Periodic Trends Flashcards
which elements are considered metalloids?
B, Si, Ge, As, Sb, Te, Po
follows a stair pattern from B to Po, excluding Al
define oxidation state
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.
what are the groups at the bottom of a periodic table called?
inner transition metals
top group: lanthanides
bottom group: actinides
which transition metals are colourless?
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!
how do transition metals and inner transition metals differ?
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
what’s the most prominent oxidation state for transition metals and inner transition metals?
transition metals: +2
inner transition metals: +3
what are diatomic atoms?
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 ??
describe the trend of metallic character on the periodic table
increases going from right to left across a period
increases going down a group
so, the most metallic is bottom left
differentiate between the characteristics/properties of metals and non-metals
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
describe the trend of atomic radius on the periodic table
increases from right to left across a period
increases going down a group
so, the biggest element is at the bottom left
describe the reasoning behind the atomic radius trend of the periodic table
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.
what is effective nuclear charge (Zeff)?
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
describe the trend of effective nuclear charge on the periodic table
increases left to right across a period
increases going up a group
so basically, top right experiences the most effective nuclear charge
describe the reasoning behind the effective nuclear charge trend of the periodic table
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
what are isoelectronic series?
atoms that have an identical number of electrons, but different numbers of protons. anions and cations also
differentiate between anions and cations in regards to radius
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)
what’s ionization energy?
energy needed to remove an electron from an atom
describe the trend of ionization energy on the periodic table
increases going from left to right across a period
increases up a group
the element that requires the most ionization energy is top right
describe the reasoning behind the ionization energy trend of the periodic table
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.
describe the concept of multiple ionization energies
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).
what are the (2) notable exceptions to the ionization trend on the periodic table?
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
what is electron affinity?
amount of energy released when an electron is added to an atom
describe the trend of electron affinity on the periodic table
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
describe the reasoning behind the electron affinity trend of the periodic table
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.
what are the (3) notable exceptions to the electron affinity trend on the periodic table?
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
what is electronegativity?
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
describe the trend of electronegativity on the periodic table
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)
describe the reasoning behind the electronegativity trend of the periodic table
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.
