Periodic table trends: flashcards, quiz and summary

Flashcards

1. 1Why does atomic radius decrease across a period?

   Each step across a period adds a proton, so the nucleus pulls harder on the outer electrons. The extra electrons sit in the same energy level and shield each other poorly, so the atom gets smaller.

2. 2Why does atomic radius increase going down a group?

   Each step down a group adds a new energy level. Outer electrons sit farther from the nucleus, and inner electrons shield them from its pull, so the atom is larger.

3. 3What is effective nuclear charge, and how does it change across a period?

   Effective nuclear charge is the pull of the nucleus felt by the outer electrons. It increases across a period because protons are added while the extra electrons share the same shell and shield poorly.

4. 4Why is ionization energy higher across a period but lower down a group?

   Across a period, electrons are held more tightly by a stronger pull. Down a group, the outer electrons are farther away and shielded, so less energy is needed to remove one.

5. 5Why is a second ionization energy always higher than the first?

   After one electron leaves, the atom has a positive charge and holds its remaining electrons more tightly, so removing another costs more energy. A very large jump, as with sodium, shows that a core electron is being removed.

6. 6How does electronegativity difference relate to bond type?

   Electronegativity measures how strongly an atom pulls shared electrons. A large difference between two atoms suggests an ionic bond, while a small difference suggests a covalent bond, which is polar if the sharing is uneven.

7. 7Why are cations smaller and anions larger than their neutral atoms?

   Cations lose electrons, sometimes a whole shell, so they shrink. Anions gain electrons that repel one another, so they expand.

8. 8Among isoelectronic ions, which is smallest, and why?

   The ion with the most protons is smallest, because its stronger nuclear pull draws the same number of electrons in more tightly. Mg2+ is smaller than Na+, which is smaller than F-.

9. 9Why do alkali metals become more reactive going down group 1?

   Each alkali metal has one valence electron. Going down the group, that electron sits farther out and is more shielded, so it is lost more easily.

10. 10How do metallic character and nonmetallic character change across a period and down a group?

    Metallic character decreases across a period and increases down a group, because atoms lose electrons more easily as you move down. Nonmetallic character does the opposite.

Practice quiz

Pick an answer to check it.

1. 1What happens to atomic radius as you move from left to right across a period?

   1. It decreases, because the nucleus pulls the electrons closer
   2. It increases, because more electrons are added
   3. It increases first and then decreases
   4. It stays the same, because the energy level does not change
   Show the answer

   A. It decreases, because the nucleus pulls the electrons closer Each step adds a proton, which increases the pull on the outer electrons. Because the electrons share an energy level and shield poorly, the atom gets smaller.

2. 2Which element has the highest electronegativity?

   1. Oxygen
   2. Cesium
   3. Fluorine
   4. Chlorine
   Show the answer

   C. Fluorine Fluorine is the most electronegative element, at about 4.0 on the Pauling scale. Cesium is among the least electronegative.

3. 3How does sodium's second ionization energy compare with its first?

   1. It is about the same, because both remove one electron
   2. It is much higher, because a core electron is being removed
   3. It is lower, because the atom is already positive
   4. It is zero, because sodium has only one electron to lose
   Show the answer

   B. It is much higher, because a core electron is being removed After the first electron leaves, the second must come from an inner shell that is much closer to the nucleus. That causes a very large jump in the required energy.

4. 4Which ion is the smallest?

   1. Na+
   2. O2-
   3. F-
   4. Mg2+
   Show the answer

   D. Mg2+ Mg2+, Na+, F-, and O2- all have 10 electrons. Mg2+ has the most protons, so its nucleus pulls the electrons in most tightly and it is the smallest.

5. 5How does reactivity change going down group 17, the halogens?

   1. It increases, because the atoms get larger
   2. It increases first, then decreases
   3. It stays constant across the group
   4. It decreases, with fluorine the most reactive
   Show the answer

   D. It decreases, with fluorine the most reactive Halogen reactivity decreases going down the group, so fluorine at the top is the most reactive. This is the opposite of the trend in the alkali metals.

6. 6Why is a cation smaller than its neutral atom?

   1. It gains a new energy level
   2. Its nucleus shrinks
   3. It loses electrons, sometimes a whole shell
   4. It gains electrons that repel each other
   Show the answer

   C. It loses electrons, sometimes a whole shell Cations lose electrons, sometimes an entire shell, so fewer electrons remain. The nucleus pulls the remaining electrons more tightly, making the ion smaller.

Test traps

What teachers like to test on this, and the mistakes students make.

• Students often say atomic radius increases across a period because more electrons are added. Radius decreases, because the extra protons pull the electrons closer.
• Students confuse shielding with nuclear pull. Electrons in the same shell shield each other poorly, so the pull across a period wins.
• Students expect ionization energy to rise going down a group. It falls, because the outer electrons are farther away and more shielded.
• Students assume anions are smaller because they are negative. Extra electrons repel each other, so anions are larger than their neutral atoms.
• When comparing isoelectronic ions, students compare charges instead of proton counts. The ion with more protons is smaller.
• Students label bonds as ionic or covalent without comparing electronegativity differences. A large difference suggests ionic, and a small difference suggests covalent.

Summary

• Elements are arranged by atomic number, and each period adds one proton and one electron to the same energy level.
• Effective nuclear charge increases across a period because the extra electrons shield each other poorly, so the nucleus pulls harder on the outer electrons.
• Atomic radius decreases across a period and increases down a group because each step down adds an energy level.
• Ionization energy and electronegativity increase across a period and decrease down a group.
• Metallic character decreases across a period and increases down a group, which explains the reactivity trends of alkali metals and halogens.
• Cations are smaller than their neutral atoms, anions are larger, and among isoelectronic ions the one with more protons is smaller.
• The size of the electronegativity difference between two atoms indicates whether a bond is ionic or covalent.

Key terms

Effective nuclear charge

The pull of the nucleus felt by the outer electrons. It increases across a period because protons are added while the extra electrons sit in the same energy level and shield poorly.

Shielding

The reduction of the nucleus pull on outer electrons by inner electrons. Shielding grows as more energy levels are added going down a group.

Ionization energy

The energy needed to remove an electron from a gaseous atom. It increases across a period and decreases down a group.

Electronegativity

How strongly an atom attracts shared electrons in a bond. It increases across a period and decreases down a group, with fluorine the highest.

Electron affinity

The energy change when an atom gains an electron. It generally becomes more negative across a period, with halogens releasing the most energy.

Isoelectronic ions

Ions that have the same number of electrons. Among them, the ion with more protons is smaller.

How to study it

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Questions