoms of an element based on the number of protons in the nucleus.

• Atomic Number (Z): The number of protons in atoms of a given element. All atoms of a particular element have the same number of protons in the nucleus.

• Mass Number (A): The sum of the protons and neutrons in an atom.

• Neutral Atoms:

  • Atoms are electrically neutral; they have no net charge.

  • The number of positively charged protons and the number of negatively charged electrons are the same in each atom.

  • Therefore, in a neutral atom, the atomic number (Z) also equals the number of electrons.

  • Example: Neutral Carbon, with $Z = 6$, has $6$ protons and $6$ electrons.

• Worked Example 2.1: Phosphorus

  • Problem: Phosphorus has the atomic number $Z = 15$. How many protons, electrons, and neutrons are there in phosphorus atoms, which have a mass number $A = 31$?

  • Analysis: The atomic number ($Z$) gives the number of protons, which is also the number of electrons in a neutral atom. The mass number ($A$) gives the total number of protons plus neutrons.

  • Solution:

    • Protons: $Z = 15$, so $15$ protons.

    • Electrons: In a neutral atom, electrons = protons = $15$, so $15$ electrons.

    • Neutrons: Subtract the atomic number from the mass number: $A - Z = 31 - 15 = 16$ neutrons.

• Worked Example 2.2: Identifying an Element

  • Problem: An atom contains $28$ protons and has $A = 60$. Give the number of electrons and neutrons in the atom, and identify the element.

  • Analysis: The number of protons is the atomic number ($Z$). In a neutral atom, the number of electrons equals the number of protons. The number of neutrons is the mass number minus the atomic number ($A - Z$).

  • Solution:

    • Protons: $28$, so $Z = 28$.

    • Electrons: $28$ (equal to protons for a neutral atom).

    • Neutrons: $A - Z = 60 - 28 = 32$ neutrons.

    • Element Identification: Looking at the periodic table, the element with atomic number $Z = 28$ is Nickel (Ni).

2.3 Isotopes and Atomic Weight

• Learning Objective: Write the symbols for different isotopes of an element, and use relative abundances and atomic masses of isotopes to calculate the average atomic weight of an element.

• Isotopes: Atoms of the same element (identical atomic numbers, $Z$) but with different mass numbers ($A$).

  • This means isotopes have the same number of protons but different numbers of neutrons.

• Example: Hydrogen Isotopes

  • Protium: One proton, no neutrons; mass number = $1$ ($_{1}^{1} ext{H}$)

  • Deuterium: One proton, one neutron; mass number = $2$ ($_{1}^{2} ext{H}$)

  • Tritium: One proton, two neutrons; mass number = $3$ ($_{1}^{3} ext{H}$)

• Isotope Notation: A specific isotope is represented by showing its mass number ($A$) as a superscript and its atomic number ($Z$) as a subscript in front of the atomic symbol.

  • Example: The symbol for tritium is $_{1}^{3} ext{H}$.

• Isotope Naming: For most elements, isotopes do not have distinctive names. Instead, the mass number ($A$) is given after the name of the element.

  • Example: Uranium-235 (or U-235) is an isotope used in nuclear reactors.

• Natural Occurrence: Most naturally occurring elements are mixtures of isotopes.

• Atomic Weight: The weighted average mass of an element's atoms.

  • Calculation: To calculate the atomic weight, the individual masses of the naturally occurring isotopes and the percentage of each (isotopic abundance) must be known.

  • Formula: $ext{Atomic weight} = ext{Σ}[( ext{isotopic abundance}) × ( ext{isotopic mass})]$

• Worked Example 2.3: Gallium Atomic Weight

  • Problem: Gallium has two naturally occurring isotopes: 60.4 ext{%} is Ga-69 (mass = $68.9257$ amu) and 39.6 ext{%} is Ga-71 (mass = $70.9248$ amu). Calculate the atomic weight for gallium.

  • Analysis: Calculate the average atomic mass by summing the contributions from each naturally occurring isotope.

  • Ballpark Estimate: The isotopic masses are $68.9$ amu and $70.9$ amu. Since slightly more than half ( 60.4 ext{%} ) is the lighter isotope (Ga-69), the average mass should be slightly less than halfway between the two. Estimate: $ext{approximately } 69.8$ amu.

  • Solution:

    • Step 1 (Known Information):

      • Ga-69: 60.4 ext{%} abundance, mass = $68.9257$ amu

      • Ga-71: 39.6 ext{%} abundance, mass = $70.9248$ amu

    • Step 2 (Unknown Answer & Units): Atomic weight for Ga (in amu) = ?

    • Step 3 (Equation): $ext{Atomic weight} = ext{Σ}[( ext{isotopic abundance}) × ( ext{isotopic mass})]$

    • Step 4 (Solve):

      • Contribution from Ga-69: $(0.604) × (68.9257 ext{ amu}) = 41.6311 ext{ amu}$

      • Contribution from Ga-71: $(0.396) × (70.9248 ext{ amu}) = 28.0862 ext{ amu}$

      • Total Atomic Weight: $41.6311 ext{ amu} + 28.0862 ext{ amu} = 69.7173 ext{ amu}$

      • Final Answer (3 significant figures): $69.7 ext{ amu}$

• Worked Example 2.4: Identifying Isotope Properties

  • Problem: Identify element X in the symbol $_{78}^{194} ext{X}$ and give its atomic number, mass number, number of protons, number of electrons, and number of neutrons.

  • Analysis: The subscript is the atomic number ($Z$), which identifies the element. The superscript is the mass number ($A$). In a neutral atom, protons = electrons = $Z$. Neutrons = $A - Z$.

  • Solution:

    • Atomic Number ($Z$): $78$.

    • Element Identity: Looking at the periodic table, element X with $Z = 78$ is Platinum (Pt).

    • Mass Number ($A$): $194$.

    • Number of Protons: $Z = 78$ protons.

    • Number of Electrons: $78$ electrons (for a neutral atom).

    • Number of Neutrons: $A - Z = 194 - 78 = 116$ neutrons.

2.4 The Periodic Table

• Learning Objective: Locate elements on the periodic table and classify them as metals, nonmetals, or metalloids based on their location.

• Periodic Table: A tabular format listing all known elements, where the atomic symbol, name of the element, and atomic mass are given in each box representing the element.

• History: Dmitri Mendeleev produced the forerunner of the modern periodic table, organizing elements by their properties.

• Classification by Physical Properties:

  • Metals:

    • Properties: Malleable, lustrous appearance, good conductors of heat and electricity.

    • Location: Occur on the left side of the periodic table.

  • Nonmetals:

    • Properties: Poor conductors of heat and electricity.

    • Location: Occur on the upper-right side of the periodic table.

  • Metalloids:

    • Properties: Exhibit properties intermediate between those of a metal and a nonmetal.

    • Location: Located in a zigzag band between the metals on the left and nonmetals on the upper-right side of the periodic table.

• Classification by Chemical Behavior: Elements are also classified by their chemical behavior, primarily based on their position in vertical columns.

  • Groups (Vertical Columns): Elements in the same vertical column (group) have similar chemical properties.

  • Periods (Horizontal Rows): Elements in the same horizontal row are in the same period.

  • Categories of Elements:

    • Main group elements (Groups 1A-8A)

    • Transition metal elements (Groups 3B-2B, or Groups 3-12)

    • Inner transition metal elements (Lanthanides and Actinides, typically below the main table)

2.5 Some Characteristics of Different Groups

• Learning Objective: Classify elements and describe chemical behavior based on group membership.

• Periodicity: A repeating rise-and-fall pattern in properties is observed as elements are arranged by increasing atomic number.

• Group 1A—Alkali Metals:

  • Elements: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr).

  • Properties: Shiny, soft metals with low melting points.

  • Reactivity: Highly reactive; react vigorously with water to form highly alkaline (basic) products.

  • Occurrence: Never found in nature in a pure, uncombined state due to their high reactivity.

• Group 2A—Alkaline Earth Metals:

  • Elements: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra).

  • Properties: Lustrous, silvery metals.

  • Reactivity: Less reactive than Group 1A neighbors, but still reactive.

  • Occurrence: Never found in nature in a pure state.

• Group 7A—Halogens:

  • Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At).

  • Properties: Colorful and corrosive nonmetals.

  • Occurrence: Found in nature only in combination with other elements (e.g., sodium chloride, NaCl, which is table salt).

  • Name Origin: The group name "halogen" comes from the Greek word hals, meaning "salt," reflecting their ability to form salts.

• Group 8A—Noble Gases:

  • Elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rn).

  • Properties: Colorless gases.

  • Reactivity: Known for their lack of chemical reactivity, earning them the label "noble" gases.

    • Helium, neon, and argon do not combine with any other elements.

    • Krypton and xenon combine with very few other elements under specific conditions.