Chapter 13 Review
Terms in this set (105)
two or more substances mixed together physically, not combined chemically. Its composition is variable, and it retains some properties of its components. Two types of mixtures are solutions and colloids.
contain only a dozen or so components (ex. glass)
i.e. humans, trees. The most complex. (6000+ components)
Solution (type of mixture)
Homogeneous mixture (exists as one phase). The particles are individual atoms, ions or small molecules. They can be gaseous, liquid, or solid.
charged particles, such as a nitrate ion (NO3-).
the smallest individual particle (barring subatomic particles), such as Al, B, C, etc (various elemental atoms).
combinations of the atoms, such as water (H2O), or carbon dioxide (CO2).
Heterogeneous mixture consisting of two or more phases. (ex. pebbles in concrete, smoke, milk). The particles are typically macromolecules or aggregations of small molecules that are dispersed so finely they don't settle out. Sometimes they are visibly heterogenous to the naked eye, sometimes they are not (fog or milk), but they appear murky.
They have a very large total surface area which attracts other particles through various intermolecular forces.
They are commonly classified by the physical state of dispersed and dispersing substances
forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the intramolecular forces, the forces which keep a molecule together, (a bond BETWEEN molecules or compounds).
attractive forces within a compound. (i.e. Ionic and Covalent bonds)
An uneven sharing of electron pairs (this is a intRAmolecular bond). All ionic compounds are polar, covalent compounds can be polar or non-polar. When 2 atoms have an electronegativity difference of .5 or greater (the electron pair is held closer to the nucleus of one of the atoms) - (there is a remaining electron pair; and uneven sharing).
A covalently bonded molecule is polar when
the difference in electronegativity between the 2 atoms is .5 or greater (this means there is a significant difference in the sharing of electrons, making the sharing unequal. This means the molecules will be drawn to another compound with higher electronegativity).
A covalently bonded molecule is non polar when
the difference in electronegativity between the 2 atoms is .4 or less (this means there is an equal sharing of electron pairs between the two molecules).
Will result when 2 nonmetal atoms have a difference in electronegativity of .4 or less. When the electron pairs are of two atoms are equidistant from each other.
Hydrogen Bond (H Bonds)
General: (IntERmolecular bond) Very strong strength between molecules/compounds. They occur between covalent POLAR molecules. Occur when hydrogen of one molecule is directly bonded to the Fluorine (F), Oxygen (O), or Nitrogen (N), atom of the neighboring molecule.
In Solutions: The principal force in solutions of polar, O- and N- containing organic and biological compounds, such as alcohols, amines, and amino acids.
General: (IntERmolecular bond) Medium strength between molecules/compounds. Occurs when there is an attraction between two polar covalent molecules, and the positive end of one polar molecule is attracted to the negative end of another polar molecule.
In Solutions: In the absence of Hydrogen Bonding, allow polar molecules to dissolve in polar solvents.
London Forces (Dispersion)
(IntERmolecular bond) Weak strength between molecules/compounds. These bonds occur between two molecules that are non-polar (there is an equal sharing of electron pairs).
In solutions: contribute to the solubility of all solutes in all solvents, but they are the principal intermolecular force in solutions of non polar substances (ex. petroleum and gasoline).
The substance that dissolves in the solvent
The substance in which the solute (s) dissolve. When the solute dissolves in the solvent, a solution is formed. The solvent is usually the most abundant component.
matter which has a specific composition and specific properties. Every pure element is a substance. Every pure compound is a substance.
A pure compound
A compound is a pure substance composed of two or more different atoms chemically bonded to one another. A compound can be destroyed by chemical means. It might be broken down into simpler compounds, into its elements or a combination of the two (it is not a mixture because they are chemically bonded, not physically).
When the solute and solvent are soluble in each other in any proportion and the solute and solvent lose their meaning (a homogeneous mixture).
What usually determines the physical state of the solution?
The physical state of the solvent.
in relation to the solute, the solubility is the maximum amount that dissolves in a fixed quantity of a given solvent at a given temperature, when an excess of the solute is present (different solutes have different solubilities).
- Sodium chloride (NaCl), S = 39.12 g/100. mL water at 100. *C (concentrated)
- Silver chloride (AgCl), S = 0.0021 g/100. mL water at 100. *C (dilute)
Solubility is a QUANTITATIVE (based on a numbers, and referring to the relative amounts of dissolves solute) term [dilute and concentration are qualitative].
Why might a solute dissolve in one solvent but not another?
Depends on the intermolecular force strengths within both the solute and the solvent as well as between them.
Like dissolves like
Used as a general rule of thumb to predict is a solute will dissolve in a solvent. It states that substances with similar types of intermolecular forces (hydrogen bond, dipole-dipole bond, or dispersion), will dissolve in each other.
Ion-dipole forces in solutions (occurs in solutions and pure substances)
(dipole-dipole forces in general: attraction between positive and negative ends of two polar covalent molecules)
In Solutions: The principle force involved when an ionic compound (balanced) dissolves in water. Two events occur simultaneously, forces compete and hydration shells.
Forces Compete (Ion-dipole forces in solutions)
(occurs in Ion-dipole forces in solutions) when the solute ions attract the oppositely charged poles of the solvent ions. These attractions compete with and over come attractions between the ions, and the crystal structure breaks down (i.e. soluble salt in water).
A chemical compound in which ions are held together in a structure by electrostatic forces termed ionic bonding.
A material that is composed of only one type of particle; examples of a pure substance include gold, oxygen and water.
A material made up of at least two different pure substances.
Hydration shells form (Ion-dipole forces in solutions).
When a solute ion separates and water molecules (the solvent) cluster around it. The number of water molecules in the innermost shell depends on the ion's size. In the innermost shell, hydrogen bonding is disrupted to form the ion-dipole forces. The innermost shell of water molecules are bonded to the ion via ion-dipole forces (positive end attracting negative end), but the innermost shell of water molecules are then H-bonded to the next shell of water molecules (weak bonding. Hydrogen bonding continues to further away shells).
Ion-induced dipole forces in solutions
a charge-induced dipole force (between positive and negative poles) that relies on polarizability. Arises when an ion's charge distorts the electron cloud of a nearby non-polar molecule.
Example: this force initiates the binding of Fe2+ ion in hemoglobin to an O2 molecule entering a red blood cell.
Dipole-induced dipole forces in solutions
based on polarizability. This happens when a polar molecule distorts the electron cloud of a non-polar molecule. Weaker than ion-induced dipole forces because the charge of each pole is less than an ion's (Coulomb's Law).
The solubility in water of atmospheric O2, N2, and noble gases is due in part to these forces.
the ability for a molecule to be polarized. It is a property of matter. Polarizabilities determine the dynamical response of a bound system to external fields, and provide insight into a molecule's internal structure.
a law of physics describing the electrostatic interaction between electrically charged particles.
The higher the charge of an ion and the smaller its radius, the closer it gets to the oppositely charged pole of an H2O molecule, and the stronger the attraction.
Intermolecular forces in Solution
There are 6:
1. Ion-dipole forces (forces compete and hydration shells)
2. Hydrogen bonding
3. Dipole-dipole forces (absence of H-bonding. Between polar molecules)
4. Ion-induced dipole forces (ion's charge to non-polar molecule)
5. Dipole-induced forces (polar molecule to non-polar molecule)
6. Dispersion forces (solutions on non-polar substances).
polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole or multipole moment. Polar molecules interact through dipole-dipole intermolecular forces and hydrogen bonds
a pair of equal and oppositely charged or magnetized poles separated by a distance.
organic compounds that have duel polarity, a polar hydroxyl group (-OH) bonded to a non polar hydrocarbon group.
(-OH) interacts through strong hydrogen bonds with water.
When a gas dissolves in a solid, it occupies the spaces between the closely packed particles.
They diffuse so little, that their mixtures are typically heterogenous. Waxes are solid-solid solutions.
mixtures of elements that have a metallic character (solid-solid solutions).
Thermochemical solution cycle
1. Solute particles need to separate (endothermic - feels cold).
2. Solvent particles need to separate (endothermic- feels cold).
3. Solute and solvent particles need to mix together (exothermic - feels hot).
- These three enthalpy changes are combined to find the overall enthalpy (heat) of the solution [the total enthalpy change that occurs when solute and solvent form a solution).
When the final solution has a LOWER heat (H) than the initial solution (less than zero). This occurs in the solution process when the sum of the endothermic terms is smaller than the exothermic term.
When the final solution has a HIGHER heat (H) than the initial solution (greater than zero). This occurs in the solution process when the sum of the endothermic terms is greater than the exothermic term. If the heat of the solution is highly positive, it may not dissolve significantly in the solvent.
the process of surrounding a solute particle with solvent particles. It is the combined process of the enthalpy change of solvent and the enthalpy change of mix.
Hydration of an ion
(solvation in water) the enthalpy change for the hydration of 1 mol of separated gaseous ions.
The ratio of the ion's charge to it's volume. An ion's charge density determines trends in its heat of hydration (the higher the charge density, the more negative the heat of hydration).
Heat of hydration
based on trends in charge density and radial size.
-When looking down a group, the charge stays the same and the size increases, therefore the charge densities decrease as well as the heat of hydration values.
-When looking across a period, the ions charges increase, and the radius decreases, increasing the heat of hydration values.
For ions, heats of hydration depend on the ion's charge density but are always negative because ion-dipole forces are strong.
the energy required (enthalpy change) to separate a mole of an ionic solid into gaseous ions. It is highly positive always.
Two factors that determine whether a solute dissolves
1. Heat of Solution. (If the heat of is highly positive, the solute may not dissolve significantly in the solvent).
The natural tendency of a system of particles to spread out, which results in the systems kinetic energy becoming more dispersed or more widely distributed.
It is directly related to the number of ways a system can distribute its energy, which involves the freedom of motion of the particles.
The three states of matter all differ dramatically in their entropy.
Order of entropy from highest to lowest for three states of matter.
Sgas > Sliquid > Ssolid
The more freedom of motion the particles have, the more ways they can distribute their kinetic energy. There is a change of entropy associated with phase change, and it can be positive or negative.
Entropy (S) change when a liquid vaporizes.
Always positive. (Svap > 0)
Entropy (S) change when a liquid freezes (fusion).
Always negative. (Sfus < 0)
Entropy of solution
usually higher than the pure solute and pure solvent because the number of ways to distribute the energy is related to the number of interactions between different molecules. (Salon > Ssolute + Ssolvent).
Solution formation involves the interplay of two factors.
1. Systems change toward a state of lower enthalpy.
2. Systems change toward a state of higher enthalpy.
The relative sizes of Heat of Solution (Enthalpy) and Entropy of Solution determine if a solution will form.
If the sum of the endothermic terms (Hlattice and Hsolvent) is much larger than the exothermic (Hmix) term (Hsoln >> 0). [in solution formation]
A solution will not form because the entropy increase from mixing solute and solvent would be much smaller than the enthalpy increase required to separate the solute.
If a substance has a positive Hsoln
it only dissolves if Ssoln is larger than Hsoln.
a measure of the quantity of solute dissolved in a given quantity of solution (or solvent). It is an intensive property, meaning it is independent of the solution volume.
The unit for concentration is Molarity (M). M = moles of solute/liters of solution.
The proportion of a substance in a mixture.
Solubility as an equilibrium process: the concentration of a solution rises as long as
the rate of dissolving solute particles is greater than the rate of recrystallizing solute particles.
Solubility as an equilibrium process: the concentration of a solution remains constant when
at a given temperature, when solid solute particles are dissolving at the same rate they are recrystallizing (undissolved solute is in equilibrium with dissolved solute).
when a solution is AT equilibrium (rate of dissolving is the same as recrystallizing) and contains the maximum amount of dissolved solute at a given temperature in the presence of undissolved solute.
If you filter off the solution, and add more solute, the solute will not dissolve because the rate of recrystallization will be greater than the rate of dissolving particles.
a solution that contains LESS that the equilibrium concentration of dissolved solute.
You are able to add more solute and it will dissolve until the rates of dissolving and recrystallizing are equal (aka at equilibrium).
a solution that contains MORE than the equilibrium concentration and is unstable relative to the standard solution.
This can be prepared if the solute is more soluble at a higher temperature. When mixing the solution at a higher temperature, you are able to put more solute in the solvent and it will dissolve. Then cool the solution and the solute will still remain dissolved until a "seed" crystal of solute is added to the solution at a standard temperature. With the addition of the "seed" crystal, the excess solute will recrystallize immediately (when the solution was heated, it was a saturated solution, when it was cooled to standard temperature, it became supersaturated).
The effect of temperature on the solubility of solids in water
In order to fully know the effect of temperature on the solubility of solids in water, you need to know the sign of enthalpy change very close to the point of saturation. It is very unpredictable.
For the most part, ionic compounds are more soluble at higher temperatures.
The effect of temperature on the solubility of gases in water
(this process takes place in water, so Hsolvent and Hmix become Hhydr. And since gas particles [the solute] are already separated, the Hsolute is similar to 0. Hhydr is then exothermic, making the Hsoln negative). Hsoln is always negative (exothermic) for the solubility of gases in water at standard temperature.
With the increase of temperature (addition of heat), gas solubility decreases, because gases have weak intermolecular forces with water.
Effect of Pressure on Solubility
Pressure has little effect on the solubility of liquids and solids because they are almost incompressible. But it has a major effect on the solubility of gases.
Effect of pressure on solubility: as the gas volume decreases..
1. Pressure (concentration) increase.
2. More gas particles enter than leave the solution per unit of time.
3. Shift to the right in the equation, until the system re-establishes equilibrium.
the solubility of gas (Sgas) is directly proportional to the partial pressure of the gas (Pgas) above the solution.
Henry's Law expresses the quantitative relationship between gas pressure and solubility.
Henry's Law constant
(kH) - it is specific for a given gas-solvent combination as a given temperature.
equation: Sgas (mol/L) = kH (mol/ L * atm) x Pgas (atm)
Molarity, Molality, Parts by mass, Parts by volume, Mole fraction
a ratio of: amount (mol) of solute / volume (L) of solution
it is the number of moles of solute dissolved in 1 L of solution. DRAWBACKS: effect of temperature and effect of mixing.
a ratio of: amount (mol) of solute / mass (kg) of solvent
it is the number of moles of solute dissolved in 1000 g (1 kg) of solvent. ADVANTAGES: effects of temperature (mass does not change with temperate) and effects of mixing (you can add masses, unlike volumes).
The preferred term when temperature and density change in a study of physical properties.
Parts by mass
a ratio of: mass of solute / mass of solution
mass percent, % (w/w). (mass of solute dissolved in 100. parts by mass of solution).
there are parts per million (ppm - multiply by 10^6) and parts per billion (ppb - multiply by 10^9) by mass.
Parts by volume
a ratio of: volume of solute / volume of solution
volume percent, % (v/v). (the volume of solute in 100. volumes of solution).
Parts by Mole: Mole fraction (X)
a ratio of: amount (mol) of solute / amount (mol) of solute + amount (mol) of solvent
gives the proportion of solute (or solvent) particles in a solution. There is also a mole percent, which is the mole fraction multiplied by 100.
a ratio of solute weight (actual mass) to solution volume.
1. Vapor Pressure Lowering
2. Boiling Point Elevation
3. Freezing Point Depression
4. Osmotic Pressure
These are all influenced by the NUMBER of solute particles, NOT their chemical identity,
Conducts a current because the solute separates into IONS as it dissolves. There are strong electrolytes and weak electrolytes.
1. Strong electrolytes dissociate completely (soluble salts, strong acids, and strong bases).
2. Weak electrolytes dissociate very little, so their solutions conduct currents poorly (weak acids and weak bases).
Solutions that do not conduct a current. The solute compounds do not dissociate into ions at all.
SO; 1 mol of compounds yields 1 mol of particles when it dissolves in solution.
Colligative properties of three types of solutes. These solutes are.
nonvolatile non-electrolytes, volatile non-electrolytes, and strong electrolytes.
Solutions that contain solutes that do not dissociate (not ionic), and they have negligible vapor pressure at the boiling point of the solvent.
Colligative property: Vapor Pressure Lowering for a nonvolatile non-electrolyte
the vapor pressure of a nonvolatile non-electrolyte solution is always lower than the vapor pressure of the pure solvent. The difference between these two pressures is called the vapor pressure lowering (P).
volatility is the tendency of a substance to vaporize. Volatility is directly related to a substance's vapor pressure. At a given temperature, a substance with higher vapor pressure vaporizes more readily than a substance with a lower vapor pressure.
nonvolatile = not vaporizing ready.
the vapor pressure of a solvent about a solution equals the mole fraction of the solvent times the vapor pressure of the pure solvent.
Solutions always follow Raoult's Law at any concentration.
Colligative Property: Boiling Point Elevation
A solution boils at a higher temperature than the pure solvent. It is proportional to the concentration of a solute.
Boiling point of a liquid (Tb)
is the temperature as which its vapor pressure equals the external pressure.
Kb and Kf
molal boiling point elevation constant and molal freezing point depression. It has units of degrees celsius per molal unit (C/m) and is specific for a given solvent.
Freezing Point Depression
A solution freezes at a lower temperature than the pure solvent, also a result of vapor pressure lowering. The freezing point of a solution is the temperature at which its vapor pressure equals that of the pure solvent, when solid solvent and solid liquid solution are in equilibrium. The freezing point depression occurs because the vapor pressure of the solution is always lower than that of the solvent, so the solution freezes at a lower temperature.
It is proportional to the molal concentration of solute.
Colligative Property: Osmotic Pressure
observed when solutions of higher and lower concentrations are separated by a semipermeable membrane, one that allows solvent, but not solute, to pass through. It is measured when the two solutions are at equilibrium. It is proportional to the number of solute particles in a given solution volume (molarity - M).
Osmosis: a net flow of solvent into he more concentrated solution, causing a pressure difference (osmotic pressure).
II = MRT
Equation for finding osmotic pressure. R = the proportionality constant (0.0821 atm
L / mol
K). T = the absolute temperature. This is similar to the ideal gas law equation (P = nRT/V) because both relate the pressure of a system to its concentration and temperature.
Themes of colligative properties for nonvolatile solutes (solutes that are not ready to vaporize).
Each colligative property is proportional to solute concentration (number of moles of solute per volume of solution).
Each colligative property arises because the solute particles cannot move between two phases.
1. Vapor pressure lowering and boiling point elevation are a result of the solute being unable to enter the gaseous phase. (when having a low vapor pressure lowering, more heat is needed to vaporize particles).
2. Freezing point depression is a result of the solute being unable to enter a solid phase. (a solution has a lower freezing point than the pure solvent because it has a higher entropy).
3. Osmotic pressure is a result of the solute being unable to cross a semipermeable membrane. (the pure solvent flows down the concentration gradient until both solutions are at equilibrium, the rate of solvent flowing in is equivalent to the rate os solvent flowing out).
In all three themes, the presence of solute decreases the mole fraction of solvent (mole fraction must equal 1, so as moles of solute increase, moles of solvent must decrease).
Colligative properties for Volatile Non-electrolyte Solutions
when the vapor consists of solute and solvent molecules. Using Raoult's Law for a volatile non-electrolyte is the mole fractions of solvent or solute in the liquid phase.
The total vapor pressure is the sum of the partial vapor pressures.
The presence of each volatile component lowers the vapor pressure of the other by making the mole fraction less than 1.
The vapor has a higher mole fraction of the more volatile component.
Colligative properties for Strong Electrolyte Solutions
the solute formula tells us the number of particles.
Van't Hoff factor (i)
multiplying factor used for strong electrolyte solutions when dealing with colligative properties ( i = measured value for electrolyte solution / expected value for non electrolyte solution), i is plugged into the same colligative property equation for nonvolatile non-electrolytes.
3 types of mixtures
1. Suspensions (heterogeneous mixture with particles large enough to be visibly distinct from the surrounding solution).
2. Solutions (a homogeneous mixture in which particles are invisible, and individual molecules are distributed equally throughout the solution).
3. Colloids (a heterogeneous mixture where the particles are not large enough to settle out but not small enough to dissolve).
When a scattered light beam appears broader than one passing through a solution. This is experienced with colloids (ex. light passing through dust particles).
an erratic change of speed and direction. This is experienced with colloids, because of the collisions of the particles with molecules of the dispersing medium.
Stabilizing and destabilizing colloids
Colloidal particles dispersed in water have charged surfaces that stabilize the colloid through ion-dipole forces.
formed by molecules with dual polarities. When the charged heads are on the exterior and the hydrocarbon tails are on the interior.
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