Aqueous Reactions and Solution Stoichiometry

1. Aqueous Reactions and Solution Stoichiometry Chapter 4 BLB 12th

2. 4.1 General Properties of Aqueous Solutions • Solution – homogeneous mixture (ch. 13) • Solvent – dissolving medium; aq-water • Solute – dissolved substance Electrolytic Properties • Electrolyte – a substance whose aqueous solution contains ions; conducts electricity • Nonelectrolyte – substance that does not form ions in solution

3. Aqueous Solutions, cont. Ionic Compounds in Water • Ionic solids dissociate (or ionize) into ions as they dissolve. • hydration – process of dissolving an ionic substance in water • solvation – dissolving in any solvent; dissolution Why? Water is polar.

4. Polarity of molecules • Electrons are shared unequally. • Results in partial charges (δ), and a… • Dipole moment

5. Hydration of NaCl(s)

6. How many ions does an ionic compound produce when it dissociates? KCl, MgCl2, or K2SO4?

7. Aqueous Solutions, cont. Molecular Compounds in Water • nonelectrolytes – contain only molecules (no ions); do not dissociate; do not conduct electricity; may dissolve in water • Most molecular compounds are nonelectrolytes. • Some may have strong interaction with water (alcohols). • Some may dissociate (acids).

8. Methanol (CH3OH) in water

9. Aqueous Solutions, cont. Strong and weak electrolytes – depend on the extent of dissociation • Strong – completely dissociate into ions - all water-soluble ionic compounds, strong acids & bases • Weak – remain mostly as neutral molecules and produce very few ions; establish chemical equilibrium - weak acids (like acetic acid) and weak bases (like amines); water Note: strong doesn’t mean soluble and vv.

10. 4.2 Precipitation Reactions • Marked by the formation of an insoluble product (precipitate) • Solubility – amount of solute that can be dissolved in a given amount of solvent at a certain temperature; g/100g or g/L or mol/L • Insoluble – solubility < 0.01 mol/L • Solubility Rules – Table 4.1, p. 121 • Metathesis (or exchange) reactions

11. Note: All common compounds of Group I metals and NH4+ are soluble in water.

12. Metathesis (or exchange) reactions Molecular: BaCl2(aq) + Na2SO4(aq) → 2 NaCl(aq) + BaSO4(s) Complete ionic: Net ionic:

13. Metathesis (or exchange) reactions Molecular: NaI(aq) + Pb(NO3)2(aq) → Complete ionic: Net ionic:

14. Metathesis (or exchange) reactions Molecular: NaOH(aq) + Co(NO3)2(aq) → Complete ionic: Net ionic:

15. 4.3 Acids, Bases, and Neutralization Reactions • Involve H+ • Acid – H+ donor Base – H+ acceptor • Neutralization: acid + base → salt + water HCl + NaOH → NaCl + H2O

18. 4.3 Acids, Bases, and Neutralization Reactions Strong and Weak • Strong – completely dissociate • Weak – only partially ionize • Neutralization: acid + base → salt + water HCl + NaOH → NaCl + H2O

22. 4.4 Oxidation-Reduction (Redox) Reactions • Involve transfer of e¯ • Oxidation – loss of e¯ Reduction – gain of e¯ • Oxidation number – a “charge” assigned to an atom to keep track of electrons transferred during redox • Displacement reaction – ion in solution is replaced through oxidation of an element.

25. 4.5 Concentrations of Solutions • Molarity (M) – mole solute/L solution M • V = mol • Dilution – adding solvent to decrease concentration M1V1 = M2V2 mol1 = mol2; only volume changes

26. Calculate the concentration (in M) if 2.50 g (NH4)2SO4 is dissolved in enough water to form 250 mL of solution.

27. How many grams of K2Cr2O7 are needed to make 50.0 mL of 0.850 M solution?

28. Ion Concentration: 0.850 M K2Cr2O7 • Concentration (M) of Cr2O72-? • Concentration (M) of K+?

29. What volume (in mL) of 6.0 M HNO3 is needed to make 250 mL of 1.0 M HNO3?

30. 4.6 Solution Stoichiometry and Chemical Analysis • Use M and volume to obtain moles • Titration – process used to determine the concentration of a solution (p. 145 ff) • Standard solution – one of precisely known concentration • Analyte – solution of unknown concentration

31. Stoichiometry Overview, p. 144