Chemistry-ch07_condensed Phases Liquids And Solids
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7Condensed phases: liquids and solids
7.1 Liquids • In a liquid, intermolecular forces are strong enough to confine the molecules to a specific volume • Molecules are able to move freely within a liquid • Liquids are fluid • Liquids cannot expand or contract significantly
7.1 Liquids • Properties of liquids – Surface tension • Measure of the resistance of a liquid to an increase in its surface area • There is a net attractive force on molecules at the surface that pulls them towards the interior of the liquid • Cohesive forces attract molecules in the liquid to one another • Adhesive forces attract molecules in the liquid to the walls of the container
7.1 Liquids
7.1 Liquids – Capillary action • The upward movement of water against the downward force of gravity
7.1 Liquids – Viscosity • A liquid’s resistance to flow • The greater the viscosity, the more slowly the liquid pours • A measure of how easily molecules slide by one another • This is affected by a combination of molecular shape and the strength of the intermolecular forces • Viscosity is affected by temperature
7.1 Liquids • Vapour pressure – The distribution of molecular energies guarantees that some of the molecules in any liquid have enough kinetic energy to overcome the intermolecular forces that confine the liquid – Whenever a liquid has an exposed surface, some of its molecules will escape into the vapour phase
7.1 Liquids – The number of molecules of a liquid that have enough energy to escape into the vapour phase depends on: • The strength of intermolecular forces • The temperature
7.1 Liquids – Vapour pressure is the pressure at which dynamic equilibrium is achieved in a closed container
7.2 Solids • Magnitudes of forces – Forces in solids range from very small to extremely large – Ions, atoms or molecules in solids can be bound together by various attractive forces: • • • •
Intermolecular forces Metallic bonding Covalent bonds Ionic interactions
7.2 Solids
7.2 Solids • Molecular solids – Aggregates of molecules bound together by intermolecular forces – The forces can be dispersion, dipolar, hydrogen bonding or a combination – Many larger molecules have sufficient dispersion forces to exist as solids at room temperature – Eg naphthalene
7.2 Solids
7.2 Solids – Dimethyl oxalate is an example of a molecular solid with dipolar interactions
– Crystals of benzoic acid contain pairs of molecules held together head to head by hydrogen bonds. These pairs then stack in planes which are held together by dispersion forces
7.2 Solids
7.2 Solids • Network solids – Have very high melting points – Held together by covalent bonds which are much harder to break than intermolecular forces – Bonding patterns determine the properties of network solids – Covalent bonds make network solids extremely durable
7.2 Solids
7.2 Solids • Metallic solids – Derives primarily from electrons in highly delocalised valence orbitals – Consists of a regular array of metal atoms embedded in a ‘sea’ of mobile valence electrons
7.2 Solids – Wide range of melting points – Strength of metallic bonding is variable – Metals are ductile and malleable
7.2 Solids
7.2 Solids • Ionic solids – Contain cations and anions strongly attracted to each other by electrostatic forces – Their stoichiometries are determined by the charges carried by the positive and negative ions – Many ionic solids contain metal cations and polyatomic anions
7.3 Phase changes • There are three different phases of matter: gas, liquid and solid • A phase change is the transition of a substance from one phase to another • Phase changes depend on temperature, pressure, and the magnitudes of intermolecular and bonding forces
7.3 Phase changes
7.3 Phase changes • As a phase change occurs, temperature remains constant • A substance must completely change phase before the temperature of the system can increase (or decrease) • Phase changes require that energy (usually in the form of heat) be either supplied to or removed from the substance undergoing the phase change
7.3 Phase changes • Molar enthalpy of vaporisation, ∆ Hvap : the heat needed to vaporise one mole of a substance at its normal boiling point. • Molar enthalpy of fusion, ∆ Hfus : the heat needed to melt one mole of a substance at its normal melting point. • Molar enthalpy of sublimation, ∆ Hsub : the heat needed to vaporise one mole of a substance from the solid phase (skips the liquid phase)
7.3 Phase changes
7.3 Phase changes • The effect of pressure is mainly seen for phase transitions involving gases • A gas at constant temperature can be liquefied by increasing the pressure • The pressure at which a gas liquefies at a specific temperature is known as the condensation point
7.3 Phase changes – Supercritical fluids • A supercritical fluid is a fluid that has certain properties of both liquids and gases • The critical temperature is the temperature at which the densities in the gas phase and liquid phase become equal and we can no longer distinguish between phases or observe a phase boundary • The associated pressure is called the critical pressure • The combination of these is the critical point
7.3 Phase changes
7.3 Phase changes • Phase diagrams – A way of showing the phase behaviour of a substance as a function of temperature and pressure
7.3 Phase changes – Characteristic features of phase diagrams: • Boundary lines between phases separate the regions where each phase is thermodynamically stable • Movement across a boundary line corresponds to a phase change • At any point along a boundary line, the two phases on either side of the line coexist in a state of dynamic equilibrium
7.3 Phase changes • Three boundary lines meet at a single point called a triple point • Above the temperature specified by the critical point, the gas cannot be liquefied under any pressure • What happens as temperature changes at constant pressure can be determined by drawing a horizontal line at the appropriate pressure on the phase diagram • The temperature for conversion between the gas phase and a condensed phase depends strongly on pressure
7.3 Phase changes • The melting temperature is almost independent of pressure, and making the boundary line between solid and liquid nearly vertical • The solid-gas boundary line extrapolates to p = 0 Pa and T = 0 K
7.3 Phase changes – All phase diagrams of pure substances share the 10 common features listed previously – The detailed appearance of a phase diagram is different for each substance, as determined by the strength of the interactions between its constituents – Phase diagrams are constructed by measuring the temperatures and pressures and which phase changes occur
7.3 Phase changes
7.3 Phase changes – Many substances have more than one solid phase
7.4 Order in solids • Molecules in gases and liquids can move relatively freely • The atoms, molecules or ions in a solid are in fixed positions • Their motions are restricted to vibrations about these fixed positions
7.4 Order in solids • Close-packed structures – Assume we can approximate a metal atom with a sphere – A hexagonal arrangement is able to fit more spheres into the same area – This hexagonal arrangement of identical spheres is the most dense packing you can achieve in one layer of spheres
7.4 Order in solids
7.4 Order in solids – When a second layer is added, to achieve the most compact arrangement, each sphere in layer B sits in one of the ‘dimples’ between a trio of spheres in layer A
7.4 Order in solids – As additional layers of spheres are added, the layer B eventually looks identical to the layer A, except offset slightly to allow the spheres to nestle in the dimples formed by the layer below
7.4 Order in solids – If spheres in the third layer lie directly above the first, the structure is called hexagonal close-packed (hcp) – If spheres in the third layer are offset from both of the layers, the structure is called cubic close-packed (ccp) – The two structure types are called closepacked structures because they achieve maximum space filling for identical spheres
7.4 Order in solids
7.4 Order in solids
7.4 Order in solids • The crystal lattice and the unit cell – A lattice is a pattern of points – Every individual point is a lattice point
– The unit cell is the smallest unique repeating unit of a lattice
7.4 Order in solids – The same lattice can be used to describe many different designs or structures
7.4 Order in solids • Cubic structures – In a simple cubic crystal, layers of atoms stack one directly above another, so that all atoms lie along straight lines at right angles
7.4 Order in solids – Primitive cubic structure: layers of atoms stacked one directly above another, so that all atoms lie along straight lines at right angles – Body-centred cubic lattice (bcc): simple cube with one entire atom in the center of the cube (in the body) – Face-centred cubic (fcc): simple cube with atoms in the center of each face of the cube
7.4 Order in solids • Ionic solids – The packing in ionic crystals requires that ions of opposite charge alternate with one another to maximise attractions and minimise repulsions – The cations and anions are usually different sizes – Adopt a variety of structures that depend on the stoichiometry and relative sizes of the ions
7.4 Order in solids
– Many ionic solids have the same unit cell structure as the sodium chloride lattice
7.4 Order in solids – Many ionic compounds have stoichiometries that differ from 1:1
7.5 X-ray diffraction • Can be used to determine the arrangement of atoms, ions or molecules in a crystalline structure • When the crystal is exposed to X-rays, intense beams, diffracted because of constructive interference, appear only in specific directions • In other directions, no X-rays appear because of destructive interference
7.5 X-ray diffraction • The X-rays coming from the crystal are recorded and form a diffraction pattern
7.5 X-ray diffraction • The Bragg equation relates λ, θ and the distance between the planes of atoms, d: nλ = 2dsinθ
7.6 Amorphous solids • When a liquid is cooled slowly, it often solidifies as a crystalline solid • When solids form rapidly, their atoms, ions or molecules may become locked into positions other than those of a regular crystal • These materials are said to be amorphous, meaning ‘without form’
7.6 Amorphous solids • Do not diffract X-rays • Glass is an entire family of amorphous solids based on silica
7.7 Crystal imperfections • Many solid materials are crystalline but contain defects • Crystalline defects can profoundly alter the properties of a solid material • Doped semiconductors are solids into which impurity ‘defects’ are introduced • Some gemstones are crystals containing impurities giving them colour
7.7 Crystal imperfections – Substitutional impurities replace one atom with another – Interstitial impurities occupy the spaces between regular atoms
7.8 Modern ceramics • Ceramics are materials composed of inorganic components that have been heat-treated • Many manufactured ceramics are made from inorganic minerals such as clay, silica (sand) and other silicates • In recent times, advanced ceramics have been prepared by chemists
7.8 Modern ceramics • Properties of ceramics – Almost all ceramics are very hard – They have very high melting points – Are often used as refractories – Most ceramics do not conduct electricity – Generally contain metals with high positive oxidation states, combined with small nonmetals with high negative oxidation states
7.8 Modern ceramics
7.8 Modern ceramics • Applications of advanced ceramics – Thin ceramic films are used as antireflective coatings on optical surfaces – Partially stabilised zirconia is used to make portions of hip-joint replacements – Boron nitride powder is used in cosmetics – Silicon nitride is used to make engine components for diesel engines
7.8 Modern ceramics • High-temperature superconductors – A superconductor is a material which offers no resistance to the flow of electricity – They can be levigated by a magnetic field
Summary • Liquids are fluids, but cannot expand or contract significantly • Viscosity, surface tension and capillary action depend mostly on the strengths of intermolecular attractions within the liquid • Vapour pressure is the pressure at which the number of molecules escaping the liquid exactly matches the number being captured
Summary • Solids are rigid • Solids may be classified as molecular, metallic, network or ionic solids • A phase change occurs when a substance undergoes a transition from one phase to another • Temperatures and pressures at which equilibria can exist between phases are shown in a phase diagram
Summary • Approximating atoms as spheres allows us to determine their possible closepacked arrangements • The overall structure of any crystalline solid can be described in terms of a repeating three-dimensional array of lattice points, which is called a lattice • X-ray diffraction can be used to determine the structure of crystalline solids
Summary • Amorphous solids do not have a regular arrangement of atoms, ions or molecules • The most important example is glass • Crystalline defects can significantly alter the properties of a solid • Ceramics are composed of inorganic components that have been heattreated
7.1 Liquids • In a liquid, intermolecular forces are strong enough to confine the molecules to a specific volume • Molecules are able to move freely within a liquid • Liquids are fluid • Liquids cannot expand or contract significantly
7.1 Liquids • Properties of liquids – Surface tension • Measure of the resistance of a liquid to an increase in its surface area • There is a net attractive force on molecules at the surface that pulls them towards the interior of the liquid • Cohesive forces attract molecules in the liquid to one another • Adhesive forces attract molecules in the liquid to the walls of the container
7.1 Liquids
7.1 Liquids – Capillary action • The upward movement of water against the downward force of gravity
7.1 Liquids – Viscosity • A liquid’s resistance to flow • The greater the viscosity, the more slowly the liquid pours • A measure of how easily molecules slide by one another • This is affected by a combination of molecular shape and the strength of the intermolecular forces • Viscosity is affected by temperature
7.1 Liquids • Vapour pressure – The distribution of molecular energies guarantees that some of the molecules in any liquid have enough kinetic energy to overcome the intermolecular forces that confine the liquid – Whenever a liquid has an exposed surface, some of its molecules will escape into the vapour phase
7.1 Liquids – The number of molecules of a liquid that have enough energy to escape into the vapour phase depends on: • The strength of intermolecular forces • The temperature
7.1 Liquids – Vapour pressure is the pressure at which dynamic equilibrium is achieved in a closed container
7.2 Solids • Magnitudes of forces – Forces in solids range from very small to extremely large – Ions, atoms or molecules in solids can be bound together by various attractive forces: • • • •
Intermolecular forces Metallic bonding Covalent bonds Ionic interactions
7.2 Solids
7.2 Solids • Molecular solids – Aggregates of molecules bound together by intermolecular forces – The forces can be dispersion, dipolar, hydrogen bonding or a combination – Many larger molecules have sufficient dispersion forces to exist as solids at room temperature – Eg naphthalene
7.2 Solids
7.2 Solids – Dimethyl oxalate is an example of a molecular solid with dipolar interactions
– Crystals of benzoic acid contain pairs of molecules held together head to head by hydrogen bonds. These pairs then stack in planes which are held together by dispersion forces
7.2 Solids
7.2 Solids • Network solids – Have very high melting points – Held together by covalent bonds which are much harder to break than intermolecular forces – Bonding patterns determine the properties of network solids – Covalent bonds make network solids extremely durable
7.2 Solids
7.2 Solids • Metallic solids – Derives primarily from electrons in highly delocalised valence orbitals – Consists of a regular array of metal atoms embedded in a ‘sea’ of mobile valence electrons
7.2 Solids – Wide range of melting points – Strength of metallic bonding is variable – Metals are ductile and malleable
7.2 Solids
7.2 Solids • Ionic solids – Contain cations and anions strongly attracted to each other by electrostatic forces – Their stoichiometries are determined by the charges carried by the positive and negative ions – Many ionic solids contain metal cations and polyatomic anions
7.3 Phase changes • There are three different phases of matter: gas, liquid and solid • A phase change is the transition of a substance from one phase to another • Phase changes depend on temperature, pressure, and the magnitudes of intermolecular and bonding forces
7.3 Phase changes
7.3 Phase changes • As a phase change occurs, temperature remains constant • A substance must completely change phase before the temperature of the system can increase (or decrease) • Phase changes require that energy (usually in the form of heat) be either supplied to or removed from the substance undergoing the phase change
7.3 Phase changes • Molar enthalpy of vaporisation, ∆ Hvap : the heat needed to vaporise one mole of a substance at its normal boiling point. • Molar enthalpy of fusion, ∆ Hfus : the heat needed to melt one mole of a substance at its normal melting point. • Molar enthalpy of sublimation, ∆ Hsub : the heat needed to vaporise one mole of a substance from the solid phase (skips the liquid phase)
7.3 Phase changes
7.3 Phase changes • The effect of pressure is mainly seen for phase transitions involving gases • A gas at constant temperature can be liquefied by increasing the pressure • The pressure at which a gas liquefies at a specific temperature is known as the condensation point
7.3 Phase changes – Supercritical fluids • A supercritical fluid is a fluid that has certain properties of both liquids and gases • The critical temperature is the temperature at which the densities in the gas phase and liquid phase become equal and we can no longer distinguish between phases or observe a phase boundary • The associated pressure is called the critical pressure • The combination of these is the critical point
7.3 Phase changes
7.3 Phase changes • Phase diagrams – A way of showing the phase behaviour of a substance as a function of temperature and pressure
7.3 Phase changes – Characteristic features of phase diagrams: • Boundary lines between phases separate the regions where each phase is thermodynamically stable • Movement across a boundary line corresponds to a phase change • At any point along a boundary line, the two phases on either side of the line coexist in a state of dynamic equilibrium
7.3 Phase changes • Three boundary lines meet at a single point called a triple point • Above the temperature specified by the critical point, the gas cannot be liquefied under any pressure • What happens as temperature changes at constant pressure can be determined by drawing a horizontal line at the appropriate pressure on the phase diagram • The temperature for conversion between the gas phase and a condensed phase depends strongly on pressure
7.3 Phase changes • The melting temperature is almost independent of pressure, and making the boundary line between solid and liquid nearly vertical • The solid-gas boundary line extrapolates to p = 0 Pa and T = 0 K
7.3 Phase changes – All phase diagrams of pure substances share the 10 common features listed previously – The detailed appearance of a phase diagram is different for each substance, as determined by the strength of the interactions between its constituents – Phase diagrams are constructed by measuring the temperatures and pressures and which phase changes occur
7.3 Phase changes
7.3 Phase changes – Many substances have more than one solid phase
7.4 Order in solids • Molecules in gases and liquids can move relatively freely • The atoms, molecules or ions in a solid are in fixed positions • Their motions are restricted to vibrations about these fixed positions
7.4 Order in solids • Close-packed structures – Assume we can approximate a metal atom with a sphere – A hexagonal arrangement is able to fit more spheres into the same area – This hexagonal arrangement of identical spheres is the most dense packing you can achieve in one layer of spheres
7.4 Order in solids
7.4 Order in solids – When a second layer is added, to achieve the most compact arrangement, each sphere in layer B sits in one of the ‘dimples’ between a trio of spheres in layer A
7.4 Order in solids – As additional layers of spheres are added, the layer B eventually looks identical to the layer A, except offset slightly to allow the spheres to nestle in the dimples formed by the layer below
7.4 Order in solids – If spheres in the third layer lie directly above the first, the structure is called hexagonal close-packed (hcp) – If spheres in the third layer are offset from both of the layers, the structure is called cubic close-packed (ccp) – The two structure types are called closepacked structures because they achieve maximum space filling for identical spheres
7.4 Order in solids
7.4 Order in solids
7.4 Order in solids • The crystal lattice and the unit cell – A lattice is a pattern of points – Every individual point is a lattice point
– The unit cell is the smallest unique repeating unit of a lattice
7.4 Order in solids – The same lattice can be used to describe many different designs or structures
7.4 Order in solids • Cubic structures – In a simple cubic crystal, layers of atoms stack one directly above another, so that all atoms lie along straight lines at right angles
7.4 Order in solids – Primitive cubic structure: layers of atoms stacked one directly above another, so that all atoms lie along straight lines at right angles – Body-centred cubic lattice (bcc): simple cube with one entire atom in the center of the cube (in the body) – Face-centred cubic (fcc): simple cube with atoms in the center of each face of the cube
7.4 Order in solids • Ionic solids – The packing in ionic crystals requires that ions of opposite charge alternate with one another to maximise attractions and minimise repulsions – The cations and anions are usually different sizes – Adopt a variety of structures that depend on the stoichiometry and relative sizes of the ions
7.4 Order in solids
– Many ionic solids have the same unit cell structure as the sodium chloride lattice
7.4 Order in solids – Many ionic compounds have stoichiometries that differ from 1:1
7.5 X-ray diffraction • Can be used to determine the arrangement of atoms, ions or molecules in a crystalline structure • When the crystal is exposed to X-rays, intense beams, diffracted because of constructive interference, appear only in specific directions • In other directions, no X-rays appear because of destructive interference
7.5 X-ray diffraction • The X-rays coming from the crystal are recorded and form a diffraction pattern
7.5 X-ray diffraction • The Bragg equation relates λ, θ and the distance between the planes of atoms, d: nλ = 2dsinθ
7.6 Amorphous solids • When a liquid is cooled slowly, it often solidifies as a crystalline solid • When solids form rapidly, their atoms, ions or molecules may become locked into positions other than those of a regular crystal • These materials are said to be amorphous, meaning ‘without form’
7.6 Amorphous solids • Do not diffract X-rays • Glass is an entire family of amorphous solids based on silica
7.7 Crystal imperfections • Many solid materials are crystalline but contain defects • Crystalline defects can profoundly alter the properties of a solid material • Doped semiconductors are solids into which impurity ‘defects’ are introduced • Some gemstones are crystals containing impurities giving them colour
7.7 Crystal imperfections – Substitutional impurities replace one atom with another – Interstitial impurities occupy the spaces between regular atoms
7.8 Modern ceramics • Ceramics are materials composed of inorganic components that have been heat-treated • Many manufactured ceramics are made from inorganic minerals such as clay, silica (sand) and other silicates • In recent times, advanced ceramics have been prepared by chemists
7.8 Modern ceramics • Properties of ceramics – Almost all ceramics are very hard – They have very high melting points – Are often used as refractories – Most ceramics do not conduct electricity – Generally contain metals with high positive oxidation states, combined with small nonmetals with high negative oxidation states
7.8 Modern ceramics
7.8 Modern ceramics • Applications of advanced ceramics – Thin ceramic films are used as antireflective coatings on optical surfaces – Partially stabilised zirconia is used to make portions of hip-joint replacements – Boron nitride powder is used in cosmetics – Silicon nitride is used to make engine components for diesel engines
7.8 Modern ceramics • High-temperature superconductors – A superconductor is a material which offers no resistance to the flow of electricity – They can be levigated by a magnetic field
Summary • Liquids are fluids, but cannot expand or contract significantly • Viscosity, surface tension and capillary action depend mostly on the strengths of intermolecular attractions within the liquid • Vapour pressure is the pressure at which the number of molecules escaping the liquid exactly matches the number being captured
Summary • Solids are rigid • Solids may be classified as molecular, metallic, network or ionic solids • A phase change occurs when a substance undergoes a transition from one phase to another • Temperatures and pressures at which equilibria can exist between phases are shown in a phase diagram
Summary • Approximating atoms as spheres allows us to determine their possible closepacked arrangements • The overall structure of any crystalline solid can be described in terms of a repeating three-dimensional array of lattice points, which is called a lattice • X-ray diffraction can be used to determine the structure of crystalline solids
Summary • Amorphous solids do not have a regular arrangement of atoms, ions or molecules • The most important example is glass • Crystalline defects can significantly alter the properties of a solid • Ceramics are composed of inorganic components that have been heattreated
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