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Avogadro - chemistry terms
Category: Sciences > Chemistry
Date & country: 14/09/2007, UK
Words: 50


Absolute Zero
This is the origin of a new temperature scale called the 'absolute' or Kelvin scale. It is the theoretical temperature at which an ideal gas would occupy zero volume if it could be cooled indefinitely without liquefaction or solidification. The absolute zero is -273.16 °C, and one degree on the absolute scale is equivalent to one Celsius degree. Convert a Celsius temperature to a Kelvin temperature by adding 273.16 (approximately 273) to the Celsius temperature.
In the graph below, the straight lines illustrate Charles' Law, in which the volume of a fixed mass of gas is measured at different temperatures, the pressure remaining constant. It is when these lines are extended to cross the temperature axis that they are found to do so at -273.16 °C. This is the origin of a new temperature scale, called the Kelvin or Absolute Temperature Scale.
Celsius temperatures are arbitrary, and absolute values must always be used in calculations involving gases.

Acid
An Acid is a substance containing hydrogen in its molecules (or ions) that it can release as H+ ions. The Bronsted-Lowry definition of an acid is that it is a 'proton donor'. By strength of an acid is meant the tendency with which it will donate H+ ions; the stronger the acid the more readily it will donate H+ ions. The terms 'strength' and 'concentration' with regard to acids must not be confused; a concentrated solution of a weak acid is still a weak acid.
Hydrochloric acid, HCl, is an example of a strong monobasic acid; it is essentially completely ionised in aqueous solution.
HCl(aq) + H2O(aq) ® H3O+(aq) + Cl-(aq)
The organic acid, ethanoic acid, CH3COOH, is an example of a weak monobasic acid; it is partially ionised in aqueous solution, an equilibrium existing between the acid and its conjugate base.
CH3COOH(aq) + H2O(l) = H3O+(aq) + CH3COO-(aq)
There is another common definition of an acid. It is the Lewis definition. A Lewis acid is an electron-pair acceptor.

Acid Dissociation Constant, Ka
The Acid Dissociation Constant, Ka, is the equilibrium constant for the reaction in which a weak acid is in equilibrium with its conjugate base in aqueous solution. Notice that in the equilibrium expression below the concentration of water is not included. This is because water is vastly in excess and the amount changes negligibly on equilibrium being established. Ka can be thought of as a modified equilibrium constant. For example,
CH3COOH(aq) + H2O(l) = CH3COO-(aq) + H3O+(aq)
Ka = [CH3COO-(aq)][H3O+(aq)] / [CH3COOH(aq)]
Therefore, the larger the value of Ka, the stronger is the acid. The value is sometimes expressed as the logarithm of its reciprocal, called pKa. Therefore,
pKa = -log Ka
The smaller the value of pKa the stronger the acid. Ka is a better measure of the strength of an acid than pH because adding more water to the acid solution will not change the value of the equilibrium constant Ka, but it will change the H+ ion concentration on which pH depends.
In the above reaction, ethanoic acid and ethanoate ions form a conjugate acid-base pair.

Activation Enthalpy
It is customary to use the two terms activation enthalpy and activation energy interchangeably, but there is, in fact, a difference between them.
With regard to the collision theory of reaction rates, we postulate that molecules react only if in a collision they possess an energy equal to or greater than a certain critical value. This is called the energy of activation, Ea.
The graph below shows that, at a particular temperature, the molecules posses a range of kinetic energies or molecular speeds. The area under each curve is proportional to the total number of molecules. The distribution of kinetic energies is at two different temperatures. At the higher temperature, the average kinetic energy is higher, and the shaded areas show that at the higher temperature more molecules on collision possess an energy equal to or greater than the minimum kinetic energy, called the 'activation energy'.
More modern theories of reaction rates, namely the transition state theory, modify the idea that molecules must 'collide' in a reaction. As reactant molecules approach each other a continuous series of changes in bond distances is visualised. Some bonds begin to lengthen and finally break, while other bonds begin to form. Energy changes accompany these continuous changes in the arrangement of atoms in molecules. Finally, the reacting molecules must achieve a specific arrangement before they can form the products of the reaction. This specific, transient, arrangement possessing a definite energy is known as the transition state. The transition state is assumed to possess properties common to real molecules. The transition state corresponds to the state of highest enthalpy. This is approximately equal to the activation energy.
The diagram below shows a reaction profile for an exothermic reaction. The 'activation enthalpy' is shown because it has become more usual to use this term.

Allotropy
Allotropes are different forms of the same element that exist in the same physical state.
Examples are: oxygen and ozone, rhombic sulphur and monoclinic sulphur, diamond and graphite. Such elements are said to show allotropy, and the various different forms are termed allotropes. In the case of diamond and graphite (and the Fullerenes, the most famous of which is Buckminsterfullerene, C60) the atoms of the element are arranged to form different molecules with different crystal shapes. In rhombic and monoclinic sulphur, discrete S8 molecules are arranged differently in the two crystal forms. The allotopes of an element have different chemical and physical properties.

Avogadro Constant, L
The Avogadro Constant, L is a constant number used to refer to atoms, molecules, ions and electrons. Its value is 6.023 x 1023 mol-1, like a dozen is 12 and a score is 20. Note the units of the Avogadro Constant.
For elements that consist of atoms (rather than molecules), the relative atomic mass expressed in grams contains the Avogadro Constant of atoms, whatever the element. Likewise, the relative molecular mass expressed in grams is the mass of the Avogadro Constant of molecules. As an example, 23 g of sodium contain the same number of atoms as there are molecules in 17 g of ammonia.
The value of the Avogadro constant can be determined experimentally.

Avogadro's Law (1811)
Equal volumes of any gases measured under the same conditions of temperature and pressure contain equal numbers of molecules (or atoms if the gas in monatomic).
This law was first known as Avogadro's Hypothesis. It was proposed in explanation of Gay-Lussac's Law of Combining Volumes, in which Avogadro also reasoned that elements could consist of 'molecules'.
The volume of 1 mole of any gas (the molar volume) is therefore a constant at a given temperature and pressure.

Base
A Base is a substance whose ions or molecules can accept H+ ions. The Bronsted-Lowry definition of a base is that it is a 'proton acceptor'. By strength of a base is meant the tendency with which it will accept H+ ions; the stronger the base the more readily it will accept H+ ions. The terms 'strength' and 'concentration' with regard to bases must not be confused; a concentrated solution of a weak base is still a weak base. There is another common definition of a base. It is the Lewis definition. A Lewis base is an electron-pair donor.

Base Dissociation Constant, Kb
The Base Dissociation Constant, Kb, is the equilibrium constant for the reaction in which a weak base is in equilibrium with its conjugate acid in aqueous solution. For example,
NH3(aq) + H2O(l) = NH4+(aq) + OH-(aq)
Kb = [NH4+(aq)][OH-(aq)] / [NH3(aq)]
Therefore, the larger the value of Kb, the stronger is the base. The value is sometimes expressed as the logarithm of its reciprocal, called pKb. Therefore
pKb = -log Kb
The smaller the value of pKb the stronger the base. Kb is a better measure of the strength of a base than pH because adding more water to the base solution will not change the value of the equilibrium constant Kb, but it will change the H+ ion concentration.
In the above reaction, ammonium ions and ammonia form a conjugate acid-base pair.

Bond Enthalpy
The Bond Enthalpy is the energy required to break a chemical bond. It is usually expressed in units of kJ mol-1, measured at 298 K. The exact bond enthalpy of a particular chemical bond depends upon the molecular environment in which the bond exists. Therefore, bond enthalpy values given in chemical data books are averaged values.
Bond breaking is an endothermic process, and the bond enthalpy involved is given a +ve sign. Bond making is an exothermic process. For the formation of a given chemical bond, the bond enthalpy has the same value, except it is given a -ve sign.
Generally, the shorter a chemical bond the stronger it is. This can be illustrated with the halogen-halogen bond enthalpies:
Bond Bond Length (nm) Bond Enthalpy (kJ mol-1)
F – F0.142158
Cl – Cl0.199242
Br – Br0.228193
I – I0.267151
Note the apparently anomolous bond enthalpy of the F-F bond. This is explained by the lone pairs of electrons on the fluorine atoms being closer together and therefore repelling each other more than is the case in the other halogen molecules.
Bond enthalpies can be used to calculate enthalpy changes for reactions, but even though standard bond enthalpy values are used, the calculated DH is likely to differ from that given in a chemical data book. There are two reasons for this. Bond enthalpy values are averaged values, and in the chemical reactions where they are applied, all the reactants and products are taken to be gaseous.
In photochemical reactions, for example, the free radical chlorination of methane, a chlorine molecule, Cl2, absorbs a photon of ultra-violet radiation sufficient in energy to break the Cl-Cl bond homolytically. Given the bond enthalpy of the Cl-Cl bond, E(Cl-Cl) = 242 kJ mol-1, the energy required to break just one such bond can be calculated.
Energy to break one Cl-Cl bond = 242 x 1000/6.023 x 1023 = 4.02 x 10-19 J

Boyle's Law
The volume of a fixed mass of gas (fixed number of molecules) is inversely proportional to the pressure, measured at constant temperature.
P a 1/V
\ P1V1 = constant
\ P1V1 = P2V2
(at constant temperature)

Buffer Solution
A Buffer Solution maintains the pH of a solution by reacting with small amounts of an added acid or base. For a buffer solution to be able to do this it must contain both an acid and a base; the acid to react with any added base and a base to react with any added acid. But these must be able to co-exist without reacting with each other. For this to be so the acid and base must be a conjugate acid-base pair. Examples of buffer solutions are a mixture of ethanoic acid and sodium ethanoate, or ammonia solution and ammonium chloride.
In aqueous solution, an acid and its conjugate base establish equilibrium. For example:
CH3COOH(aq) + H2O(l) = CH3COO-(aq) + H3O+(aq)
Ka = [CH3COO-(aq)][H3O+(aq)] / [CH3COOH(aq)]
Rearranging the above expression,
Ka / [H3O+(aq)] = [CH3COO-(aq)] / [CH3COOH(aq)]
Therefore a buffer solution of a desired pH can be prepared. From the pH, the H3O+(aq) concentration can be found, and the buffer solution prepared by adjusting the ratio of the weak acid and conjugate base.

Charles' Law
The volume of a fixed mass of gas (fixed number of molecules) is directly proportional to the absolute temperature, measured at constant pressure.
V a T
\ V = constant x T
\ V/T = constant
\ V1/T1 = V2/T2
(at constant pressure)

Covalent Bond
A pair of electrons shared between two bonded atoms, attracting each of the positively charged nuclei of the atoms and so bonding them together, constitutes a single covalent bond. Each of the bonded atoms provides one electron of the shared pair. Sometimes, both electrons of the shared pair can be provided by just one of the bonded atoms forming a covalent bond that is also known as a 'dative' or 'co-ordinate' bond. Once a dative bond is formed it is indistinguishable from any other covalent bonds between the same atoms. For example, in the ammonium ion, NH4+, the non-bonding pair (lone pair) of electrons on the nitrogen atom is donated to an H+ ion.
:NH3 + H+ ® NH4+
Two shared pairs of electrons between bonded atoms consitutes a double covalent bond, and three shared pairs forms a triple covalent bond.

Electronegativity
Electronegativity is a measure of the net tendency of a bonded atom in a molecule to attract electrons.
There have been several calculations of electronegativity values. Those of Linus Pauling are one example. Fluorine, the most electronegative element, is given an arbitrary value of 4.0 and the electronegativities of the atoms of other elements are related to it.
Electronegativites can be used to make rough predictions about the type of bonding to be found in a compound. The greater the difference in electronegativity values of the two bonding elements, the greater is the chance for ionic bonding. Also, the electronegativity difference between two covalently bonded atoms determines the polarity of the bond: The larger the difference, the more polar is the covalent bond with the more electronegative element bearing the partial negative charge, d-. If the difference is zero or very small, the bond is essentially non-polar.

Element, Compound and Mixture
An Element is a substance that cannot be broken down into anything simpler by chemical means. It is made up of atoms of one kind only. Some elements can consist of single atoms, whereas as others can consist of molecules, such as O2, S8 and P4.
A Compound is formed when two or more elements become chemically combined together. The elements react chemically and chemical bonds exist between the atoms involved. The elements react in definite amounts, and therefore a compound has a definite chemical composition. A compound often does not resemble the elements from which it was formed. Often it cannot easily be converted back to the elements from which it is formed.
A Mixture is formed when two or more different substances are mixed together. There is no chemical combination. The substances can be present in any proportions. Often they can be identified visually within the mixture. Often the components of a mixture can be easily separated, e.g. a magnet can be used to separate iron filings from a mixture with sulphur.

Empirical Formula
An Empirical Formula shows the simplest whole number ratio of atoms in a molecule of a substance. For example, C2H6 is the 'molecular formula' of ethane. Its 'empirical formula' is CH3.

Equilibrium Constant
The Equilibrium Constant, Kc, relates to a chemical reaction at equilibrium. It can be calculated if the equilibrium concentration of each reactant and product in a reaction at equilibrium is known. The equilibrium expression below, formed from the general chemical equilibrium, is universally true. The chemical components happen to be gases.
aA(g) + bB(g) = cC(g) + dD(g)
The Ideal Gas Equation shows that the pressure of a gas is proprtional to its concentration.
pV = nRT
where p is pressure of a particular gas (its partial pressure) in an equilibrium mixture, V is the total volume, n is the number of moles of the particular gas, R is the general gas constant, and T is the absolute temperature.
In the above equation where the temperature is also constant,
P a n/V
The equilibrium constant can therefore also be expressed in terms of partial pressures, and is denoted as Kp:
There are several types of equilibrium constants. Each is constant at a constant temperature. For example, consider the following ionic equilibrium involving the aqueous solution of a weak acid:
CH3COOH(aq) + H2O(l) = CH3COO-(aq) + H3O+(aq)
Ka is called the 'acid dissociation constant'. Note that water is omitted from the expression because it is present in such vast excess that its concentration changes negligibly on the formation of equilibrium and is therefore effectively constant. The concentration of the water is included in the equilibrium constant, and Ka can be thought of as a modified equilibrium constant.

Faraday Constant
A mole of electrons is given a special name: 1 Faraday. The charge of 1 mole of electrons is called the Faraday Constant, F.
It is calculated by multiplying the charge of one electron by the Avogadro constant.
1 Faraday = 9.648 70 x 104, coulombs/faraday (C mol-1).
In electrolysis, Michael Faraday in 1832 - 1833 discovered the relationship between the charge and the quantity of matter liberated in an electrode reaction. He did not know about partial reactions at electrodes, electrons, or a unit (to be) named after him. He did know something about atomic and molecular masses. His results are summarised in two laws, now famous as Faraday's laws.
The mass of a substance produced or consumed is proportional to the quantity of charge (current x time) that has passed through the circuit.
The number of faradays that must pass through a circuit when one mole of a substance is produced or consumed is a whole number.
What the whole number is depends on the substance and the reaction. For example:
2 faradays produce 1 mol of H2:
2H+ + 2e- ® H2
1 faraday produces 1 mol of Ag:
Ag+ + e- ® Ag
3 faradays produce 1 mol of Al:
Al3+ + 3e- ® Al
That one mole of atoms or molecules always requires an integral number of faradays suggests a once novel but now familiar idea: the existence of a natural unit of charge, the charge of an electron.

Half-Life
The Half-life, t½, is the time required for the activity of a radioactive isotope to fall to one-half of its original value. This may also be expressed in terms of the amount of the material in grams. It is a constant for a given isotope and does not depend on the particular compound in which it is present. It is impossible to know when any one particular atom will undergo decay, if ever. But even an extremely small sample of a radioactive isotope contains trillions of atoms, and when dealing with such large numbers scientists can make predictions about the number which will undergo decay in a given period of time. The time taken for half the original number to decay is chosen.

Hess's Law
The overall enthalpy change that accompanies a chemical reaction is independent of the route by which the reaction takes place, provided the initial and final states are the same. It is a statement of the 'Law of Conservation of Energy' since, for example, if the reverse of a reaction was less endothermic than the forward reaction was exothermic, energy could be created.
The precise value of DH depends on the temperature and pressure at which it is measured, and therefore standard conditions have to be defined so that DH values can be compared. Standard temperature is 298 K and standard pressure is 1 atm (101 325 Nm-2). Standard values are given in chemical data books.

Ideal Gas
The gas laws (Boyle's, Charles' and Avogadro's, and the Ideal Gas Law) have a mathematical basis. Real gases do not obey these laws exactly. For example, in a real gas two molecules upon collision can stick together for a while, in which time they behave as one molecule. An ideal gas, or perfect gas, only exists in the imagination. Its molecules are volumeless points in space, have mass and velocity, do not attract or repel each other or the walls of their vessel, collide with each other and with the walls of the vessel perfectly elastically without loss of total energy. A real gas approximates to ideal behaviour at high temperatures and low pressures. The gas laws are mostly adequate for dealing with real gases. The van der Waal's equation is an example of an equation derived to give better approximations of the behaviour of real gases than the ideal gas law.

Ideal Gas Law
The Ideal Gas Law is a combination of Boyle's Law, Charles' Law and Avogadro's Law. It can be expressed by a single equation, PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the general gas constant, and T is the absolute temperature.
The value of the gas constant, R is 8.3143 J K-1 mol-1. Note that these are SI units, and therefore the volume used in this equation must be in m3and the pressure in N-2. In all gas law calculations, absolute values of temperature (Kelvin) must be used.
Another commonly encountered value for R is 0.08205 atm dm3 K-1 mol-1. If this value is used then the volume must be in dm3 and the pressure in atm.

Indicators
An Indicator is a substance that undergoes a change in colour when the end-point of a titration is reached.
Acid-base indicators are perhaps the most common types, for example, phenolphthalein and methyl orange. but different types of indicators are used in precipitation reactions, and in reactions where there is a colour change an indicator need not be added.
Acid-base indicators are used to signal the end of acid-base titrations. These indicators change colour within a characteristic pH range. For example:
IndicatorAcidBasepH range
Methyl OrangeRedYellow3 - 4.4
PhenolphthaleinColourlessPink8 - 10
Indicators are also commonly used in precipitation reactions. For example, in the determination of Cl- by titration with AgNO3, Chromate(VI) ion, CrO42-, is used as the indicator. A brick-red precipitate of silver chromate(VI) signals the end-point.
Colour changes may also occur as part of the reaction, and in these cases an indicator may not be needed. For example, in manganate(VII) titrations a tiny drop of the manganate(VII) solution from the burette is sufficient to impart a faint permanent pink colour to the solution at the end-point of a titration.

Ionic Bond
An Ionic Bond is an electrostatic force of attraction between oppositely charged ions. In a giant ionic crystalline lattice, the oppositely charged ions are packed together in a highly regular way. Ionic Bonding exists in every direction, three-dimensionally throughout the entire crystalline structure. Ionic bonding is not uni-directional. Any particular ion is not associated with just one other. For example, in sodium chloride each Na+ ion is surrounded by six Cl- ions, and each Cl- by six Na+ ions. The separate entity, NaCl, does not exist. The substance is not molecular it is ionic. Relating to their structure, ionic compounds typically are hard and have high melting and boiling points, and they can conduct electricity when their ions are free to move about, as in aqueous solutions and when molten.
The picture below illustrates the sodium chloride crystalline ionic lattice structure. (Sir William Henry Bragg (1862-1942), with his son Sir William Lawrence Bragg, was awarded the Nobel Prize for Physics in 1915 for work using the X-ray spectrometer on the arrangement of particles in crystals.)

Ionic Product of Water
The Ionic Product of Water, Kw, is the equilibrium constant for the reaction in which water undergoes an acid-base reaction with itself. That is, water is behaving simultaneously as both an acid and a base.
H2O(l) + H2O(l) = H3O+(aq) + OH-(aq)

Kw = [H3O+(aq)][OH-(aq)]
At 298 K, the value of Kw is 1 x 10-14 mol2 dm-6. This makes the concentration of H+ ions equal to 1 x 10-7 mol dm-3, and therefore the pH is 7. This is defined as 'neutral'.
From the above equilibrium expression, taking -log10 throughout
pKw = pH + pOH = 14
Ionic Product does not apply only to water. It applies, for example, to the equilibrium in liquid ammonia:
NH3 + NH3 = NH2- + NH4+

Isomerism
Isomers have the same molecular formula, but their molecules are different. Their differences arise in a number of ways. Structural isomerism occurs when the atoms are bonded together in a different order. Stereoisomerism arises when the molecules are not superimposable because their atoms are orientated differently in space. Stereoisomers have the same sequence of chemical bonds. There are two types of stereoisomerism: cis-trans geometric isomerism arises owing to the lack of free rotation about a double covalent bond; optical isomerism arises when an atom is tetrahedrally bonded to four different atoms or groups of atoms.

Isotope
Isotopes are different atoms of the same element, therefore having the same number of protons, but containing a different number of neutrons. Their relative masses are different because it is essentially the protons and neutrons in the nucleus that make up the mass of the atom. Chlorine, for example, has two naturally occurring isotopes, chlorine-35 and chlorine-37.

Le Chatelier's Principle
Le Chatelier's Principle states that if a change is imposed upon a chemical system at equilibrium, the equilibrium will respond in such a way as to undo, in part, the effect of the change imposed upon it.
In doing so, the system will re-establish equilibrium but with a new 'equilibrium position'. The changes that can be imposed on a chemical system include changes to the amount of a reactant or product, temperature, and pressure/volume. Each of these can be increased or decreased. Do not confuse the equilibrium constant with equilibrium position. Changing the amount of a reactant or product, or the pressure/volume, will not change the value of the equilibrium constant, but changing the temperature will.

Molar Volume
This is the volume of 1 mole of any gas. Its value depends upon the temperature and pressure at which it is measured. The Molar Volume at s.t.p. (273 K, 1 atm.) is 22.4 dm3. The Molar Volume at room temperature and pressure (298 K, 1 atm.) is 24 dm3.

Molecular Formula
The Molecular Formula shows the exact number of each type of atom in a molecule of a substance. For example, C2H6 is the 'molecular formula' of ethane. One molecule of ethane contains two atoms of carbon and six atoms of hydrogen. Its 'empirical formula' is CH3.

Molecule
A Molecule is a group of two or more atoms covalently bonded together to form a particle that is a separate entity. Molecules do not carry an overall charge and therefore do not conduct electricity. A molecule consisting of a small number of atoms is referred to as a small molecule or simple molecule, e.g. H2O, CO2. Typically these substances have low melting and boiling points. Boiling point increases with increasing relative molecular mass, Mr. Giant molecular structures have covalent bonding extending throughout. Examples are diamond and silicon dioxide. These are very hard substances with high melting and boiling points.

Nucleophile (Nucleophilic reagent)
Nucleophilic means 'neucleus-loving'. A Nucleophile is a chemical species that is electron-rich. It will seek out an electron-deficient site in an organic molecule.
Examples of nucleophiles are: H2O, NH3, OH-, Cl-, Br-, CN-.
The diagram below illustrates a nucleophilic subsitution reaction in which a primary halogenoalkane is refluxed with dilute aqueous potassium hydroxide solution.

pH
pH is related to hydrogen ion concentration. Since the H+ ion concentration in solution is often small, the concentration is generally expressed as the logarithm of its reciprocal, which is called a pH value. Therefore, pH is defined as
pH = -log[H+]
For a ten times increase in H+ ion concentration there is a decrease in the pH value of one unit.
Given the pH of a solution, its H+ concentration can be found:
[H+(aq)] = antilog -pH
or
[H+(aq)] = 10-pH

Relative Atomic Mass, Ar
To consider the masses of atoms measured in grams, for example, would be to deal inconveniently with extremely small numbers. Rather, the mass of an atom is compared with that of an atom of carbon-12. The relative atomic mass of carbon-12 is taken to be 12. Relative masses have no units because they have cancelled in their calculation.
Some elements have isotopes. In calculating the relative atomic mass of an element with isotopes, the relative mass and proportion of each is taken into account. For example, naturally occurring chlorine consists of atoms of relative isotopic masses 35 (75%) and 37 (25%). Its relative atomic mass is 35.5.
Ar = (75/100 x 35) + (25/100 x 37) = 35.5
The relative masses of atoms are measured using an instrument called a mass spectrometer, invented by the English physicist Francis William Aston (1877-1945) when he was working in Cambridge with J. J. Thomson. It was in his use of this instrument that the existence of isotopes of elements was discovered. Aston eventually discovered many of the naturally occurring isotopes of non-radioactive elements. He was awarded the Nobel Prize for Chemistry in 1922.
Briefly, the mass spectrometer works by bombarding gaseous atoms with fast-moving electrons which knock out an electron from the atom. The cations formed are brought down on to a detector in turn according to their mass. The instrument provides a measure of the relative mass (compared to 12C) and the relative number of each isotope.
The diagrams below represent the mass spectrum of naturally occurring chlorine.
The above right spectrum has been represented so that the most abundant isotope has a relative abundance of 100%, with the other mass peaks scaled in relation to this. The relative atomic mass of chlorine is now calculated as shown below:
Ar = (100/133 x 35) + (33/133 x 37) = 35.5

Relative Molecular Mass, Mr
To consider the masses of molecules measured in grams, for example, would be to deal inconveniently with extremely small numbers. Rather, the mass of a molecule is compared with that of an atom of carbon-12. The relative atomic mass of carbon-12 is taken to be 12. Relative masses have no units because they have cancelled in their calculation. A relative molecular mass can be calculated easily by adding together the relative atomic masses of the constituent atoms. For example, ethanol, CH3CH2OH, has a Mr of 46.
Some care is needed. The term 'relative molecular mass' is sometimes used to refer to ionic compounds such as sodium chloride, NaCl. Such compounds are not molecular, and more accurately the term 'relative formula mass' can be used.
The mass spectrometer is also used to measure relative molecular masses (see relative atomic mass). The molecular ions formed in the instrument can often fragment, and it is from the the relative masses and abundances of these fragments that information about molecular structure can be deduced.

Room Temperature and Pressure
Room Temperature and Pressure is 298 K and 1 atm. pressure (101325 Nm-2).

Solubility Product
The Solubility Product, Ks, is the equilibrium constant for the equilibrium that exists between a slightly soluble salt and its ions in a saturated solution. For example,
AgCl(s) = Ag+(aq) + Cl-(aq)
Ks = [Ag+(aq)][Cl-(aq)] Ks = 1.74 x 10-10 mol2 dm-6
Solubility product and 'solubility' are different and must not be confused. Unlike the solubility of a substance, the solubility product is independent of what else is dissolved in solution. The solubility of a substance can be calculated from its solubility product.

Stability Constant
The Stability Constant, Kstab, is the equilibrium constant for the equilibrium that exists between a transition metal ion surrounded by water molecule ligands and the same transition metal ion surrounded by ligands of another kind in a ligand displacement reaction.
The individual ligands are displaced stepwise, and an equilibrium expression could be written for each step, but it is more common to write an overall expression for the overall ligand displacement reaction. Like all equilibrium constants, stability constants vary with temperature. For example,
[Cu(H2O)6]2+(aq) + 4NH3(aq)          [Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l)
The value is often expressed as its logarithm, log Kstab, making the numbers easier to handle. The larger the value of log Kstab, the more powerfully stabilising are the ligands of the transition metal ion. Since for a given transition metal ion, the different types of ligands always displace water molecule ligands in the equilibrium reaction, the numerical values can be used to decide upon the order in which the ligands are stabilising of a particular transition metal ion. The higher the value of the stability constant, the more stable the complex.

Standard Enthalpy of Combustion
The Standard Enthalpy of Combustion, DH°c, 298, is the heat evolved when 1 mole of a substance burns completely in oxygen, measured under standard conditions of temperature and pressure. For example,
CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l) DH°c, 298 = -890 kJ mol-1

Standard Enthalpy of Formation
The Standard Enthalpy of Formation, DH°f, 298, is the heat change that takes place when 1 mole of a substance is formed from its elements in their standard states, measured under standard conditions of temperature and pressure. For example,
C(s) + 2H2(g) ® CH4(g) DH°f, 298 = -75 kJ mol-1
According to this definition, the enthalpy of formation of an element in its standard state is zero.
There is a very useful equation, which follows from the definition for enthalpy of formation. This can be used to calculate heat changes, DH, when the enthalpies of formation of the reactants and products are given. It is
DH = SHf(products) - SHf(reactants).

Standard Enthalpy of Hydration
The Standard Enthalpy of Hydration, DH°hyd, 298, is the heat evolved when 1 mole of gaseous ions become hydrated (surrounded by water molecules), measured under standard conditions.
Al3+(g) + aq ® Al3+(aq) DH°hyd, 298 = -4613 kJ mol-1
The higher the charge on the ions and the smaller their size, the more exothermic the hydration enthalpy.
When the solvent is other than water the process is referred to as solvation, with the term enthalpy of solvation, DH°solv, 298.

Standard Enthalpy of Ionisation
The first ionisation enthalpy, DH°IE, 298, is the minimum energy required to remove completely the outermost electron from a 'gaseous atom' in its 'ground state', measured under standard conditions. The value is usually expressed in kJ mol-1.
K(g) ® K+(g) + e- DH°IE, 298 = +425 kJ mol-1
To remove a second electron is referred to as the second ionisation enthalpy, and so on.
Ionisation Enthalpies are illustrated in the diagram below:

Standard Enthalpy of Neutralisation
The Standard Enthalpy of Neutralisation, DH°neut, 298, is the heat evolved when 1 mole of water forms from the reaction between an acid and an alkali, measured under standard conditions of temperature and pressure. Since the essential reaction is H+ ions and OH- ions reacting to form water, the value is approximately -57.2 kJ mol-1.
H+(aq) + OH-(aq) ® H2O(l) DH°neut, 298 = -57.2 kJ mol-1
Where the acid is weak, for example ethanoic acid, the value will be slightly less exothermic owing to the absorption of some energy in bringing about ionisation of the acid molecules.

Standard Enthalpy of Solution
The Standard Enthalpy of Solution, DH°sol, 298, is the heat change which takes place when one mole of a solute is completely dissolved in a solvent to form a solution of concentration 1 mol dm-3, measured under standard conditions.
Enthalpy of Solution can be measured experimentally. It can also be calculated; it is the sum of two imaginary steps: the reverse of the lattice enthalpy plus the sum of the hydration enthalpies of the ions.
DH°sol, 298® -DH°LE, 298 + (DH°hyd, 298 (anion) + DH°hyd, 298 (cation))
This relationship is illustrated in the enthalpy diagram below for sodium chloride:
DH°sol, 298 = --776 + -771 = +5 kJ mol-1

Standard Lattice Enthalpy
The Standard Lattice Enthalpy, DH°LE, 298, is the heat evolved when one mole of a crystalline ionic compound forms from its gaseous ions, measured under standard conditions.
Mg2+(g) + 2Br-(g) ® MgBr2(s) DH°LE, 298 = -2440kJ mol-1
The magnitude of the lattice enthalpy depends on how closely the ions pack together in the crystal structure. The more closely they pack the more exothermic the lattice enthalpy. The higher the charge on the ions and the smaller their size, the more closely they will pack in the crystal structure.

Standard Temperature and Pressure
Standard Temperature and Pressure, S.T.P., is 273 K and 1 atm. pressure (101325 Nm-2).

Structural Formula
The Structural Formula shows how the atoms which make up a molecule are bonded together. For example, given the formula C4H10, there are two structural formulae relating to this molecular formula. These are CH3CH2CH2CH3 and CH3CHCH3CH3.

The Mole
A Mole is the mass of a substance containing the Avogadro Constant of particles. For example, 23g of sodium contain the Avogadro Constant of atoms, and the mass of this is 1 mole. 17g of ammonia contain the Avogadro Constant of NH3 molecules.
One Mole is the amount of any substance that contains as many particles as there are carbon atoms in 12 g of carbon-12. Some care is needed. 58.5 g of sodium chloride is the mass of one mole of this substance. However, it is ionic, and 1 mole consists of 1 mole of Na+ ions and 1 mole of Cl- ions. This is two moles of ions in total.

Transition Metal
A Transition Metal is defined as an element that forms at least one ion with a partially filled set of d electron atomic orbitals. In the first row of the d block only the eight elements from titanium to copper are classed as transition metals. Scandium forms only the Sc3+ ion with no 3d electrons; zinc forms only the Zn2+ with a full set of 3d electrons.
The electronic structures of Chromium, [Ar]4s13d5, and Copper, [Ar]4s13d10, are exceptional. Although the 4s atomic orbital fills with electrons before the 3d atomic orbitals, the 4s electrons are the first to be removed in ion formation. The Cu+ could be referred to as non-transitional.