Presentation on theme: "UNIT 7: REACTIONS IN SOLUTION. Chemical Reactions in Solutions Most reactions occur in aqueous solution SOLUTE is the substance to be dissolved in solution."— Presentation transcript:
UNIT 7: REACTIONS IN SOLUTION
Chemical Reactions in Solutions Most reactions occur in aqueous solution SOLUTE is the substance to be dissolved in solution SOLVENT is the substance (often a liquid) the solute dissolves in The concentration of the solution is Molarity (M) = moles solute L solution
Chemical Reactions in Solutions The maximum amount of solute that can dissolve in a fixed amount of solvent is the solubility. Salt (NaCl) and sugar (C 12 H 22 O 11 ) both dissolve easily in water; salt and sugar are soluble in water Oil does not dissolve in water at all, so it is insoluble in water.
Chemical Reactions in Solutions The composition of a solution is given by its concentration. molarity = M= moles solute L solution This tells us what the solution is made of, but what happens to the solute after it is added to the solvent?
Dissolving Process Resulting solutions have unique properties: Taste Conductivity (ability to conduct electricity) Non-electrolyte Electrolyte
How do we explain these behaviors? Remember… Molecular substances are made of molecules, share electrons (covalent bond), composed of nonmetals C 12 H 22 O 11 (s) C 12 H 22 O 11 (aq) Ionic substances are made of cations (+) and anions (-), transfer electrons, ionic substances are salts NaCl (s) Na 1+ (aq) + Cl 1– (aq) Ions that make up crystal separate (dissociate); move freely throughout liquid
Are solids the only things that are soluble? Nope. Gases are molecular substances that may dissolve in water: CO 2 (g) CO 2 (aq) NOTE: Most gases are not very soluble in water, but you can increase gas solubility by increasing the pressure of the gas. Think Soda. Think Scuba Divers.
Trends in dissolving Most solid substances are more soluble at higher temperature. Increasing Temp.= Increasing KE= faster dissolving Gases show the opposite trend: gases are less soluble at higher temperature. Increasing Temp.= Increasing KE= Evaporation
Trends in dissolving When you have reached to limit of a solutes solubility, and no more solute can dissolve, the solution is saturated.
Colligative Properties of aqueous solutions Aqueous solutions have different physical characteristics than just pure water. Electrolyte vs. Non-electrolyte (conductivity) depends on what kind of solute is dissolved. Colligative properties- the effect of adding solute depends on how many solute particles are present, but not what kind of particles they are. These properties include, melting pt., boiling pt., freezing pt.
Freezing Point Depression Aqueous solutions always freeze or melt at a lower temperature than pure water does. Presence of solute interferes w/ crystal formation Identity of the solute does not matter –– only the number of solute particles dissolved in the water. The freezing pt. of H 2 O decreases by 1.86 °C for every mole of particles dissolved in a liter of water.
Heres a scenario… You dissolve 1 mole of sugar (a molecular solute, C 12 H 22 O 11 ) in a liter of water, the solution freezes at –1.86 °C, You dissolve 1 mole of salt (an ionic solute, NaCl) in 1 liter of water, the solution freezes at –3. 6°C. What happened here? Why doesnt the water freeze at the same temp? Remember…freezing pt. decreases 1.86 °C for every 1 mole.
Boiling Point Elevation A non-volatile solute (one that doesnt evaporate easily, like salt or sugar) lowers the vapor pressure of the solution. So, a solution needs to reach a higher temperature than a solvent does in order to boil. This is BP elevation. The boiling pt. of H 2 O increases by 0.54 °C for every mole of particles in a liter of water.
New Freezing pt New Boiling pt
Precipitation Reactions and Net Ionic Equations In precipitation rxns., solutions that contain dissolved ions are mixed and at least one new, insoluble salt (a precipitate) forms, which settles out of the mixture. Example: Pb(NO 3 ) 2 (aq) and KI (aq) are both clear, colorless solutions.
Precipitation Reactions and Net Ionic Equations Precipitation reaction is also known as a double replacement reaction, in which the substances change partners. Pb(NO 3 ) 2 (aq) + 2KI (aq) PbI 2 (s) + 2 KNO 3 (aq) How does this happen?
Precipitation Reactions and Net Ionic Equations When each salt dissolves in water, its ions dissociate and move about freely in the solution: Pb(NO 3 ) 2 (s) Pb 2+ (aq) + 2 NO 3 1– (aq) KI (s) K 1+ (aq) + I 1– (aq) Pb 2+ (aq) + 2 NO 3 1– (aq) + 2 K 1+ (aq) + 2 I 1– (aq) PbI 2 (s) + 2 NO 3 1– (aq) + 2 K 1+ (aq) Pb 2+ (aq) + 2 I 1– (aq) PbI 2 (s)
Lets Practice…But first, an example BaCl 2 (aq) + Na 2 SO 4 (aq) BaSO 4 (s) +2 NaCl (aq) Each aqueous salt is dissociated into its component ions. The solid precipitate is not dissociated: Ba 2+ (aq) + 2 Cl 1– (aq) + 2 Na 1+ (aq) + SO 4 2– (aq) BaSO 4 (s) + 2 Na 1+ (aq) + 2 Cl 1– (aq) The Na 1+ and Cl 1– ions did not change during the reaction and are spectators. We cross them out: Ba 2+ (aq) + SO 4 2– (aq) BaSO 4 (s)
Lets Practice… Write the net ionic equation for this reaction: AgNO 3 (aq) + KCl (aq) KNO 3 (aq) + AgCl (s)
Rules You should be able to predict the products of a precipitation reaction and write its net ionic equation: 1. Precipitation is a double replacement reaction. To form the products, switch the ion partners. 2. When writing products, be attentive to ion charges. When Na 1+ and Cl 1– combine, the formula is NaCl, not Na 2 Cl 2. Also, dont make up impossible combinations like Na 1+ with Ba 2+ ! Ions with like charges will not stick together. 3. It is not necessary to balance the equation before identifying the precipitate and crossing out spectators, and it is easier to balance after removing spectator ions. 4. Salts containing NH 4 1+, Na 1+, K 1+, or NO 3 1– ions (water babies) are soluble, so the precipitate is not going to be a salt containing any of these ions. Mark the combination that is soluble (aq). The other combination must be the precipitate; mark it (s). 5. Identify the spectator ions, cross them out, and write the net ionic equation. Balance the net ionic equation if necessary.