The periodic table reveals consistent patterns in atomic properties, known as periodic trends, which arise due to the arrangement of electrons and the structure of atomic nuclei.
Organization of the Periodic Table
Understanding periodic trends starts with understanding the structure of the periodic table. The table is organized in a way that reflects recurring chemical properties. These patterns stem from how electrons are arranged in atoms and how the nucleus interacts with these electrons.
Periods: Horizontal Rows
The periodic table has seven horizontal rows called periods. Elements within the same period have the same number of electron shells but increasing atomic numbers from left to right. The atomic number, which is the number of protons in the nucleus, determines the identity and basic chemical behavior of the element.
Increasing Atomic Number: As you move across a period, the number of protons increases. Each successive element has one more proton than the last.
Electron Configuration: Electrons are added to the same principal energy level, meaning that the shielding effect remains relatively constant across a period.
Same Shells, Different Properties: Despite having the same number of electron shells, elements in the same period have different physical and chemical properties due to the increase in nuclear charge.
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Example:
Sodium (Na), with 11 protons, and Argon (Ar), with 18 protons, both belong to Period 3. While they have the same number of electron shells (three), Argon’s stronger nuclear charge pulls electrons closer, altering its properties significantly compared to Sodium.
Groups: Vertical Columns
Groups are the vertical columns on the periodic table, totaling eighteen. Each group contains elements with the same number of valence electrons, which gives rise to similar chemical behaviors.
Same Valence Electrons: Elements in a group share similar bonding patterns and reactivity because they all have the same number of electrons in their outermost shell.
Increasing Shells: Moving down a group, each element has an additional electron shell, which increases atomic size and affects many other trends.
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Example:
Neon (Ne) and Xenon (Xe) are both in Group 18, the noble gases. Ne has two electron shells while Xe has five, but both have eight valence electrons, making them stable and chemically inert.
Effective Nuclear Charge (Z_eff)
The effective nuclear charge is the net positive charge experienced by an electron in an atom. It accounts for both the total positive charge from the nucleus and the shielding effect caused by inner electrons.
Actual Nuclear Charge (Z): The total number of protons in the nucleus.
Shielding Constant (S): The repelling effect of inner electrons that reduces the pull on outer electrons.
Effective Nuclear Charge Equation: Z_eff = Z − S
Across a period, Z increases while S remains relatively stable, so Z_eff increases. This stronger attraction pulls electrons closer to the nucleus. Down a group, both Z and S increase, but S increases more significantly, so Z_eff doesn’t increase as much.
This concept helps explain the behavior of electrons and the patterns in atomic properties that recur throughout the periodic table.
Atomic Radius
The atomic radius is defined as the distance from the nucleus of an atom to the outermost electron shell. It provides a sense of the size of an atom.
Across a Period → Decreases
As you move from left to right, protons are added to the nucleus.
Electrons are added to the same energy level, so no additional shielding occurs.
The increased nuclear charge pulls the electron cloud inward.
This causes the size of the atom to shrink gradually across a period.
Example:
In Period 2, lithium (Li) is larger than fluorine (F) because fluorine’s greater nuclear charge draws electrons in more tightly.
Down a Group ↓ Increases
Each successive element adds a new electron shell.
This increases the distance between the nucleus and the outer electrons.
Even though the nuclear charge increases, the added shielding prevents a stronger pull.
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Example:
In Group 1, lithium (Li) has 2 shells, sodium (Na) has 3, and cesium (Cs) has 6, making Cs the largest.
Ionic Radius
The ionic radius is the distance from the nucleus to the outermost electron in an ion. When atoms form ions, they either lose or gain electrons, which affects their size.
Cations (Positive Ions)
Metals typically lose electrons to form cations.
Losing electrons reduces electron-electron repulsion.
Often results in the loss of an entire outer shell.
Cations are smaller than their neutral atoms.
Example:
Sodium (Na) becomes Na⁺ by losing one electron. The result is a smaller radius because it now has only two occupied shells instead of three.
Anions (Negative Ions)
Nonmetals gain electrons to form anions.
The increase in electron-electron repulsion expands the electron cloud.
Anions are larger than their neutral atoms.
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Example:
Chlorine (Cl) becomes Cl⁻ by gaining an electron, causing the radius to increase due to added repulsion in the outer shell.
Electronegativity
Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
Across a Period → Increases
Atoms have a higher Z_eff, attracting bonding electrons more strongly.
Atomic radius decreases, so bonding electrons are closer to the nucleus.
Down a Group ↓ Decreases
Atomic radius increases, so bonding electrons are farther from the nucleus.
Increased shielding weakens the attraction.
Most Electronegative Element:
Fluorine (F) has the highest electronegativity with a value of 4.0.
Note:
Noble gases typically do not have assigned electronegativity values due to their lack of bonding.
Ionization Energy
Ionization energy is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms.
First Ionization Energy
Refers to the energy required to remove the outermost electron.
Across a Period → Increases
Higher Z_eff holds electrons more tightly.
Smaller atomic radius means electrons are closer to the nucleus.
Down a Group ↓ Decreases
Outer electrons are farther from the nucleus.
Shielding reduces the effective nuclear pull.
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Example:
Helium (He) has the highest ionization energy, while cesium (Cs) has one of the lowest among stable elements.
Higher Ionization Energies
Second Ionization Energy: Higher than the first because electrons are removed from a closer shell.
Large Energy Jumps: Occur when removing core (non-valence) electrons.
Diagnostic Tip:
Given ionization energies:
I1 = 500
I2 = 1500
I3 = 7000
I4 = 9000
The big jump from I2 to I3 suggests the element has 2 valence electrons.
Exceptions
Some deviations occur due to electron configurations:
Group 15 elements have higher first I.E. than Group 16.
Example: Nitrogen (N) > Oxygen (O) because N has a half-filled p orbital, which is relatively stable.
Quantum Tunneling
Explains why elements like boron (B) have lower ionization energy than expected.
A 2p electron in B is more easily removed than a 2s electron in Be due to its greater distance from the nucleus.
Electron Affinity
Electron affinity is the energy change when an atom gains an electron in the gaseous state.
Across a Period → More Negative
Increased Z_eff causes stronger attraction to added electrons.
More energy is released when electrons are added to atoms on the right side of the table.
Down a Group ↓ Less Negative (More Positive)
Atoms are larger, so added electrons are farther from the nucleus.
Less energy is released.
Note:
Electron affinity values are usually negative (exothermic process).
Exception: Chlorine (Cl) has a more negative electron affinity than fluorine (F).
In fluorine, the small size causes crowding in the valence shell, increasing repulsion.
Overview Image of Periodic Trends
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Key Terms and Definitions
Atomic Number: Number of protons in an atom’s nucleus; determines the identity of the element.
Atomic Radius: Distance from the nucleus to the outermost electron shell.
Coulomb’s Law: Describes force between two charged particles:
F = (k × q1 × q2) / r²
where F is force, q1 and q2 are charges, r is the distance, and k is a constant.Effective Nuclear Charge (Z_eff): Net positive charge on an electron:
Z_eff = Z − SElectron Affinity: Energy change when a neutral atom gains an electron.
Electronegativity: Atom’s ability to attract shared electrons in a chemical bond.
Groups: Columns on the periodic table; share number of valence electrons.
Ionization Energy: Energy needed to remove an electron from a gaseous atom or ion.
Ionic Radius: Size of an ion after gaining or losing electrons.
Noble Gases: Group 18 elements with full valence shells; inert and stable.
Periodic Table: Organized layout of elements by increasing atomic number and recurring chemical properties.
Periods: Horizontal rows on the periodic table; correspond to number of electron shells.
Protons: Positively charged particles in the atomic nucleus.
Quantum Tunneling: Phenomenon where electrons can pass through potential barriers due to wave-like behavior.
Valence Electrons: Electrons in the outermost shell involved in bonding.
These periodic trends form the foundation for predicting chemical reactivity, bonding behavior, and physical properties of elements in AP Biology and beyond.
FAQ
Transition metals often deviate from standard periodic trends due to the involvement of d-orbitals in electron configurations. These d-electrons do not shield outer electrons as effectively as s- and p-electrons, which affects atomic size and nuclear attraction.
As electrons are added to the (n-1)d subshell, the increasing nuclear charge pulls electrons closer without a proportional increase in shielding.
This leads to a relatively constant atomic radius across the transition series, unlike the more pronounced decrease in main group elements.
Ionization energies of transition metals may not increase consistently across a period because d-electron repulsion and half-filled subshell stability can lower the energy needed to remove an electron.
Some transition elements exhibit variable oxidation states due to the similar energies of s- and d-orbitals, which also influences ionization behavior.
Synthetic elements (transuranic elements) follow periodic trends theoretically, but their extreme instability makes experimental validation difficult. However, their predicted properties align with periodic behavior based on electron configurations.
They are placed in the actinide or transactinide series, with increasing atomic numbers.
As atomic number increases, relativistic effects (involving high-speed electrons) influence electron behavior, causing deviations from lighter element trends.
For example, element 114 (flerovium) was predicted to behave like lead, but shows noble-gas-like properties due to closed-shell stabilization effects.
Trends like ionization energy and atomic radius can be distorted due to these relativistic effects, making theoretical models more important for predicting properties.
Hydrogen and noble gases display unusual electronegativity values due to their unique atomic structures and bonding behavior.
Hydrogen, despite being in Group 1, has a relatively high electronegativity (2.1) because it only has one electron and one proton, resulting in a strong attraction to shared electrons in covalent bonds.
It does not follow group trends because it does not form cations like alkali metals; instead, it often shares or gains electrons.
Noble gases, particularly helium, neon, and argon, traditionally had no electronegativity values due to their full valence shells and lack of bonding behavior.
However, heavier noble gases like xenon can form compounds (e.g., XeF4), and are now assigned low but measurable electronegativity values due to their ability to participate in chemical bonding under specific conditions.
The shielding effect differs significantly depending on the type of orbital occupied by electrons, influencing periodic trends like atomic size and ionization energy.
s-electrons are closest to the nucleus and have the highest shielding ability.
p-electrons are farther from the nucleus than s-electrons in the same shell and are less effective at shielding outer electrons.
d- and f-electrons are even less effective at shielding because they are more diffused and lie closer to the nucleus in lower energy levels.
In transition metals (d-block), inner d-electrons don’t shield the outer s-electrons well, leading to increased effective nuclear charge and smaller-than-expected atomic radii.
In lanthanides and actinides (f-block), the poor shielding by f-electrons causes a phenomenon known as the lanthanide contraction, where atomic and ionic radii decrease more than expected across the series.
Elements in the same group have the same number of valence electrons, which determines how they interact in chemical reactions. While atomic size increases down the group, this does not significantly alter the type of chemical reactions they undergo.
The identical valence electron configuration in a group leads to similar bonding patterns and reactivity.
For example, all Group 1 elements (alkali metals) have one valence electron, making them highly reactive and prone to losing that electron to form +1 ions.
Though cesium is much larger than lithium, both react vigorously with water because the outer electron is easily lost.
Reactivity may increase down a group despite larger size because the outer electron is further from the nucleus and less tightly held, making it easier to remove.
This trend reinforces the concept that electron configuration, especially valence electrons, plays a greater role in reactivity than atomic size alone.
Practice Questions
Explain why atomic radius decreases across a period from left to right, even though the number of electrons increases.
As you move from left to right across a period on the periodic table, the atomic number increases, meaning more protons are added to the nucleus. This increase in positive nuclear charge pulls the negatively charged electrons closer to the nucleus. Although electrons are also being added, they enter the same principal energy level, so there is minimal additional shielding. The effective nuclear charge increases across the period, resulting in a stronger attraction between the nucleus and the electron cloud, which causes the atomic radius to decrease despite the added electrons.
Compare and contrast the ionization energies of lithium and cesium, and explain the trend based on their positions on the periodic table.
Lithium has a higher first ionization energy than cesium because its outermost electron is closer to the nucleus and experiences less shielding. In contrast, cesium has more electron shells, so its valence electron is farther from the nucleus and more weakly held due to greater shielding by inner electrons. Both elements are in Group 1 and have one valence electron, but the larger atomic size and reduced effective nuclear charge in cesium make its outer electron easier to remove. Therefore, ionization energy decreases as you move down a group due to increasing atomic radius and electron shielding.
