4A2 Trends and Patterns of the Periodic Table

Question 1

Define:

atomic radius

Answer 1

The size of a neutral atom, measured from the nucleus to the outermost electron.

Atomic radius increases down a group due to additional orbitals and decreases across a period due to stronger nuclear pull.

Question 2

Define:

ionic radius

Answer 2

The radius of an ion, which changes when an atom gains or loses electrons to form an ion.

Gaining electrons increases ionic radius (anion), while losing electrons decreases it (cation).

Question 3

True or False:

Ionic radius is the same as atomic radius.

Answer 3

False

Ionic radius changes based on electron gain/loss, while atomic radius pertains to neutral atoms.

Question 4

Why does ionic radius increase down a group?

Answer 4

Because the number of electron orbitals increases, resulting in a larger ion size.

Moving down a group means entering higher periods with additional electron shells.

Question 5

Which has a smaller atomic radius:

sodium (Na) or chlorine (Cl)

Answer 5

Chlorine (Cl)

Chlorine has more protons, pulling electrons closer and decreasing its radius.

Question 6

How does ionic radius differ between cations and anions?

Answer 6

Cations have smaller ionic radii, while anions have larger ionic radii compared to their atoms.

This is due to the loss of electrons in cations and the gain of electrons in anions.

Extra electrons in anions repel each other, expanding the electron cloud.

Question 7

Why does atomic radius decrease across a period?

Answer 7

Because the number of protons increases, pulling electrons closer to the nucleus.

This stronger nuclear pull reduces the size of the electron cloud.

Question 8

List two factors affecting the ionic radius.

Answer 8

1. Electron gain or loss.
2. Spin and coordination number.

More surrounding bonds or higher spin increases ionic radius.

Question 9

How does the number of protons affect ionic radius in isoelectronic ions?

Answer 9

More protons result in a smaller ionic radius.

Protons pull electrons closer, reducing the size of the ion.

Question 10

True or False:

Fluorine (F⁻) has a larger ionic radius than Oxygen (O²⁻).

Answer 10

False

Oxygen has fewer protons than fluorine, so it exerts less pull on its electrons, leading to a larger ionic radius.

Question 11

What is the relationship between atomic radius and electronegativity?

Answer 11

Atoms with smaller atomic radii tend to have higher electronegativity.

A tighter hold on electrons leads to a greater ability to attract electrons in a bond.

Question 12

Why does fluorine (F) have a smaller atomic radius than lithium (Li)?

Answer 12

It has more protons, pulling electrons closer to the nucleus.

The increased nuclear charge reduces the size of the electron cloud.

Question 13

Define:

diagonal relationship on the periodic table

Answer 13

It is the similarities in properties between certain pairs of diagonally adjacent elements in the second and third periods of the periodic table.

It is caused by similar ionic potentials and opposing trends in atomic properties across periods and down groups.

For example, atomic radii increase down a group but decrease across a period, creating similarities diagonally.

Question 14

List five element pairs that exhibit diagonal relationships.

Answer 14

1. Lithium-Magnesium
2. Beryllium-Aluminum
3. Boron-Silicon
4. Carbon-Phosphorus
5. Nitrogen-Sulfur

These pairs share properties such as:

• Atomic radii
• Amphoteric behavior
• Covalent character
• Electronegativity

Question 15

List three properties shared by Beryllium and Aluminum.

Answer 15

1. Amphoteric oxides and hydroxides.
2. Covalent bonding tendencies.
3. Solubility in organic solvents.

Beryllium and Aluminum form compounds with low melting points and exhibit amphoteric behavior.

Question 16

Define:

amphoteric behavior

Answer 16

A substance’s ability to act as both an acid and a base.

Beryllium and Aluminum oxides exhibit this property.

Question 17

Define:

metallic character

Answer 17

It is an element’s tendency to lose electrons and exhibit properties like malleability and electrical conductivity.

Metallic character depends on atomic radii and ionization energy.

Question 18

How does atomic radius trend affect metallic character?

Answer 18

As atomic radius increases, metallic character increases.

Larger atomic radii reduce the nuclear pull on valence electrons, making it easier to lose electrons.

Question 19

How does metallic character vary across a period?

Answer 19

It decreases from left to right.

As you move across a period, atomic radii decrease, and ionization energy increases, making it harder for elements to lose electrons.

Question 20

Fill in the blank:

Metallic character increases when moving _____ a group.

Answer 20

down

As you move down a group, atomic radii increase, and ionization energy decreases, making it easier for elements to lose electrons.

Question 21

True or False:

Boron and Silicon are semiconductors.

Answer 21

True

Both Boron and Silicon are nonmetals, semiconductors, and exhibit allotropy.

Allotropy refers to the ability of an element to exist in different physical forms.

Question 22

Which elements form triple bonds and exhibit allotropy?

Answer 22

Carbon and phosphorus

Both elements show strong sigma and pi bonding and exist in different forms.

Question 23

How does boiling point trend across a period?

Answer 23

It increases from left to the middle of a period and decreases sharply in the nonmetals.

Metals have stronger bonds, leading to higher boiling points, while nonmetals have weaker intermolecular forces.

Question 24

Fill in the blank:

Nitrogen and Sulfur both exhibit _______.

Answer 24

diamagnetism

Diamagnetism is a property where substances are weakly repelled by a magnetic field.

Question 25

# Define: electronegativity

Answer 25

The ability of an atom to **attract and hold electrons** in a chemical bond. ## Footnote Electronegativity is measured using the Pauling scale and varies across elements in the periodic table.

Question 26

How does electronegativity **change down a group**?

Answer 26

It decreases. ## Footnote Due to the increase in the number of electron shells, which increases nuclear shielding.

Question 27

# True or False: Electronegativity **decreases across a period** in the periodic table.

Answer 27

False ## Footnote Electronegativity increases across a period due to the increase in nuclear charge and a decrease in atomic radius.

Question 28

# Which element has a **lower** electronegativity: potassium or phosphorus

Answer 28

potassium ## Footnote Electronegativity decreases down a group and increases across a period; potassium is farther left than phosphorus.

Question 29

Why does fluorine have a **high electronegativity**?

Answer 29

* Its small atomic radius. * Its high nuclear charge. ## Footnote Fluorine’s strong pull on electrons is due to minimal shielding and a high proton count. It has the highest electronegativity of any element with a value of 3.98 on the Pauling scale.

Question 30

# Define: nuclear shielding

Answer 30

The reduction in **effective nuclear charge** on outer electrons due to inner electron shells. ## Footnote Greater shielding reduces electronegativity by lessening the nucleus's pull on outer electrons.

Question 31

# Fill in the blank: The **least electronegative** element is \_\_\_\_\_\_\_.

Answer 31

Francium ## Footnote Francium has an electronegativity value close to 0, indicating its inability to hold electrons tightly.

Question 32

# True or False: Noble gases typically have electronegativity values **greater than 0**.

Answer 32

False ## Footnote Most noble gases have a value of 0 due to their full electron shells, which do not require attracting more electrons.

Question 33

How does the number of electron shells **affect electronegativity**?

Answer 33

More shells reduce electronegativity due to **increased shielding**. ## Footnote The greater the number of electron shells, the weaker the attraction between the nucleus and the outermost electrons.

Question 34

# Fill in the blank: Electronegativity **increases across a period** due to increasing \_\_\_\_\_\_ \_\_\_\_\_\_.

Answer 34

nuclear charge ## Footnote A higher number of protons increases the pull on bonding electrons.

Question 35

# Which element has a **greater** electronegativity: sulfur or oxygen

Answer 35

oxygen ## Footnote Oxygen is farther to the right on the periodic table, and electronegativity increases across a period.

Question 36

# Define: ionization energy

Answer 36

The minimum energy required to **remove the outermost electron** from a gaseous atom of an element. ## Footnote Ionization energy is measured in kilojoules per mole and is influenced by atomic structure.

Question 37

List *three* factors **affecting** ionization energy.

Answer 37

1. Number of protons 1. Number of electron shells 1. Number of electrons ## Footnote These factors influence how strongly the nucleus attracts the outermost electrons.

Question 38

How does the number of protons **affect ionization energy**?

Answer 38

More **protons increase ionization energy** by creating a stronger positive charge that attracts electrons. ## Footnote A higher proton count in the nucleus increases the hold on electrons.

Question 39

What is the relationship between **ionization energy and atomic radius**?

Answer 39

Ionization energy **decreases as atomic radius increases**. ## Footnote A larger radius places electrons farther from the nucleus, reducing nuclear attraction.

Question 40

How does the number of **electron shells affect ionization** energy?

Answer 40

**More electron shells decrease ionization energy** because outer electrons are farther from the nucleus. ## Footnote Additional electron shells weaken the nucleus's hold on valence electrons.

Question 41

# Fill in the blank: Ionization energy \_\_\_\_\_\_\_ **across a period** from left to right.

Answer 41

increases ## Footnote This is due to an increase in the number of protons.

Question 42

# True or False: Ionization energy decreases as you **move down a group**.

Answer 42

True ## Footnote Due to an increase in the number of electron shells.

Question 43

What happens to ionization energy as **atomic radius increases**?

Answer 43

It decreases. ## Footnote A larger atomic radius increases the distance between the nucleus and valence electrons, weakening the attraction.

Question 44

Why does cesium have **lower ionization energy than lithium**?

Answer 44

Cesium has **more electron shells**, increasing the distance between the nucleus and its valence electrons. ## Footnote The outermost electrons in cesium experience less nuclear attraction than in lithium.

Question 45

# Which element has **higher** ionization energy: fluorine or lithium

Answer 45

fluorine ## Footnote Fluorine has more protons, resulting in a stronger nuclear pull on its electrons.

Question 46

# Define: second ionization energy

Answer 46

The energy required to remove an electron from a **singly charged gaseous cation**. ## Footnote It is always greater than the first ionization energy due to increased nuclear attraction.

Question 47

# True or False: Nonmetals have **low ionization energy**.

Answer 47

False ## Footnote Nonmetals have high ionization energy, meaning they strongly attract electrons and do not lose them easily.

Question 48

# Fill in the blank: The more \_\_\_\_\_\_\_ an ion becomes, the harder it is to **remove additional electrons**.

Answer 48

positive ## Footnote The increasing positive charge strengthens the hold on remaining electrons, raising ionization energy.