Chemistry Video 6

Question 1

Li, Na, K

Answer 1

All shiny, conduct heat and electricity. Reacts with oxygen atoms in 2:1 ratio

Question 2

Ca, Sr, Ba

Answer 2

All soft and shiny, reacts with oxygen atoms in 1:1 ratio

Question 3

Dmitri Mendeleev

Answer 3

Organized the final periodic table. His table made undiscovered elements

Question 4

Periodic table element block

Answer 4

Symbol of element in centre, under symbol is atomic mass, atomic number in corner

Question 5

Group 1

Answer 5

Alkali metals; shiny metals, conduct heat, conduct electricity, react 2:1 with oxygen

Question 6

Group 2

Answer 6

Alkaline earth metals; metallic properties, react with oxygen in 1:! ration

Question 7

Groups 3-12

Answer 7

Transition metals; most of them are easily oxidized.

Question 8

Group 15

Answer 8

Pnictogens

Question 9

Group 16

Answer 9

Chalcogens

Question 10

Group 17

Answer 10

Halogens

Question 11

Group 18

Answer 11

Inert Gases, noble gases, unreactive. Monoatomic.

Question 12

Atomic radius/covalent radius

Answer 12

One half the distance between the nuclei of 2 identical atoms that are bonded to each other.

Increases down the group. Decreases going left to right on a period.

Question 13

Bond length

Answer 13

Distance between the nuclei of 2 atoms forming a bond. Sum of the atomic radii for the 2 atoms that are participating in the bond.

Question 14

Effective nuclear charge, (Zeff) and trend

Answer 14

Pull exerted by the nucleus on an electron. Core electrons shield valence electrons from the pull of the nucleus.

Zeff = number of protons - core electrons

Increases from left to right across a period, decreases going down a group.

Reasoning: Across a period, the numbers of protons are increasing with no increase in a shielding effect, which results in electrons being pulled closer to the nucleus due to a stronger attraction. Going down a group, more shielding causes the effective nuclear charge to decrease. As electrons get further away from the nucleus, the attractive force between protons and electrons naturally lessens.

Question 15

Ionic radius

Answer 15

Cation = smaller ionic radius compared to non-ion version

Anion = larger ionic radius compared to non-ion version

When a neutral atom loses an electron and becomes a cation, its radius will decrease due to an increase in the effective nuclear charge allowing the protons to pull in the electrons closer to the nucleus.

Conversely, if an atom gains electrons and becomes an anion, its atomic radius will increase. This is due to a decrease in effective nuclear charge and causing a decrease in the pull from the protons in the nucleus.

Metals typically form cations resulting in their ionic radius to be less than their atomic radius and non-metals typically form anions resulting in their ionic radius to be greater than their atomic radius.

Question 16

Isoelectronic species

Answer 16

Atoms/ions with the same electron configuration, but differing numbers of protons

the most positively charged atom will have the smallest radius because it has the most protons with the same number of electrons, and therefore the greatest attraction to the nucleus.

Question 17

Ionization energy

Answer 17

Energy required to remove an electron (gas phase, ground state configuration). Higher ionization energy = more difficult to remove the electron. Successive ionization energies are greater than the previous. Always remove the outermost electron

Increases going left to right on a period. Decreases going down a group

He has the highest. Fr has the smallest

Exceptions: Orbitals are most stable when half full or completely full.

Reasoning: As you go from left to right across the periodic table, the number of electrons and protons increases. As the valence shell continues to fill, the electrons become harder to remove (require more energy) due to an increase in effective nuclear charge. Fluorine, for example, has a high ionization energy because its electrons are strongly attracted to the more positively charged atomic nucleus.

Note: Going down a group, the shielding effect is increasing. Outer electrons are not experiencing the full pull of the positive charge of protons due to increasing inner electrons shielding them. This results in lower ionization energy as you go down a group because electrons become easier to expel with increasing shells. Iodine, for example, has a lower ionization energy than fluorine because it contains more electron shells, which makes it easier to extract an electron from the atom’s valence shell.

Question 18

Electron affinity

Answer 18

Energy associated with adding an electron (gas phase, ground state configuration); The amount of energy released when an electron is added to
an atom. Can be favourable (negative) or unfavourable (positive).

The harder it is to remove an electron or the higher the Zeff, the easier it is to add an electron, and thus a greater electron affinity. Going down a group, the attraction of an electron to the nucleus decreases due to shielding; Electron affinity, therefore, decreases.

Trend: Electron affinity increases going from left to right across a period and
decreases down a group.

F and Cl have large electron affinity.

Exceptions: noble gases have full shell. Half-stable and full-stable orbitals.

i.e. Silicon has a greater affinity for an electron than phosphorus because accepting an electron will give it half-filled subshell configuration.

There are successive electron affinities.

Question 19

Electronegativity

Answer 19

How well an atom can attract electron density towards itself. The more strongly it can attract electrons, the greater its electronegativity.

Increases going left to right on a period. Decreases down a group. F has greatest and Fr has the lowest. Exclude noble gases from trend.

Question 20

Inner transition metals

Answer 20

The two rows (which belong in periods 6 and 7) beneath the periodic table are known as the inner transition metals. Period 6 elements are the lanthanides and Period 7 elements are the actinides.

Question 21

Diatomic atoms

Answer 21

hydrogen, nitrogen, fluorine, oxygen, iodine, chlorine, and bromine.

Have No Fear Of Ice Cold Beer

Question 22

Metallic character

Answer 22

increases going from right to left across a period and increases going down a group.

Question 23

Metals vs non metals

Answer 23

Metals:

• Malleable, lustrous
• Good conductors of electricity/heat
• Form basic oxides
• Lose electrons to form cations
• Usually solid at room temperature, with the exception of Hg (liquid)
• Generally, high melting and boiling points

Non-metals:

• Brittle, dull
• Poor conductors of electricity/heat
• Form acidic oxides
• Gain electrons to form anions
• Gas or solid at room temperature, with the exception of Br (liquid)
• Generally, low melting and boiling points

Question 24

Atomic radius trend

Answer 24

Trend: The atomic radius decreases
from left to right across a period and
increases going down a group.

Reasoning: Across a period, the number
of protons in an atom increases. Increasing protons results in greater nuclear attraction between the protons and electrons, which results in shells being pulled closer to the nucleus (which equals a smaller radius). Going down a group, the number of electrons increases. Each additional energy level gets farther and farther away from the nucleus, which causes the atomic radius to increase.

Question 25

Multiple ionization energies

Answer 25

The first ionization energy is the energy required to remove the outermost electron. Following removal of the first electron, elements can have second, third, fourth, etc. ionization energies. These values are always larger than the first ionization energy because subsequent electrons are more difficult to remove

Question 26

Exceptions to multiple ionization energies concept

Answer 26

1. The first is that alkaline earth metals have greater ionization energy than group 13 elements. This is because alkaline earth metals have completely filled orbitals and require more energy to remove an electron than group 13 elements
2. The second exception is that group 15 elements have greater ionization energy than group 16 elements. Group 15 elements have half-filled orbitals, which is a more stable configuration than that of group 16 elements and thus require more energy to remove an electron