IB Chemistry/Periodicity

Periodicity Revision Notes

3.1 The Periodic Table

3.1.1 : Elements increase in atomic number across each period, and down each group.

3.1.2 : Group - the columns proceeding vertically. Period - the rows proceeding horizontally.

3.1.3 : Group = number of valence electrons in the atom. Period = number of main electron shells...s, p , d and f blocks as described above.

3.2 Physical Properties

3.2.1 : Electron structure

• As a period is crossed, valency electrons increase
• As a group is descended, a shell is added (with the same number of valence electrons)

Atomic radii

• Across a period, radii decreases due to an increasing nuclear charge and reduced screening (more electrons and protons in the outer shell)
• Down a group, radii increases due to additional shells blocking nuclear charge and increased screening (shielding)

Ionic radii

• Positive ions (Na ‡ Na+ + e-) have smaller radii due to 1 less shell and reduced screening
• Negative ions (F + e- ‡ F-) have larger radii because the positive charge of the nucleus has less positive pulling power (PPP) due to the presence of an extra electron in this ionic state.

Ionisation energy Ionization energy is the energy required to remove 1MOLE of electrons from 1MOLE of gaseous atoms (e.g. M ‡ M+ + e-)

Successive ionization energies are 3rd > 2nd > 1st, because of smaller radii and greater charge.

Electron affinity Electron affinity is the energy change on the addition of 1 mole of electrons to 1 mole of gaseous atoms (e.g. X(g) + e- ‡ X-(g))

Electronegativity Electronegativity is the measure of the ability of an element in a bond to attract electrons. The most electronegative elements are (in decreasing) F, O, and N.

• Across a period, electronegativity increases because the increase in nuclear charge makes the nucleus more attracted to electrons
• Down a group, electronegativity decreases due to increased screening and less nuclear attraction

Melting Points

• Melting point decrease down group 1 due to the elements metallic structure, which is held together by attractive forces between de-localized electrons.
• Melting point increases down group 7 due to the strong intermolecular forces that increase in strength with more electrons.

Periodic Trends - Moving Across the Periodic Table

Because of the way in which the periodic table is organized, many trends become apparent.

Moving across the periodic table, proton and electrons are added. Because all of these electrons are added into the same shell, the ionization energy increases. The ionization energy is the amount of energy required to remove the outermost electrons. Because of the overall greater attraction between more protons and more electrons across the periodic table, it becomes more and more difficult to remove an electron. The atomic radius of the elements decreases. Also, metallic character decreases (electronegativity increases; electronegativity is the tendency of an element to gain electrons). This traces the tendency of atoms of metals to lose electrons while nonmetals gain them.

Periodic Trends - Moving Down the Periodic Table

Because elements in the same family (column) contain the same valence shell electron configuration, they tend to behave very similarly. Moving down the periodic table, electrons are added to successively higher energy levels. Because of this, atomic radius increases. Also, it is easier to remove the outermost electron of larger atoms, so ionization energy decreases as well as electro negativity.

Noble Gases 0/8A/18

helium, neon, argon, krypton, xenon, radon

Noble gases except helium all share a full valence shell electron configuration ( 8 outer electrons, 2 for helium). Because this configuration is extremely stable as well as symmetrical, the noble gases are very unreactive and will only react under extremely rigorous conditions.

Halogens 7A/17

fluorine, chlorine, bromine, iodine, astatine

Halogens have an electron configuration of 7, meaning they require only one more electron for a noble configuration. They have very negative electron affinities, are extremely reactive, and form ions with a -1 charge. They are so reactive that in their homogenous state UV light will catalyze a radical reaction.

Fluorine is the most electronegative of the elements and is so reactive that it attacks almost any other element (noble gases, oxygen, nitrogen, and gold are the exceptions) to form fluorides. Chlorine is somewhat less reactive, bromine somewhat less reactive than chlorine, and iodine even less, but even iodine is a formidable iodizer. Strong radioactivity masks the chemical properties of astatine.

All except astatine form gaseous compounds with hydrogen: hydrogen fluoride HF, hydrogen chloride HCl, hydrogen bromide HBr, and hydrogen iodide HI; these are acidic, strongly reactive substances; except for hydrogen fluoride these substances are among the strongest known acids. They free halogens and the hydrogen halides react with most metals, halides of metals being known as salts, of which sodium chloride NaCl is best known as "salt".

Electronegativity decreases in this group with increasing atomic mass, and oxygen is more electronegative than any element except fluorine; it acts much like a halogen except for its -2 oxidation state. Fluorine and oxygen oxidize these elements to the +6 oxidation state, resulting in such substances (for sulfur) as sulfur hexafluoride SF₆ and sulfur trioxide SO₃ and its derivative sulfuric acid H₂SO₄, one of the most heavily-used industrial chemicals. Fluorine, oxygen, chlorine, and bromine oxidize all of the elements of this group except oxygen to the +4 oxidation state.

Oxides of sulfur, selenium, and tellurium are acidic. Strong radioactivity largely masks the chemical properties of polonium.

Alkaline Earth Metals group 2A/2

(beryllium, magnesium, calcium, strontium, barium, radium)

These substances are all metals having two electrons in the outer shell, or the ns² configuration. With increasing mass these elements become softer, have lower melting and boiling points, and become more reactive. None appear uncombined in nature, and all are separated from their compounds with difficulty. All react with halogens and, except for beryllium, with water and oxygen (magnesium at temperatures higher than those of living things)

All oxidize to the +2 state that represents an ion in the stable configuration of an inert gas.

Beryllium is the least reactive; it is the hardest of these elements. Its oxide is amphoteric; it reacts with both strong acids and bases.

Magnesium burns in hot air or steam to form alkaline oxides or hydroxides. Powders or thin slices of this metal can be ignited with a match. Magnesium hydroxide is a strong base, although its low solubility masks its alkalinity. At room temperature magnesium metal is light and strong, but it can be ignited in air. Halogens and acids attack magnesium violently with the formation of salts like magnesium chloride (MgCl₂) or magnesium sulfate (MgSO₄). Although magnesium would react with atmospheric oxygen, an oxide layer protects the metal from oxidation.

With calcium and the heavier elements of this group, the reactions are even more violent and take place at lower temperatures; they react not only with halogens and acids but also with water, oxygen, and even nitrogen. They are such strong reducing agents that they can reduce carbon dioxide to carbon as well as many metal oxides to metals.

The strong radioactivity of radium masks its chemical properties, but to the extent that its chemistry is known it fits well into the group and its pattern.

Alkali Metals group 1A/1

(lithium, sodium, potassium, rubidium, caesium (US spelling cesium), francium)

These elements are best marked by their reactivity. Physically they are soft, shiny (when freshly prepared) solids with low melting points; they conduct electricity well. They all have one outermost electron that they lose easily to almost any electronegative substance other than nitrogen (lithium reacts with nitrogen, unlike the others). None is ever found uncombined on earth, and none is ever put to use as a structural metal. They must be kept under inert liquids such as kerosene or in inert gases (nitrogen suffices for any of these elements other than lithium.

All oxidize easily to the +1 oxidation state. High reactivity masks the chemistry of francium.

They react with atmospheric oxygen to form various oxides and react violently with water, halogens, and acids. A typical reaction between one of these metals and water is as such:

Na_(s) + H₂O → Na⁺_(aq) + OH^-_(aq) + 1/2 H_{2 (g)}

(WARNING: this reaction generates much heat. One should not cast any alkali metal into water or acids except with the cautions that professional chemists use. The hydrogen gas from this reaction may itself ignite with atmospheric oxygen in a dangerous flame. These metals should never allowed to touch flesh because they react with any water upon them and yield corrosive hydroxides that burn flesh. Strong solutions of alkali metal hydroxides are destructive to flesh).

Hydroxides of these elements dissociate completely in water to form some of the strongest bases known. These are as strongly alkaline as any acids and react violently with acids to form halides and water in neutralization:

NaOH_(s) + HCl _(g) → NaCl_(s) + H₂O _(s)

The results of any reactions of these metals and any acids are salts. Almost all salts of these elements are highly soluble in water and form conducting solutions, proving their ionic nature. The best known of these substances is sodium chloride, NaCl, a substance known as common salt. Salts of these elements and strong acids are neutral (for example, potassium nitrate KNO₃); salts with weak acids such as acetic acid are alkaline (sodium carbonate, Na₂CO₃).

Hydrogen, an element like no other

Hydrogen is the most common element in the universe; as a gas, it is too light for the Earth's gravitation to hold. It is by far the most common constituent of the Sun and all other stars and of the gas giant planets of our solar system. It exists on or just under the surface of the Earth as a component of water and in innumerable compounds of carbon, many essential to life.

The heat and light from the sun (or any other star) arises largely from the nuclear fusion of hydrogen into helium. Nuclear reactions are outside of the normal scope of the discussion of chemistry, so that can be largely ignored here.

Hydrogen, although having one outermost electron does not fit into the alkali metals or any other group. It deserves its own treatment. It forms compounds analogous to those of the alkali metals, but such hydrogen compounds are much less alkaline (or more acidic), much less ionic, and more volatile. Sodium chloride, the stereotypical salt, is neutral and clearly ionic; hydrogen chloride is a non-ionic gas under normal conditions and is a strong acid. The hydrogen analog of sodium hydroxide is a volatile liquid under normal situations: unlike the strongly alkaline and solid sodium hydroxide, water is feebly ionic and effectively neutral.

Hydrogen is a non-metal, forming a diatomic gas which results from the sharing of the single electrons of hydrogen atoms. It can achieve a stable ionic structure (no electrons!) by losing an electron or by gaining an electron and achieving the completed shell configuration of helium. The hydrogen molecule is best described as sharing the two electrons between two hydrogen atoms. This structure is highly stable and has little inclination to form bonds between other hydrogen molecules; hydrogen is a gas down to some of the lowest temperatures known. It is also the lightest of gases, weighing less even than helium.

Hydrogen readily shares its electron with a strongly-electronegative element like any halogen, oxygen, or sulfur. The combination with fluorine is particularly violent and possible down to very low temperatures:

1/2 H_{2 (g)} + 1/2 F_2(g) → HF_(g)

Light is enough to force combustion between hydrogen and chlorine, and a spark is enough to cause combustion between hydrogen and oxygen. In view of the Hindenburg disaster of 1936, helium has long replaced hydrogen in lighter-than-air aircraft. That reaction is as such:

H_{2 (g)} + 1/2 O_{2 (g)} → H₂O _(g)

(Note that at the temperatures associated with such a combustion, water is in the gaseous state!)

Somewhat analogously (but not very well), hydrogen can act somewhat like a halogen, forming hydrides with some metals. Most of these react violently with water to form hydrogen gas and the metal hydroxide. Hydrogen compounds with non-metals are typically among the most volatile substances of those elements.

Under pressure, in aqueous solution or non-solid acids, hydrogen is a good reducing agent. Strong acids attack most metals:

Zn _(s) + 2 H⁺ _(aq) → Zn ²⁺ _(aq) + H _{2 (g)} .

In the atmospheres of gas giants, gaseous hydrogen under great pressure reduces nitrogen to ammonia, carbon compounds to methane and other hydrocarbons, and oxides to water.

Hydrogen forms more chemical compounds than any other element including carbon (almost all carbon compounds are compounds of hydrogen as well and vice-versa, but more substances containing hydrogen but not carbon exist than do compounds of carbon but not hydrogen). Hydrogen forms bonds with most non-metals, including oxygen, nitrogen, and carbon. Although a hydrogen atom can bond with only one other element, and then only in a single bond, hydrogen allows very long chains of carbon atoms to form. Most of the hydrogen compounds with carbon alone are combustible gases or volatile liquids or waxy solids that can be vaporized and burned to produce water, carbon dioxide, and much heat. Natural gas, gasoline (a mixture of liquid hydrocarbons), and waxes as found in candles make suitable fuels. With such other elements as oxygen, nitrogen, sulfur, and in some cases metals hydrogen allows the formation of substances necessary for life, including carboxylic acids, sugars, proteins, nucleic acids, hemoglobin, and chlorophyll.

Such complex compounds are ordinarily discussed in organic chemistry, a study associated more obviously with carbon.