S3.1 The periodic table: SL Part 1

Question 1

What do the columns in the periodic table represent, and how many are there?

Answer 1

The columns represent groups, and there are 18 groups in the periodic table.

Question 2

How are elements arranged in the periodic table, and what does each row signify?

Answer 2

Elements are arranged horizontally by atomic number. Each row is called a period, showing similar properties among the elements.

Question 3

Identify the property that separates metals from non-metals in the periodic table.

Answer 3

Metals are on the left-hand side and exhibit properties like malleability and high conductivity, whereas non-metals are on the right, typically being brittle and with low conductivity.

Question 4

What are metalloids, and where are they located in the periodic table?

Answer 4

Metalloids have properties of both metals and non-metals and are located along the diagonal staircase between metals and non-metals.

Question 5

How does the electron configuration determine the block an element belongs to in the periodic table?

Answer 5

The block (s, p, d, f) an element belongs to is determined by the sublevel (s, p, d, f) that the outermost valence electrons occupy.

Question 6

How can you predict the electron configuration of an element up to atomic number 36 using the periodic table?

Answer 6

By using the element’s position within the periodic table, you can deduce its electron configuration, considering the period number indicates the outermost energy level.

Question 7

What is the significance of group numbers in the periodic table for electron configuration?

Answer 7

Group numbers help identify the number of valence electrons an element has, crucial for predicting chemical reactivity and bonding behavior.

Question 8

How do the alkali metals (Group 1) differ in their electron configurations?

Answer 8

Alkali metals have a single electron in their outermost s-sublevel, which is indicated by their position in the s-block of the periodic table.

Question 9

What pattern is observed in the electron configurations of Group 17 elements (halogens)?

Answer 9

Halogens, located in Group 17, have a total of seven valence electrons: two in the s-sublevel and five in the p-sublevel. This configuration is a reflection of their position in the p-block and is crucial for their high reactivity.

Question 10

How can the block method help determine the electron configuration of an element?

Answer 10

The block method involves identifying the element’s period and block (s, p, d, f) in the periodic table, then deducing its electron configuration based on its position.

Question 11

Atomic Radius

Answer 11

The distance from an atom’s nucleus to the outermost orbital of one of its electrons. Usually measured as the half-distance between the nuclei of two of the same atoms bonded together.

Question 12

Explain the periodicity of atomic radii

Answer 12

Atomic radius increases down a group due to additional electron shells, and decreases across a period due to increased effective nuclear charge pulling electrons closer.

Question 13

Explain the periodicity of ionic radii

Answer 13

Ionic radii are affected by the gain or loss of electrons, altering electron-electron repulsion and attraction to the nucleus

Question 14

Why the ionic radii decrease as the nuclear charge (number of protons) increases for isoelectronic species?

Answer 14

This is because a higher nuclear charge exerts a stronger attractive force on the same number of electrons, pulling them closer to the nucleus and resulting in a smaller ionic radius

Question 15

Ionisation energy

Answer 15

The energy required with remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Question 16

Electron affinity

Answer 16

The amount of energy released when a neutral atom gains an electron to form a negatively charged ion.

Question 17

Electronegativity

Answer 17

A measure of how much an atomic nucleus attracts the shared electrons that are involved in a covalent bond.

Question 18

The periodicity of Ionization Energy

Answer 18

It increases across a period due to higher nuclear charge attracting electrons more strongly, and decreases down a group as electrons are further from the nucleus and more shielded by inner electrons.

Question 19

The periodicity of Electron Affinity

Answer 19

It generally increases across a period, indicating a greater tendency to gain electrons due to a higher effective nuclear charge, and decreases down a group due to increased distance of the added electron from the nucleus and weaker attraction.

Question 20

The periodicity of Electronegativity

Answer 20

Like ionization energy and electron affinity, electronegativity increases across a period due to increased nuclear charge, and decreases down a group as atomic radius increases and the effective attraction between the nucleus and shared electrons diminishes.

Question 21

What metals are these: hydrogen (H), lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr)?

Answer 21

Alkali metals (group 1)

Question 22

What are the properties of the elements of halogens

Answer 22

Highly reactive and form ionic salts when they react with metals.

Question 23

These elements can form ions with different ionic charges. Metals have high melting points and high conductivities.

Answer 23

Transition  elements

Question 24

Colourless monoatomic gases that have little to no reactivity with other elements.

Answer 24

Noble gases