Atoms and the Periodic Table Notes

Section 2.4 Subatomic Particles, Isotopes, and Ions

• Learning Objectives:

  • Identify the mass number of an atom using the count of each type of subatomic particle (protons, electrons, and neutrons) present in that atom, and vice versa.

  • Identify which specific element an atom represents using the mass number of that atom, or vice versa.

• Subatomic Particles and the Atomic Nucleus:

  • Subatomic Particles: Atoms are composed of three primary particles: neutrons, protons, and electrons.

  • The Nucleus:

    • The atomic nucleus is incredibly small, possessing a radius only $1/10,000$ of the radius of the cell.

    • The nucleus contains almost all of the mass of the atom.

  • Electrons:

    • Electrons possess very little mass.

    • Despite their low mass, electrons occupy almost the entire volume of the atom.

• Properties of Subatomic Particles (Table 2.2):

  • Proton (p):

    • Relative Charge ($e$): $+1$.

    • Mass ($u$): $1.0073\,\text{u}$.

    • Location: Inside the nucleus.

  • Neutron (n):

    • Relative Charge ($e$): $0$.

    • Mass ($u$): $1.0087\,\text{u}$.

    • Location: Inside the nucleus.

  • Electron (e):

    • Relative Charge ($e$): $-1$.

    • Mass ($u$): $0.000549\,\text{u}$.

    • Location: Outside the nucleus.

  • Fundamental Unit of Charge: The relative charges are based on the charge of the electron, represented as $e$, where $1\,e = -1.60 \times 10^{-19}\,\text{C}$.

  • Atomic Mass Units: Masses are measured in atomic mass units, abbreviated as $\text{u}$.

• Relative Sizes and Masses Metaphors:

  • Mass Comparison: If an electron had the mass of a marble, a proton or neutron would have the mass of a bowling ball.

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

Atomic Number, Mass Number, and Isotopes

• Atomic Number ($Z$):

  • The atomic number, symbol $Z$, is defined by the number of protons in an atom's nucleus: $Z = p$.

  • The atomic number determines the identity of the element.

  • On the periodic table, every element is assigned a whole number $Z$. For example, $Z = 1$ is hydrogen and $Z = 8$ is oxygen.

  • In an electrically neutral atom, the number of protons equals the number of electrons: $\text{number of } p = \text{number of } e$.

• Isotopes:

  • Atoms of the same element can have different numbers of neutrons in their nuclei.

  • Definition: Two atoms with the same number of protons but different numbers of neutrons are called isotopes of the same element.

  • Isotopes share the same atomic number but differ in mass due to the differing neutron count.

• Mass Number ($A$):

  • The mass number, symbol $A$, is the sum of protons and neutrons in the atom: $A = p + n = Z + n$.

  • The number of neutrons can be calculated as the difference between the mass number and the atomic number: $n = A - Z$.

• Symbology for Isotopes:

  • Isotopes are named using the element name or symbol followed by the mass number (e.g., hydrogen-1).

  • Isotope Symbols: Typically written as ${}^{A}_{Z}\text{E}$, where $A$ is the mass number and $Z$ is the atomic number.

  • Hydrogen Examples:

    • Hydrogen-1: ${}^{1}\text{H}$ (contains 1 proton and 0 neutrons).

    • Hydrogen-2: ${}^{2}\text{H}$ (contains 1 proton and 1 neutron).

Examples: Particles and Isotope Notation

• Example 2.7: Subatomic Particle Counting:

  • Question: How many protons, neutrons, and electrons are in a neutral atom of each of the following isotopes?

    • a. ${}^{32}\text{S}$

    • b. ${}^{19}\text{F}$

    • c. ${}^{20}\text{Ne}$

    • d. ${}^{238}\text{U}$

  • Solution:

    • a. Sulfur ($Z=16$): $16\,p$, $16\,e$, $16\,n$ (since $32-16=16$).

    • b. Fluorine ($Z=9$): $9\,p$, $9\,e$, $10\,n$ (since $19-9=10$).

    • c. Neon ($Z=10$): $10\,p$, $10\,e$, $10\,n$ (since $20-10=10$).

    • d. Uranium ($Z=92$): $92\,p$, $92\,e$, $146\,n$ (since $238-92=146$).

• Example 2.8: Identifying Correct Symbols:

  • Question: Assuming the element symbols are correct, which of the following isotopic symbols are correct? Explain why.

  • (Note: Identifying specific correct symbols requires verifying if the atomic number matches the element symbol on the periodic table).

Ions: Cations and Anions

• Definition: Atoms can form charged particles called ions by gaining or losing electrons.

• Cations:

  • Formed when an atom loses one or more electrons, resulting in a positive charge.

  • Example: Sodium reacts to lose an electron: $\text{Na} \rightarrow \text{Na}^{+} + e^{-}$.

  • This process is termed oxidation.

• Anions:

  • Formed when an atom gains one or more electrons, resulting in a negative charge.

  • Example: Oxygen reacts to gain electrons: $\text{O} + 2e^{-} \rightarrow \text{O}^{2-}$.

  • This process is termed reduction.

• Example 2.11: Identifying Particles in Ions:

  • Question: How many protons and electrons are in each of the following ions?

    • a. $\text{F}^{-}$

    • b. $\text{Mg}^{2+}$

    • c. $\text{N}^{3-}$

    • d. $\text{W}^{6+}$

  • Solution:

    • a. $\text{F}^{-}$: $9\,p$ and $10\,e$.

    • b. $\text{Mg}^{2+}$: $12\,p$ and $10\,e$.

    • c. $\text{N}^{3-}$: $7\,p$ and $10\,e$.

    • d. $\text{W}^{6+}$: $74\,p$ and $68\,e$.

Section 2.5 Atomic Masses

• Learning Objectives:

  • Calculate the atomic mass of any element using the masses and relative abundances of its naturally occurring mixture of isotopes.

• The Atomic Mass Scale:

  • Atomic masses are measured on a relative scale because they are extremely small. One isotope is assigned a specific value, and all others are measured relative to it.

  • Historical Standard: Initially, oxygen was the standard, with the naturally occurring mixture of oxygen isotopes assigned a value of exactly $16\,\text{atomic mass units}$ ($\text{u}$ or $\text{amu}$).

  • Current Standard (1961): Carbon-12 ($\text{C-12}$) was designated as the new standard, assigned a mass of exactly $12\,\text{u}$.

• Modern Determination of Atomic Mass:

  • The atomic mass displayed on the periodic table is a weighted average of the actual masses of all naturally occurring isotopes of that element.

  • Weighted averages account for the relative abundance (the number of each type of isotope present in nature).

• Example 2.12: Weighted Average Metaphor (Course Grading):

  • Scenario: Exams = 50%, Quizzes = 25%, Homework = 20%, Participation = 5%.

  • Student Scores: Exams = 80.2%, Quizzes = 77.3%, Homework = 87.8%, Participation = 100.0%.

  • Calculation: Overall Grade = $(0.50 \times 80.2) + (0.25 \times 77.3) + (0.20 \times 87.8) + (0.05 \times 100.0)$.

• Calculation of Atomic Mass Formula:

  • $\text{atomic mass} = \sum(\text{fraction of isotope} \times \text{mass of isotope})$

• Example 2.13: Calculating Atomic Mass of Copper:

  • Data:

    • ${}^{63}\text{Cu}$: 69.17% abundance, mass = $62.9396\,\text{u}$.

    • ${}^{65}\text{Cu}$: 30.83% abundance, mass = $64.9278\,\text{u}$.

  • Steps:

    • Convert percentages to fractions: 0.6917 and 0.3083.

    • Calculation: $(0.6917 \times 62.9396) + (0.3083 \times 64.9278)$.

• Artificial Isotopes and Elements:

  • Artificial Elements: Elements with Z > 92 are artificial and do not have standard atomic masses based on natural abundance.

  • Periodic Table Representation: For artificial elements, the mass number of the most stable isotope is provided in parentheses.

  • Artificial Isotopes: Artificial isotopes of naturally occurring elements are not included in the calculation of the element's atomic mass.

Section 2.6 The Periodic Table

• Learning Objectives:

  • Discuss the historical contributions of Mendeleev and Meyer.

  • Locate and classify elements based on their position in the table.

  • Identify periods, groups, and sections by name and number.

• Development of the Periodic Table:

  • History: The table resulted from refining the grouping of elements with similar chemical characteristics over many years.

  • Early Ordering: Initially, elements were ordered by atomic mass. After reaching Dalton's hypothesis, scientists focused on determining relative atomic masses.

  • Mendeleev and Meyer: Dimitri Mendeleev and Lothar Meyer independently produced periodic tables combining internal characteristics and atomic masses.

  • Mendeleev's Innovation: Mendeleev used the table to predict the existence and properties of undiscovered elements. He also prioritized chemical properties over mass when they conflicted, notably swapping Iodine ($\text{I}$) and Tellurium ($\text{Te}$). His belief that their masses were incorrectly determined was later validated.

• Modern Arrangement:

  • Elements are arranged horizontally by increasing atomic number ($Z$) and vertically by similar properties.

  • Example 2.15: Atomic Mass Inversions: Besides Iodine and Tellurium, other naturally occurring pairs out of order by mass (but correctly ordered by atomic number) include:

    • Argon ($\text{Ar}$) and Potassium ($\text{K}$).

    • Cobalt ($\text{Co}$) and Nickel ($\text{Ni}$).

• Structure of the Modern Table:

  • Periods: Horizontal rows. There are 7 periods. Periods 6 and 7 include the inner transition elements.

  • Groups/Families: Vertical columns. Elements in a group share similar chemical properties.

  • Numbering Systems:

    • Older system: Groups 1–8 with A/B designations.

    • Modern system: Groups 1–18 (left to right).

• Group Names (Figure 2.17):

  • Group 1 (1A): Alkali metals.

  • Group 2 (2A): Alkaline earth metals.

  • Group 11: Coinage metals.

  • Group 15 (5A): Pnictogens.

  • Group 16 (6A): Chalcogens.

  • Group 17 (7A): Halogens.

  • Group 18 (8A): Noble gases (colorless, odorless, mostly inert; discovered after the initial periodic table development).

Classifications and Categories

• Larger Group Classifications:

  • Main Group Elements: Found in groups 1, 2, and 13–18. These are the most common elements and exhibit predictable chemistry.

  • Transition Elements: Found in groups 3–12; all are metals.

  • Inner Transition Elements: Displayed in two rows below the main table.

    • Lanthanoids: Elements 58–71.

    • Actinoids: Elements 90–103.

• Metals, Nonmetals, and Metalloids (Figure 2.19):

  • Metals: Make up the majority of elements; located on the left side. Properties: Shiny, malleable, ductile, good conductors of heat and electricity.

  • Nonmetals: Includes Hydrogen plus elements on the upper-right side. Properties: Brittle as solids.

  • Metalloids: Border the division between metals and nonmetals; they share properties with both categories.

• Example Identifications (Examples 2.16 – 2.20):

  • Period 4: Starts with Potassium ($\text{K}$), ends with Krypton ($\text{Kr}$), contains 18 elements.

  • Location Identification:

    • Neon ($\text{Ne}$): Period 2, Group 18.

    • Rubidium ($\text{Rb}$): Period 5, Group 1 (corrected from transcript slide 41 list).

    • Tungsten ($\text{W}$): Period 6, Group 6.

    • Francium ($\text{Fr}$): Period 7, Group 1.

  • Similar Chemical Properties: Rubidium ($\text{Rb}$) and Francium ($\text{Fr}$) (both are in Group 1).

  • Classification:

    • Cesium ($\text{Cs}$): Main group element.

    • Indium ($\text{In}$): Main group element.

    • Uranium ($\text{U}$): Inner transition element (Actinoid).

  • Element Trivia (Example 2.20):

    • Fifth element in the first row of transition metals: Manganese ($\text{Mn}$).

    • Element in Period 4, Group 6: Chromium ($\text{Cr}$).

    • Seventh transition element: Cobalt ($\text{Co}$).

    • First element of Group 8: Iron ($\text{Fe}$).

    • Third halogen: Bromine ($\text{Br}$).

    • First alkaline earth metal: Beryllium ($\text{Be}$).

    • First coinage metal: Copper ($\text{Cu}$).