Atomic Properties and the Periodic Table Flashcards
Topic 2, Lectures 5-8 - Jon Rourke
Mendeleev (1869)
Endeavoured to order elements in a logical sequence using chemical properties. ordered the elements using increasing relative atomic masses.
Benefits of Mendeleev’s periodic table
Allowed predictions for elements such as gallium before they were even invented.
Moseley (1913)
Found that the proper sequence criterion was not determined by relative atomic mass but by atomic number (no of electrons). Also found that hroups in the periodic table were not only chemically similar but also electronically. Ordering the elements based on electronic configuration creates Mendeleev’s arrangement.
What can be found from electronic configuration?
The element’s chemistry can be predicted.
How were the groups in the periodic table arranged?
The noble gases were placed at the far right of the table, as they have full ns2np6 configurations. Alkaline metals (ns1) and alkaline earth metals (ns2) became groups 1 and 2 on the far left. Halogens (ns2np5), calcogens (ns2np4), and pnictogens (ns2np3) became groups 17, 16, and 15, respectively.
Main group elements
Outer shells consisting of only s and p electrons.
Transition elements
An element with an incomplete d orbital in either a metal compound or complex—zinc is an exception as it has a complete d orbital
Lanthanides
Strip at the bottom of the periodic table contain an incomplete f orbital
Multi-electron system assumptions
Assume the orbitals in ME systems look exactly as they do in hydrogen. Quantum numbers and angular functions (shape) are the same. Radial functions are similar but are contracted due to a higher nuclear charge in ME systems.
Valence electrons
Outermost electrons, highest energy, partake in chemical reactions.
Core electrons
Low-energy electrons in filled shells. Shield the nucleus
Shielding
In ME systems, electrons are attracted to the positive nucleus and repelled by negative electrons; repulsion between electrons causes shielding from nuclear charge. The lessened nuclear charge is known as effective nuclear charge. Extent of shielding is dependent on n
Penetration
The relative electron density of an electron to the nucleus. Electrons in different orbitals have different wavefunctions and therefore different RDFs. E.g., a 2s orbital has more electron density near the nucleus in comparison to a 2p orbital, therefore is said to be more penetrating. The extent of penetration is dependent on n and l.
Trend of electron penetration in an orbital
Provided the n value is the same: s > p > d > f. S orbitals are spherically symmetrical, therefore protect the nucleus in all directions, and have high electron density near the nucleus. P, d and f orbitals, however, have more complex shapes and therefore do not protect in all directions. They also have more nodes, resulting in lower electron density near the nuclear.
Why is the 2s orbital filled before the 2p?
Through observing the RDF graphs, the 2s orbital has a maximum peak closer to the nucleus than the 2p, therefore has lower energy/more stability and is filled first.
Why is the 4s orbital filled before the 3d?
Although on the RDF graph the 3d orbital has a maximum closer to the nucleus, the 4s orbital penetrates the core electrons. The 4s orbitals experience lower shielding effects than 3d orbitals, making them more stable when filled first.
Effective nuclear charge equation
Zeff = Z - S (shielding effects)
“Perfect shield”
If an electron were to screen the effects of one positive charge in the nucleus. Would contribute a value of 1 to the shielding constant. However, shielding is not always perfect.
Criteria for calculating shielding constant
Electrons with higher values of n contribute 0 to the constant. Those with the same principal quantum number as the observed electron contribute 0.35. Electrons with a n value of one below the observed contribute 0.85. Electrons with low n values contribute 1.
Criteria for calculating shielding constant when the observed electron is nf or nd
All electrons with lower n values contribute a value of 1 because d and f orbitals are completely shielded by lower shells.
General trend of Zeff
Increases down a group and across a row.
Ionisation enthalpy: noble gases
High ionisation energy due to complete shells, therefore very stable. Full outer shells also have high attraction to the nucleus due to opposite charges.
Ionisation enthalpy: alkaline metals
Low ionisation energy due to the singular valence electron, which experiences high levels of shielding from core electrons.
Ionisation enthalpy: elements in the same group
Analogous elements tend to have similar IEs. With alkaline metals, elements all with 1 s electron outside of a noble gas configuration, it would be expected that the IE would decrease at a large rate with increasing n value, however, extra shielding does not cause a drastic effect to the Zeff, and has a lessened impact with each added shell.
Ionisation enthalpy: elements in the same period
General increase in IE across a period. Even though the number of electrons increases, as does the number of protons in the nucleus - therefore Zeff does not increase by a huge amount.
Exceptions: B and Be
Boron has a lower IE than beryllium despite being further along across the period. This occurs because the first ionisation energy of boron removes a singular electron from the 2p orbital forming a complete 2s orbital, which requires much less energy than removing a paired electron from a complete 2s orbital
Exceptions: O and N
The IE of oxygen is lower than that of nitrogen because N’s 2p orbital is made from 3 unpaired electrons, whereas O’s is made from that plus one paired electron orbital. It requires less energy to remove electrons from an orbital with a repelling pair of electrons.
Electron affinity values
Can be positive or negative. Negative is more common, as when electrons are added, they become attracted to the nucleus and make the element more stable - releasing energy. In the rare event that adding an electron makes an element less stable e.g. noble gases, it will require energy to occur, thus causing a positive value. Same for group 2 elements, which have full 2s orbitals.
Electron affinity: G14 v G15
Group 15 has a lower negative EA because each element has 3 half-occupied orbitals, which would require energy to add to due to electron repulsion; therefore, less energy is released. Group 14, however, has only 2 half-occupied orbitals; therefore, less energy is required to add an electron and more energy is released.
Electron affinity: first row elements
Tends to be lower than expected due to their small size, leading to stronger electron repulsion.
Pauling electronegativity scale
Used bond dissociation energies (energy required to break a bond) to calculate electronegativities. Mathematical formula: electronegativity difference = (bond energy of compound - sum of individual bond energies)^1/2. More difference = more electronegativity. Fluorine was given the highest electronegativity of 4 and EN tends to increase diagonally to the right
Mulliken electronegativity
Argued that the arithmetic mean of the IE and EA should be a good measure of the electronegativity.
Alfred-Rochow electronegativity
Calculated based off Zeff and covalen radius.
What can electronegativity show
Can be used to explain electron distribution in a bond. If two atoms have the same electronegativity, electrons are shared equally, and the bond is non-polar. If electronegativity differs, the polarity is proportional to the difference in electronegativity. Thus, atoms with extreme electronegativities are more likely to form ionic compounds.
Differing sizes of atoms
For the same atoms, atomic radii can differdepending upon the nature of the bond e.g. metallic, covalent etc.
Why is it impossible to measure the size of an atom?
The probability of finding an electron at any given distance from the nucleus is never zero. Even if you measured the distance at which the probability was 95%, it would still differ between compounds. Only distances between nuclei can be measured.
Measuring covalent radii
Experimentally measured bond lengths can be used to find atomic radii if divided by 2. If elements do not form single bonds in the elemental state, the radii can be found from bond lengths in compounds with known bond lengths.
Unsaturated compounds covalent atomic radii
Bond lengths in double or triple bonds tend to be shortened, so calculated radii must be adjusted accordingly. e.g., C-C = 154pm C=C = 135pm.
Trends in atomic radius
Decreases across a period due to a higher number of electrons, therefore increased Zeff. Increases down a gorup due to electrons being placed into orbitals of a higher n value.
Covalent: Group 13 atomic radii
The atomic radii for Al and Ga are the same despite the general increasing trend. Al has valence electrons in the 3p orbitals and Ga has valence electrons in the 4s orbitals; however, because the 3d orbital has quite poor shielding abilities overall, the Zeff of Ga remains quite similar to Al regardless of the extra shell.
Metallic atomic radii
Half of the internuclear distance in a metal. Dependent on how closely packed the atoms are and their coordination number
Van der Waal’s atomic radii
One half of the internuclear distance between non-bonded atoms in a solid or liquid state. This is only known for some elements and can vary considerably. Tend to be larger on average because atoms aren’t bonded.
Ionic atomic radii
Determined from crystal structures where the sum of the cation and anion radii is equal to the internuclear distance. Half are the atomic radii. The only difficulty comes from knowing where the radius of the cation ends and the anion begins.
Trends in ionic radii
The radius decreases with increasing positive charge and with increasing oxidation states.
