Solution Chemistry

Solutions Solution chemistry- where water is the solvent (called aqueous solutions). Solutions—homogeneous mixture of two or more substances in which dissolved substance is ionized, its composition is the same throughout its volume. Usually a solid is dissolved in a liquid, but liquids can be dissolved in other liquids and gases can be dissolved in a liquid as well.

Solutions Solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent.

Solubility? The amount of a substance that dissolves in a given volume of solvent at a given temperature. Common units: grams solute/ 100 grams of solvent or grams solute/ 100mL of water

Solutions How does a solid dissolve into a liquid? What ‘drives’ the dissolution process? What are the energetics of dissolution?

How Does a Solution Form? 1.Solvent molecules attracted to surface ions. 2.Each ion is surrounded by solvent molecules. 3.Enthalpy (  H) changes with each interaction broken or formed. Ionic solid dissolving in water

How Does a Solution Form? 1.Solvent molecules attracted to surface ions. 2.Each ion is surrounded by solvent molecules. 3.Enthalpy (  H) changes with each interaction broken or formed.

How Does a Solution Form The ions are solvated (surrounded by solvent). If the solvent is water, the ions are hydrated. The intermolecular force here is ion-dipole.

Factors Affecting Solubility Solute-solvent interactionSolute-solvent interaction - Like dissolves like Temperature Factor -Temperature Factor - i) Solids/Liquids- Solubility increases with Temperature. Increase K.E. increases motion and collision between solute / solvent. ii) gas - Solubility decreases with Temperature Increase K.E. result in gas escaping to atmosphere. Pressure Factor -Pressure Factor - i) Solids/Liquids - Very little effect. Solids and Liquids are already lose together, extra pressure will not increase solubility. ii) gas - Solubility increases with Pressure.Increase pressure squeezes gas solute into solvent.

Properties of Water 2-2-  + Water is polar Water forms hydrogen bonds When water freezes it expands

hydrogen bonding is the weak intermolecular bond between the H end of one molecule and the O, N, or F end of another molecule 2-2-  + 2-2-

Polar liquids tend to dissolve in polar solvents. “Likes dissolve likes” Miscible liquids: mix in any proportions. water and ethanol are miscible- broken hydrogen bonds in both pure liquids are re-established in the mixture. Immiscible liquids: do not mix. The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water. Solute-Solvent Interactions

The number of -OH groups within a molecule increases solubility in water. The more polar bonds in the molecule, the better it dissolves in a polar solvent. The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non- polar solvent. The greater the contact between the solute and solvent, the faster the substance will dissolve. Therefore, stirring and crushing, which increases the surface area, increase the rate of solution Solute-Solvent Interactions Cont.

Ionic Solutes

Polar water molecules interacting with positive and negative ions of a salt

+- Positive end of the water molecule attracts to the anion (negative ion) Negative end of the water molecule attracts to the cation (positive ion). Making it a strong electrolyte..

Electrolytes Compounds- dissolved in water conduct an electric current. Strength correlates with number of ions in solution. Strong electrolytes conduct a much stronger current because have more ions Nonelectrolytes Do not conduct an electric current because do not contain ions.

Electrolytes Strong Weak Non Electrolyte Electrolyte Electrolyte

Polar Molecules

Sugar is a molecular solid with weak bonds. It does not form ions and therefore does not ionize in solution. Meaning sugar cannot act as an electrolyte.

In order for a substance to dissolve in water, the water molecules must be more attracted to the new substance added. Remember water molecules are weakly bonded to each other by hydrogen bonds. Will this substance dissolve in water? Yes, because water will be attracted to it.

Substances that don’t dissolve are called insoluble E.g. Petroleum (crude oil), which are non-polar So if you want to dissolve grease which is non-polar, you need to use a non-polar solvent. Petroleum in a non-polar organic molecule

If solution process absorbs energy- solubility will be INCREASED as the temperature is increased. If the solution process releases energy- solubility will DECREASE with increasing temperature. Temperature-

Absorption of energy and increase in solubility: sugar dissolves better in warm water than cold. As temperature increases, solubility of solids generally increases. Exception: Ce 2 (SO 4 ) 3 Temperature Cont.

Gases and liquids: gases get less soluble as temperature increases. Example: carbonated beverages Thermal pollution: lakes get too warm, CO 2 and O 2 become less soluble and are not available for plants or animals. Temperature Effects:

Pressure:

Higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution. –higher the pressure, the greater the solubility. Pressure Effects:

Carbonated beverages are bottled with a partial pressure of CO 2 > 1 atm. Opening bottle, the partial pressure of CO 2 decreases and the solubility of CO 2 decreases. bubbles of CO 2 escape from solution. Pressure Effects:

Degree of saturation Saturated solution  Solvent holds as much solute as is possible at that temperature.  Undissolved solid remains in flask.  Dissolved solute is in dynamic equilibrium with solid solute particles.

Degree of saturation Unsaturated Solution  Less than the maximum amount of solute for that temperature is dissolved in the solvent.  No solid remains in flask.

Degree of saturation Supersaturated  Solvent holds more solute than is normally possible at that temperature.  These solutions are unstable; crystallization can often be stimulated by adding a “seed crystal” or scratching the side of the flask.

Degree of saturation Unsaturated, Saturated or Supersaturated?  How much solute can be dissolved in a solution?

Saturated Solutions

Saturated/unsaturated solutions

To dissolve 120 g the temp must be raised to 80 o C at 50 o 88.0 g of KNO 3 will dissolve

1. Which substance’s solubility increases the most with temperature? 2. Which substance’s solubility changes the least with temperature? 3. What is the solubility of KI at 8 o C? 4. What temperature is needed to dissolve 160 g of potassium iodide, KI, in 100 g of water?

Colloids: Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity.

Tyndall Effect Colloidal suspensions can scatter rays of light. This phenomenon is known as the Tyndall effect.

Colloids in Biological Systems Some molecules have a polar, hydrophilic (water- loving) end and a nonpolar, hydrophobic (water-hating) end.

Colloids in Biological Systems Sodium stearate is one example of such a molecule.

Colloids in Biological Systems These molecules can aid in the emulsification of fats and oils in aqueous solutions.

Ways of Expressing Concentrations of Solutions

Concentration Mass Percent or Volume Percent Mole Fraction Molarity (M) Molality (M)

Mass Percentage Mass % of A = mass of A in solution total mass of solution  100 Volume of A in solution Total volume of solution Volume % of A =  100

Parts per Million and Parts per Billion (for mass or volume) ppm = mass of A in solution total mass of solution  10 6 Parts per Million (ppm) Parts per Billion (ppb) ppb = mass of A in solution total mass of solution  10 9

moles of A total moles in solution X A = Mole Fraction (X) A could be either solvent or solute – make sure you find the quantity you need!

mol of solute L of solution M = Molarity (M) Concentration of solution; or its strength Can change with temperature.

mol of solute kg of solvent m = Molality (m) Not temperature dependent- Mass and moles are not dependent on temperature

Changing Molarity to Molality

Problem Give the concentration (molarity of grams of NaCl in 545 mL of water) Molar mass of NaCl = g/mol

Honors: Problem Give the concentration (in molality) of grams of NaCl in 545 mL of water Density of 25°C = g/mL MW of NaCl = g/mol

Honors Problem A 10.7 molal solution of NaOH has a density of 1.33 g/cm 3 at 20°C. molar mass of NaOH = g/mol and molar mass of H 2 O = g/mol. – Calculate the mole fraction of NaOH, the weight percentage of NaOH and the molarity of the solution.

Honors practice An aqueous solution of NaCl is created using 133 g of NaCl diluted to a total solution volume of 1.00 L. – Calculate the molarity, molality, and mass percent of the solution – Density of 1.08 g/mL – Molar mass of NaCl = g/mol.

Acid/Base Chemistry

Acids from the Latin word acere  “sour” taste sour (but you wouldn’t taste an acid to see) change litmus paper red corrosive to some metals (reacts to create hydrogen gas – H 2 ) Donates a hydrogen ion (H + ) to another substance Create a hydrogen ion (H + ) or hydronium ion (H 3 O + ) when dissolved in water HCl  H + + Cl - Hydrochloric Acid Hydrogen ion Chloride ion Examples: hydrochloric acid, vinegar, lemon juice, rainwater H2OH2O Notice how the hydrogen ion is released when the acid is in water

Bases (Alkalis) taste bitter feel slippery or soapy change litmus paper blue react with oils and grease- soaps Accept a hydrogen ion (H + ) create a hydroxide ion (OH - ) when dissolved in water Examples: sodium hydroxide, Drano, Tums, baking soda NaOH  Na + + OH - Sodium Hydroxide Sodium ion Hydroxide ion H2OH2O Notice how the hydroxide ion is released when the base is in water; this ion can accept a hydrogen ion (H + )

Neutralization Reaction occurs when acids and bases react with each other to produce water and salt – acids release a hydrogen ion (H + ) and bases release a hydroxide ion (OH - )  water (H 2 O) – the negative ion from the acid joins with the positive ion of a base  salt HCl + NaOH  H 2 O + NaCl Hydrochloric Acid (acid) Sodium Hydroxide (base) Water Sodium Chloride (salt) Both the salt and water are neutral substances; therefore, that is why this is referred to as a neutralization reaction.

Acid, Base, or Neutralization? Zn + 2H +  Zn 2+ + H 2 NH 3 + H 2 O  NH OH - HClO + LiOH  LiClO + H 2 O HCl + H 2 O  H 3 O + + Cl - Acid – because H 2 gas was given off Acid – because H 3 O + is present in the products Base – because OH - is present in the products Neutralization – because of the salt and water in the products

Who Theory: Acid= When Arrheniusincreases H ’s Brønstedproton donor1923 Lowry ditto1923 Lewis Electron-pair acceptor 1923 Three definitions of acid

Some Definitions Arrhenius acids and bases – Acid:Substance that, when dissolved in water, increases the concentration of hydrogen ions (protons, H + ). – Base:Substance that, when dissolved in water, increases the concentration of hydroxide ions.

Some Definitions Brønsted–Lowry: must have both 1. an Acid:Proton donor and 2. a Base:Proton acceptor

The Brønsted-Lowry acid donates a proton (H + ion), while the Brønsted-Lowry base (H + ion) accepts it. Brønsted-Lowry Acids and Bases: Which is the acid and which is the base in each of these rxns?

A Brønsted–Lowry acid… …must have a removable (acidic) proton. HCl, H 2 O, H 2 SO 4 A Brønsted–Lowry base… …must have a pair of nonbonding electrons. NH 3, H 2 O

Types of Proton acceptors and donators: Monoprotic acid- Donates 1 proton (H + ) – HCl, HF, HI, HClO 3 Diprotic acid- Donates 2 protons (2H + ) – H 2 S, H 2 SO 4 Triprotic acid- Donates 3 protons (3 H + ) – H 3 PO 4

If it can be either…...it is amphiprotic. HCO 3 – HSO 4 – H2OH2O

What Happens When an Acid Dissolves in Water? Water acts as a Brønsted–Lowry base and abstracts a proton (H + ) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.

Conjugate Acids and Bases: From the Latin word conjugare, meaning “to join together.” Reactions between acids and bases always yield their conjugate bases and acids.

Acid and Base Strength Strong acids are completely dissociated in water. – Their conjugate bases are quite weak. Weak acids only dissociate partially in water. – Their conjugate bases are weak bases.

Acid and Base Strength Substances with negligible acidity do not dissociate in water. – Their conjugate bases are exceedingly strong.

Autoionization of Water As we have seen, water is amphoteric. In pure water, a few molecules act as bases and a few act as acids. This process is called autoionization.

Ion-Product Constant The equilibrium expression for this process is K c = [H 3 O + ] [OH – ] This special equilibrium constant is referred to as the ion-product constant for water, K w. At 25°C, K w = 1.0 

pH pH is defined as the negative base-10 logarithm of the hydronium ion concentration. pH = –log [H 3 O + ]

pH In pure water, K w = [H 3 O + ] [OH – ] = 1.0  Because in pure water [H 3 O + ] = [OH - ], [H 3 O + ] = (1.0  ) 1/2 = 1.0  10 -7

pH Therefore, in pure water, pH = –log [H 3 O + ] = –log (1.0  ) = 7.00 An acid has a higher [H 3 O + ] than pure water, so its pH is <7 A base has a lower [H 3 O + ] than pure water, so its pH is >7.

pH These are the pH values for several common substances.

Other “p” Scales The “p” in pH tells us to take the negative log of the quantity (in this case, hydronium ions). Some similar examples are – pOH –log [OH - ] – pK w –log K w

Watch This! Because [H 3 O + ] [OH − ] = K w = 1.0  , we know that –log [H 3 O + ] + – log [OH − ] = – log K w = or, in other words, pH + pOH = pK w = 14.00

How Do We Measure pH? – Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5 – An indicator Compound that changes color in solution.

How Do We Measure pH? pH meters measure the voltage in the solution