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Chapter 5: The Nature and Properties of Solutions

Chapter 5: The Nature and Properties of Solutions. Solution:. A homogenous mixture that looks the same throughout. Solute:. Substance being dissolved. Present in smaller amount. Solvent:. Substance doing the dissolving. P resent in greater amount. Demo!!. Solute - KMnO 4.

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Chapter 5: The Nature and Properties of Solutions

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  1. Chapter 5: The Nature and Properties of Solutions

  2. Solution: A homogenous mixture that looks the same throughout. Solute: Substance being dissolved. Present in smaller amount. Solvent: Substance doing the dissolving. Present in greater amount.

  3. Demo!! Solute - KMnO4 Solvent - H2O Substance doing the dissolving. Present in greater amount. Substance being dissolved. Present in smaller amount.

  4. The process of dissolving. solute particles are surrounded by solvent particles • Solvation: First... solute particles are separated and pulled into solution Then...

  5. Conducting Electricity When a solute (solid or liquid) is placed in a solvent 1 of two things can happen. 1. The solute can sink to the bottom and be insoluble (slightly soluble). 2. The solute can dissolve in the solvent and do MANY things.

  6. Conducting Electricity If a compound dissolves in the solvent it can either break apart and create positive and negative ions. NaCl (Salt) Na+ Na+ Cl- Cl- Na+ Cl- Or it can stay together and NOT form ions. Whether or not a substance breaks apart depends on what type of compound it is.

  7. Molecular Compounds(Non-Metal + Non-Metal) Molecular Compounds do not break up and form ions. Thus, cannot conduct electricity. Sugar(s) Sugar(aq)

  8. Molecular Acids(Weak Acids only) Weak molecular acids only break up slightly, only a few ions are created. Thus, poor conductors of electricity. Acetic Acid (Vinegar) (CH3COOH) H+ CH3COO-

  9. Ionic Compounds(Metal + Non-Metal) Ionic Compounds break apart when they dissolve and produce lots of ions. Thus are excellent conductors of electricity. Salt (NaCl) Na+ Cl- Na+ Cl- Cl- Na+ Cl- Cl- Na+ Na+ Na+ Cl-

  10. - + - - + + acetic acid salt sugar Conducting Electricity Non- Electrolyte Weak Electrolyte Strong Electrolyte solute exists as ions and molecules solute exists as ions only solute exists as molecules only DISSOCIATION IONIZATION

  11. Dissociation (Ionic): • separation of an ionic solid into aqueous ions • separation of ions that already existed before being added to water Dissociation Equation NaCl(s)  Na+(aq) + Cl–(aq) Back

  12. Ionization (Molecular Acids) • breaking apart of some polar molecules into aqueous ions • the production of new ions, specifically hydrogen/hydronium ions Ionization Equation HNO3(aq) + H2O(l)  H3O+(aq) + NO3–(aq)

  13. Molecular Solvation (Molecular) • Polar molecules stay intact. Dissociation Equation C12H22O11(s)  C12H22O11(aq)

  14. NONPOLAR NONPOLAR POLAR POLAR Molecular Solvation “Like Dissolves Like”

  15. Substances in Water • Ionic Substances: Solubility table allows you to predict if substance will dissociate 100% or not (s or aq) • Molecular Compounds: require the understanding of intermolecular forces to predict entities in water (s/l/g or aq) • Polar substances dissolve in water (like dissolves like) • Hydrogen bonding makes them even more soluble • Elements: Most will remain the same (s,l, or g) • Exception: Group 1 elements like sodium will dissolve in water to form basic solutions (aqueous ions)

  16. Universal Solvent • Water is often referred to as the “universal solvent” due to its ability to dissolve many things. • However, there is no actual universal solvent. • How would you store such a liquid?

  17. Acids • Empirical Definitions (observable properties): • Acids: • taste sour • Electrolytes (Weak Acids Low…..Strong Acids High) • React with metals to produce hydrogen gas • Turn blue litmus red • Neutralize bases

  18. Acids • Properties of both ionic and molecular compounds. • Strong acids (6) are strong electrolytes (complete ionization). 1. HCl 2.HBr 3.HNO3 4.H2SO4 5.____________ 6. _____________ Strong Acids are Strong electrolytes (Complete dissociation).

  19. Weak Acids are Weak electrolytes (Partial dissociation).

  20. Bases • Empirical Definitions (observable properties): • Bases: • taste bitter (coffee) • Electrolytes (Weak Bases Low…..Strong Bases High) • Feel greasy/slippery • Turn Red litmus blue • Neutralize acids

  21. Bases • Properties of both ionic and molecular compounds. • Strong bases are strong electrolytes (complete dissociation). • Strong bases are any compounds containing OH (hydroxide). -NaOH -Ba(OH)2 -Ca(OH)2 -Li2(OH) -Mg(OH) -KOH

  22. H H – + O O Cl Cl H H H H Definitions • Arrhenius - In aqueous solution… • Acidsform hydronium ions (H3O+) HCl + H2O  H3O+ + Cl– acid

  23. H H – + N O O N H H H H H H H H Definitions • Arrhenius - In aqueous solution… • Bases form hydroxide ions (OH-) NH3 + H2O  NH4+ + OH- base

  24. Qualitative vs. Quantitative Qualitative: what is present in a solution • Is the solution acidic/basic/ionic etc.. Quantitative: how much is present in a solution • What is the concentration of the solution

  25. Concentration • The amount of solute in a solution. • Describing Concentration • % by mass - medicated creams • % by volume - rubbing alcohol • ppm, ppb - water contaminants • molarity - used by chemists

  26. Concentration SAWS Water Quality Report - June 2000

  27. % by Volume • Vinegar is usually labeled as 5% acetic acid by volume • This means there is 5mL of acetic acid in every 100mL of vinegar solution

  28. % Weight per volume • Hydrogen Peroxide used as antiseptic is 3% W/V • This means there are 3g of hydrogen peroxide for every 100mL of solution

  29. % weight by weight • This form is used when making a solution of two solids like in the production of a metal alloy such as sterling silver Merely a flesh wound

  30. ppm • Parts per million is often used for very dilute solutions and is very important in environmental studies • This is a special case of the W/W ratio or

  31. Molarity/amount concentration

  32. MASS IN GRAMS MOLES NUMBER OF PARTICLES LITERS OF SOLUTION 6.02  1023 (particles/mol) molar mass (g/mol) Molarity (mol/L)

  33. Example: • Find the molarity of a solution containing 75 g of MgCl2 in 250 mL of water. Xmol MgCl2 1 mol MgCl2 95.21 g MgCl2 Xmol MgCl2 = = 75 g MgCl2 0.25 L water = 3.2mol/L MgCl2

  34. Example: • How many grams of NaCl are req’d to make a 1.54mol/L solution using 0.500 L of water? 1.54 mol NaCl 1 L water X mol NaCl = = 0.77 mol 0.500 L water

  35. 58.44 g NaCl 1 mol NaCl X g NaCl 0.77 mol NaCl = = 45 g

  36. Concentration of Ions • Balanced dissociation or ionization equations for soluble ionic substances and strong acids allows us to determine the amount concentration of ions or compounds in solution

  37. Concentration of Ions • Use the mole ratio from the balanced equation to set up the need over got conversion factor. • ex: What is the amount concentration of aluminum ions and sulfate ions in a 0.40 mol/L solution of aluminum sulfate? G N Al2(SO4)3(aq) 2Al3+(aq) + 3SO42-(aq) C = 0.40 mol/L C = 0.80 mol/L C = ? () N G 2 1 CAl3+(aq) = (0.40 mol/L)(2/1) = 0.80 mol/L

  38. Concentration of Ions N G Al2(SO4)3(aq) 2Al3+(aq) + 3SO42-(aq) C = 0.40 mol/L C=1.20 mol/L C = ? () N G 3 1 1.20 mol/L CSO42-(aq) = (0.40 mol/L)(3/1) =

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