Chemistry 20 LG 5

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Chemistry 20 LG 5
2010-11-29 11:15:14
Chemistry Solutions

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  1. Define "solution".
    homogeneous mixture of two substances
  2. Define "solvent".
    the substance present in a larger quantity in the mixture
  3. Define "solute".
    substance present in the smaller quantity in the mixture.
  4. Define "concentration".
    perportion of solute to solvent.
  5. Define "concentrated solution".
    solution with a large amount of solute compared to solvent.
  6. Define "diluted solution".
    solution with a relatively small amount of solute compared to solvent.
  7. What two tests should be performed when using water in an experiment?

  8. What does "ppm" stand for?
    parts per million.
  9. What is one ppm equivalent to?
    one milligram/ litre
  10. What does "ppb" stand for?
    parts per billion.
  11. What is one ppb equivalent to?
    one µg/ litre
  12. What is the formula for calculating amount concentration?
  13. _______compounds are conductive.

    ionic and acidic compounds are conductive.
  14. ________compounds are non-conductive

    molecular compounds are non-conductive
  15. What are the steps involved in preparing a solution?
    1.Calculate the moles required to make the volume and concentration.

    2.Calculate the mass required in moles.

    3. Weigh out the mass required.
  16. What formula is used to calculate the moles required for a specific volume and concentration?
    C=n/v or n=Cv
  17. What formula is used to calculate the mass in the moles required for a specific concentration of a solution?
  18. What kind of flask is used when creating a solution?
    A volumetric flask
  19. What line is used to take a measurment from on a volumetric flask?
    The lower line.
  20. In a dilution, the initial concentration of the solution is _________ or _________ than that of the final solution.
    • i) higher
    • ii) stronger
  21. What formula is used when calculation dilutions?
    CiVi = CfVf
  22. If a substance dissolves in water and the ions separate, becoming surrounded by water, the solution created is _________.
    an ionic solution.
  23. If a polar molecule dissolves in water and molecular bonds are broken and ionic bonds are not, the solution is ___________.
    a molecular solution.
  24. If ions in an ionic substance stay bonded together and form a solid, what should you do?
    Check the periodic table for solubility of the substance.
  25. If ions do not separate, it is said they do not __________.
  26. If ions in a solution separate, it is said that they __________.
  27. What are the four things substances can do when added to a solution?



    4.Remain unchanged
  28. Ionic compounds that are soluble _________.
  29. Molecular compounds that are soluble __________.
  30. Substances that ionize in solution form _____.
  31. Compounds with a H+ ion are often ______.
  32. Explain why some substances may not change when put into a solution.
    A very small amount of the solute dissociates and the rest remains solid.
  33. What happens when slightly soluble ionic solute is added to solvent?
    no change occurs.
  34. Would Sr(OH)2(s) be ionic, molecular or acidic?
    Ionic, as it has both a metal and a non-metal.
  35. Would Sr(OH)2(S) be electrolytic or non-electrolytic?
    Electrolytic, as it is an ionic compound.
  36. What entities would be present in a solution containing water and Sr(OH)2(s)?
    Sr2+(aq) + 2 OH-(aq)
  37. Would NH3(g) be ionic, acidic or molecular?
    Molecular, as it contains two nonmetals.
  38. Would NH3(g) be electrolytic or non-electrolytic?
    Non-electrolytic, as it is molecular.
  39. Would NH3(g) undergo separation dissociation or no change?
    Separation, as it is polar.
  40. Define the term "saturated solution".
    A solution with the maximum amount of solute at a given temperature.
  41. Define the term "unsaturated solution".
    A solution where the solvent is able to dissolve more solute at a given temperature.
  42. Define the term "supersaturated solution".
    A solution in which there is more dissolved solute than the solubility indicates for a given temperature.
  43. Solids have a _______ solubility in water that is at higher tempteratures.
  44. Gases have a higher solubility in water at ________ temperatures.
  45. Why do gases have lower solubility in hot water?
    • Hot water has more kinetic energy than cold water.
    • Colder water = less movement= a
    • better ability to hold gases.
    • Gases are able to escape less quickly from cold water than hot water.
  46. __________ generally have a low solubility in water.
  47. In polar molecules, solubility increases when the water temperature ___________.
  48. Polar liquids that do not dissolve are said to be ___________.
  49. What kind of polar molecules tend to dissolve completely in water to any proportion?
    Small polar molecules (short chain) with hydrogen bonds will completely dissolve in water to any proportion.
  50. Polar liquids that dissolve are said to be __________.
  51. Solution equilibrium occurs when:
    The concentration becomes constant/ the solute stops dissolving and the solvent is saturated.
  52. Would BaCO3(s) form an ionic, molecular or acid solution?
    Ionic, as it contains both a metal and a non-metal.
  53. Would BaCO3(s) form an electrolytic or non-electrolytic solution?
    Electrolytic, as it is ionic.
  54. Would BaCO3(S) undergo dissociation, deparation, ionization or no change when in solution?
    No change