Periodic Trends

Question 1

which elements are considered metalloids?

Answer 1

B, Si, Ge, As, Sb, Te, Po

follows a stair pattern from B to Po, excluding Al

Question 2

define oxidation state

Answer 2

related to the number of electrons that an atom loses, gains, or appears to use when bonding with another atom

almost all transition metals have
multiple oxidation states because they have several electrons with similar energies, meaning that one or all of
them can be removed, depending on the
circumstances. manganese, for example, shows oxidation states from +2 to +7.

Question 3

what are the groups at the bottom of a periodic table called?

Answer 3

inner transition metals

top group: lanthanides
bottom group: actinides

Question 4

which transition metals are colourless?

Answer 4

although most transition metals have color, row 4 transition metals are an exception as they are colorless

row 4 includes zinc, copper, cobalt, and so on!

Question 5

how do transition metals and inner transition metals differ?

Answer 5

inner transition metals are far less abundant on earth compared to transition metals

transition metal atoms have their valence electrons in the outermost d-orbital, whereas inner transition metal atoms have their valence electrons in the f-orbital. hence, d-block and f-block

Question 6

what’s the most prominent oxidation state for transition metals and inner transition metals?

Answer 6

transition metals: +2

inner transition metals: +3

Question 7

what are diatomic atoms?

Answer 7

atoms that are usually found paired due to their unstable nature

H, N, F, O, I, Cl, and Br

Have No Fear Of Ice Cold Beer ??

Question 8

describe the trend of metallic character on the periodic table

Answer 8

increases going from right to left across a period

increases going down a group

so, the most metallic is bottom left

Question 9

differentiate between the characteristics/properties of metals and non-metals

Answer 9

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 mercury (Hg), which is liquid
- generally high melting/boiling points

NON-METALS
- brittle & dull
- poor conductors of electricity/heat
- form acid oxides
- gains electrons to form anions
- gas or solid at room temperature, with the exception of bromine (Br), which is liquid
- generally low melting/boiling points

Question 10

describe the trend of atomic radius on the periodic table

Answer 10

increases from right to left across a period

increases going down a group

so, the biggest element is at the bottom left

Question 11

describe the reasoning behind the atomic radius trend of the periodic table

Answer 11

increases right to left across a period: the number of protons in an atom decreases moving from right to left. decreasing protons results in a weaker nuclear attraction between the protons and electrons, which results in electron shells
being further apart from the nucleus, therefore increasing the radius.

increases going down a group: the number of electron shells increases moving down the group. each additional electron level gets further and further away from the nucleus, which causes the atomic radius to increase.

Question 12

what is effective nuclear charge (Zeff)?

Answer 12

effective nuclear charge (Zeff) is the amount of positive charge experienced by an electron. the shielding effect of lower orbital electrons prevents higher orbital electrons from experiencing a strong attraction to the nucleus. this effect explains why valence electrons are more easily removed.

the effective nuclear charge is calculated given the following equation:

Zeff = Z - S
where Z = number of protons
S = number of shielding
(non-valence) electrons

so basically, Zeff = valence electrons

Question 13

describe the trend of effective nuclear charge on the periodic table

Answer 13

increases left to right across a period

increases going up a group

so basically, top right experiences the most effective nuclear charge

Question 14

describe the reasoning behind the effective nuclear charge trend of the periodic table

Answer 14

increases left to right across a period: the numbers of protons increase (with no increase in electron shells) from left to right, and thus no increase in shielding effect. results in electrons being pulled closer to the nucleus due to a stronger attraction

increases going up a group: the number of electron shells decreases, moving up the table, which brings outer shell electrons closer to the positively charged nucleus

Question 15

what are isoelectronic series?

Answer 15

atoms that have an identical number of electrons, but different numbers of protons. anions and cations also

Question 16

differentiate between anions and cations in regards to radius

Answer 16

anions are ions that have
gained electrons and have more electrons than protons, making them negatively charged

when a neutral atom gains electrons and becomes an anion, the increased electron number results in increased electron-electron repulsions. this
expands the size of the electron cloud,
resulting in a LARGER RADIUS (for anions)

cations are ions that have lost electrons and have more protons than electrons, making them positively charged

when a neutral atom loses electrons and becomes a cation, the decreased electron number results in decreased electron-electron repulsions. this
reduces the size of the electron cloud, resulting in a SMALLER RADIUS (for cations)

Question 17

what’s ionization energy?

Answer 17

energy needed to remove an electron from an atom

Question 18

describe the trend of ionization energy on the periodic table

Answer 18

increases going from left to right across a period

increases up a group

the element that requires the most ionization energy is top right

Question 19

describe the reasoning behind the ionization energy trend of the periodic table

Answer 19

increases left to right across a period: the number of protons increases from left to right. as the valence shell continues to fill, the electrons become harder to remove (require more energy) due to an increase in effective nuclear charge. note that noble gases will require lots of energy due to their filled octet

increases going up a group: moving up a
group, there are fewer electron shells and subsequently less of a shielding effect from the inner electrons. this creates difficulty in removing the electrons from valence shells. the distance decreases between the nucleus and the highest-energy electron, strengthening the nuclear attraction to that electron, and therefore requiring more energy.

Question 20

describe the concept of multiple ionization energies

Answer 20

the first ionization energy is the energy required to remove the outermost electron. following the removal of the first electron, elements can have second, third, fourth, etc. ionization energies. the energy associated with removing each successive electron from an atom or ion
increases.

e.g. the first ionization energy of sodium is 496 kJmol-1. following that, there is a huge jump to 4562 kJmol-1 for the second ionization energy because we would be removing an electron from a stable configuration (a full outer shell).

Question 21

what are the (2) notable exceptions to the ionization trend on the periodic table?

Answer 21

alkaline earth metals have filled orbitals, which gives them greater stability, leading to their higher ionization energy compared to Group 13 elements in the same period. this is why Be has higher ionization energy compared to B

group 15 elements have half-filled orbitals, which gives it greater stability, leading to its higher ionization energy compared to Group 16 elements in the same period. this is why N has a higher ionization energy compared to O

Question 22

what is electron affinity?

Answer 22

amount of energy released when an electron is added to an atom

Question 23

describe the trend of electron affinity on the periodic table

Answer 23

increases going from left to right across a period

increases going up a group

the element that is most electron affinitive is the top right, fluorine

Question 24

describe the reasoning behind the electron affinity trend of the periodic table

Answer 24

increases left to right across a period: across a period, as the atom’s valence shell gets filled, there is increased attraction between the nucleus and the electrons of the atom. this creates a stronger affinity for electrons.

increases going up a group: moving up a
group, there are fewer electron shells, leading to decreased electron shielding and an increased proximity between the nucleus and valence electrons, increasing the nuclear attraction and thereby increasing electron affinity.

Question 25

what are the (3) notable exceptions to the electron affinity trend on the periodic table?

Answer 25

group 2 elements have filled s-orbitals, so their electron affinities are very low

group 15 elements have half-filled orbitals p-orbitals, so their electron affinities are lower than group 14 elements of the same period

noble gases have filled electron shells, so their electron affinities are negligible

Question 26

what is electronegativity?

Answer 26

measurement of an atom’s ability to attract electrons in a bond. the higher the electronegativity of an atom, the greater ability to attract an electron pair

Question 27

describe the trend of electronegativity on the periodic table

Answer 27

increases going from left to right across a period

increases moving up a group

the most electronegative element is the top right, fluorine

(noble gases negligible)

Question 28

describe the reasoning behind the electronegativity trend of the periodic table

Answer 28

increases left to right across a period: with increasing protons as you go from left to right across a period, the ability of an atom to attract an electron pair is increased. this is similar to the electron affinity trend; however, this trend is
specific to electron pairs in a bond, not the addition of a single electron.

increases going up a group: moving up a
group, as the atomic radius decreases, and the valence electrons experience less shielding, the ability of an atom to attract an electron pair is increased.