Periodic Trends Flashcards: Atomic Radius, Ionization Energy & Electronegativity

Question 1

What is a periodic trend?

Accepted Answer

A recurring pattern in element properties as atomic number increases across periods and down groups.

Question 2

What does covalent radius measure?

Accepted Answer

Half the distance between the nuclei of two identical atoms joined by a covalent bond.

Question 3

What is first ionization energy?

Accepted Answer

The minimum energy needed to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

Question 4

What is electronegativity?

Accepted Answer

An atom's ability to attract shared electrons toward itself in a chemical bond.

Question 5

What is a period on the periodic table?

Accepted Answer

A horizontal row. For main-group elements, moving across a period fills orbitals in the same principal electron shell.

Question 6

What is ionic radius?

Accepted Answer

A measure of an ion's size, usually inferred from distances between ions in crystals.

Question 7

What does electron affinity describe?

Accepted Answer

The energy change when an electron is added to an isolated gaseous atom to form a gaseous anion.

Question 8

What does metallic character describe?

Accepted Answer

How readily an element shows metallic behavior, especially losing valence electrons and forming cations.

Question 9

What is a group on the periodic table?

Accepted Answer

A vertical column. Main-group elements in one group usually share a valence-electron pattern and similar chemistry.

Question 10

What is effective nuclear charge?

Accepted Answer

The net positive pull an electron feels from the nucleus after shielding and electron–electron repulsion are taken into account.

Question 11

What are successive ionization energies?

Accepted Answer

The energies needed to remove electrons one after another from the same atom, then from its increasingly positive ions.

Question 12

How does electronegativity difference relate to bond polarity?

Accepted Answer

A larger difference generally produces a more uneven electron distribution and a more polar bond.

Question 13

Why do main-group elements in one group often behave similarly?

Accepted Answer

They have the same general number and arrangement of valence electrons, which drive much of their bonding and reactivity.

Question 14

What is electron shielding?

Accepted Answer

The reduction in nuclear attraction felt by an electron because other electrons lie between it and the nucleus and repel it.

Question 15

Why does the definition of ionization energy specify a gaseous atom?

Accepted Answer

It isolates the atom from bonding and intermolecular effects, so the energy reflects electron removal from that species itself.

Question 16

Where are metals and nonmetals generally found on the periodic table?

Accepted Answer

Metals occupy the left and center; nonmetals cluster toward the upper right, with metalloids near the boundary.

Question 17

What changes in the electron arrangement across a main-group period?

Accepted Answer

Electrons are added to the same principal shell while the nucleus gains one proton from one element to the next.

Question 18

How does atomic radius generally change from left to right across a period?

Accepted Answer

It decreases.

Question 19

How does first ionization energy generally change from left to right across a period?

Accepted Answer

It increases, although a few recurring subshell and electron-pairing exceptions interrupt the rise.

Question 20

How does electronegativity generally change from left to right across a period?

Accepted Answer

It increases for the elements normally assigned electronegativity values.

Question 21

What changes in the electron arrangement down a main-group group?

Accepted Answer

Each step adds a higher principal electron shell while preserving a similar valence-electron pattern.

Question 22

How does atomic radius generally change down a group?

Accepted Answer

It increases.

Question 23

How does first ionization energy generally change down a group?

Accepted Answer

It decreases.

Question 24

How does electronegativity generally change down a group?

Accepted Answer

It decreases.

Question 25

How does effective nuclear charge generally change across a main-group period?

Accepted Answer

It increases because nuclear charge rises while added electrons enter the same principal shell and do not fully shield one another.

Question 26

How does a cation's radius compare with its neutral parent atom?

Accepted Answer

The cation is smaller.

Question 27

How does electron addition generally change across a period?

Accepted Answer

It generally becomes more energetically favorable toward the right, but electron affinity has substantial exceptions and depends on the sign convention used.

Question 28

How does metallic character generally change from left to right across a period?

Accepted Answer

It decreases.

Question 29

Why are valence electrons generally farther from the nucleus down a group?

Accepted Answer

They occupy shells with higher principal quantum numbers, so the electron cloud extends farther outward.

Question 30

How does an anion's radius compare with its neutral parent atom?

Accepted Answer

The anion is larger.

Question 31

How does favorable electron addition generally change down a group?

Accepted Answer

It generally becomes less favorable as the added electron enters a larger, more shielded shell, though electron-affinity irregularities are common.

Question 32

How does metallic character generally change down a group?

Accepted Answer

It increases.

Question 33

Why does shielding change less than nuclear charge across a main-group period?

Accepted Answer

The added electrons enter the same principal shell, so they do not shield one another as effectively as inner-shell electrons do.

Question 34

Which has the larger atomic radius, Li or Na?

Accepted Answer

Na. It lies below Li and has an additional occupied electron shell.

Question 35

Which has the larger atomic radius, Na or Mg?

Accepted Answer

Na. Atomic radius generally decreases from left to right across Period 3.

Question 36

How does greater electron–nucleus distance affect electrostatic attraction?

Accepted Answer

It weakens the attraction, all else being equal.

Question 37

How do successive ionization energies for one element compare?

Accepted Answer

Each successive ionization energy is higher than the one before it because an electron is removed from an increasingly positive species.

Question 38

Which has the higher first ionization energy, Li or Na?

Accepted Answer

Li. Its valence electron is closer to the nucleus and less shielded.

Question 39

Which has the higher first ionization energy, Na or Mg?

Accepted Answer

Mg. Its greater effective nuclear charge holds the valence electrons more tightly.

Question 40

Why does forming a cation usually shrink an atom?

Accepted Answer

Electron loss reduces electron–electron repulsion and increases the nuclear pull per remaining electron; losing the outer shell can shrink it sharply.

Question 41

Which element is most electronegative on the Pauling scale?

Accepted Answer

Fluorine.

Question 42

Which is more electronegative, Li or Na?

Accepted Answer

Li. Electronegativity generally decreases down Group 1.

Question 43

Which is more electronegative, Na or Mg?

Accepted Answer

Mg. Electronegativity generally increases across Period 3.

Question 44

What does a large jump between successive ionization energies reveal?

Accepted Answer

The next electron would come from a lower, core shell; the number removed before the jump indicates the valence-electron count for a main-group atom.

Question 45

Why is Cl⁻ larger than neutral Cl?

Accepted Answer

The added electron increases repulsion within the valence shell while the nuclear charge stays the same.

Question 46

How do you compare the radii of isoelectronic species?

Accepted Answer

The species with more protons is smaller because the same number of electrons feels a stronger nuclear attraction.

Question 47

Why does an atom have no single sharp physical radius?

Accepted Answer

Its electron cloud has no hard edge, so atomic size depends on a defined measurement such as covalent, metallic, or van der Waals radius.

Question 48

Why are noble-gas electronegativities often omitted in introductory tables?

Accepted Answer

Electronegativity describes attraction in a bond, and many noble gases form too few ordinary bonds for a standard value to be useful on common scales.

Question 49

Why does forming an anion usually expand an atom?

Accepted Answer

The extra electron increases electron–electron repulsion and lowers the nuclear pull available per electron.

Question 50

Order O²⁻, F⁻, and Ne from largest to smallest radius.

Accepted Answer

O²⁻ > F⁻ > Ne. All have 10 electrons, and increasing proton count pulls that electron cloud inward.

Question 51

After which removal does Na show its first large ionization-energy jump?

Accepted Answer

After the first electron. Removing one valence electron leaves a stable core, so the second removal reaches that core.

Question 52

Does electronegativity difference create a universal ionic-versus-covalent cutoff?

Accepted Answer

No. A larger difference usually means more bond polarity, but bonding lies on a continuum and context matters.

Question 53

How do noble gases generally differ from halogens in electron affinity?

Accepted Answer

Adding an electron to a noble gas is generally unfavorable because it must begin a higher-energy shell; halogens usually gain one much more favorably.

Question 54

Which is smaller, Na⁺ or Mg²⁺?

Accepted Answer

Mg²⁺. Both have 10 electrons, but Mg²⁺ has one more proton.

Question 55

After which removal does Mg show its first large ionization-energy jump?

Accepted Answer

After the second electron. Mg has two valence electrons, so the third removal reaches a core shell.

Question 56

How does electronegativity differ from electron affinity?

Accepted Answer

Electronegativity is a relative measure of attraction for shared electrons in a bond; electron affinity is an energy change for adding an electron to an isolated gaseous species.

Question 57

How are atomic radius and first ionization energy generally related?

Accepted Answer

A larger radius usually means a lower first ionization energy because the valence electron is farther from the nucleus and easier to remove.

Question 58

Which has the more exothermic first electron affinity, F or Cl?

Accepted Answer

Cl. Fluorine's very compact 2p shell creates stronger electron–electron repulsion for the incoming electron, so this pair breaks the simple down-group expectation.

Question 59

Which has the higher first ionization energy, Be or B?

Accepted Answer

Be. B loses a higher-energy 2p electron, while Be loses a more penetrating 2s electron from a filled 2s subshell.

Question 60

Why does Group 1 metal reactivity generally increase down the group?

Accepted Answer

The valence electron is farther out and more shielded, so its first ionization energy falls and electron loss becomes easier.

Question 61

How are atomic radius and electronegativity generally related?

Accepted Answer

Smaller atoms usually attract bonding electrons more strongly, so electronegativity tends to rise as radius falls.

Question 62

Why is Na⁺ much smaller than neutral Na?

Accepted Answer

Na loses its entire third-shell valence level, leaving the smaller neon-like electron configuration.

Question 63

Which has the higher first ionization energy, Mg or Al?

Accepted Answer

Mg. Al's removed electron is a higher-energy 3p electron, while Mg loses a more penetrating 3s electron from a filled 3s subshell.

Question 64

Why does halogen reactivity generally decrease down Group 17?

Accepted Answer

Larger radius and greater shielding weaken attraction for an incoming electron, so oxidizing ability generally falls. Particular reactions still depend on bond energies and conditions.

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Periodic Trends Flashcards: Atomic Radius, Ionization Energy & Electronegativity

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