Topic 4 States of Matter

4 States of Matter

4.1 Gases

Candidates should be able to:

(a) explain the pressure and behaviour of ideal gas using the kinetic theory;
(b) explain qualitatively, in terms of molecular size and intermolecular forces, the conditions necessary for a gas approaching the ideal behaviour;
(c) define Boyle’s law, Charles’ law and Avogadro’s law;
(d) apply the pV nRT equation in calculations, including the determination of the relative molecular mass, Mr;
(e) define Dalton’s law, and use it to calculate the partial pressure of a gas and its composition;
(f) explain the limitation of ideality at very high pressures and very low temperatures.

4.2 Liquids

Candidates should be able to:

(a) describe the kinetic concept of the liquid state;
(b) describe the melting of solid to liquid, vaporisation and vapour pressure using simple kinetic theory;
(c) define the boiling point and freezing point of liquids.

4.3 Solids

Candidates should be able to:

(a) describe qualitatively the lattice structure of a crystalline solid which is:
(i) ionic, as in sodium chloride,
(ii) simple molecular, as in iodine,
(iii) giant molecular, as in graphite, diamond and silicon(IV) oxide,
(iv) metallic, as in copper;

(b) describe the allotropes of carbon (graphite, diamond and fullerenes), and their uses.

4.4 Phase diagrams

Candidates should be able to:

(a) sketch the phase diagram for water and carbon dioxide, and explain the anomalous behaviour of water;
(b) explain phase diagrams as graphical plots of experimentally determined results;
(c) interpret phase diagrams as curves describing the conditions of equilibrium between phases and as regions representing single phases;
(d) predict how a phase may change with changes in temperature and pressure;
(e) discuss vaporisation, boiling, sublimation, freezing, melting, triple and critical points of H2O and CO2;
(f) explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent, and hence, on its melting point and boiling point (colligative properties);
(g) state the uses of dry ice.

4 States of Matter

4.1 Gases

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Candidates should be able to:

(a) explain the pressure and behaviour of ideal gas using the kinetic theory;
(b) explain qualitatively, in terms of molecular size and intermolecular forces, the conditions necessary for a gas approaching the ideal behaviour;
(c) define Boyle’s law, Charles’ law and Avogadro’s law;
(d) apply the pV nRT equation in calculations, including the determination of the relative molecular mass, Mr;
(e) define Dalton’s law, and use it to calculate the partial pressure of a gas and its composition;
(f) explain the limitation of ideality at very high pressures and very low temperatures.

 

Candidates should be able to:

(a) explain the pressure and behaviour of ideal gas using the kinetic theory;

(b) explain qualitatively, in terms of molecular size and intermolecular forces, the conditions necessary for a gas approaching the ideal behaviour;

(c) define Boyle’s law,

 

Charles’ law

and Avogadro’s law;

(d) apply the pV nRT equation in calculations, including the determination of the relative molecular mass, Mr;

(e) define Dalton’s law, and use it to calculate the partial pressure of a gas and its composition;

(f) explain the limitation of ideality at very high pressures and very low temperatures.

4.2 Liquids

Back

Candidates should be able to:

(a) describe the kinetic concept of the liquid state;
(b) describe the melting of solid to liquid, vaporisation and vapour pressure using simple kinetic theory;
(c) define the boiling point and freezing point of liquids.

 

Candidates should be able to:
(a) describe the kinetic concept of the liquid state;

(b) describe the melting of solid to liquid, vaporisation and vapour pressure using simple kinetic theory;

(c) define the boiling point and freezing point of liquids.

4.3 Solids

Back

Candidates should be able to:

(a) describe qualitatively the lattice structure of a crystalline solid which is:
(i) ionic, as in sodium chloride,
(ii) simple molecular, as in iodine,
(iii) giant molecular, as in graphite, diamond and silicon(IV) oxide,
(iv) metallic, as in copper;
(b) describe the allotropes of carbon (graphite, diamond and fullerenes), and their uses.

Candidates should be able to:

(a) describe qualitatively the lattice structure of a crystalline solid which is:

(i) ionic, as in sodium chloride,

(ii) simple molecular, as in iodine,

(iii) giant molecular, as in graphite, diamond and silicon(IV) oxide,

(iv) metallic, as in copper;

(b) describe the allotropes of carbon (graphite, diamond and fullerenes), and their uses.

4.4 Phase diagrams

Back

Candidates should be able to:

(a) sketch the phase diagram for water and carbon dioxide, and explain the anomalous behaviour of water;
(b) explain phase diagrams as graphical plots of experimentally determined results;
(c) interpret phase diagrams as curves describing the conditions of equilibrium between phases and as regions representing single phases;
(d) predict how a phase may change with changes in temperature and pressure;
(e) discuss vaporisation, boiling, sublimation, freezing, melting, triple and critical points of H2O and CO2;
(f) explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent, and hence, on its melting point and boiling point (colligative properties);
(g) state the uses of dry ice.

 

Candidates should be able to:
(a) sketch the phase diagram for water and carbon dioxide, and explain the anomalous behaviour of water;

A phase diagram in physical chemistry, engineering, mineralogy, and materials science is a type of chart used to show conditions (pressure, temperature, volume, etc.) at which thermodynamically distinct phases occur and coexist at equilibrium.

Phase diagram for water and carbon dioxide

Triple points mark conditions at which three different phases can coexist. For example, the water phase diagram has a triple point corresponding to the single temperature and pressure at which solid, liquid, and gaseous water can coexist in a stable equilibrium (273.16 K and a partial vapor pressure of 611.657 Pa).

Convert from Pascal to Atmosphere

Phase diagram for carbon diaoxide

The only thing special about this phase diagram is the position of the triple point which is well above atmospheric pressure. It is impossible to get any liquid carbon dioxide at pressures less than 5.11 atmospheres. That means that at 1 atmosphere pressure, carbon dioxide will sublime at a temperature of -78°C.

(b) explain phase diagrams as graphical plots of experimentally determined results;

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(c) interpret phase diagrams as curves describing the conditions of equilibrium between phases and as regions representing single phases;

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(d) predict how a phase may change with changes in temperature and pressure;

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(e) discuss vaporisation, boiling, sublimation, freezing, melting, triple and critical points of H2O and CO2;

 

Vaporization

(or vapourisation) of an element or compound is a phase transition from the liquid phase to vapor. There are two types of vaporization: evaporation and boiling. Evaporation is a surface phenomenon, whereas boiling is a bulk phenomenon.

Boiling,

Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point, the temperature at which the vapour pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding atmosphere.

Sublimation of dry ice

Sublimation,

Sublimation is the phase transition of a substance directly from the solid to the gas phase without passing through the intermediate liquid phase. Sublimation is an endothermic process that occurs at temperatures and pressures below a substance’s triple point in its phase diagram.

Freezing,

Freezing, or solidification, is a phase transition in which a liquid turns into a solid when its temperature is lowered below its freezing point. For most substances, the melting and freezing points are the same temperature; however, certain substances possess differing solid–liquid transition temperatures.

Melting,

Melting, or fusion, is a physical process that results in the phase transition of a substance from a solid to a liquid. This occurs when the internal energy of the solid increases, typically by the application of heat or pressure, which increases the substance’s temperature to the melting point.

Triple point,

all three phases (solid, liquid and gas) can coexist simultaneously in the equilibrium

Critical point,

highest temperature and pressure at which a liquid may be observed.

(f) explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent, and hence, on its melting point and boiling point (colligative properties);

 

explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent,

and hence, on its melting point and boiling point ;

Because the change in vapor pressure is a colligative property, which depends only on the relative number of solute and solvent particles, the changes in the boiling point and the melting point of the solvent are also colligative properties.

Notes for elevation of boiling point

The freezing point of a solution (solute + solvent) is always lower than that of a pure solvent.

Freezing-point depression describes the process in which adding a solute to a solvent decreases the freezing point of the solvent.

Examples include salt in water, alcohol in water, or the mixing of two solids such as impurities in a finely powdered drug. In the last case, the added compound is the solute, and the original solid is thought of as the solvent. The resulting solution or solid-solid mixture has a lower freezing point than the pure solvent or solid. This phenomenon is what causes sea water, (a mixture of salt (and other things) in water) to remain liquid at temperatures below 0 °C (32 °F), the freezing point of pure water.

Notes for freezing point depressions

(g) state the uses of dry ice.

Dry ice, sometimes referred to as “cardice” (chiefly by British chemists), is the solid form of carbon dioxide. It is used primarily as a cooling agent. Its advantages include lower temperature than that of water ice and not leaving any residue (other than incidental frost from moisture in the atmosphere).

 

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