periodic system

In the SparkNote on the Periodic table we discussed a number of simple periodic trends. In this section we will discuss a number of more complex trends, the understanding of which relies on knowledge of atomic structure.

Before getting into these trends, we should engage a quick review and establish some terminology. As seen in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group of atoms with a net charge (whether positive or negative) is called an ion. A positively charged ion is a cation while a negatively charged ion is an anion.

Now we are ready to discuss the periodic trends of atomic size, ionization energy, electron affinity, and electronnegativity.

Atomic Size (Atomic Radius)

The atomic size of an atom, also called the atomic radius, refers to the distance between an atom's nucleus and its valence electrons. Remember, the closer an electron is to the nucleus, the lower its energy and the more tightly it is held.

Moving Across a Period

Moving from left to right across a period, the atomic radius decreases. The nucleus of the atom gains protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the nucleus upon the electrons. True, electrons are also added as the elements move from left to right across a period, but these electrons reside in the same energy shell and do not offer increased shielding.

Moving Down a Group

The atomic radius increases moving down a group. Once again protons are added moving down a group, but so are new energy shells of electrons. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount of the protons' positive charge.

Cations and Anions

Cations and anions do not actually represent a periodic trend in terms of atomic radius, but they do affect atomic radius, and so we will discuss them here.

A cation is positively charged, meaning that it is an atom that has lost an electron or electrons. The positive charge of the nucleus is thus distributed over a smaller number of electrons and electron-electron repulsion is decreased, meaning that the electrons are held more tightly and the atomic radius is smaller than in the normal neutral atom. Anions, conversely, are negatively charged ions: atoms that have gained electrons. In anions, electron-electron repulsion increases and the positive charge of the nucleus is distributed over a large number of electrons. Anions have a greater atomic radius than the neutral atom from which they derive.

Ionization Energy and Electron Affinity

The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity.

Ionization Energy

The ionization energy is the energy it takes to fully remove an electron from the atom. When several electrons are removed from an atom, the energy that it takes to remove the first electron is called the first ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so on. In general, the second ionization energy is greater than first ionization energy. This is because the first electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted to the nucleus. If a particular ionization energy follows a previous electron loss that emptied a subshell, the next ionization energy will take a rather large leap, rather than follow its normal gently increasing trend. This fact helps to show that just as electrons are more stable when they have a full valence shell, they are also relatively more stable when they at least have a full subshell.

Ionization Energy Across a Period

Ionization energy predictably increases moving across the periodic table from left to right. Just as we described in the case of atomic size, moving from left to right, the number of protons increases. The electrons also increase in number, but without adding new shells or shielding. From left to right, the electrons therefore become more tightly held meaning it takes more energy to pry them loose. This fact gives a physical basis to the octet rule, which states that elements with few valence electrons (those on the left of the periodic table) readily give those electrons up in order to attain a full octet within their inner shells, while those with many valence electrons tend to gain electrons. The electrons on the left tend to lose electrons since their ionization energy is so low (it takes such little energy to remove an electron) while those on the right tend to gain electrons since their nucleus has a powerful positive force and their ionization energy is high. Note that ionization energy does show a sensitivity to the filling of subshells; in moving from group 12 to group 13 for example, after the d shell has been filled, ionization energy actually drops. In general, though, the trend is of increasing ionziation energy from left to right.

Ionization Energy Down a Group

Ionization energy decreases moving down a group for the same reason atomic size increases: electrons add new shells creating extra shielding that supersedes the addition of protons. The atomic radius increases, as does the energy of the valence electrons. This means it takes less energy to remove an electron, which is what ionization energy measures.

Electron Affinity

An atom's electron affinity is the energy change in an atom when that atom gains an electron. The sign of the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its potential energy decreases: upon gaining an electron the atom gives off energy and the electron affinity is negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is positive. An atom with a negative electron affinity is far more likely to gain electrons.

Electron Affinities Across a Period

Electron affinities becoming increasingly negative from left to right. Just as in ionization energy, this trend conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence shells will tend to gain electrons. Such atoms are located on the right of the periodic table and have very negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and become more stable. Be careful, though: the nobel gases, located in the extreme right hand column of the periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not want to add more electrons: noble gas electron affinities are positive. Similarly, atoms with full subshells also have more positive electron affinities (are less attractive of electrons) than the elements around them.

Electron Affinities Down a Group

Electron affinities change little moving down a group, though they do generally become slightly more positive (less attractive toward electrons). The biggest exception to this rule are the third period elements, which often have more negative electron affinities than the corresponding elements in the second period. For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity.

Electronegativity

Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. Electronegativity is based on an atom's ionization energy and electron affinity. For that reason, electronegativity follows similar trends as its two constituent measures.

Electronegativity generally increases moving across a period and decreases moving down a group. Flourine (F), in group VIIa and period 2, is the most powerfully electronegative of the elements. Electronegativity plays a very large role in the processes of Chemical Bonding.

 

Atomic Structure :  Periodic Trends

Atomic Radii

1) As you move down a group, atomic radius increases.

WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.  Each subsequent energy level is further from the nucleus than the last.  Therefore, the atomic radius increases as the group and energy levels increase.

2) As you move across a period, atomic radius decreases.

WHY? - As you go across a period, electrons are added to the same energy level.  At the same time, protons are being added to the nucleus.  The concentration of more protons in the nucleus creates a "higher effective nuclear charge."  In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.

View a periodic table with atomic radii.
 

 

Ionic Radii

1) Anions (negative ions) are larger than their respective atoms.

WHY?

Electron-electron repulsion forces them to spread further apart.
Electrons outnumber protons; the protons cannot pull the extra electrons as tightly toward the nucleus.

2) Cations (positive ions) are smaller than their respective atoms.

WHY?

There is less electron-electron repulsion, so they can come closer together.
Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly.
If the electron that is lost is the only valence electron so that the electron configuration of the cation is like that of a noble gas, then an entire energy level is lost.  In this case, the radius of the cation is much smaller than its respective atom.

First Ionization Energy

Definition:  The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.

1) As you move down a group, first ionization energy decreases.

WHY?

Electrons are further from the nucleus and thus easier to remove the outermost one.
"SHIELDING" - Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus.  It therefore becomes easier to remove the outer electron

2) As you move across a period, first ionization energy increases.

WHY? - As you move across a period, the atomic radius decreases, that is, the atom is smaller.  The outer electrons are closer to the nucleus and more strongly attracted to the center.  Therefore, it becomes more difficult to remove the outermost electron.

Exceptions to First Ionization Energy Trends

1) Xs2 > Xp1  e.g. 4Be > 5B WHY? - The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital.  Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital. 2) Xp3 > Xp4  e.g.  7N > 8O WHY? - After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired.  The electron-electron repulsion makes it easier to remove the outermost, paired electron. (See Hund's Rule)

Veiw a periodic table with first ionization energies.

Second and Higher Ionization Energies

Definition:  Second Ionization Energy is the energy required to remove a second outermost electron from a ground state atom.
Subsequent ionization energies increase greatly once an ion has reached the state like that of a noble gas.  In other words, it becomes extremely difficult to remove an electron from an atom once it loses enough electrons to lose an entire energy level so that its valence shell is filled.

Ionization Energies (kJ/mol)
Element Na Mg Al		1st IE 495.8 737.7 577.6		2nd IE 4562.4 1450.6 1816.6		3rd IE 6912 7732.6 2744.7		4th IE 9543 10,540 11,577

Electron Affinity

Definition:  The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.

1) As you move down a group, electron affinity decreases.

2) As you move across a period, electron affinity increases.

	Exceptions Among nonmetals, however, the elements in the first period have lower electron affinities than the elements below them in their respective groups.  Elements with electron configurations of Xs2, Xp3, and Xp6 have electron affinities less than zero because they are unusually stable.  In other words instead of energy being given off, these elements actually require an input of energy in order to gain electrons.  e.g.  Be, N, Ne  WHY? - Electron affinities are all much smaller than ionization energies.  Xs2 < 0:  Stable, diamagnetic atom with no unpaired electrons.  Xp3 < 0: Stable atom with 3 unpaired p-orbital electrons each occupying its own subshell.  Xp6 < 0: Stable atom with filled valence (outermost) shell.

Lattice Energy

Definition:  The energy given off when oppositely charged ions in the gas phase come together to form a solid.

The strength of a bond between ions of opposite charge can be calculated using Coulomb's Law.

Coulomb's Law - The force of attraction between oppositely charged particles is directly proportional to the product of the charges of the particles (q₁ and q₂) and inversely proportional to the square of the distance between the particles.

1) As you move down a group, lattice energy decreases.

WHY? - The atomic radius increases as you move down a group.  Since the square of the distance is inversely proportional to the force of attraction, lattice energy decreases as the atomic radius increases.

2) As you increase the magnitude of the charge (becomes more positive or more negative), lattice energy increases.

WHY? - The product of the charges of the particles is directly proportional to the force of attraction. Therefore, lattice energy increases as the charges increase.

Lattice Energies of Alkali Metals with Halides (kJ/mol) Li+ Na+ K+ Rb+ Cs+ F- 1036 923 821 785 740 Cl- 853 787 715 689 659 Br- 807 747 682 660 631 I- 757 704 649 630 604

Lattice Energies of Salts of OH- and O2- with Cations of varying charge (kJ/mol) Na+ Mg2+ Al3+ OH- 900 3006 5627 O2- 2481 3791 15916

    Metals

    Common characteristics:

        Metallic luster (shine)

        Generally solids at room temperature

        Malleable

        Ductile

        Conduct heat and electricity

        Exist as extended planes of atoms

        Combine with other metals to form alloys which have metallic characteristics

        Form positive ions, e.g.  Na+, Mg2+, and Al3+

    Nonmetals

    Common characteristics:

        Rarely have metallic luster (shine)

        Generally gases at room temperature

        Neither malleable nor ductile

        Poor conductors of heat and electricity

        Usually exist as molecules in thier elemental form

        Combine with other nonmetals to form covalent

        Generally form negative ions, e.g.  Cl-, SO42-, and N3-

    The differences in the characteristics of metals and nonmetals can be explained by the following:

        Metals have relatively few electrons in their valence shells.

        Metals have lower ionization energies than nonmetals.

        Metals have smaller electron affinities than nonmetals.

        Metals have larger atoms than nonmetals.

    1) As you move across a period, metallic character decreases and nonmetallic character increases.

    2) As you move down a group, metallic character increases and nonmetallic character decreases.

    Semimetals (Metalloids)

        A class of 8 elements that have properties of both metals and nonmetals.

B             Si             Ge          As           Sb           Te           Po           At

    Common characteristics:

        Generally look metallic but are brittle (not malleable or ductile)

        Neither good conductors or insulators; instead they are semiconductors.

Bonding :  IntrBonding :  Intramolecular Forces of AttractionamolBonding :  Intraaecular Forces of Attraction

1902-1916 - G.N. Lewis

Described atoms with a cube model, explaining that the electron in an atom were arranged in positions at the corners of a cube.  This model was based upon the following assumptions:  The number of electrons in an atom's outermost cube is equal to the number of electrons lost when an atom forms a positive ion.  Each successive neutral atom on the periodic table (as the atomic number increases) has one more electron it its outermost cube.  To fill the corners of the cube, it takes eight electrons, or an octet.  When a cube has filled its octet, it becomes the center about which another cube can be built.

Lewis explained simple ionic compounds by assuming that atoms with less than four electrons in their outermost cube transferred those electrons to another atom which needed to gain that many electrons to fill its octet.  Both atoms, then, would end up with an outer cube that was either completely filled or completely empty.

e.g.  Na loses one electron to Cl to form Na⁺ and Cl^-.

It was later understood why Lewis's "octet" theory was valid.  It requires eight electrons to fill the s (1) and p (3) orbitals in one shell which are the outermost orbitals in a shell.  The outermost shell of an atom came to be known as the valence shell, and the electrons which reside in the valence shell are now known as valence electrons.

• Valence electrons - The electrons on an atom that are not present in the previous atom with a filled-shell electron configuration, ignoring filled d and f orbitals.

e.g.  Bromine has the electron configuration:  [Ar] 4s²3d¹⁰4p⁵.  If we add up the electrons in the s and p orbitals only, we find that bromine has 7 valence electrons.

Lewis later found that atoms could fill their outermost cubes by sharing one or more pairs of electrons.

e.g.  One Cl atom can share one of its outermost electrons with another Cl atom to form the diatomic molecule Cl₂.

e.g.  One O atom can share two of its outermost electrons with another O atom to form the diatomic molecule O₂.

Since atoms always seemed to share pairs of electrons when they formed bonds, Lewis changed his cube model with the eight electrons in the corners to a model with pairs of electrons.  These models are known as Lewis Structures (or Lewis Dot Structures).

e.g.  The bonding of two Cl atoms for form Cl₂ would now be written as:

These Lewis Structures models are still in use today.  The only major difference between Lewis's system, and the system in use today is that the covalent bonds between the atoms today are represented by a line instead of two dots.  Double bonds are represented by a double line and triple bonds are represented by triple lines.

Intramolecular Forces of Attraction  - The forces of attraction that exist between bonds within a molecule.

• Ionic bond - A bond between two ions which are formed from the TRANSFER or electrons from one (which forms a positive ion) to another (which forms a negative ion).  The bond is held together by the force of attraction from the opposite charges of the ions formed.
• Ionic compounds dissociate into their ions when they dissolve in water.  This makes them GOOD CONDUCTORS of electricity in aqueous solution.

NaCl_(s) Na⁺_(aq) + Cl^-_(aq)

• Covalent bond - A bond between two atoms formed by the SHARING of electrons.
• Covalent compounds cannot dissociate when they dissolve in water.  This makes them POOR CONDUCTORS of electricity in aqueous solution.

C₆H₁₂O_6(s) C₆H₁₂O_6(aq)

Electronegativity and Polarity

• Electronegativity - The tendency of an atom to draw electrons in a bond toward itself.
• There are two periodic trends concerning electronegativity.
• As you move down a group, electronegativity decreases.
• As you move across a period, electronegativity increases.

H 2.20

He  *

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne *

Na 0.93

Mg 1.31

Al 1.61

Si 1.90

P 2.19

S 2.58

Cl 3.16

Ar *

K 0.82

Ca 1.00

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.90

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr *

Rb 0.82

Sr 1.95

Y 1.22

Zr 1.33

Nb 1.6

Mo 2.16

Tc 1.9

Ru 2.2

Rh 2.28

Pd 2.20

Ag 1.93

Cd 1.69

In 1.78

Sn 1.96

Sb 2.05

Te 2.1

I 2.66

Xe *

Cs 0.79

Ba 0.89

La 1.1

Hf 1.3

Ta 1.3

W 2.36

Re 1.9

Os 2.2

Ir 2.20

Pt 2.28

Au 2.54

Hg 2.00

Tl 2.04

Pb 2.33

Bi 2.02

Po 2.0

At 2.2

Rn *

Fr 0.7

Ra 0.9

Ac 1.1

Unq

Unp

Unh

Uns

Uno

Une

Bonds can be classified according to the difference in electronegativities of the atoms (EN).  Bonds are:

• ionic ifEN > 1.8
• polar covalent if 1.8 EN  0.4
• nonpolar covalent if EN < 0.4

The larger the difference in the electronegativities of the atoms in a bond, the stronger the strength of the bond.  As the bond becomes stronger, melting and boiling points generally increase as well.

When comparing bonds with the same atoms but with different oxidation states, we have to consider another factor besides the difference in the electronegativities of the atoms since this will be the same if the atoms are the same.  As the oxidation number on an atom increase, its ability to draw electrons in a bond toward itself increases as well.  For example, the electronegativity of an Mn atom in Mn₂O₇ (oxidation number of +7) is much greater than the electronegativity of an Mn atom in MnO (oxidation number of +2).

Polar Versus Nonpolar Bonds

Nonpolar covalent - Covalent bond in which the electrons are shared equally.

e.g.  The diatomic elements (Br, I, N, Cl, H, O, and F) are all examples of nonpolar covalent bonds.

Two general relationships between single, double and triple bonds

1.      BOND STRENGTH increases from single to double to triple
 

SINGLE < DOUBLE < TRIPLE

2.      BOND LENGTH decreases from single to double to triple

SINGLE > DOUBLE > TRIPLE

Polar covalent - Covalent bond in which the electrons are NOT shared equally.  The charged ends are called dipoles and are represented by the symbol.

·  The end of the bond with the larger electronegativity is the slightly charged negative end.

·  The end of the bond with the smaller electronegativity is the slightly charged positive end.

In this H-Cl bond, for example, the bond is covalent since the EN < 1.8, but the electrons are shared unequally, so the bond is polar.  The Cl atom has the higher EN so the electrons are drawn more toward the Cl atom and away from the H atom.  Since electrons carry a negative charge, this gives the Cl atom a slightly negative charge.  The H atom, consequently, has a slightly positive charge.

Polar Versus Nonpolar Molecules

Molecules that contain polar covalent bonds may or may not be polar molecules.  The polarity of a molecule is determined by measuring the dipole moment which is represented by the symbol .  This depends on two factors:

·  The degree of the overall separation of charge between the atoms in the bond

·  The distance between the positive and negative poles

If there are equal polar bonds that balance each other around the central atom, then the overall molecule will be NONPOLAR with no dipole moment, even though the bonds within the molecule may be polar.

In this molecule, CH4, the C atom has a higher EN than the surrounding H atoms.  This makes the individual C-H bonds polar with a slightly negative charge on the C atom and a slightly positive charge on the H atoms.  However, the polarities on these bonds balance each other out around the central C atom.  The overall dipole moment of the molecule is therefore 0 and the molecule is nonpolar.

If there are unequal polar bonds around the central atom, then the overall molecule will be POLAR with an a dipole moment.

In the this molecule, CH3Cl, all of the bonds around the central C atom are not equal.  The the Cl atom has a higher EN than the  central C atom, so the electrons are more drawn toward the Cl atom.  The H atoms have a lower EN than the central C atom and cannot balance out the large EN of the Cl atom.  Therefore, the overall molecule has a dipole moment with a slightly negative charge on the Cl end and a slightly positive charge on the H ends, making the molecule polar.

In this molecule, H2O, the two O-H bonds are polar but are equal to each other.  However, the central O atom has a higher EN than the H atoms so the electrons are drawn toward the O atom.  The O atom also has two pairs of nonbonding electrons, and combined the nonbonding electrons and the electrons drawn toward the O from the O-H bonds result in an overall dipole moment.  Therefore, this molecule is polar.