Premed · Premed · General Chemistry 1

Lecture 9: Periodic Properties of the Elements

General Chemistry I

Learning Objectives

By the end of this lecture, students will be able to:

1. Explain the organization of the periodic table in terms of electron configurations
2. Classify elements by group characteristics (alkali metals, alkaline earth metals, halogens, noble gases, transition metals)
3. Describe and explain periodic trends in atomic radius, ionic radius, ionization energy, and electron affinity
4. Predict relative sizes of atoms and ions across the periodic table
5. Relate trends in metallic character and chemical reactivity to electron configuration

Lecture Content

I. Organization of the Periodic Table

The modern periodic table arranges elements in order of increasing atomic number (Z). The horizontal rows, called periods, group elements that share the same highest principal quantum number (n). Period 1 has n = 1 and contains 2 elements, Period 2 has n = 2 with 8 elements, Period 3 has n = 3 with 8 elements, and Period 4 has n = 4 with 18 elements. The vertical columns, called groups, collect elements with similar valence electron configurations, which in turn gives them similar chemical properties. The table can also be divided into blocks based on which subshell is being filled: the s-block (Groups 1-2 and He) fills ns orbitals, the p-block (Groups 13-18) fills np orbitals, the d-block (Groups 3-12, the transition metals) fills (n-1)d orbitals, and the f-block (lanthanides and actinides, the inner transition metals) fills (n-2)f orbitals.

II. Important Group Families

The Group 1A elements (alkali metals: Li, Na, K, Rb, Cs, Fr) share an ns^1 configuration, making them very reactive metals that form +1 cations and react vigorously with water. The Group 2A elements (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra) have an ns^2 configuration and form +2 cations; they are reactive but less so than the alkali metals. The Group 7A elements (halogens: F, Cl, Br, I, At) possess an ns^2 np^5 configuration, making them very reactive nonmetals that form -1 anions and exist naturally as diatomic molecules. The Group 8A elements (noble gases: He, Ne, Ar, Kr, Xe, Rn) have completely filled valence shells (ns^2 np^6, except for He which is 1s^2) and are generally unreactive and remarkably stable. The transition metals of the d-block display variable oxidation states, form colored compounds, often produce complex ions, and frequently serve as catalysts.

III. Effective Nuclear Charge (Z_eff)

The effective nuclear charge experienced by an electron is given by Z_eff = Z - sigma, where sigma is the shielding or screening constant. Electrons in inner shells shield outer electrons from the full positive charge of the nucleus, while electrons within the same shell provide some shielding but much less effectively. Core electrons are far more effective at shielding than valence electrons. Moving across a period from left to right, Z increases by 1 with each element, but shielding increases only slightly because the added electrons enter the same shell. The result is a steady increase in Z_eff across a period. Moving down a group, Z increases substantially, but electrons are added to new, more distant shells, so the Z_eff experienced by valence electrons increases only modestly. This concept of effective nuclear charge is the key factor underlying most periodic trends.

<image>A diagram illustrating effective nuclear charge. The left side shows a simplified atom with a nucleus labeled "Z = 11 (Na)" surrounded by concentric electron shells. Inner electrons (1s^2, 2s^2 2p^6) are shown as a shielding cloud labeled "10 core electrons, shielding." The outer electron (3s^1) is shown experiencing a reduced nuclear pull. An arrow shows Z_eff ≈ +11 - 10 = +1 for the valence electron. The right side shows a periodic table heat map where each element's square is color-coded by Z_eff of the outermost electron, ranging from cool colors (low Z_eff, left side) to warm colors (high Z_eff, right side), demonstrating the increase across each period.</image>

IV. Atomic Radius

Atomic radius is defined as half the distance between the nuclei of two bonded identical atoms (the covalent radius) or half the distance between nearest neighbors in a metallic crystal. Across a period from left to right, atomic radius decreases because Z_eff increases, pulling electrons closer to the nucleus while they remain in the same principal energy level with no additional shielding layers. Down a group, atomic radius increases because electrons are added to higher-n shells at greater average distances from the nucleus, and increased shielding from additional inner electron shells weakens the nuclear pull on the outermost electrons. To illustrate: across Period 2, Li (152 pm) is larger than C (77 pm), which is larger than F (64 pm). Down Group 1A, Li (152 pm) is smaller than Na (186 pm), which is smaller than K (227 pm).

V. Ionic Radius

Cations are smaller than their parent atoms because the loss of electrons reduces electron-electron repulsion, often removes an entire outer shell, and increases Z_eff per remaining electron. Sodium, for instance, shrinks from 186 pm as a neutral atom to 95 pm as Na+. Anions are larger than their parent atoms because the extra electrons increase electron-electron repulsion and decrease Z_eff per electron. Chlorine grows from 99 pm as a neutral atom to 181 pm as Cl-.

An isoelectronic series consists of ions that all have the same number of electrons. Consider O^2-, F-, Na+, Mg^2+, and Al^3+, each with 10 electrons. As nuclear charge increases across this series (8, 9, 11, 12, 13 protons), ionic radius decreases because more protons pull on the same number of electrons. Within a single group, ionic radius increases going down, following the same trend as atomic radius.

VI. Ionization Energy (IE)

The first ionization energy (IE_1) is the energy required to remove the outermost electron from a neutral gaseous atom: X(g) -> X+(g) + e-. This process is always endothermic, so IE_1 is always positive. Successive ionization energies increase (IE_2 > IE_1 > IE_3, etc.), with a dramatic jump occurring when a core electron must be removed, breaking into a noble gas configuration.

Across a period, ionization energy generally increases from left to right because increasing Z_eff holds electrons more tightly. Two notable exceptions appear: there is a slight dip at Group 3A, where the first p electron (higher in energy and easier to remove than a paired s electron) is being removed, and another at Group 6A, where one of the paired p electrons is removed and the electron-electron repulsion within the pair makes removal easier. Down a group, ionization energy decreases because electrons are farther from the nucleus, experience more shielding, and are therefore easier to remove.

<image>A plot of first ionization energy (y-axis, in kJ/mol) versus atomic number (x-axis, elements 1-36). The graph shows a general sawtooth pattern: peaks at noble gases (He, Ne, Ar, Kr) and troughs at alkali metals (Li, Na, K). Within each period, the overall trend is increasing, but notable dips occur at B (Group 3A, removing a 2p electron) and O (Group 6A, paired electron repulsion in 2p^4). Elements are labeled at their data points. Annotations explain the dips: "B: first electron in higher-energy 2p subshell" and "O: removing paired electron from 2p^4 is easier due to electron-electron repulsion."</image>

VII. Electron Affinity (EA)

Electron affinity is the energy change when a gaseous atom gains an electron: X(g) + e- -> X-(g). By convention, a more negative electron affinity indicates a more favorable (more exothermic) process. Across a period, electron affinity generally becomes more negative as atoms approach a filled valence shell and have a stronger tendency to gain electrons. Exceptions include Group 2A and Group 8A, where the filled s subshell or complete octet leaves no energetic benefit to adding an electron (EA is approximately zero or positive), and Group 5A, where the half-filled p subshell is already relatively stable. Down a group, electron affinity generally becomes less negative (weaker attraction for an added electron). A notable exception is fluorine, which has a less negative electron affinity than chlorine because fluorine's compact 2p orbitals create strong electron-electron repulsion when an additional electron is forced into such a small space.

VIII. Metallic Character

Metallic character refers to an element's tendency to lose electrons and form cations. It increases going down a group (lower ionization energy makes it easier to lose electrons) and decreases going across a period from left to right (higher ionization energy makes it harder). Consequently, metals cluster in the lower left of the periodic table while nonmetals dominate the upper right. Along the staircase line separating the two regions lie the metalloids or semimetals (B, Si, Ge, As, Sb, Te), which have properties intermediate between metals and nonmetals.

IX. Summary of Periodic Trends

X. Electronegativity (Preview)

Electronegativity measures the ability of an atom in a bond to attract shared electrons toward itself. On the Pauling scale, fluorine has the highest electronegativity at 4.0, while cesium is among the lowest at 0.7. The trend follows ionization energy: electronegativity increases across a period and decreases down a group. The difference in electronegativity between two bonded atoms determines the polarity of the bond, a topic explored in depth in the bonding lectures.

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