Knowt
Note
Studied by 4 people
Home
CHAPTER 7: PERIODIC PROPERTIES OF THE ELEMENTS
7.1 ∣ Development of the Periodic Table
Russian Dmitri Mendeleev (1834–1907) and German Lothar Meyer (1830–1895) proposed similar classification schemes in 1869. When elements are grouped by atomic weight, identical chemical and physical properties occur periodically. Atomic numbers were unknown to scientists. However, atomic weights rise with atomic number, therefore Mendeleev and Meyer organized the elements roughly correctly.
He boldly predicted their existence and properties, referring to them as eka-aluminum (“under” aluminum) and eka-silicon (“under” silicon), respectively, after the elements under which they appeared in his table.
In 1913, two years after Rutherford proposed the nuclear model of the atom, English physicist Henry Moseley (1887–1915) developed the concept of atomic numbers.
He arranged the X-ray frequencies in order by assigning a unique whole number, called an atomic number, to each element.
7.2 ∣ Effective Nuclear Charge
Coulomb’s law tells us that the strength of the interaction between two electrical charges depends on the magnitudes of the charges and on the distance between them.
Each electron in a many-electron atom is screened from the nucleus by the other electrons.
Effective nuclear charge. The effective nuclear charge experienced by the 3s electron in a sodium atom depends on the 11+ charge of the nucleus and the 10- charge of the core electrons.
The effective nuclear charge increases from left to right across any period of the periodic table.
The valence electrons added to counterbalance the increasing nuclear charge screen one another ineffectively.
Going down a column, the effective nuclear charge experienced by valence electrons changes far less than it does across a period.
Effective nuclear charge increases slightly as we go down a column because the more diffuse core electron cloud is less able to screen the valence electrons from the nuclear charge.
7.3 ∣ Sizes of Atoms and Ions
According to the quantum-mechanical model, however, atoms do not have sharply defined boundaries at which the electron distribution becomes zero.
The bonding atomic radius for any atom in a molecule is equal to half of the bond distance d.
The bonding atomic radius (also known as the covalent radius) is smaller than the nonbonding atomic radius.
Periodic Trends in Atomic Radii
Within each group, bonding atomic radius tends to increase from top to bottom.
Within each period, bonding atomic radius tends to decrease from left to right.
Periodic Trends in Ionic Radii
Ionic compounds can be used to calculate ionic radii from interatomic distances, much like molecules can. Ion size depends on its nuclear charge, amount of electrons, and valence electron orbitals, just as atom size. A neutral atom becomes a cation by removing electrons from the most distant atomic orbitals. Cations diminish electron–electron repulsions.
Cations are smaller than their parent atoms. Anions are larger than their parent atoms.
As the principal quantum number of the outermost occupied orbital of an ion increases, the radius of the ion increases.
An isoelectronic series is a group of ions all containing the same number of electrons.
In any isoelectronic series, we can list the members in order of increasing atomic number; therefore, the nuclear charge increases as we move through the series.
7.4 ∣ Ionization Energy
The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion.
In general, the first ionization energy, is the energy needed to remove the first electron from a neutral atom.
The second ionization energy, is the energy needed to remove the second electron, and so forth, for successive removals of additional electrons.
Variations in Successive Ionization Energies
The magnitude of the ionization energy tells us how much energy is required to remove an electron; the greater the ionization energy, the more difficult it is to remove an electron.
The sharp increase in ionization energy that occurs when an inner-shell electron is removed.
Every element exhibits a large increase in ionization energy when the first of its inner-shell electrons is removed.
Periodic Trends in First Ionization Energies
Smaller atoms ionize more. Atomic size impacts ionization energy. Energy to remove an electron from the outermost inhabited shell depends on the effective nuclear charge and average electron distance from the nucleus. Electron-nucleus attraction increases with effective nuclear charge or distance. This attraction makes electron removal harder, raising ionization energy. Ionization energy rises as effective nuclear charge and atomic radius decrease. As we descend, atomic radius increases, but effective nuclear charge increases slowly. Increased radius decreases nucleus-electron attraction, lowering ionization energy.
Electron Configurations of Ions
When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest principal quantum number, n.
If there is more than one occupied subshell for a given value of n, the electrons are first removed from the orbital with the highest value of l.
Electrons added to an atom to form an anion are added to the empty or partially filled orbital having the lowest value of n.
7.5 ∣ Electron Affinity
The first ionization energy of an atom is a measure of the energy change associated with removing an electron from the atom to form a cation.
The positive value of the ionization energy means that energy must be put into the atom to remove the electron.
The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity.
For most atoms, energy is released when an electron is added.
Ionization energy measures the energy change when an atom loses an electron.
Electron affinity measures the energy change when an atom gains an electron.
The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity.
7.6 ∣ Metals, Nonmetals, and Metalloids
Characteristic Properties of Metals and Nonmetals | |
|---|---|
Metals | Nonmetals |
Have a shiny luster; various colors, although most are silvery | Do not have a luster; various colors |
Solids are malleable and ductile | Solids are usually brittle; some are hard, and some are soft |
Good conductors of heat and electricity | Poor conductors of heat and electricity |
Most metal oxides are ionic solids that are basic | Most nonmetal oxides are molecular substances that form acidic solutions |
Tend to form cations in aqueous solution | Tend to form anions or oxyanions in aqueous solution |
The more an element exhibits the physical and chemical properties of metals, the greater its metallic character.
Metals
Metals are shiny, malleable, and ductile.
Metals conduct heat and electricity.
Metals tend to have low ionization energies and therefore tend to form cations relatively easily.
Compounds made up of a metal and a nonmetal tend to be ionic substances.
Most metal oxides are basic.
Nonmetals
can be solid, liquid, or gas.
They are not lustrous and generally are poor conductors of heat and electricity.
Because of their relatively large, negative electron affinities, nonmetals tend to gain electrons when they react with metals.
Compounds composed entirely of nonmetals are typically molecular substances that tend to be gases, liquids, or low melting point solids at room temperature.
Most nonmetal oxides are acidic, which means that those that dissolve in water form acids.
Metalloids
They may have some characteristic metallic properties but lack others.
One of the reasons metalloids can be used for integrated circuits is that their electrical conductivity is intermediate between that of metals and that of nonmetals.
Very pure silicon is an electrical insulator, but its conductivity can be dramatically increased with the addition of specific impurities called dopants.
7.7 ∣ Trends for Group 1A and Group 2A Metals
Group 1A: The Alkali Metals
The alkali metals are soft metallic solids.
The name alkali comes from an Arabic word meaning “ashes.”
Sodium, like the other alkali metals, is soft enough to be cut with a knife.
The alkali metals exist in nature only as compounds.
All alkali metals combine directly with most non-metals.
Group 2A: The Alkaline Earth Metals
Compared with the alkali metals, the alkaline earth metals are harder and denser, and melt at higher temperatures.
The alkaline earth metals are less reactive than their alkali metal neighbors.
The heavier alkaline earth ions give off characteristic colors when heated in a hot flame.
Like their neighbors sodium and potassium, magnesium and calcium are relatively abundant on Earth and in seawater and are essential for living organisms as cations in ionic compounds.
7.8 ∣ Trends for Selected Nonmetals
Hydrogen
Hydrogen is a nonmetal that occurs as a colorless diatomic gas, H2(g), under most conditions.
Unlike the alkali metals, hydrogen reacts with most nonmetals to form molecular compounds in which its electron is shared with, rather than completely transferred to, the other nonmetal.
The ability of hydrogen to form bonds with carbon is one of the most important aspects of organic chemistry.
The fact that hydrogen can gain an electron further illustrates that it behaves much more like a nonmetal than an alkali metal.
Group 6A: The Oxygen Group
Oxygen, sulfur, and selenium are typical nonmetals.
Tellurium is a metalloid, and polonium, which is radioactive and quite rare, is a metal.
Oxygen is a colorless gas at room temperature; all of the other members of group 6A are solids.
Oxygen has a great tendency to attract electrons from other elements (to oxidize them).
After oxygen, the most important member of group 6A is sulfur. This element exists in several allotropic forms, the most common and stable of which is the yellow solid having the molecular formula S8.
Below sulfur in group 6A is selenium, Se. This relatively rare element is essential for life in trace quantities, although it is toxic at high doses.
The next element in the group is tellurium, Te. Its elemental structure is even more complex than that of Se, consisting of long, twisted chains of Te—Te bonds.
From O to S to Se to Te, the elements form larger and larger molecules and become increasingly metallic.
Group 7A: The Halogens
Astatine, which is both extremely rare and radioactive, is omitted because many of its properties are not yet known. Even less is known about the recently discovered Tennessine.
Fluorine and chlorine are more reactive than bromine and iodine.
Chlorine is the most industrially useful of the halogens.
The halogens react directly with most metals to form ionic halides.
Group 8A: The Noble Gases
Gases, are all nonmetals that are gases at room temperature.
They are all monatomic (that is, they consist of single atoms rather than molecules).
In fact, until the early 1960s the elements were called the inert gases because they were thought to be incapable of forming chemical compounds.
In 1962, Neil Bartlett (1932–2008) at the University of British Columbia reasoned that the ionization energy of Xe might be low enough to allow it to form compounds.
In 2000, Finnish scientists reported the first neutral molecule that contains argon, the HArF molecule, which is stable only at low temperatures.