C2 The Periodic Table

2.1 Development of the Periodic Table

  • Learning Objectives:

    • Understand the historical development of the periodic table.
    • Recognize how testing predictions can validate or challenge scientific ideas.

  • Challenges in Understanding Elements in the Early 1800s:

    • Limited knowledge of atomic structure.
    • Misidentification of some compounds as elements.
    • Incomplete listing of all elements.

  • 19th Century Chemistry:

    • Chemists actively discovered new elements.
    • Focused on identifying behavioral patterns among elements for organization and improved understanding.

  • John Dalton's Contribution (1808):

    • Arranged elements by atomic weights.
    • Published a table of elements in "A New System of Chemical Philosophy".

  • John Newlands' "Law of Octaves" (1864):

    • Organized elements by mass, noting similar properties in every eighth element.
    • Assumed all elements were discovered, leading to inaccuracies.
    • The pattern held only up to calcium; faced ridicule from other scientists.

Mendeleev's Breakthrough

  • Dmitri Mendeleev (1869):

    • Arranged ~50 known elements by atomic weights, organized to show periodic property patterns.
    • Left gaps for undiscovered elements, predicting their properties.
    • Subsequent discoveries validated Mendeleev's predictions.

  • Challenges to Mendeleev's Pattern:

    • Argon's higher atomic mass than potassium posed a problem.
    • Mendeleev adjusted element order to maintain property-based grouping.
    • The reason was a mystery for decades.

  • 20th Century Resolution:

    • Discovery of atomic structure explained the discrepancies.
    • Elements ordered by the number of protons (atomic number).
    • Isotopes accounted for unexpected atomic weights.

  • Key Points:

    • The periodic table arranges elements to group similar ones.
    • The table is named for regular property patterns.
    • Mendeleev's table was accepted for predicting undiscovered element properties.

C2.2 Electronic Structures and the Periodic Table

  • Learning Objectives:

    • Link atomic structure to the periodic table layout.
    • Differentiate metals and non-metals based on electronic structure and table position.
    • Explain the unreactive nature of noble gases.

  • Periodic Table Organization:

    • Elements are arranged by atomic (proton) number.
    • Elements with similar properties are aligned vertically in groups.
    • There are eight main groups.

  • Electronic Structures and Group Numbers:

    • Elements in the same group behave similarly due to having the same number of electrons in their outermost shell.
    • The group number indicates the number of outer shell electrons.
    • Example: Group 2 elements have two outer electrons; Group 6 elements have six.

  • Metals vs. Non-metals:

    • Metals conduct electricity while non-metals are generally insulators.
    • Metals typically have higher melting and boiling points.
    • Metals are ductile and malleable; non-metal solids are brittle.

  • Location in the Periodic Table:

    • Non-metals are in the top right corner.
    • Group 5, 6, and 7 elements gain electrons to form negative ions.
    • Metals are on the left and center; Groups 1, 2, and 3 lose electrons to form positive ions.

  • Group 0 - Noble Gases:

    • Noble gases have eight outer electrons (except helium, which has two).
    • This stable electron configuration makes them unreactive.

  • Noble Gas Properties:

    • They exist as monatomic gases.
    • They do not easily form molecules.
    • A few compounds with larger noble gases exist, containing fluorine or oxygen (e.g., XeF and XeO).
    • Boiling points increase down the group (Helium: -269°C, Radon: -62°C).

  • Key Points:

    • Atomic number determines position.
    • Outer shell electrons dictate chemical properties.
    • Group number = number of outer shell electrons.
    • Metals lose electrons; non-metals gain electrons.
    • Noble gases are unreactive due to stable electron configurations.

C2.3 Group 1 – The Alkali Metals

  • Learning Objectives:

    • Understand the behavior of Group 1 elements.
    • Know how their properties change down the group.

  • List of Alkali Metals:

    • Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Francium (Fr).
    • Rubidium, Cesium, and Francium are highly reactive or unstable and typically not used in schools.

  • Properties:

    • Very reactive; must be stored in oil to prevent reaction with air.
    • Reactivity increases down the group (Lithium is least reactive, Francium is most).
    • Low density; Lithium, Sodium, and Potassium float on water (density < 1 g/cm³).
    • Soft, can be cut with a knife; silvery, shiny surface when freshly cut, but quickly dulls due to oxide formation.
    • Example reaction: 4Na(s) + O2(g) \rightarrow 2Na2O(s)

  • Electronic Structure:

    • Atoms have one electron in their outermost shell, making them reactive.
    • They lose this electron to achieve a stable noble gas configuration.
    • Form 1+ ions (e.g., Na^+, K^+$), always creating ionic compounds.

  • Melting and Boiling Points:

    • Low melting and boiling points for metals.
    • These decrease going down the group; Cesium is liquid at just 29°C.

  • Reaction with Water:

    • Metals float, move around, and fizz.
    • Fizzing due to hydrogen gas formation.
    • Potassium reacts vigorously; hydrogen ignites with a lilac flame (due to potassium ions).
    • Metal hydroxide is also produced, making the solution alkaline (high pH).
    • Example: 2Na(s) + 2H2O(l) \rightarrow 2NaOH(aq) + H2(g)
    • Example: 2K(s) + 2H2O(l) \rightarrow 2KOH(aq) + H2(g)

  • Other Reactions:

    • React vigorously with non-metals like chlorine to form metal chlorides (white solids that dissolve in water to form colorless solutions).
    • Reactivity increases down the group due to easier loss of the outer electron.
    • Example: 2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)
    • Similar reactions occur with fluorine, bromine, and iodine; all produce white, water-soluble, colorless solutions.

  • Key Points:

    • Group 1 elements are the alkali metals.
    • Melting and boiling points decrease down the group.
    • They react with water to produce hydrogen and alkaline solutions containing metal hydroxides.
    • They form 1+ ions in reactions to make ionic compounds which are generally white, dissolve in water and produce colorless solutions.
    • Reactivity increases down the group.

C2.4 Group 7 - The Halogens

  • Learning Objectives:

    • Understand the behavior of Group 7 elements.
    • Recognize how their properties change down the group.

  • Properties of Halogens:

    • Toxic non-metals with colored vapors.
    • Low melting and boiling points that increase down the group.
    • Poor conductors of heat and electricity.
    • Exist as diatomic molecules (e.g., F2, Cl2, Br2, I2).

  • Reactions of Halogens:

    • They have seven electrons in their outermost shell.
    • They gain one electron by sharing with non-metals (e.g., hydrogen).

  • Reactions with Hydrogen:

    • F2(g) + H2(g) \rightarrow 2HF(g) (Explosive, even at -200°C in the dark).
    • Cl2(g) + H2(g) \rightarrow 2HCl(g) (Explosive in sunlight, slow in the dark).
    • Br2(g) + H2(g) \rightarrow 2HBr(g) (Over 300°C with platinum catalyst).
    • I2(g) + H2(g) \rightarrow 2HI(g) (Over 300 °C; very slow, reversible).

  • Halogen reactivity decreases down the group

  • Reactions with metals:

    • Halogens gain a single electron to form ions with a 1- charge (e.g., F^-, Cl^-, Br^-, I^-).
    • Examples: Sodium chloride (NaCl) and iron(III) bromide (FeBr_3) - ionic compounds.

  • Displacement Reactions:

    • A more reactive halogen displaces a less reactive one from its salt solutions.
    • Cl2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br2(aq)$$
    • Fluorine violently reacts in aqueous solutions and is thus an exception.

  • Key Points:

    • Halogens form ions with a single negative charge in ionic compounds.
    • Halogens form covalent compounds by sharing electrons with other non-metals.
    • A more reactive halogen displaces a less reactive one from a solution of its salts.
    • Halogen reactivity decreases down the group.

C2.5 Explaining Trends

  • Reactivity Trends:

    • Group 1: Reactivity increases down the group (Li < Na < K < Rb < Cs).
    • Group 7: Reactivity decreases down the group (F > Cl > Br > I > At).

  • Explanation based on Electronic Structure:

    • Down a group, the number of electron shells increases, making atoms larger.
    • Larger atoms lose electrons more easily (Group 1).
    • Larger atoms gain electrons less easily (Group 7).

  • Reasons for these Effects:

    • Outer electrons are farther from the nucleus, reducing attraction (electrons are negatively charged, and the nucleus is positively charged).
    • Inner electron shells shield outer electrons from the nucleus's positive charge.

  • Group 1 Trends Explained:

    • Reactivity increases down Group 1 because the outermost electron is less strongly attracted to the nucleus.
    • Increased distance and shielding outweigh the increasing nuclear charge.

  • Group 7 Trends Explained:

    • Reactivity decreases down Group 7.
    • The same factors are in play: atomic size, shielding, and nuclear charge.
    • Down the group, it's harder to attract an extra electron.

  • Key Points:

    • Reactivity trends depend on the attraction between outer electrons and the nucleus.

  • Electrostatic attraction depends on:

    • Distance between outer electrons and nucleus.
    • Number of occupied inner shells (shielding effect).
    • Size of the positive charge on the nucleus (nuclear charge).
    • Increased nuclear charge down the group is outweighed by increased distance and shielding.
    • Electrons are easier to lose and harder to gain for larger atoms down a group.