4: The Periodic Table and Some Atomic Properties
- Page ID
- 518182
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 4.1: Classifying the Elements- The Periodic Law and the Periodic Table
- The periodic table came from the efforts of Döbereiner (triads), Newlands (law of octaves), Meyer (atomic volume trends), and Mendeleev, who left spaces for undiscovered elements like Ga and Ge. Mendeleev's accurate predictions confirmed his table. Then, noble gases were included by Ramsay, and atomic number (not mass) was brought in by Moseley as the organizational basis! All this led to the modern periodic table that organizes elements according to properties and electron configuration.
- 4.2: Metals and Nonmetals and their Ions
- Elements are classified as metals (lustrous, conductive, malleable), nonmetals (dull, brittle, insulators), or metalloids (intermediate properties). Metals form cations, have high luster, and are good conductors. Nonmetals form anions, exist as gases/solids, and form acidic oxides. Metalloids like Si have intermediate properties and are semiconductors. Metallic character diminishes from left to right and enhances top to bottom on the periodic table, and nonmetallic character is vice versa.
- 4.3: Rules Governing Ground State Electron Configurations
- The Aufbau Principle (also called the building-up principle or the Aufbau rule) states that, in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels.
- 4.4: Sizes of Atoms and Ions
- Atomic sizes decrease left-to-right across periods (increasing nuclear charge) and increase top-to-bottom down groups (increasing nuclear shells). Cations are smaller than the parent atom (missing electrons reduce shielding), and anions are larger (gain electrons enhance repulsion). For isoelectronic series (same number of electrons), radius decreases with increasing nuclear charge (e.g., N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺). These trends enable us to estimate relative sizes of atoms/ions.
- 4.5: Ionization Energy
- Ionization energy (I) is the energy required to remove an electron from an atom. Ionization energies increase (I₁ < I₂ < I₃) as electrons are removed from successively more positive ions. Across periods, I₁ generally increases due to greater nuclear charge, with dips at group 13 and group 16 arrangements. Down groups, I₁ decreases since valence electrons are less strongly attracted to the nucleus. Transition metals show very small I₁ variations due to d-electron shielding.
- 4.6: Electron Affinity
- Electron affinity (EA) measures energy change when an electron is added to an atom, they can be negative (energy released), positive (energy absorbed), or zero. Trends show that EA becomes more negative across periods (increasing nuclear charge) and less negative down groups (larger atomic size). Second EAs are always positive due to strong electron-electron repulsion in dianions. It's good to note that Cl has the most negative EA (-346 kJ/mol), while noble gases have EAs ≥0.
- 4.7: Magnetic Properties
- Magnetic behavior is dependent on electron configuration, paramagnetic materials contain unpaired electrons that are parallel to magnetic fields, diamagnetic materials contain paired electrons and weakly repel magnetic fields. To determine magnetism: (1) Write e⁻ configuration, (2) Draw orbitals, (3) Count unpaired e⁻, (4) Paramagnetic if unpaired e⁻, diamagnetic if paired.
- 4.8: Periodic Properties of the Elements
- A summary of major periodic trends including, Atomic Radii, Electron Affinity, First Ionization Energy, and Electronegativity. These trends help build an intuitive understanding of why elements behave the way they do, both at an individual atom's scale and in a larger compound scale. It is important to note that each of these trends have been discussed in previous sections, to truly understand the trends refer to the previous corresponding sections.
- 4.10: Key Equations
- These are key equations. You should know how and when to use them. To practice their applications, check out 4.11.
- 4.11: Summary
- This section provides a brief summary of all the sections covered. It is important to understand that this should not be seen as a summary that will help you understand everything; it will only get you familiar with big concepts.
- 4.12: Exercises
- These homework exercises are designed to help you practice the content from this chapter. Feel free to work on them on your own or with a study group.