Chemistry: The Central Science - Chapter 7 Study Notes

Development of the Periodic Table

• Dmitri Mendeleev and Lothar Meyer independently proposed a grouping method for elements.

Mendeleev and the Periodic Table

• Table 7.1 provides a comparison between Mendeleev’s predictions for eka-silicon (now known as germanium) and the actual properties of germanium.

  • Predicted vs. Observed Properties of Germanium:

  • Atomic weight: 72 (predicted) vs. 72.59 (observed)

  • Density (g/cm³): 5.5 (predicted) vs. 5.35 (observed)

  • Specific heat (J/g·K): 0.305 (predicted) vs. 0.309 (observed)

  • Melting point (°C): High (predicted) vs. 947 (observed)

  • Color: Dark gray (predicted) vs. grayish white (observed)

  • Formula of oxide: XO₂ (predicted) vs. GeO₂ (observed)

  • Density of oxide (g/cm³): 4.7 (predicted) vs. 4.70 (observed)

  • Formula of chloride: XCl₄ (predicted) vs. GeCl₄ (observed)

  • Boiling point of chloride (°C): A little under 100 (predicted) vs. 84 (observed)

• Mendeleev is credited significantly due to his use of chemical properties for organizing the table and successfully predicting properties of missing elements such as germanium.

Atomic Number

• Mendeleev's initial table utilized atomic masses as the primary property of the elements.

• About 35 years after Mendeleev, Ernest Rutherford discovered the nuclear atom.

• Henry Moseley introduced the atomic number concept, with protons defining the periodic property of elements.

Periodicity

• Periodicity is defined as a repetitive pattern of element properties related to atomic number.

• Key properties discussed include:

  • Sizes of atoms and ions

  • Ionization energy

  • Electron affinity

  • Group chemical property trends

• The chapter begins with an examination of effective nuclear charge, a fundamental property influencing many trends.

Effective Nuclear Charge

General Concept (1 of 2)

• Properties of elements are influenced by the attractions between valence electrons and the nucleus.

• Electrons are attracted to the nucleus but also repel one another.

• The net forces experienced by an electron arise from both attractive and repulsive interactions.

Effective Nuclear Charge Calculation (2 of 2)

• The effective nuclear charge ($Z<em>{eff}$) is calculated using the formula: $Z</em>{eff} = Z - S$

  • Where $Z$ represents the atomic number & $S$ denotes a screening constant (typically close to the count of inner electrons).

• Effective nuclear charge trends:

  • Increases across a period.

  • Slight increase down a group.

Atomic Sizing

Size Concept

• The nonbonding atomic radius, or van der Waals radius, measures half of the shortest distance between two nuclei during atomic collision.

• The bonding atomic radius, or covalent radius, is defined as half the distance between two nuclei in a bond.

Trends in Sizes of Atoms

• Trends in bonding atomic radius indicate:

  • Generally decreases from left to right across a period ($Z_{eff}$ increases).

  • Generally increases from top to bottom within a group (energy level $n$ increases).

Ionic Sizes

Determinants of Ionic Size (1 of 2)

• Ionic size is influenced by:

  • Nuclear charge

  • Number of electrons

  • Electron orbital configurations

Cations vs. Anions (2 of 2)

• Cations are smaller than their parent atoms due to:

  • The removal of the outermost electron which reduces electron-electron repulsion.

• Anions are larger than their parent atoms because:

  • Electrons are added, increasing electron-electron repulsion.

Isoelectronic Series

• An isoelectronic series consists of ions with identical electron counts.

• In this series, ionic size decreases as nuclear charge increases.

• Example of trends:

  • Increasing nuclear charge leads to a decreasing ionic radius.

Ionization Energy (I)

• Ionization energy (IE) is the energy necessary to detach an electron from a gaseous atom or ion under ground-state conditions.

  • The first ionization energy (I₁) refers to the energy for removing the initial electron.

  • The second ionization energy (I₂) refers to the energy for removing the second electron.

• A higher ionization energy indicates increased difficulty for electron removal.

Trends in Ionization Energy

• It becomes progressively harder to remove each subsequent electron.

• After all valence electrons are removed, ejecting core electrons requires significantly more energy.

• Table 7.2 lists successive values of ionization energies for elements sodium through argon (in kJ/mol).

Periodic Trends in First Ionization Energy (I₁)

• Ionization energy tends to increase across a period.

• Ionization energy tends to decrease down a group.

• S- and P-block elements exhibit a broader range of I₁ values.

  • D-block shows minor increases across a period.

  • F-block elements demonstrate minimal variations.

Influencing Factors of Ionization Energy

• Smaller atomic size correlates with higher ionization energy values.

• Ionization energy values are influenced by effective nuclear charge and the average distance of the electron from the nucleus.

Irregularities in General Trend

• General trends may not hold when:

  • The next valence electron enters a new energy sublevel.

  • The next electron pairs in the same orbital (causing electron repulsion that lowers energy).

Electron Configurations of Ions

Cations

• Cations appear when electrons are lost from the highest energy level (denoted by the highest principal quantum number, n).

  • Example: Losing a 2s electron.

  • Example: Losing two 4s electrons.

Anions

• Anions are formed when electron configurations are filled, such as:

  • Gaining one electron in the 2p orbital.

Electron Affinity

• Electron affinity is the energy change when an electron is added to a gaseous atom.

  • Typically an exothermic process; hence, the value is often negative.

General Trends in Electron Affinity

• Electron affinity displays minimal variation down a group.

• Generally increases across a period.

  • Notable exceptions include elements in:

  • Group 2A: s sublevel is full.

  • Group 5A: p sublevel is half-full.

  • Group 8A: p sublevel is full.

  • Important to note: Many of these exceptions exhibit a positive electron affinity.

Metals, Nonmetals, and Metalloids

• The classification of elements shows increasing metallic character starting from nonmetals to metals within the periodic table.

• Key groups include:

  • 1A: Alkali metals

  • 2A: Alkaline earth metals

  • 6A: Oxygen group

  • 7A: Halogens

  • 8A: Noble gases

  • Discussion on why hydrogen is classified as a nonmetal.

Metals vs. Nonmetals

Metals

• Metals typically form cations.

• Characteristics include:

  • Shiny luster

  • Good conductivity of heat and electricity

  • Malleable and ductile

  • Predominantly solids at room temperature (except for mercury)

  • Low ionization energies allowing easy cation formation

Metal Chemistry

• Compounds formed between metals and nonmetals are often ionic.

• Metal oxides usually exhibit basic properties and react with acids.

Nonmetals

• Nonmetals are primarily located on the right side of the periodic table.

• Key properties include:

  • Existing as solids, liquids, or gases (element-dependent)

  • Solid nonmetals generally appear dull, brittle, and are poor conductors

  • Possess high negative electron affinities, facilitating anion formation

Nonmetal Chemistry

• Molecules solely composed of nonmetals form molecular compounds.

• Most nonmetal oxides are acidic in nature.

Comparison of Metal and Nonmetal Properties

• Table 7.3 outlines distinct properties between metals and nonmetals:

  • Metals have shiny luster; nonmetals do not.

  • Metals are malleable and ductile; nonmetals are typically brittle.

  • Metals conduct heat and electricity effectively; nonmetals are poor conductors.

  • Most metal oxides are basic; most nonmetal oxides are acidic.

  • Metals tend to form cations; nonmetals form anions or oxyanions in solution.

Metalloids

• Metalloids possess mixed characteristics of both metals and nonmetals.

• Some metalloids function as electrical semiconductors (used in computer chips).