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Solubility of Ionic Compounds

Solubility of Ionic Compounds. Solution a homogeneous mixture of two or more substances Solute the substance we dissolve in a solvent Solvent the liquid we use to dissolve a substance. Solubility of Ionic Compounds. Aqueous solution a solution where solute is dissolved in water.

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Solubility of Ionic Compounds

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  1. Solubility of Ionic Compounds • Solution • a homogeneous mixture of two or more substances • Solute • the substance we dissolve in a solvent • Solvent • the liquid we use to dissolve a substance

  2. Solubility of Ionic Compounds • Aqueous solution • a solution where solute is dissolved in water

  3. Solubility of Ionic Compounds • Electrolytes • dissolves in water • the ability to dissociate • formation of ions • conducts electricity

  4. Solubility of Ionic Compounds • Nonelectrolytes • dissolves in water • does not dissociate • does not conduct electricity

  5. Solubility Rules • The common salts of the alkali metals (group IA) and the ammonium ion (NH4+) are soluble. • Salts of nitrate (NO3–), chlorate (ClO3–), perchlorate (ClO4–), and acetate (CH3COO–) anions are soluble. • All chlorides, bromides, and iodides are soluble, except those of Ag+, Pb2+, and Hg22+. • Sulfates (SO42–) are soluble except those of Ba2+, Pb2+, Hg2+, and Hg22+. • Most metal hydroxides are insoluble. The exceptions are the hydroxides of the alkali metals and the heavier alkaline earth metals (Ca, Sr, Ba). • All carbonates (CO32– ) and phosphates (PO43– ) are insoluble, except those of the group IA metals and NH4+ ions.

  6. Precipitation Reactions • A precipitation reaction involves the formation of an insoluble product or products from the reaction of soluble reactants.

  7. Net Ionic Equations • The overall equation shows all of the reactants and products in undissociated form. • The complete ionic equation shows separately all species – both ions and molecules – as they are present in solution.

  8. Net Ionic Equations • The net ionic equation shows only those species in the solution that actually undergo a chemical change. • Spectator ions do not participate in the reaction and undergo any change.

  9. Net Ionic Equations • Write the net ionic equation for the reaction in water of hydrobromic acid, HBr, and potassium hydroxide, KOH. Write the overall equation for the reaction HBr(aq) + KOH(aq) H2O(l) + KBr(aq)

  10. Net Ionic Equations Write the complete ionic equation H+(aq) + Br–(aq) + K+(aq) + OH–(aq) H2O(l) + K+(aq) + Br–(aq) Write the complete ionic equation H+(aq) + OH–(aq) H2O(l)

  11. Solubility & Net Ionic Equations • Write the net ionic equation that describes the reaction that occurs when solutions of barium nitrate and sodium sulfate are mixed. Write the overall equation for the reaction Ba2(NO)3(aq) + Na2(SO)4(aq) NaNO3(aq) + BaSO4(s)

  12. Net Ionic Equations Write the complete ionic equation Ba2+ (aq) + 2NO3–(aq) + Na+(aq) + SO42–(aq) Na+(aq) + 2NO3–(aq) + BaSO4(s) Write the complete ionic equation Ba2+(aq) + SO42–(aq) BaSO4(s)

  13. Thermochemistry • Thermochemistry is the study of the relationship between heat and chemical reactions. • The system is the matter composed of the substances of the reaction. • The surroundings is all the other matter outside of the reaction.

  14. Thermochemistry • The Law of Conservation of Energy states that the total energy of the universe(the system + surroundings) is constant during a chemical change. • An exothermic reaction gives off heat to the surroundings. DH - • An endothermic reaction absorbs heatfrom the surroundings. DH +

  15. Enthalpy • Enthalpy (DH) is the quantity of heat absorbed by the system, under conditions of constant pressure. DH  the chemical reaction gives off heat DH + the chemical reaction absorbs heat

  16. Calorimetry • Calorimetry • determination in change in enthalpy • measure heat flow in chemical reaction • Calorimeter • a device in which calorimetry is measured

  17. Calorimetry • Molar Heat Capacity • quantity of heat required to raise the temperature of the object 1 kelvin or °C • J/K or J/°C • Specific Heat • quantity of heat required to raise the temperature of a one-gram sample by1 kelvin or °C • J/g-K or J/g-°C

  18. Calorimetry q = m csDT where q = heat in joules m = mass in grams C = specific heat DT = Tfinal -Tinitial

  19. Heat Change Calculations What quantity of heat is needed to raise the temperature of 120 g of water by 6.0 °C? q = m csDT and q = (120 g H2O) (4.184) (6) = 3.0 x 103 J or 3.0 kJ

  20. Heat Change Calculations A chemical reaction releases enough heat to raise the temperature of 49.9 g of water from 17.82°C to 19.72 °C. Calculate the heat evolved q = m csD T q = (49.9g H2O)(4.184)(19.72 - 17.82) q = 397 J

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