Difference Between Reactivity Series And Electrochemical Series

Reactivity And Electrochemical Series - Metals

The reactivity series can be used to compare the relative reactivities of different metals. It lists metals in order of general chemical reactivity. Metals generally react by losing electrons to form positive ions. The more readily a metal loses electrons, the more reactive it is – and the greater its strength as a reductant. Metals higher up in the series can reduce the ions of those lower down.

The standard electrode potential of a metal also indicates its strength as a reductant. The more negative the value of the standard electrode potential of a metal, the greater is its strength as a reductant. Hence, you might expect the metal reactivity series and standard electrode potentials to list metals in the same order. However, you must remember that the metal reactivity series is based on observing a range of reactions, such as displacement reactions between solid metals and solid metal oxides. Standard electrode potentials refer specifically to reactions taking place in aqueous solution.

Below is a comparison of the reactivity series with the electrochemical series, which ranks metals according to their standard electrode potentials.

The obvious discrepancy is the relative positions of sodium and calcium. Calcium is a stronger reductant than sodium according to Eo (standard electrode potential) values, but the metal reactivity series suggests that calcium is less reactive than sodium. This discrepancy arises because calcium reacts at a much slower rate, in displacement reactions for example, which in turn happens because two electrons must be removed, not one as for sodium.

Note also that aluminium reacts readily with oxygen in the air, forming a layer of stable aluminium oxide on its surface. This impervious oxide coat often causes aluminium to exhibit lower reactivity than its position in the metal reactivity series indicates.

Group IV Elements - CAPE Chemistry Unit 1

Elements - Structure And Bonding

Main Points To Be Used In Answering A Past Paper Question:

• Down the group there is a change in structure from giant molecular to giant metallic and a change in bonding from covalent to metallic.
• From C to Ge elements exhibit a giant molecular structure.
• Sn and Pb exhibit a giant metallic lattice structure.
• From C to Ge bond length between group IV atoms increases and bond energy/strength decreases.

Elements - Electrical Conductivity

Trend: Going down the group there is a general increase in electrical conductivity.

Reason: Down the group there is a gradual increase in metallic character due to an increase in delocalisation of electrons throughout the structure.

Note: There is an increase in electrical conductivity from C(diamond) to Sn. There is a decrease from Sn to Pb which has never been addressed in the past papers. Hence we use the phrase "general increase" when talking about the trend in electrical conductivity down the group. 

Unit 1 Module 1 SS 2 Notes

Global Chemistry Lessons

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2.1

State the various forces of attraction between particles.

Ionic bonds, covalent bonds, metallic bonds, van der Waals' forces.

Both bonding (intramolecular) forces and intermolecular forces arise from electrostatic attractions between opposite charges. Bonding forces are due to the attraction between cations and anions (ionic bonding), nuclei and electron pairs (covalent bonding), or metal cations and delocalized valence electrons (metallic bonding). Intermolecular forces, on the other hand, are due to the attraction between molecules as a result of partial charges, or the attraction between ions and molecules. The two types of forces differ in magnitude, and Coulomb's law explains why:

• Bonding forces are relatively strong because they involve larger charges that are closer together.
• Intermolecular forces are relatively weak because they typically involve smaller charges that are farther apart.

Ion-Dipole Forces

When an ion and a nearby polar molecule (dipole) attract each other, an ion-dipole force results. The most important example takes place when an ionic compound dissolves in water. The ions become separated because the attractions between the ions and the oppositely charged poles of the H2O molecules overcome the attractions between the ions themselves.

Dipole-Dipole Forces

When polar molecules lie near one another, as in liquids and solids, their partial charges act as tiny electric fields that orient them and give rise to dipole-dipole forces: the positive pole of one molecule attracts the negative pole of another (diagram below).

Polar molecules and dipole-dipole forces. In a solid or a liquid, the polar molecules are close enough for the partially positive pole of one molecule to attract the partially negative pole of a nearby molecule. The orientation is more orderly in the solid (left) than in the liquid (right) because, at the lower temperatures required for freezing, the average kinetic energy of the particles is lower. (Interparticle spaces are increased for clarity.)

For molecular compounds of approximately the same size and molar mass, the greater the dipole moment, the greater the dipole-dipole forces between the molecules are, and so the more energy it takes to separate them. Consider the boiling points of the compounds in the next diagram. Methyl chloride, for instance, has a smaller dipole moment than acetaldehyde, so less energy is needed to overcome the dipole-dipole forces between its molecules and it boils at a lower temperature.

Dipole moment and boiling point. For compounds of similar molar mass, the boiling point increases with increasing dipole moment. (Note the increasing color intensities in the electron density models.) The greater dipole moment creates stronger dipole-dipole forces, which require higher temperatures to overcome.

The Hydrogen Bond

A special type of dipole-dipole force arises between molecules that have an H atom bonded to a small, highly electronegative atom with lone electron pairs. The most important atoms that fit this description are N,0, and F. The H-N, H-O,
and H- F bonds are very polar, so electron density is withdrawn from H. As a result, the partially positive H of one molecule is attracted to the partially negative lone pair on the N, 0, or F of another molecule, and a hydrogen bond (H bond) forms. Thus, the atom sequence that allows an H bond (dotted line) to form is -B:····H-A-, where both A and B are N, O, or F. Three examples are

The small sizes of N, O, and F are essential to H bonding for two reasons:

1. It makes these atoms so electronegative that their covalently bonded H is
    highly positive.


2. It allows the lone pair on the other N, O, or F to come close to the H.

The Significance of Hydrogen Bonding

Hydrogen bonding has a profound impact in many systems. Here we'll examine one major effect on physical properties and preview its enormous importance in biological systems.