Periodic Trends
Related Examples and Practice Problems
Topic Summary & Highlights
and Help Videos
Core Concept
Atomic trends arise from the periodic table's structure and describe how properties like atomic radius, ionization energy, electronegativity, electron affinity, ion size, metallic character, and effective nuclear charge change across periods and down groups, influenced by the balance between nuclear charge, electron shielding, and energy levels.
Practice Tips
Use the periodic table as a visual aid to see where each trend increases or decreases.
Practice predicting properties of unknown elements based on their position relative to known elements.
Understand the "why" for each trend—focus on how nuclear charge and electron shells influence each trend.
Practice with examples: Compare the properties of elements in the same group and period, like comparing sodium (Na) and chlorine (Cl) for trends across a period, or lithium (Li) and potassium (K) for trends down a group.
Topic Overview Podcast
Topic Related Resources
|
LABORATORY
|
DEMONSTRATIONS
|
ACTIVITIES
|
VIRTUAL SIMULATIONS
|
| Property | Across a Period | Down a Group |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionization Energy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Electron Affinity | Becomes more negative | Becomes less negative |
| Ion Size | Decreases (for same charge) | Increases |
| Metallic Character | Decreases | Increases |
| Effective Nuclear Charge (Z_eff) | Increases | Remains relatively constant |
Coulomb’s Law
Imagine you have two tiny magnets—that's kind of what electric charges are like. Coulomb's Law tells you how those magnets (or charges) will push or pull on each other.
Rule 1: Opposites Attract, Likes Repel
Opposite Charges attract each other (pull closer). Like Charges will repel each other (push away).
In an atom, the positive nucleus and the negative electrons always attract!
Rule 2: More Charge = Stronger Force
This rule is about how hard the push or pull is.
If you have a small charge, the force (push or pull) will be weak.
If you have a big charge (lots of protons or lots of electrons), the force will be super strong.
Simply put: If you double the amount of charge, the strength of the electric push or pull will double, too!
ule 3: Distance Makes a BIG Difference
It explains how the distance between the charges affects the force. The farther apart two charges are, the weaker the force is—but it weakens very, very fast.
Imagine shining a flashlight 🔦:
If you put the flashlight right next to a wall, the spot of light is bright and strong. (Close distance = Strong Force)
If you walk 10 feet away, the light on the wall is much weaker and more spread out. (Far distance = Weak Force)
In electric terms: If you double the distance between a nucleus and an electron, the electric force doesn't just get half as weak; it gets four times (1/4) weaker! That's why the electrons in the outer shells of an atom aren't held as tightly as the ones close to the nucleus.
Atomic Radius
Definition: The distance from the nucleus to the outermost electron shell.
Trend:
Across a Period (Left to Right): Atomic radius decreases.
Why? Increased nuclear charge (more protons) pulls electrons closer to the nucleus without adding additional electron shells.
Down a Group (Top to Bottom): Atomic radius increases.
Why? Each row down adds a new electron shell, increasing the size of the atom.
Practice Example
Indicate which has the larger atomic radius and explain using atomic structure: Na or Mg
Worked out Solution:
Both Sodium (Na) and Magnesium (Mg) are in Period 3, meaning their valence electrons are in the same electron shell (n = 3). As you move from Na (11 protons) to Mg (12 protons), the nuclear charge increases (11 > 12). Since the shielding by inner electrons is essentially the same, this results in a higher Effective Nuclear Charge (Zeff) for Mg. The stronger positive charge in the Mg nucleus pulls its electron cloud closer, causing the atom to have a smaller atomic radius than Na.
Ionization Energy
Definition: The energy required to remove an electron from an atom in its gaseous state.
Trend:
Across a Period: Ionization energy generally increases.
Why? As atomic radius decreases, electrons are closer to the nucleus and more tightly held, requiring more energy to remove.
Down a Group: Ionization energy decreases.
Why? Electrons are further from the nucleus and experience less attraction, making them easier to remove.
Exceptions: There are slight decreases in ionization energy within periods due to electron repulsion in p-orbitals and half-filled orbital stability.
Practice Example
Indicate which has the smaller ionization energy and explain using atomic structure: Na or K
Worked out Solution:
Potassium (K) is located below Sodium (Na) in Group 1 of the periodic table, meaning it has one more electron shell. This difference in atomic structure explains the lower ionization energy:
Greater Distance: K's outermost electron is in the 4s shell, making it farther from the positive nucleus than Na's 3s electron. According to Coulomb's Law, the attractive force weakens rapidly with distance.
Greater Shielding: K has more inner electron shells (core electrons) which effectively block, or shield, the outermost electron from the nucleus's full positive charge.
Both the increased distance and the greater shielding make the outermost electron in K much less tightly held, requiring less energy to remove it.
Electronegativity
Definition: A measure of an atom's ability to attract and bond with electrons when in a compound.
Trend:
Across a Period: Electronegativity increases.
Why? Nonmetals on the right side need only a few electrons to achieve a stable electron configuration, so they attract electrons more strongly.
Down a Group: Electronegativity decreases.
Why? Increased atomic size reduces the pull of the nucleus on bonding electrons.
Practice Example
Indicate which has the smaller electronegativity and explain using atomic structure: C or Si?
Worked out Solution:
Electronegativity decreases as you move down a group. Since both Carbon (C) and Silicon (Si) are in Group 14, and Si is below C, Si has a lower ability to attract bonding electrons. This is due to:
Greater Distance: Si's bonding electrons are in the 3rd shell, making them farther from the nucleus than C's 2nd shell electrons. The attraction weakens dramatically with this increase in distance.
Increased Shielding: Si has a full n=2 shell (8 core electrons) that effectively shields its bonding electrons from the nucleus's positive charge.
This weaker attraction means Si is less able to pull in a shared pair of electrons, resulting in a lower electronegativity.
Electron Affinity (EA)
Definition: The energy change that occurs when an atom gains an electron, typically resulting in a negative value (energy is released).
Trend:
Across a Period: Electron affinity becomes more negative (favorable) in general.
Why? Atoms on the right side of the periodic table (nonmetals) have a stronger tendency to gain electrons to complete their valence shells.
Down a Group: Electron affinity becomes less negative.
Why? Larger atomic size reduces the attraction between the nucleus and added electrons.
Note: Noble gases do not follow the trend as they have full valence shells and generally do not gain electrons.'
Practice Example
Indicate which has the smaller electronegativity and explain using atomic structure: C or Si?
Worked out Solution:
Both
Metallic Character
Definition: The tendency of an element to lose electrons and form cations (typical of metals).
Trend:
Across a Period: Metallic character decreases.
Why? Moving across a period, elements more readily gain electrons rather than lose them, as they approach a full valence shell.
Down a Group: Metallic character increases.
Why? Larger atoms with lower ionization energies more easily lose electrons, a characteristic of metals.
Ionic Size
Definition: The tendency of an element to lose electrons and form cations (typical of metals).
When an atom (neutral) becomes an ion (charged) it either loses or gains electrons.
Atoms (trend to be metals) that LOSE electrons become CATIONS (positively charged). Atoms get SMALLER when an electron is LOST.
Why? xx
Atoms (tend to be non-metals) that GAIN electrons become ANIONS (negatively charged). Atoms get LARGER when an electron is GAINED.
Why? xx
Effective Nuclear Charge (Zeff)
CHECK OUT EFFECTIVE NUCLEAR CHARGE TOPIC.
Definition: The net positive charge experienced by valence electrons after accounting for shielding from inner electrons.
Trend:
Across a Period: Z_eff increases.
Why? With each additional proton, the nuclear charge increases, but added electrons are in the same shell and don’t shield each other effectively.
Down a Group: Z_eff slightly decreases or stays roughly constant.
Why? New shells reduce the nuclear pull on outer electrons despite the increase in nuclear charge.
