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Chapter 1

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Chapter 6

Chapter 7

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Chapter 12

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Chapter 14

Chapter 15

Chapter 16

Chapter 17

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Chapter 21

       

 

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CHAPTER 14 Study Guide ==> Here!
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CHAPTER 8/9 Study Guide on Bonding ==> Here!

   

  Other Books -- Please understand that these are to be used only for reference as they were given to me for only that purpose...

5 Steps to a 5 AP Chemistry
AP Chemistry For Dummies
Modern Chemistry
Chemistry Chang 10th Edition
Zumdahl Chemistry 8th



 

Chapter Study Guides

Chapters 8 & 9 - Bonding

* Understand that when forming chemical bonds atoms are attempting to form more stable electronic configurations
* Understand the concept of ionic bonding and the nature of the ionic bond
* Understand the concept of covalent bonding and nature of the covalent bond
* Be able to draw Lewis structures
* Understand the concept of resonance related to Lewis structures
* Understand the concept of formal charge related to Lewis structures
* Be able to predict the shape of, and bond angles in, simple molecules and ions using VSEPR theory
* Understand the concept of the dative (co-ordinate) bond related to Lewis structures
* Understand that ionic bonding and covalent bonding are at two ends of a sliding scale of bond type
* Understand the concept of electronegativity as it pertains to bonding theory
* Understand that polarization caused by small highly charged cations leads to ionic compounds exhibiting some covalent character
* Understand that differences in electronegativity in covalent molecules causes dipoles and some ionic character in covalent compounds
* Understand when molecules exhibit polarity
* Be able to predict the shapes of simple molecules and ions using Lewis structures
* Understand the occurrence, relative strength and nature of dipole-dipole interactions, London dispersion forces and hydrogen bonds
* Understand how solid structure influences properties
* Understand the nature of liquids
* Understand the nature of sigma and pi bonds
* Understand and be able to identify different types of orbital hybridization

Chapter 10 - Gases

* Be able to convert between different units of pressure
* Be able to convert between different units of temperature
* Recall and be able to use Boyle's law in calculations
* Recall and be able to use Charles' law in calculations
* Recall and be able to use Gay-Lussac's law in calculations
* Recall and be able to use Avogadro's law in calculations
* Recall and be able to use the Combined gas law and the General gas law in calculations
* Recall and be able to use the Ideal gas law in calculations
* Understand and be able to use the van der Waals equation (modified ideal gas law) in calculations
* Recall and be able to use Dalton's law of partial pressures in calculations
* Recall the conditions that are used as standard in calculations
* Be able to use molar gas volume in calculations
* Understand the Kinetic theory as applied to gases
* Understand the concept of, and be able to perform calculations involving, the root-mean-square-speed of gases
* Understand the terms effusion and diffusion and be able to perform calculations relating to those concepts

Chapter 11 - Intermolecular Forces

11.1 A Molecular Comparison of Gases, Liquids, and Solids
• Physical properties of substances are understood in terms of kinetic-molecular theory:
• Gases are highly compressible and assume the shape and volume of their container.
• Gas molecules are far apart and do not interact much with one another.
• Liquids are almost incompressible, assume the shape but not the volume of the container.
• Liquids molecules are held together more closely than gas molecules but not so rigidly that the molecules cannot slide past each other.
• Solids are incompressible and have a definite shape and volume.
• Solid molecules are packed closely together.
• The molecules are so rigidly packed that they cannot easily slide past each other.
• Solids and liquids are condensed phases.
• Solids with highly ordered structures are said to be crystalline.
• Converting a gas into a liquid or solid requires the molecules to get closer to each other.
• We can accomplish this by cooling or compressing the gas.
• Converting a solid into a liquid or gas requires the molecules to move further apart.
• We can accomplish this by heating or reducing the pressure on the gas.
• The forces holding solids and liquids together are called intermolecular forces.
• Physical properties of liquids and solids are due to intermolecular forces.
• These are forces between molecules.

11.2 Intermolecular Forces
• The attraction between molecules is an intermolecular force.
• Intermolecular forces are much weaker than ionic or covalent bonds.
• When a substance melts or boils, intermolecular forces are broken.
• When a substances condenses, intermolecular forces are formed.
• Boiling points reflect intermolecular force strength.
• A high boiling point indicates strong attractive forces.
• Melting points also reflect the strength of attractive forces.
• A high melting point indicates strong attractive forces.
• van der Waals forces are the intermolecular forces that exist between neutral molecules.
• These include London dispersion forces, dipole–dipole forces, and hydrogen-bonding forces.
• Ion–dipole interactions are important in solutions.
• These are all weak (<15% as strong as a covalent or ionic bond) electrostatic interactions.

Ion-Dipole Forces
• An ion–dipole force is an interaction between an ion (e.g., Na+) and the partial charge on the end of a polar molecule/dipole (e.g., water).
• It is especially important for solutions of ionic substances in polar liquids.
• Example: NaCl(aq)

Dipole–Dipole Forces
• Dipole–dipole forces exist between neutral polar molecules.
• Polar molecules attract each other.
• The partially positive end of one molecule attracts the partially negative end of another.
• Polar molecules need to be close together to form strong dipole–dipole interactions.
• Dipole–dipole forces are weaker than ion–dipole forces.
• If two molecules have about the same mass and size, then dipole–dipole forces increase with increasing polarity.
• For molecules of similar polarity, those with smaller volumes often have greater dipole–dipole attractions.

London-Dispersion Forces
• These are the weakest of all intermolecular forces.
• It is possible for two adjacent neutral molecules to affect each other.
• The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).
• For an instant, the electron clouds become distorted.
• In that instant a dipole is formed (called an instantaneous dipole).
• One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).
• These two temporary dipoles attract each other.
• The attraction is called the London dispersion force, or simply a dispersion force.
• London dispersion forces exist between all molecules.
• What affects the strength of a dispersion force?
• Molecules must be very close together for these attractive forces to occur.
• Polarizability is the ease with which an electron distribution can be deformed.
• The larger the molecule (the greater the number of electrons) the more polarizable it is.
• London dispersion forces increase as molecular weight increases.
• London dispersion forces depend on the shape of the molecule.
• The greater the surface area available for contact, the greater the dispersion forces.
• London dispersion forces between spherical molecules are smaller than those between more cylindrically shaped molecules.
• Example: n-pentane vs. neopentane.

Hydrogen Bonding
• Experiments show that the boiling points of compounds with H–F, H–O, and H–N bonds are abnormally high.
• Their intermolecular forces are abnormally strong.
• Hydrogen bonding is a special type of intermolecular attraction.
• This is a special case of dipole–dipole interactions.
• H-bonding requires:
• H bonded to a small electronegative element (most important for compounds of F, O, and N).
• an unshared electron pair on a nearby small electronegative ion or atom (usually F, O, or N on another molecule).
• Electrons in the H-X bond (X is the more electronegative element) lie much closer to X than H.
• H has only one electron, so in the H-X bond, the H+ presents an almost bare proton to the X-.
• Bond energies of hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol.
• They are much weaker than ordinary chemical bonds.
• Intermolecular and intramolecular hydrogen bonds have exceedingly important biological significance.
• They are important in stabilizing protein structure, in DNA structure and function, etc.
• An interesting consequence of H-bonding is that ice floats.
• The molecules in solids are usually more closely packed than those in liquids.
• Therefore, solids are usually more dense than liquids.
• Ice is ordered with an open structure to optimize H-bonding.
• Water molecules in ice are arranged in an open, regular hexagon.
• Each d+ H points towards a lone pair on O.
• Therefore, ice is less dense than water.
• Ice floats, so it forms an insulating layer on top of lakes, rivers, etc. Therefore, aquatic life can survive in winter.
• Water expands when it freezes.
• Frozen water in pipes may cause them to break in cold weather.

Comparing Intermolecular Forces
• Dispersion forces are found in all substances.
• Their strength depends on molecular shapes and molecular weights.
• Dipole–dipole forces add to the effect of dispersion forces.
• They are found only in polar molecules.
• H-bonding is a special case of dipole–dipole interactions.
• It is the strongest of the intermolecular forces involving neutral species.
• H-bonding is most important for H compounds of N, O, and F.
• If ions are involved, ion–dipole (if a dipole is present) and ionic bonding are possible.
• Ion–dipole interactions are stronger than H-bonds.
• Keep in mind that ordinary ionic or covalent bonds are much stronger than these interactions!

FORWARD REFERENCES
• Molecular materials held together by intermolecular forces will be called “soft” materials in Ch. 12.
• Intermolecular forces between polymer chains and in liquid crystals will be mentioned in Ch. 12 (sections 12.6 and 12.8).
• Breaking of solute–solute and solvent–solvent intermolecular forces and replacing them with solute–solvent interactions will take place in the solution process (Ch. 13).
• The binding between the substrate and the active site in the enzyme action thanks to the intermolecular forces will be discussed in Ch. 14 (section 14.7).
• Hydrogen bonding and the formation of hydrated hydronium ions will be discussed in Ch. 16 (section 16.2).
• Hydrogen bonding will be partially responsible for the relative weakness of HF compared to the strength of other binary acids involving halides (section 16.10).
• Hydrogen bonding and high heat capacity, high melting, and boiling points of water will be mentioned again in Ch. 18 (section 18.5).
• Intermolecular attractions in ice will be discussed in Ch. 19 (section 19.3).
• Hydrogen bonding in alcohols will be discussed in Ch. 25 (section 25.5).
• Hydrogen bonding in the a helix of a protein will be discussed in Ch. 25 (section 25.9).


11.3 Some Properties of Liquids

Viscosity
• Viscosity is the resistance of a liquid to flow.
• A liquid flows by sliding molecules over one another.
• Viscosity depends on:
• the attractive forces between molecules.
• The stronger the intermolecular forces, the higher the viscosity.
• the tendency of molecules to become entangled.
• Viscosity increases as molecules become entangled with one another.
• the temperature.
• Viscosity usually decreases with an increase in temperature.

Surface Tension
• Bulk molecules (those in the liquid) are equally attracted to their neighbors.
• Surface molecules are only attracted inward towards the bulk molecules.
• Therefore, surface molecules are packed more closely than bulk molecules.
• This causes the liquid to behave as if it had a “skin.”
• Surface tension is the amount of energy required to increase the surface area of a liquid by a unit amount.
• Stronger intermolecular forces cause higher surface tension.
• Water has a high surface tension (H-bonding)
• Hg(l) has an even higher surface tension (there are very strong metallic bonds between Hg atoms).
• Cohesive and adhesive forces are at play.
• Cohesive forces are intermolecular forces that bind molecules to one another.
• Adhesive forces are intermolecular forces that bind molecules to a surface.
• Illustrate this by looking at the meniscus in a tube filled with liquid.
• The meniscus is the shape of the liquid surface.
• If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g., water in glass).
• If cohesive forces are greater than adhesive forces, the meniscus is curved downwards (e.g., Hg(l) in glass)
• Capillary action is the rise of liquids up very narrow tubes.
• The liquid climbs until adhesive and cohesive forces are balanced by gravity.

FORWARD REFERENCES
• Viscosity of liquid crystals will be discussed in Ch. 12 (section 12.8).
• Viscosity of organic compounds, such as polyhydroxyl alcohols, will be mentioned in Ch. 25 (section 25.5).

11.4 Phase Changes
• Phase changes are changes of state.
• Matter in one state is converted into another state.
• Sublimation: solid ==> gas.
• Melting or fusion: solid ==> liquid.
• Vaporization: liquid ==> gas.
• Deposition: gas ==> solid.
• Condensation: gas ==> liquid.
• Freezing: liquid ==> solid.

Energy Changes Accompanying Phase Changes
• Energy changes of the system for the above processes are:
melting or fusion: ΔHfus > 0 (endothermic).
• The enthalpy of fusion is known as the heat of fusion.
• vaporization: ΔHvap > 0 (endothermic).
• The enthalpy of vaporization is known as the heat of vaporization.
• sublimation: ΔHsub > 0 (endothermic).
• The enthalpy of sublimation is called the heat of sublimation.
• deposition: ΔHdep < 0 (exothermic).
• condensation: ΔHcon < 0 (exothermic).
• freezing: ΔHfre < 0 (exothermic).
• Generally the heat of fusion (enthalpy of fusion) is less than heat of vaporization.
• It takes more energy to completely separate molecules than to partially separate them.
• All phase changes are possible under the right conditions (e.g., water sublimes when snow disappears without forming puddles).
• The following sequence is endothermic:
heat solid ==> melt ==> heat liquid ==> boil ==> heat gas
• The following sequence is exothermic:
cool gas ==> condense ==> cool liquid ==> freeze ==> cool solid

Heating Curves
• Plot of temperature change versus heat added is a heating curve.
• During a phase change adding heat causes no temperature change.
• The added energy is used to break intermolecular bonds rather than cause a temperature change.
• These points are used to calculate ΔHfus and ΔHvap.
• Supercooling: When a liquid is cooled below its freezing point and it still remains a liquid.

Critical Temperature and Pressure
• Gases may be liquefied by increasing the pressure at a suitable temperature.
• Critical temperature is the highest temperature at which a substance can exist as a liquid.
• Critical pressure is the pressure required for liquefaction at this critical temperature.
• The greater the intermolecular forces, the easier it is to liquefy a substance.
• Thus the higher the critical temperature

FORWARD REFERENCES
• Supercritical fluids in green chemistry will be discussed in Ch. 18 (section 18.7).
• Thermodynamics of phase changes will be further discussed throughout Ch. 19.

11.5 Vapor Pressure
Explaining Vapor Pressure on the Molecular Level
• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid.
• These molecules move into the gas phase.
• As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid.
• After some time the pressure of the gas will be constant.
• A dynamic equilibrium has been established.
• Dynamic equilibrium is a condition in which two opposing processes occur simultaneously at equal rates.
• In this case, it is the point when as many molecules escape the surface as strike the surface.
• Vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor are in dynamic equilibrium.
• The pressure of the vapor at this point is called the equilibrium vapor pressure.
Volatility, Vapor Pressure, and Temperature ,
• If equilibrium is never established the vapor continues to form.
• Eventually, the liquid evaporates to dryness.
• Liquids that evaporate easily are said to be volatile.
• The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

Vapor Pressure and Boiling Point
• Liquids boil when the external pressure at the liquid surface equals the vapor pressure.
• The normal boiling point is the boiling point at 760 mm Hg (1 atm).
• The temperature of the boiling point increases as the external pressure increases.
• Two ways to get a liquid to boil increase temperature or decrease pressure.
• Pressure cookers operate at high pressure.
• At high pressure the boiling point of water is higher than at 1 atm.
• Therefore, food is cooked at a higher temperature.

FORWARD REFERENCES
• Vapor pressure reduction of the solvent in a solution—a colligative property—will be discussed in Ch. 13 (section 13.5).

11.6 Phase Diagrams
• A phase diagram is a plot of pressure vs. temperature summarizing all equilibria between phases.
• Phase diagrams tell us which phase will exist at a given temperature and pressure.
• Features of a phase diagram include:
• vapor-pressure curve: generally as temperature increases, vapor pressure increases.
• critical point: critical temperature and pressure for the gas.
• normal melting point: melting point at 1 atm.
• triple point: temperature and pressure at which all three phases are in equilibrium.
• Any temperature and pressure combination not on a curve represents a single phase.

Phase Diagrams of H2O and CO2
• Water:
• In general, an increase in pressure favors the more compact phase of the material.
• This is usually the solid.
• Water is one of the few substances whose solid form is less dense than the liquid form.
• The melting point curve for water slopes to the left.
• The triple point occurs at 0.0098 C and 4.58 mm Hg.
• The normal melting (freezing) point is 0 C.
• The normal boiling point is 100 C.
• The critical point is 374 C and 218 atm.
• Carbon Dioxide:
• The triple point occurs at -56.4 C and 5.11 atm.
• The normal sublimation point is -78.5 C. (At 1 atm CO2 sublimes, it does not melt.)
• The critical point occurs at 31.1 C and 73 atm.
• Freeze drying: Frozen food is placed in a low pressure (< 4.58 torr) chamber.
• The ice sublimes.

FORWARD REFERENCES
• Phase diagrams for a pure solvent and for a solution of a nonvolatile solute will be discussed in Ch. 13 (section 13.5).
• Phase equilibria at melting and boiling points will be further analyzed in Ch. 19.

11.7 Structures of Solids
• A crystalline solid has a well-ordered, definite arrangements of molecules, atoms, or ions.
• Examples are quartz, diamond, salt, and sugar.
• The intermolecular forces are similar in strength.
• Thus they tend to melt at specific temperatures.
• In an amorphous solid, molecules, atoms or ions do not have an orderly arrangement.
• Examples are rubber and glass.
• Amorphous solids have intermolecular forces that vary in strength.
• Thus they tend to melt over a range of temperatures.

Unit Cells
• Crystalline solids have an ordered, repeating structure.
• The smallest repeating unit in a crystal is a unit cell.
• The unit cell is the smallest unit with all the symmetry of the entire crystal.
• The three-dimensional stacking of unit cells is the crystal lattice.
• Each point in the lattice is a lattice point which represents an identical environment within the solid.
• There are three types of cubic unit cells.
• Primitive cubic.
• The lattice points are at the corners of a simple cube with each atom shared by eight unit cells.
• Body-centered cubic (bcc).
• Lattice points occur at the corners of a cube and in addition there is a lattice point in the center of the body of the cube. The corner lattice points are shared by eight unit cells, and the center atom is completely enclosed in one unit cell.
• Face-centered cubic (fcc).
• There are lattice points at the corners of a cube plus one lattice point in the center of each face of the cube. Eight unit cells share the corner lattice points and two unit cells share the face lattice points.

The Crystal Structure of Sodium Chloride
• It has a face-centered cubic lattice.
• There are two equivalent ways of defining this unit cell:
• Cl– (larger) ions at the corners of the cell, or
• Na+ (smaller) ions at the corners of the cell.
• The cation to anion ratio in a unit cell is the same for the crystal.
• In NaCl each unit cell contains the same number of Na+ and Cl– ions.
• Note that the unit cell for CaCl2 needs twice as many Cl– ions as Ca2+ ions.

Close Packing of Spheres
• Crystalline solids have structures that maximize the attractive forces between particles.
• Their particles can be modeled by spheres.
• Each atom or ion is represented by a sphere.
• Molecular crystals are formed by close packing of the molecules.
• Maximum intermolecular forces in crystals are achieved by the close packing of spheres.
• A crystal is built up by placing close packed layers of spheres on top of each other.
• There is only one place for the second layer of spheres.
• There are two choices for the third layer of spheres:
• The third layer eclipses the first (ABAB arrangement).
• This is called hexagonal close packing (hcp).
• The third layer is in a different position relative to the first (ABCABC arrangement).
• This is called cubic close packing (ccp).
• Note that the unit cell of a ccp crystal is face-centered cubic.
• In both close-packed structures, each sphere is surrounded by 12 other spheres (6 in one plane, 3 above, and 3 below).
• Coordination number is the number of spheres directly surrounding a central sphere.
• When spheres are packed as closely as possible, there are small spaces between adjacent spheres (interstitial holes).
• If unequally sized spheres are used, the smaller spheres are placed in the interstitial holes.
• For example: Li2O
• The larger O2– ions assume the cubic close-packed structure with the smaller Li+ ions in the holes.

FORWARD REFERENCES
• Coordination numbers will be used in Ch. 23 for transition metal complexes.

11.8 Bonding in Solids
• The physical properties of crystalline solids depend on the:
• attractive forces between particles and on
• the arrangement of the particles.

Molecular Solids
• Molecular solids consist of atoms or molecules held together by intermolecular forces.
• Weak intermolecular forces give rise to low melting points.
• Intermolecular forces include dipole–dipole, London dispersion and H-bonds.
• Molecular solids are usually soft.
• They are often gases or liquids are room temperature.
• Efficient packing of molecules is important (since they are not regular spheres).
• Molecular solids show poor thermal and electrical conductivity.
• Examples: Ar(s), CH4(s), CO2(s), sucrose.

Covalent-Network Solids
• Covalent-network solids consist of atoms held together, in large networks or chains, with covalent bonds.
• They have much higher melting points and are much harder than molecular solids.
• This is a consequence of the strong covalent bonds that connect the atoms.
• Examples are diamond, graphite, quartz (SiO2), and silicon carbide (SiC).
• In diamond:
• each C atom has a coordination number of 4.
• each C atom is tetrahedral.
• there is a three-dimensional array of atoms.
• Diamond is hard, and has a high melting point (3550C).
• In graphite:
• each C atom is arranged in a planar hexagonal ring.
• layers of interconnected rings are placed on top of each other.
• the distance between adjacent C atoms in the same layer is close to that seen in benzene (1.42 vs. 1.395 in benzene).
• electrons move in delocalized orbitals (good conductor).
• the distance between layers is large (3.41 ).
• the layers are held together by weak dispersion forces.
• They slide easily past each other.
• Graphite is a good lubricant.

Ionic Solids
• Ionic solids consist of ions held together by ionic bonds.
• They are hard, brittle and have high melting points.
• Ions (spherical) are held together by electrostatic forces of attraction.
• Recall:
• The larger the charges (q1, q2) and smaller the distance (d) between ions, the stronger the ionic bond.
• The structure of the ionic solid depends on the charges on the ions and on the relative sizes of the atoms.
• Examples of some ionic lattice types are:
• NaCl structure.
• Each ion has a coordination number of six.
• It has a face-centered cubic lattice.
• The cation to anion ratio is 1:1.
• Other similar examples: LiF, KCl, AgCl, and CaO.
• CsCl structure.
• Cs+ has a coordination number of eight.
• It is different from the NaCl structure (Cs+ is larger than Na+).
• The cation to anion ratio is 1:1.
• Zinc blende (ZnS) structure.
• S2– ions adopt a face-centered cubic arrangement.
• Zn2+ ions have a coordination number of four.
• The S2– ions are placed in a tetrahedron around the Zn2+ ions.
• Another example is CuCl.
• Fluorite (CaF2) structure.
• Ca2+ ions are in a face-centered cubic arrangement.
• There are twice as many F– ions as Ca2+ ions in each unit cell.
• Other examples are BaCl2 and PbF2.

Metallic Solids
• Metallic solids consist entirely of metal atoms.
• Metallic solids are soft or hard.
• They have high melting points.
• They show good electrical and thermal conductivity.
• They are ductile and malleable.
• Examples are all metallic elements (i.e., Al, Cu, Fe, Au)
• Metallic solids have metal atoms in hexagonal close-packed, face-centered cubic or body-centered cubic arrangements.
• Thus the coordination number for each atom is either 8 or 12.
• A problem that needs to be explained:
• The bonding is too strong to be explained by London dispersion forces and there are not enough electrons for covalent bonds.
• Resolution:
• The metal nuclei float in a sea of delocalized valence electrons.
• Metals conduct heat and electricity because the valence electrons are delocalized and are mobile.

X-Ray Diffraction
• When waves are passed through a narrow slit they are diffracted.
• When waves are passed through a diffraction grating (many narrow slits in parallel) they interact to form a diffraction pattern (areas of light and dark bands).
• Efficient diffraction occurs when the wavelength of light is close to the size of the slits.
• The spacing between layers in a crystal is 2–20 , which is the wavelength range for X-rays.
• X-ray diffraction (X-ray crystallography):
• X-rays are passed through the crystal and are detected on a photographic plate.
• The photographic plate has one bright spot at the center (incident beam) as well as a diffraction pattern.
• Each close-packing arrangement produces a different diffraction pattern.
• Knowing the diffraction pattern, we can calculate the positions of the atoms required to produce that pattern.
• We calculate the crystal structure based on a knowledge of the diffraction pattern.

Chapter 13 - Properties of Solutions

13.1 The Solution Process
• A solution is a homogeneous mixture of solute and solvent.
• Solutions may be gases, liquids, or solids.
• Each substance present is a component of the solution.
• The solvent is the component present in the largest amount.
• The other components are the solutes.

The Effect of Intermolecular Forces

• Intermolecular forces become rearranged in the process of making solutions with condensed phases.
• Consider NaCl (solute) dissolving in water (solvent):
• Water molecules orient themselves on the NaCl crystals.
• H-bonds between the water molecules have to be broken.
• NaCl dissociates into Na+ and Cl–.
• Ion–dipole forces form between the Na+ and the negative end of the water dipole.
• Similar ion–dipole interactions form between the Cl– and the positive end of the water dipole.
• Such an interaction between solvent and solute is called solvation.
• If water is the solvent, the interaction is called hydration.

Energy Changes and Solution Formation
• There are three steps involving energy in the formation of a solution:
• Separation of solute molecules (ΔH1),
• Separation of solvent molecules (ΔH2), and
• Formation of solute–solvent interactions (ΔH3).
• We define the enthalpy change in the solution process as:
ΔHsoln = ΔH1 + ΔH2 + ΔH3
• ΔHsoln can either be positive or negative depending on the intermolecular forces.
• To determine whether ΔHsoln is positive or negative, we consider the strengths of all solute–solute, solvent–solvent, and solute–solvent interactions:
• Breaking attractive intermolecular forces is always endothermic.
• ΔH1 and ΔH2 are both positive.
• Forming attractive intermolecular forces is always exothermic.
• ΔH3 is always negative.
• It is possible to have either ΔH3 > (ΔH1 + ΔH2) or ΔH3 < (ΔH1 + ΔH2).
• Examples:
• MgSO4 added to water has ΔHsoln = –91.2 kJ/mol.
• NH4NO3 added to water has ΔHsoln = + 26.4 kJ/mol.
• MgSO4 is often used in instant heat packs and NH4NO3 is often used in instant cold packs.
• How can we predict if a solution will form?
• In general, solutions form if the ΔHsoln is negative.
• If ΔHsoln is too endothermic a solution will not form.
• “Rule of thumb”: polar solvents dissolve polar solutes.
• Nonpolar solvents dissolve nonpolar solutes.
• Consider the process of mixing NaCl in gasoline.
• Only weak interactions are possible because gasoline is nonpolar.
• These interactions do not compensate for the separation of ions from one another.
• Result: NaCl doesn't dissolve to any great extent in gasoline.
• Consider the process of mixing water in octane (C8H18).
• Water has strong H-bonds.
• The energy required to break these H-bonds is not compensated for by interactions between water and octane.
• Result: water and octane do not mix.

Solution Formation, Spontaneity, and Disorder
• A spontaneous process occurs without outside intervention.
• When the energy of the system decreases (e.g., dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous.
• Spontaneous processes tend to be exothermic.
• However, some spontaneous processes do not involve the movement of the system to a lower energy state (e.g., an endothermic reaction).
• Some endothermic processes occur spontaneously.
• The amount of randomness or disorder of the system is given by its entropy.
• In most cases, solution formation is favored by the increase in entropy that accompanies mixing.
• Example: A mixture of CCl4 and C6H14 is less ordered than the two separate liquids.
• Therefore, they spontaneously mix even though ΔHsoln is very close to zero.
• A solution will form unless the solute–solute or solvent–solvent interactions are too strong relative to solute–solvent interactions.

Solution Formation and Chemical Reactions
• Some solutions form by physical processes and some by chemical processes.
• Consider:
Ni(s) + 2HCl(aq) ‡ NiCl2(aq) + H2(g)
• Note that the chemical form of the substance being dissolved has changed during this process (Ni ‡ NiCl2)
• When all the water is removed from the solution, no Ni is found, only NiCl2•6H2O remains.
• Therefore, the dissolution of Ni in HCl is a chemical process.
• By contrast:
NaCl(s) + H2O (l) ‡ Na+(aq) + Cl–(aq).
• When the water is removed from the solution, NaCl is found.
• Therefore, NaCl dissolution is a physical process.

FORWARD REFERENCES
• Hydrolysis of metal ions will be brought up again in Ch. 16 (section 16.11).
• Thermodynamics of processes will be further discussed throughout Ch. 19.
• Rust is a hydrate (Ch. 20, section 20.8).
• Solution alloys will be further described in Ch. 23 (section 23.6).

13.2 Saturated Solutions and Solubility
• As a solid dissolves, a solution forms:
• Solute + solvent ==> solution
• The opposite process is crystallization.
• Solution ==> solute + solvent
• If crystallization and dissolution are in equilibrium with undissolved solute, the solution is saturated.
• There will be no further increase in the amount of dissolved solute.
• Solubility is the amount of solute required to form a saturated solution.
• A solution with a concentration of dissolved solute that is less than the solubility is said to be unsaturated.
• A solution is said to be supersaturated if more solute is dissolved than in a saturated solution.

FORWARD REFERENCES
• Solubility equilibria will be further discussed in Ch. 17 (section 17.4).
• Various crystalline substances (anhydrous and hydrated) for transition metal compounds will be described throughout Ch. 23.

13.3 Factors Affecting Solubility
• The tendency of a substance to dissolve in another depends on:
• the nature of the solute.
• the nature of the solvent.
• the temperature.
• the pressure (for gases).

Solute–Solvent Interactions
• Intermolecular forces are an important factor in determining solubility of a solute in a solvent.
• The stronger the attraction between solute and solvent molecules, the greater the solubility.
• For example, polar liquids tend to dissolve in polar solvents.
• Favorable dipole–dipole interactions exist (solute–solute, solvent–solvent, and solute–solvent).
• Pairs of liquids that mix in any proportions are said to be miscible.
• Example: Ethanol and water are miscible liquids.
• In contrast, immiscible liquids do not mix significantly.
• Example: Gasoline and water are immiscible.
• Consider the solubility of alcohols in water.
• Water and ethanol are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture.
• However, not all alcohols are miscible with water.
• Why?
• The number of carbon atoms in a chain affects solubility.
• The greater the number of carbons in the chain, the more the molecule behaves like a hydrocarbon.
• Thus, the more C atoms in the alcohol, the lower its solubility in water.
• Increasing the number of –OH groups within a molecule increases its solubility in water.
• The greater the number of –OH groups along the chain, the more solute-water H-bonding is possible.
• Generalization: “like dissolves like”.
• Substances with similar intermolecular attractive forces tend to be soluble in one another.
• The more polar bonds in the molecule, the better it dissolves in a polar solvent.
• The less polar the molecule the less likely it is to dissolve in a polar solvent and the more likely it is to dissolve in a nonpolar solvent.
• Network solids do not dissolve because the strong intermolecular forces in the solid are not reestablished in any solution.

Pressure Effects
• The solubility of a gas in a liquid is a function of the pressure of the gas over the solution.
• Solubilities of solids and liquids are not greatly affected by pressure.
• With higher gas pressure, more molecules of gas are close to the surface of the solution and the probability of a gas molecule striking the surface and entering the solution is increased.
• Therefore, the higher the pressure, the greater the solubility.
• The lower the pressure, the smaller the number molecules of gas close to the surface of the solution resulting in a lower solubility.
• The solubility of a gas is directly proportional to the partial pressure of the gas above the solution.
• This statement is called Henry's law.
• Henry's law may be expressed mathematically as: Sg=kPg
• Where Sg is the solubility of gas, Pg the partial pressure, k = Henry’s law constant.
• Note that the Henry's law constant differs for each solute–solvent pair and differs with temperature.
• An application of Henry's law is the preparation of carbonated soda.
• Carbonated beverages are bottled under PCO2 > 1 atm.
• As the bottle is opened, PCO2 decreases and the solubility of CO2 decreases.
• Therefore, bubbles of CO2 escape from solution.

Temperature Effects
• Experience tells us that sugar dissolves better in warm water than in cold water.
• As temperature increases, solubility of solids generally increases.
• Sometimes solubility decreases as temperature increases [e.g., Ce2(SO4)3].
• Experience tells us that carbonated beverages go flat as they get warm.
• Gases are less soluble at higher temperatures.
• An environmental application of this is thermal pollution.
• Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available for plants or animals.
• Fish suffocate.

FORWARD REFERENCES
• The dynamic equilibrium between a solid solute and its solution will be mentioned in Ch. 14 (section 14.7).
• Factors affecting solubility will be discussed in detail in Ch. 17 (section 17.5).
• Temperature effects on solubility of NaCl will be mentioned in Ch. 19 (section 19.7).
• Reactions involving CO2, HCO3- and H2CO3 will be discussed in Ch. 22 (section 22.9).
• Solubility of organic substances in polar solvents will be mentioned in Ch. 25 (section 25.1).

13.4 Ways of Expressing Concentration
• All methods involve quantifying the amount of solute per amount of solvent (or solution).
• Concentration may be expressed qualitatively or quantitatively.
• The terms dilute and concentrated are qualitative ways to describe concentration.
• A dilute solution has a relatively small concentration of solute.
• A concentrated solution has a relatively high concentration of solute.
• Quantitative expressions of concentration require specific information regarding such quantities as masses, moles, or liters of the solute, solvent, or solution.
• The most commonly used expressions for concentration are:
• mass percentage.
• mole fraction.
• molarity.
• molality.

Mass Percentage, ppm, and ppb
• Mass percentage is one of the simplest ways to express concentration.
• By definition: (wouldn't copy -- in book)
• Similarly, parts per million (ppm) can be expressed as the number of mg of solute per kilogram of solution.
• By definition: (wouldn't copy -- in book)
• The density of a very dilute aqueous solution is similar to that of pure water (1g/ml)
• If the density of the solution is 1g/ml, then 1 ppm = 1 mg solute per liter of solution.
• We can extend this again!
• Parts per billion (ppb) can be expressed as the number of g of solute per kilogram of solution.
• By definition: (wouldn't copy -- in book)
• If the density of the solution is 1g/ml, then 1 ppb = 1 g solute per liter of solution.

Mole Fraction, Molarity, and Molality
• Common expressions of concentration are based on the number of moles of one or more components.
• Recall that mass can be converted to moles using the molar mass.
• Recall:

• Note that mole fraction has no units.
• Note that mole fractions range from 0 to 1.
• Recall:
• Note that molarity will change with a change in temperature (as the solution volume increases or decreases).
• We can define molality (m), yet another concentration unit:
• Molality does not vary with temperature.
• Note that converting between molarity (M) and molality (m) requires density.
• The molarity and molality of dilute solutions are often very similar.

FORWARD REFERENCES
• Molar concentrations will be used in rate law expressions in Ch. 14.
• Molar concentrations will be used in equilibrium constant and reaction quotient expressions in Ch. 15, 16, 17, 19, and 20.
• Concentrations (ppm) of trace constituents in mixtures will be mentioned in Ch. 18 (section 18.1).

13.5 Colligative Properties
• Colligative properties depend on number of solute particles.
• There are four colligative properties to consider:
• vapor pressure lowering (Raoult's Law).
• boiling point elevation.
• freezing point depression.
• osmotic pressure.

Lowering the Vapor Pressure
• Consider a volatile liquid in a closed container.
• After a period of time an equilibrium will be established between the liquid and its vapor.
• The partial pressure exerted by the vapor is the vapor pressure.
• Nonvolatile solutes (with no measurable vapor pressure) reduce the ability of the surface solvent molecules to escape the liquid.
• Therefore, vapor pressure is lowered.
• The amount of vapor pressure lowering depends on the amount of solute.
• Raoult’s law quantifies the extent to which a nonvolatile solute lowers the vapor pressure of the solvent.
• If PA is the vapor pressure with solute, PA is the vapor pressure of the pure solvent, and CA is the mole fraction of A, then
• An ideal solution is one that obeys Raoult’s law.
• Real solutions show approximately ideal behavior when:
• the solute concentration is low.
• the solute and solvent have similarly sized molecules.
• the solute and solvent have similar types of intermolecular attractions.
• Raoult’s law breaks down when the solvent–solvent and solute–solute intermolecular forces are much greater or weaker than solute–solvent intermolecular forces.

Boiling-Point Elevation
• A nonvolatile solute lowers the vapor pressure of a solution.
• At the normal boiling point of the pure liquid, the solution has a has a vapor pressure less than 1 atm.
• Therefore, a higher temperature is required to reach a vapor pressure of 1 atm for the solution (ΔTb).
• The molal boiling-point-elevation constant, Kb, expresses how much ΔTb changes with molality, m: ΔTb = Kbm
• The nature of the solute (electrolyte vs. nonelectrolyte) will impact the colligative molality of the solute.

Freezing-Point Depression
• When a solution freezes, crystals of almost pure solvent are formed first.
• Solute molecules are usually not soluble in the solid phase of the solvent.
• Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution.
• The melting-point (freezing-point) curve is a vertical line from the triple point.
• Therefore, the solution freezes at a lower temperature (ΔTf) than the pure solvent.
• The decrease in freezing point (ΔTf) is directly proportional to molality.
• Kf is the molal freezing-point-depression constant.
ΔTf = Kfm
• Values of Kf and Kb for most common solvents can be found in Table 13.4.

Osmosis
• Semipermeable membranes permit passage of some components of a solution.
• Often they permit passage of water but not larger molecules or ions.
• Examples of semipermeable membranes are cell membranes and cellophane.
• Osmosis is the net movement of a solvent from an area of low solute concentration to an area of high solute concentration.
• Consider a U-shaped tube with a two liquids separated by a semipermeable membrane.
• One arm of the tube contains pure solvent.
• The other arm contains a solution.
• There is movement of solvent in both directions across a semipermeable membrane.
• As solvent moves across the membrane, the fluid levels in the arms become uneven.
• The vapor pressure of solvent is higher in the arm with pure solvent.
• Eventually the pressure difference due to the difference in height of liquid in the arms stops osmosis.
• Osmotic pressure, p, is the pressure required to prevent osmosis.
• Osmotic pressure obeys a law similar in form to the ideal-gas law.
• For n moles, V= volume, M= molarity, R= the ideal gas constant, and an absolute temperature, T, the osmotic pressure is:
pV = nRT
• Two solutions are said to be isotonic if they have the same osmotic pressure.
• Hypotonic solutions have a lower p, relative to a more concentrated solution.
• Hypertonic solutions have a higher p, relative to a more dilute solution.
• We can illustrate this with a biological system: red blood cells.
• Red blood cells are surrounded by semipermeable membranes.
• If red blood cells are placed in a hypertonic solution (relative to intracellular solution), there is a lower solute concentration in the cell than the surrounding tissue.
• Osmosis occurs and water passes through the membrane out of the cell.
• The cell shrivels up.
• This process is called crenation.
• If red blood cells are placed in a hypotonic solution, there is a higher solute concentration in the cell than outside the cell.
• Osmosis occurs and water moves into the cell.
• The cell bursts (hemolysis).
• To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic relative to the intracellular fluids of cells.
• Everyday examples of osmosis include:
• If a cucumber is placed in NaCl solution, it will lose water to shrivel up and become a pickle.
• A limp carrot placed in water becomes firm because water enters via osmosis.
• Eating large quantities of salty food causes retention of water and swelling of tissues (edema).
• Water moves into plants, to a great extent, through osmosis.
• Salt may be added to meat (or sugar added to fruit) as a preservative.
• Salt prevents bacterial infection: A bacterium placed on the salt will lose water through osmosis and die.
• Active transport is the movement of nutrients and waste material through a biological membrane against a concentration gradient.
• Movement is from an area of low concentration to an area of high concentration.
• Active transport is not spontaneous.
• Energy must be expended by the cell to accomplish this.

Determination of Molar Mass
• Any of the four colligative properties may be used to determine molar mass.

FORWARD REFERENCES
• Desalination via reverse osmosis will be described in Ch. 18 (section 18.5).

13.6 Colloids
• Colloids or colloidal dispersions are suspensions in which the suspended particles are larger than molecules but too small to separate out of the suspension due to gravity.
• Particle size: 5 to 1000 nm.
• There are several types of colloids:
• aerosol: gas + liquid or solid (e.g., fog and smoke),
• foam: liquid + gas (e.g., whipped cream),
• emulsion: liquid + liquid (e.g., milk),
• sol: liquid + solid (e.g., paint),
• solid foam: solid + gas (e.g., marshmallow),
• solid emulsion: solid + liquid (e.g., butter),
• solid sol: solid + solid (e.g., ruby glass).
• The Tyndall effect is the ability of colloidal particles to scatter light.
• The path of a beam of light projected through a colloidal suspension can be seen through the suspension.

Hydrophilic and Hydrophobic Colloids
• Focus on colloids in water.
• Water-loving colloids are hydrophilic.
• Water-hating colloids are hydrophobic.
• In the human body, large biological molecules such as proteins are kept in suspension by association with surrounding water molecules.
• These macromolecules fold up so that hydrophobic groups are away from the water (inside the folded molecule).
• Hydrophilic groups are on the surface of these molecules and interact with solvent (water) molecules.
• Typical hydrophilic groups are polar (containing C–O, O–H, N–H bonds) or charged.
• Hydrophobic colloids need to be stabilized in water.
• One way to stabilize hydrophobic colloids is to adsorb ions on their surface.
• Adsorption: when something sticks to a surface we say that it is adsorbed.
• If ions are adsorbed onto the surface of a colloid, the colloid appears hydrophilic and is stabilized in water.
• Consider a small drop of oil in water.
• Add a small amount of sodium stearate.
• Sodium stearate has a long hydrophobic hydrocarbon tail and a small hydrophilic head.
• The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface.
• The hydrophilic heads then interact with the water and the oil drop is stabilized in water.
• A soap acts in a similar fashion.
• Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water.
• Most dirt stains on people and clothing are oil based.
• Biological application of this principle:
• The gallbladder excretes a fluid called bile.
• Bile contains substances (bile salts) that form an emulsion with fats in our small intestine.
• Emulsifying agents help form an emulsion.
• Emulsification of dietary fats and fat-soluble vitamins is important in their absorption and digestion by the body.

Removal of Colloid Particles
• We often need to separate colloidal particles from the dispersing medium.
• This may be problematic.
• Colloid particles are too small to be separated by physical means (e.g., filtration).
• However, colloid particles often may be coagulated (enlarged) until they can be removed by filtration.
• There are various methods of coagulation.
• Colloid particles move more rapidly when the colloidal dispersion is heated, increasing the number of collisions. The particles stick to each other when they collide.
• Adding an electrolyte neutralizes the surface charges on the colloid particles.
• A biological application of another approach to separating colloidal particles from the suspending medium is dialysis.
• In dialysis a semipermeable membrane is used to separate ions from colloidal particles.
• In kidney dialysis, the blood is allowed to pass through a semipermeable membrane immersed in a washing solution.
• The washing solution is isotonic in ions that must be retained.
• The washing solution does not have the waste products that are found in the blood.
• Wastes therefore dialyze out of the blood (move from the blood into the washing solution).
• The "good" ions remain in the blood.

FORWARD REFERENCES
• Adsorption and absorption in heterogeneous catalysis will be mentioned in Ch. 14 (section 14.7).
• A model of hemoglobin in blood will be provided in chapters 24 and 25 (sections 24.2 and 25.9, respectively).
• Surfactant organic molecules will be mentioned again in Ch. 25 (section 25.1).

Chapter 11 Charts, Diagrams, and Illustrations







Chapter 13 Charts, Diagrams, and Illustrations


   


Ka is the equilibrium constant for the dissociation reaction of a weak acid.
The value of Ka is used to calculate pH of weak acids.
The pKa value is used to choose a buffer when needed.
Choosing an acid or base where pKa is close to the pH needed gives the best results.

Ka of Weak Acids

Name Formula Ka pKa
acetic HC2H3O2 1.8 x 10-5 4.7
ascorbic (I) H2C6H6O6 7.9 x 10-5 4.1
ascorbic (II) HC6H6O6- 1.6 x 10-12 11.8
benzoic HC7H5O2 6.4 x 10-5 4.2
boric (I) H3BO3 5.4 x 10-10 9.3
boric (II) H2BO3- 1.8 x 10-13 12.7
boric (III) HBO32- 1.6 x 10-14 13.8
carbonic (I) H2CO3 4.5 x 10-7 6.3
carbonic (II) HCO3- 4.7 x 10-11 10.3
citric (I) H3C6H5O7 3.2 x 10-7 6.5
citric (II) H2C6H5O7- 1.7 x 105 4.8
citric (III) HC6H5O72- 4.1 x 10-7 6.4
formic HCHO2 1.8 x 10-4 3.7
hydrazidic HN3 1.9 x 10-5 4.7
hydrocyanic HCN 6.2 x 10-10 9.2
hydrofluoric HF 6.3 x 10-4 3.2
hydrogen peroxide H2O2 2.4 x 10-12 11.6
hydrogen sulfate ion HSO4- 1.2 x 10-2 1.9
hypochlorous HOCl 3.5 x 10-8 7.5
lactic HC3H5O3 8.3 x 10-4 3.1
nitrous HNO2 4.0 x 10-4 3.4
oxalic (I) H2C2O4 5.8 x 10-2 1.2
oxalic (II) HC2O4- 6.5 x 10-5 4.2
phenol HOC6H5 1.6 x 10-10 9.8
propanic HC3H5O2 1.3 x 10-5 4.9
sulfurous (I) H2SO3 1.4 x 10-2 1.85
sulfurous (II) HSO3- 6.3 x 10-8 7.2
uric HC5H3N4O3 1.3 x 10-4 3.9

 

What is an acid-base indicator?
An acid-base indicator is a weak acid or a weak base. The undissociated form of the indicator is a different color than the iogenic form of the indicator. An Indicator does not change color from pure acid to pure alkaline at specific hydrogen ion concentration, but rather, color change occurs over a range of hydrogen ion concentrations. This range is termed the color change interval. It is expressed as a pH range.

How is an indicator used?
Weak acids are titrated in the presence of indicators which change under slightly alkaline conditions. Weak bases should be titrated in the presence of indicators which change under slightly acidic conditions.

What are some common acid-base indicators?
Several acid-base indicators are listed below, some more than once if they can be used over multiple pH ranges. Quantity of indicator in aqueous (aq.) or alcohol (alc.) solution is specified. Tried-and-true indicators include: thymol blue, tropeolin OO, methyl yellow, methyl orange, bromphenol blue, bromcresol green, methyl red, bromthymol blue, phenol red, neutral red, phenolphthalein, thymolphthalein, alizarin yellow, tropeolin O, nitramine, and trinitrobenzoic acid. Data in this table are for sodium salts of thymol blue, bromphenol blue, tetrabromphenol blue, bromcresol green, methyl red, bromthymol blue, phenol red, and cresol red.

Indicator pH Range Quantity per 10 ml Acid Base
Thymol Blue 1.2-2.8 1-2 drops 0.1% soln. in aq. red yellow
Pentamethoxy red 1.2-2.3 1 drop 0.1% soln. in 70% alc. red-violet colorless
Tropeolin OO 1.3-3.2 1 drop 1% aq. soln. red yellow
2,4-Dinitrophenol 2.4-4.0 1-2 drops 0.1% soln. in 50% alc. colorless yellow
Methyl yellow 2.9-4.0 1 drop 0.1% soln. in 90% alc. red yellow
Methyl orange 3.1-4.4 1 drop 0.1% aq. soln. red orange
Bromphenol blue 3.0-4.6 1 drop 0.1% aq. soln. yellow blue-violet
Tetrabromphenol blue 3.0-4.6 1 drop 0.1% aq. soln. yellow blue
Alizarin sodium sulfonate 3.7-5.2 1 drop 0.1% aq. soln. yellow violet
a-Naphthyl red 3.7-5.0 1 drop 0.1% soln. in 70% alc. red yellow
p-Ethoxychrysoidine 3.5-5.5 1 drop 0.1% aq. soln. red yellow
Bromcresol green 4.0-5.6 1 drop 0.1% aq. soln. yellow blue
Methyl red 4.4-6.2 1 drop 0.1% aq. soln. red yellow
Bromcresol purple 5.2-6.8 1 drop 0.1% aq. soln. yellow purple
Chlorphenol red 5.4-6.8 1 drop 0.1% aq. soln. yellow red
Bromphenol blue 6.2-7.6 1 drop 0.1% aq. soln. yellow blue
p-Nitrophenol 5.0-7.0 1-5 drops 0.1% aq. soln. colorless yellow
Azolitmin 5.0-8.0 5 drops 0.5% aq. soln. red blue
Phenol red 6.4-8.0 1 drop 0.1% aq. soln. yellow red
Neutral red 6.8-8.0 1 drop 0.1% soln. in 70% alc. red yellow
Rosolic acid 6.8-8.0 1 drop 0.1% soln. in 90% alc. yellow red
Cresol red 7.2-8.8 1 drop 0.1% aq. soln. yellow red
a-Naphtholphthalein 7.3-8.7 1-5 drops 0.1% soln. in 70% alc. rose green
Tropeolin OOO 7.6-8.9 1 drop 0.1% aq. soln. yellow rose-red
Thymol blue 8.0-9.6 1-5 drops 0.1% aq. soln. yellow blue
Phenolphthalein 8.0-10.0 1-5 drops 0.1% soln. in 70% alc. colorless red
a-Naphtholbenzein 9.0-11.0 1-5 drops 0.1% soln. in 90% alc. yellow blue
Thymolphthalein 9.4-10.6 1 drop 0.1% soln. in 90% alc. colorless blue
Nile blue 10.1-11.1 1 drop 0.1% aq. soln. blue red
Alizarin yellow 10.0-12.0 1 drop 0.1% aq. soln. yellow lilac
Salicyl yellow 10.0-12.0 1-5 drops 0.1% soln. in 90% alc. yellow orange-brown
Diazo violet 10.1-12.0 1 drop 0.1% aq. soln. yellow violet
Tropeolin O 11.0-13.0 1 drop 0.1% aq. soln. yellow orange-brown
Nitramine 11.0-13.0 1-2 drops 0.1% soln in 70% alc. colorless orange-brown
Poirrier's blue 11.0-13.0 1 drop 0.1% aq. soln. blue violet-pink
Trinitrobenzoic acid 12.0-13.4 1 drop 0.1% aq. soln. colorless orange-red
These are molar heats of formation for anions and cations in aqueous solution.
In all cases, the heats of formation are given in kJ/mol at 25C for 1 mole of the ion.

Cations DHf (kJ/mol)
Anions DHf (kJ/mol)
Ag+ (aq) +105.9
Br- (aq) -120.9
Al3+ (aq) -524.7
Cl- (aq) -167.4
Ba2+ (aq) -538.4
ClO3- (aq) -98.3
Ca2+ (aq) -543.0
ClO4- (aq) -131.4
Cd2+ (aq) -72.4
CO32- (aq) -676.3
Cu2+ (aq) +64.4
CrO42- (aq) -863.2
Fe2+ (aq) -87.9
F- (aq) -329.1
Fe3+ (aq) -47.7
HCO3- (aq) -691.1
H+ (aq) 0.0
H2PO4- (aq) -1302.5
K+ (aq) -251.2
HPO42- (aq) -1298.7
Li+ (aq) -278.5
I- (aq) -55.9
Mg2+ (aq) -462.0
MnO4- (aq) -518.4
Mn2+ (aq) -218.8
NO3- (aq) -206.6
Na+ (aq) -239.7
OH- (aq) -229.9
NH4+ (aq) -132.8
PO43- (aq) -1284.1
Ni2+ (aq) -64.0
S2- (aq) +41.8
Pb2+ (aq) +1.6
SO42- (aq) -907.5
Sn2+ (aq) -10.0


Zn2+ (aq) -152.4


Reference: Masterton, Slowinski, Stanitski, Chemical Principles, CBS College Publishing, 1983.

This is an alphabetical list of alloys grouped according to the base metal of the alloy. Some alloys are listed under more than one element, since the composition of the alloy may vary such that one element is present in a higher concentration than the others.


Aluminum Alloys
* AA-8000: used for building wire
* Al-Li (aluminum, lithium, sometimes mercury)
* Alnico (aluminum, nickel, copper)
* Duralumin (copper, aluminum)
* Magnalium (aluminum, 5% magnesium)
* Magnox (magnesium oxide, aluminum)
* Nambe (aluminum plus seven other unspecified metals)
* Silumin (aluminum, silicon)
* Zamak (zinc, aluminum, magnesium, copper)
* Aluminum forms other complex alloys with magnesium, manganese, and platinum
Bismuth Alloys
* Wood's metal (bismuth, lead, tin, cadmium)
* Rose metal (bismuth, lead, tin)
* Field's metal
* Cerrobend
Cobalt Alloys
* Megallium
* Stellite (cobalt, chromium, tungsten or molybdenum, carbon)
o Talonite (cobalt, chromium)
* Ultimet (cobalt, chromium, nickel, molybdenum, iron, tungsten)
* Vitallium
Copper Alloys
* Arsenical copper
* Beryllium copper (copper, beryllium)
* Billon (copper, silver)
* Brass (copper, zinc)
o Calamine brass (copper, zinc)
o Chinese silver (copper, zinc)
o Dutch metal (copper, zinc)
o Gilding metal (copper, zinc)
o Muntz metal (copper, zinc)
o Pinchbeck (copper, zinc)
o Prince's metal (copper, zinc)
o Tombac (copper, zinc)
* Bronze (copper, tin, aluminum or any other element)
o Aluminum bronze (copper, aluminum)
o Arsenical bronze (copper, arsenic)
o Bell metal (copper, tin)
o Florentine bronze (copper, aluminum or tin)
o Glucydur (beryllium, copper, iron)
o Guani-n (likely a manganese bronze of copper, manganese, with iron sulfides and other sulfides)
o Gunmetal (copper, tin, zinc)
o Phosphor bronze (copper, tin and phosphorus)
o Ormolu (Gilt Bronze) (copper, zinc)
o Speculum metal (copper, tin)
* Constantan (copper, nickel)
* Copper-tungsten (copper, tungsten)
* Corinthian bronze (copper, gold, silver)
* Cunife (copper, nickel, iron)
* Cupronickel (copper, nickel)
* Cymbal alloys (Bell metal) (copper, tin)
* Devarda's alloy (copper, aluminum, zinc)
* Electrum (copper, gold, silver)
* Hepatizon (copper, gold, silver)
* Heusler alloy (copper, manganese, tin)
* Manganin (copper, manganese, nickel)
* Nickel silver (copper, nickel)
* Nordic gold (copper, aluminum, zinc, tin)
* Shakudo (copper, gold)
* Tumbaga (copper, gold)

Gallium Alloys
* Galinstan (gallium, indium, tin)
Gold Alloys
* Electrum (gold, silver, copper)
* Tumbaga (gold, copper)
* Rose gold (gold, copper)
* White gold (gold, nickel, palladium, or platinum)
Indium Alloys
* Field's metal (indium, bismuth, tin)
Iron or Ferrous Alloys
* Steel (carbon)
o Stainless steel (chromium, nickel)
+ AL-6XN
+ Alloy 20
+ Celestrium
+ Marine grade stainless
+ Martensitic stainless steel
+ Surgical stainless steel (chromium, molybdenum, nickel)
o Silicon steel (silicon)
o Tool steel (tungsten or manganese)
o Bulat steel
o Chromoly (chromium, molybdenum)
o Crucible steel
o Damascus steel
o HSLA steel
o High speed steel
o Maraging steel
o Reynolds 531
o Wootz steel
* Iron
o Anthracite iron (carbon)
o Cast iron (carbon)
o Pig iron (carbon)
o Wrought iron (carbon)
* Fernico (nickel, cobalt)
* Elinvar (nickel, chromium)
* Invar (nickel)
* Kovar (cobalt)
* Spiegeleisen (manganese, carbon, silicon)
* Ferroalloys
o Ferroboron
o Ferrochrome (chromium)
o Ferromagnesium
o Ferromanganese
o Ferromolybdenum
o Ferronickel
o Ferrophosphorus
o Ferrotitanium
o Ferrovanadium
o Ferrosilicon
Lead Alloys
* Antimonial lead (lead, antimony)
* Molybdochalkos (lead, copper)
* Solder (lead, tin)
* Terne (lead, tin)
* Type metal (lead, tin, antimony)
Magnesium Alloys
* Magnox (magnesium, aluminum)
* T-Mg-Al-Zn (Bergman phase)
* Elektron
Mercury Alloys
* Amalgam (mercury with just about any metal except platinum)
Nickel Alloys
* Alumel (nickel, manganese, aluminum, silicon)
* Chromel (nickel, chromium)
* Cupronickel (nickel, bronze, copper)
* German silver (nickel, copper, zinc)
* Hastelloy (nickel, molybdenum, chromium, sometimes tungsten)
* Inconel (nickel, chromium, iron)
* Monel metal (copper, nickel, iron, manganese)
* Mu-metal (nickel, iron)
* Ni-C (nickel, carbon)
* Nichrome (chromium, iron, nickel)
* Nicrosil (nickel, chromium, silicon, magnesium)
* Nisil (nickel, silicon)
* Nitinol (nickel, titanium, shape memory alloy)

Potassium Alloys
* KLi (potassium, lithium)
* NaK (sodium, potassium)
Rare Earth Alloys
* Mischmetal (various rare earths)
Silver Alloys
* Argentium sterling silver (silver, copper, germanium)
* Billon (copper or copper bronze, sometimes with silver)
* Britannia silver (silver, copper)
* Electrum (silver, gold)
* Goloid (silver, copper, gold)
* Platinum sterling (silver, platinum)
* Shibuichi (silver, copper)
* Sterling silver (silver, copper) Tin Alloys
* Britannium (tin, copper, antimony)
* Pewter (tin, lead, copper)
* Solder (tin, lead, antimony)
Titanium Alloys
* Beta C (titanium, vanadium, chromium, other metals)
* 6al-4v (titanium, aluminum, vanadium)
Uranium Alloys
* Staballoy (depleted uranium with titanium or molybdenum)
* Uranium may also be alloyed with plutonium
Zinc Alloys
* Brass (zinc, copper)
* Zamak (zinc, aluminum, magnesium, copper)
Zirconium Alloys
* Zircaloy (zirconium and tin, sometimes with niobium, chromium, iron, nickel)

Pressure Units
 
pascal
(Pa)

bar
(bar)
technical atmosphere
(at)

atmosphere
(atm)

torr
(Torr)
pound-force per
square inch
(psi)
1 Pa = 1 N/m2 10-5 1.0197 x 10-5 9.8692 x 10-6 7.5006 x 10-3 145.04 x 10-6
1 bar 100,000 = 106 dyn/cm2 1.0197 0.98692 750.06 14.5037744
1 at 98,066.5 0.980665 = 1 kg/cm2 0.96784 735.56 14.223
1 atm 101,325 1.01325 1.0332 = 1 atm 760 14.696
1 torr 133.322 1.3332 x 10-3 1.3595 x 10-3 1.3158 x 10-3 =1 Torr
 1 mmHg
19.337 x 10-3
1 psi 6.894 x 103 68.948 x 10-3 70.307 x 10-3 68.046 x 10-3 51.715 = 1 lb/in2

Boiling Point vs. Atmoic Number

Vapor Pressure of Selected Compounds

pH Scale

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