Topic 3 Periodic Table

Question 1

What is the periodic trend for atomic radius across a period?

Answer 1

The atomic radius decreases from left to right due to the increasing effective nuclear charge (Zeff), which pulls electrons closer to the nucleus.

Question 2

What is electronegativity, and which element has the highest value?

Answer 2

Electronegativity is the ability of an atom in a compound to attract electrons. Fluorine has the highest electronegativity value of 4.0

Question 3

How does the atomic radius change down a group in the periodic table?

Answer 3

Atomic radius increases down a group because of the addition of electron shells, which increases the distance between the nucleus and the outermost electrons

Question 4

Write the shorthand electron configuration for Magnesium

Answer 4

[Ne] 3s²

Question 5

What is the trend for ionization energy across a period?

Answer 5

Ionization energy increases across a period due to a higher effective nuclear charge, making it harder to remove an electron.

Question 6

Define the octet rule

Answer 6

Atoms are most stable when they have a full outer shell of electrons, typically eight, except for the first shell, which holds two electrons

Question 7

Why is Cesium (Cs) more electropositive than Sodium (Na)?

Answer 7

Cesium’s outer electron is in the 6s orbital, much further from the nucleus than Sodium’s 3s orbital electron, making it easier to lose and thus more electropositive.

Question 8

What are the characteristic charges of ions formed by Group 1 and Group 7 elements?

Answer 8

Group 1 forms +1 ions, and Group 7 forms -1 ions

Question 9

How do you calculate the effective nuclear charge (Zeff)?

Answer 9

Zeff = Z - S, where Z is the atomic number, and S is the number of shielding (non-valence) electrons.

Question 10

What is the difference between ionic and covalent bonding?

Answer 10

Ionic bonding occurs when electrons are transferred between atoms, creating ions. Covalent bonding involves the sharing of electrons between atoms.

Question 11

What is Hund’s Rule?

Answer 11

Hund’s Rule states that electrons will occupy orbitals singly before pairing up to maximize the number of unpaired electrons.

Question 12

Which groups in the periodic table contain elements that typically do not form ions?

Answer 12

Group 4 (does not form ions due to high energy requirements) and Group 8 (noble gases, stable with a full outer shell).

Question 13

What is the general trend for electronegativity down a group?

Answer 13

Electronegativity decreases down a group because the increasing atomic radius reduces the attraction between the nucleus and bonding electrons.

Question 14

Define first ionization energy.

Answer 14

The first ionization energy is the minimum energy required to remove one electron from a gaseous neutral atom in its ground state.

Question 15

Why do Group 2 elements form +2 ions?

Answer 15

Group 2 elements lose two electrons from their outer s-orbital to achieve a stable electron configuration.

Question 16

What is the electron configuration of Neon?

Answer 16

1s² 2s² 2p⁶

Question 17

Describe the trend for melting points in transition metals.

Answer 17

Transition metals generally have high melting points because of strong metallic bonding due to d-electron delocalization.

Question 18

What are p-orbitals, and how are they shaped?

Answer 18

P-orbitals are regions where electrons are likely to be found. They have a dumbbell shape and can hold up to six electrons.

Question 19

Explain why ionic radius decreases for cations compared to their parent atoms.

Answer 19

When atoms lose electrons to form cations, there is less electron-electron repulsion, and the remaining electrons are pulled closer to the nucleus.

Question 20

What is the trend for ionization energy down a group?

Answer 20

Ionization energy decreases down a group because the outermost electron is further from the nucleus, reducing the nuclear attraction.

Question 21

What is noble gas shorthand in electron configuration?

Answer 21

It uses the configuration of the nearest preceding noble gas as a base, followed by the remaining electron configuration. Example: [Ne] 3s² for Magnesium.

Question 22

What is the significance of the Aufbau principle?

Answer 22

It states that electrons fill orbitals starting with the lowest energy levels before moving to higher ones.

Question 23

Why do halogens form -1 ions?

Answer 23

Halogens gain one electron to complete their valence shell, achieving a stable noble gas configuration.

Question 24

What is the periodic trend for metallic character?

Answer 24

Metallic character increases down a group and decreases across a period from left to right.

Question 25

Which element is a metalloid and used in semiconductors?

Answer 25

Silicon (Si), because it partially conducts electricity under certain conditions.

Question 26

Define bond polarity

Answer 26

Bond polarity refers to the unequal sharing of electrons in a covalent bond, resulting in a partial positive charge on one atom and a partial negative charge on the other.

Question 27

Why does Boron (B) form covalent rather than ionic bonds?

Answer 27

Boron’s small size and high ionization energy make it difficult to lose three electrons, so it shares electrons instead.

Question 28

What is a key chemical property of Group 1 elements?

Answer 28

Alkali metals are highly reactive, losing one electron to form +1 ions

Question 29

Why do reactivity and metallic character increase down Group 1?

Answer 29

The outer electron is further from the nucleus in larger atoms, making it easier to lose.

Question 30

Write a balanced equation for the reaction of Sodium (Na) with water.

Answer 30

2Na + 2H₂O → 2NaOH + H₂

Question 31

Why are Group 2 elements less reactive than Group 1?

Answer 31

They have two electrons to lose, requiring more energy to form +2 ions.

Question 32

Write the electron configuration for Magnesium (Mg).

Answer 32

1s² 2s² 2p⁶ 3s² or [Ne] 3s²

Question 33

What is a typical reaction of Group 2 elements with oxygen?

Answer 33

2Mg + O₂ → 2MgO

Question 34

Why do halogens form -1 ions?

Answer 34

Halogens gain one electron to complete their valence shell.

Question 35

What happens to the reactivity of halogens down the group?

Answer 35

Reactivity decreases because the atomic radius increases, reducing the attraction for electrons.

Question 36

Write a balanced equation for the reaction of Chlorine (Cl₂) with Potassium Bromide (KBr)

Answer 36

Cl₂ + 2KBr → 2KCl + Br₂

Question 37

Why are noble gases chemically inert?

Answer 37

They have a full outer electron shell, making them stable and unreactive

Question 38

Name two applications of noble gases

Answer 38

Helium in balloons and Neon in lighting

Question 39

Which noble gas is used in high-energy lasers?

Answer 39

Argon

Question 40

What is a unique property of transition metals?

Answer 40

They can exist in multiple oxidation states, such as Fe²⁺ and Fe³⁺

Question 41

Why do transition metals form colored compounds?

Answer 41

Their d-electrons absorb specific wavelengths of light, causing transitions between energy levels

Question 42

Give an example of a transition metal forming a complex ion

Answer 42

[Cu(H₂O)₆]²⁺

Question 43

Give an example of an ionic compound formed between Group 2 and Group 6 elements

Answer 43

Mg + O₂ → 2MgO (Magnesium Oxide)

Question 44

Why is Beryllium (Be) unique among the alkaline earth metals?

Answer 44

Its outer electrons are close to the nucleus, making it harder to lose them, so it often forms covalent bonds

Question 45

Write the balanced equation for the reaction of Iron(III) ion (Fe³⁺) with Chlorine (Cl⁻) to form an ionic compound

Answer 45

Fe³⁺ + 3Cl⁻ → FeCl₃