Liquids and Solids Gas Liquid and Solid Gas

  • Slides: 57
Download presentation
Liquids and Solids

Liquids and Solids

Gas, Liquid, and Solid Gas Liquid Solid Highly Compressible Slightly Compressible Very slightly compressible

Gas, Liquid, and Solid Gas Liquid Solid Highly Compressible Slightly Compressible Very slightly compressible Low Density High Density Fills container completely Does not expand to fill container Rigidly retains its volume Assumes shape of container Retains its own shape Rapid diffusion Slow diffusion Extremely slow diffusion, surface only High expansion on heating Low expansion on heating Total Disorder Ordered arrangement Disordered

Intermolecular Forces § Forces of attraction between neighboring particles § Much weaker than bonding

Intermolecular Forces § Forces of attraction between neighboring particles § Much weaker than bonding forces § Responsible for state of matter and some physical properties § e. g. , The stronger the attractive forces, the higher the melting and boiling points § Also involved in change of state

Three Types § London Dispersion forces § Dipole-dipole forces § Hydrogen bonds

Three Types § London Dispersion forces § Dipole-dipole forces § Hydrogen bonds

London Dispersion Forces § The motion of electrons can create an instantaneous dipole moment

London Dispersion Forces § The motion of electrons can create an instantaneous dipole moment on an atom § For example, if at any one time both of a helium atom’s electrons are on the same side of the atom at the same time § A temporary dipole on one atom can cause, or induce, a temporary dipole on an adjacent atom

London Dispersion Forces § These forces are significant only when molecules are very close

London Dispersion Forces § These forces are significant only when molecules are very close together, as in a compressed gas § These forces are found only in nonpolar compounds § Molecules and atoms will lose their spherical shape

 • More compact molecules have smaller surface areas, weaker London dispersion forces, and

• More compact molecules have smaller surface areas, weaker London dispersion forces, and lower boiling points. • Flatter, less compact molecules have larger surface areas, stronger London dispersion forces, and higher boiling points.

Dipole-Dipole Forces § Polar molecules have a positive end a negative end § Dipole-dipole

Dipole-Dipole Forces § Polar molecules have a positive end a negative end § Dipole-dipole forces occur when the positive end of one molecule is attracted to the negative end of another § Only effective when polar molecules are very close together § For molecules of about the same size, dipole forces increase with increasing polarity

If two neutral molecules, each having a permanent dipole moment, come together such that

If two neutral molecules, each having a permanent dipole moment, come together such that their oppositely charged ends align, they will be attracted to each other.

Hydrogen Bonds § Type of dipole-dipole force § Not a true bond! § Occurs

Hydrogen Bonds § Type of dipole-dipole force § Not a true bond! § Occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons (e. g. , N, O & F) § The hydrogen in one molecule will be attracted to the electronegative atom in another molecule

Hydrogen Bonds § Hydrogen has no inner core of electrons, so a dipole will

Hydrogen Bonds § Hydrogen has no inner core of electrons, so a dipole will expose its concentrated charge on the proton, its nucleus. § Hydrogen can approach an electronegative atom very closely and interact strongly with it.

 • Electron shell around a hydrogen atom is rather thin, giving the hydrogen

• Electron shell around a hydrogen atom is rather thin, giving the hydrogen atom a small positive charge. • Electron shell round an oxygen atom is quite thick, and so oxygen carries an extra bit of negative charge. • These opposite charges attract, although quite weakly. • This weak force is called a hydrogen bond. The hydrogen atoms of one water molecule stick to the oxygen atoms of nearby water molecules.

Properties of Liquids § Have much greater densities than their vapors § Only slightly

Properties of Liquids § Have much greater densities than their vapors § Only slightly compressible; not a discernable difference when compressed § Fluidity: ability to flow § Liquids can diffuse through one another, but at a much slower rate than gases

Properties of Liquids § Viscosity: resistance to flow § Determined by the type of

Properties of Liquids § Viscosity: resistance to flow § Determined by the type of intermolecular forces involved, the shape of the particle, and the temperature § The stronger the attractive forces, the higher the viscosity § The larger the particles, the higher the viscosity § Increases as temp decreases

Properties of Liquids § Surface Tension: the imbalance of forces at the surface of

Properties of Liquids § Surface Tension: the imbalance of forces at the surface of a liquid § The uneven forces make the surface behave as if it has a tight film stretched across it § The stronger the intermolecular forces, the higher the surface tension

Properties of Liquids § Surfactants: compounds that lower the surface tension of water §

Properties of Liquids § Surfactants: compounds that lower the surface tension of water § § Frequently added to detergents Capillary action: movement of a liquid through narrow spaces

Properties of Solids § Have extremely strong intermolecular forces in order for solids to

Properties of Solids § Have extremely strong intermolecular forces in order for solids to have definite shape and volume § Particle arrangement causes solids to almost always have higher densities than liquids § Ice is an exception: it expands when it freezes because of the way the particles arrange themselves during the freezing process

Properties of Solids § Particle arrangements cause different types of solids: § Crystalline solids

Properties of Solids § Particle arrangements cause different types of solids: § Crystalline solids § § Molecular solids Covalent network solids Ionic solids Metallic solids § Amorphous solids

Crystalline Solids § Has atoms, ions, or molecules arranged in an orderly, geometric, 3

Crystalline Solids § Has atoms, ions, or molecules arranged in an orderly, geometric, 3 -D structure § Individual pieces of a crystalline solid are called crystals § Smallest arrangement of connected points that can be repeated in 3 directions to form a lattice is called a unit cell § There are 7 different crystal systems based on shape

Molecular Solids § Held together by dispersion forces, dipole-dipole forces or hydrogen bonds §

Molecular Solids § Held together by dispersion forces, dipole-dipole forces or hydrogen bonds § NOT held together by genuine bonds (ionic and covalent) § Most are NOT solids at room temperature § Poor conductors of heat and electricity (don’t contain ions) § Examples are sucrose and ice

–Molecular such as sucrose or ice whose constituent particles are molecules held together by

–Molecular such as sucrose or ice whose constituent particles are molecules held together by the intermolecular forces.

Arrangement of molecules in liquid water Arrangement of molecules in ice

Arrangement of molecules in liquid water Arrangement of molecules in ice

Covalent Network Solids § Atoms that can form multiple covalent bonds § Form a

Covalent Network Solids § Atoms that can form multiple covalent bonds § Form a network of atoms that do not have a unit cell § Most allotropes exist in this form § Allotropes are forms of the same element that have different bonding patterns of arrangement § Examples include diamonds and graphite, quartz

Graphite Diamond

Graphite Diamond

Covalent network solids such as quartz where atoms are held together by 3 -D

Covalent network solids such as quartz where atoms are held together by 3 -D networks of covalent bonds. Here the hexagonal pattern of Si (violet) and O (red) atoms in structure matches the hexagonal crystal shape

Ionic Solids § Type of crystalline solid § Type and ratio of ions determine

Ionic Solids § Type of crystalline solid § Type and ratio of ions determine the structure of the lattice and the shape of the structure § The network of attractions that extend through an ionic compound gives these compounds their high melting points and hardness

Ionic Solids § Strong but brittle § When struck, cations and anions are shifted,

Ionic Solids § Strong but brittle § When struck, cations and anions are shifted, which causes repulsion that in turn shatter the crystal § Poor conductors of heat and electricity in solid form

 • Ionic solids are an orderly pattern of one ion, generally the anion,

• Ionic solids are an orderly pattern of one ion, generally the anion, with cations positioned in 'holes' between the anions • The occupation of these 'holes' depends on the formula of the ionic compound

Sodium chloride Cupric chloride

Sodium chloride Cupric chloride

Metallic Solids § Consist of positive metal ions surrounded by a sea of mobile

Metallic Solids § Consist of positive metal ions surrounded by a sea of mobile electrons § Mobile electrons make metals malleable, ductile, and good conductors of heat and electricity

 • A series of metals atoms that have all donated their valence electrons

• A series of metals atoms that have all donated their valence electrons to an electron cloud that permeates the structure • This electron cloud is referred to as an electron sea • Visualize the electron sea model as if it were a box of marbles that are surrounded by water. The marbles are the metal atoms and the water represents the electron sea.

 • The marbles can be pushed anywhere within the box and the water

• The marbles can be pushed anywhere within the box and the water will follow them, always surrounding the marbles. • This unique property, allows metallic bonds to be maintained when pushed and pulled in all sorts of ways. • As a result, they are malleable and ductile.

Gold Copper Silver

Gold Copper Silver

Amorphous Solids § Solid in which the particles are not arranged in a regular,

Amorphous Solids § Solid in which the particles are not arranged in a regular, repeating pattern, but still retain rigidity § Examples include glass, rubber, many plastics, tar and wax § Particles are trapped in a disordered arrangement that is characteristic of liquids

Phase Changes § Always involve a change in energy § Energy is needed either

Phase Changes § Always involve a change in energy § Energy is needed either to overcome or form attractive forces between particles

Melting and Freezing § Melting point/freezing point: temp at which solid and liquid forms

Melting and Freezing § Melting point/freezing point: temp at which solid and liquid forms exist in equilibrium § Melting is endothermic § Freezing is exothermic

Vaporization § The change of state from a liquid to a gas § Endothermic

Vaporization § The change of state from a liquid to a gas § Endothermic process § Two methods of vaporization: § Evaporation § Boiling

Evaporation § Occurs at the surface of a liquid § Occurs because molecules close

Evaporation § Occurs at the surface of a liquid § Occurs because molecules close to the surface have enough energy to overcome the attractions of neighboring molecules and escape § Slower molecules stay in the liquid state § Rate of evaporation increases as temp increases

Boiling § Occurs within the liquid § Boiling point: temp at which vapor pressure

Boiling § Occurs within the liquid § Boiling point: temp at which vapor pressure equals atmospheric pressure § If vapor pressure is less than atmospheric pressure, bubbles do not form

Condensation § Change of a gas to a liquid § Exothermic process § Molecules

Condensation § Change of a gas to a liquid § Exothermic process § Molecules of vapor can return to the liquid state by colliding with the liquid surface § The particles become trapped by the intermolecular attractions of the liquid

Sublimation and Deposition § Sublimation: solid goes directly to a gas without passing through

Sublimation and Deposition § Sublimation: solid goes directly to a gas without passing through the liquid phase § Deposition is the reverse process § Sublimation is endothermic § Deposition is exothermic

Heating Curves § Graphic illustrations of phase changes § Plot of temp of a

Heating Curves § Graphic illustrations of phase changes § Plot of temp of a sample as a function of time § Notice temp remains constant during phase changes while amount of energy varies

Heating Curve of Water A: Rise in temperature as ice absorbs heat. B: Absorption

Heating Curve of Water A: Rise in temperature as ice absorbs heat. B: Absorption of heat of fusion. C: Rise in temperature as liquid water absorbs heat. D: Water boils and absorbs heat of vaporization. E: Steam absorbs heat and thus increases its temperature. The above is an example of a heating curve. One could reverse the process, and obtain a cooling curve. The flat portions of such curves indicate the phase changes.

Phase Diagrams § Diagram that relates the states of a substance to temp and

Phase Diagrams § Diagram that relates the states of a substance to temp and pressure § State depends on temp and pressure § 2 states can exist simultaneously at certain temps and pressures § Triple point: the temp and pressure when all three states exist at the same time

 • TRIPLE POINT - The temperature and pressure at which the solid, liquid,

• TRIPLE POINT - The temperature and pressure at which the solid, liquid, and gas phases exist simultaneously. • CRITICAL POINT - The temperature above which a substance will always be a gas regardless of the pressure. • FREEZING POINT - The temperature at which the solid and liquid phases of a substance are in equilibrium at atmospheric pressure. • BOILING POINT - The temperature at which the vapor pressure of a liquid is equal to the pressure on the liquid. • Normal (Standard) Boiling Point - The temperature at which the vapor pressure of a liquid is equal to standard pressure (1. 00 atm = 760 mm. Hg = 760 torr = 101. 325 k. Pa) • NOTE – • The line between the solid and liquid phases is a curve of all the freezing/melting points of the substance. • The line between the liquid and gas phases is a curve of all the boiling points of the substance.