A Mixture Is a Combination of Two Or More Substances in Which Each Substance Retains Its

A Mixture Is a Combination of Two Or More Substances in Which Each Substance Retains Its

Solutions

Mixtures

  • A mixture is a combination of two or more substances in which each substance retains its properties
  • Most materials we see daily are mixtures
  • Mixture of elements, compounds or both
  • The components of a mixture can be separated from one another by taking advantage of differences in the components’ physical properties
  • Solids and liquids can be separated by filter paper (Filtration)
  • If a material is pure, it consists of only a single element or compound
  • If the material is impure, then it is a mixture or contains two or more elements or compounds
  • Do not be fooled by labels that claim their product is pure… “pure 100% orange juice”
  • Mixtures are termed either heterogeneous or homogeneous
  • Remember what the prefixes mean…
  • Hetero: different
  • Homo: same
  • Heterogeneous mixtures are when the different components can be seen as individuals
  • e.g. sand in water, pulp in orange juice, salad dressings
  • Homogeneous mixtures have the same composition throughout, therefore any region of the mixture has the same ratio of substances

Heterogeneous Mixtures

  • Suspension: a mixture containing particles that settle out if left undisturbed for a period of time
  • These particles are larger than those found in solvated particles and gravity has greater force on them
  • ***Particles in a solution are atomic-scale size***
  • Colloids:
  • Mixtures of medium sized particles (between suspensions and solutions)

Homogenous Mixtures

  • To describe a solution correctly, one must know the lingo or proper terms
  • The component that is present in the largest amount is the solvent
  • The other component is the solute
  • Yes you can have more than one solute
  • Think of like this… you pour the chocolate syrup (SOLUTE) into the milk (SOLVENT)
  • In other words, you usually mix the solute into the solvent
  • Anything that does not mix in the solvent is said to be INSOLUBLE
  • Have you ever seen something come out of the liquid solution when it cools?
  • This substance is the solute and is known as the PRECIPITATE
  • Lastly, the term for mixing these components into the solution is known as DISSOLVING!!!

Dissolving

  • How do they dissolve?
  • The solvent forces the solute apart into individual components and completely surrounds them
  • This forms the solution and is called solvation
  • Factors that affect the rate of dissolving:
  • Increasing surface area
  • Agitating the solution
  • Heating the solvent
  • A substance that dissolves in water to give a solution that conducts electric current and is called an electrolyte
  • These are salts (metal – nonmetal)
  • Some of the most important chemicals in our body our electrolytes (they are in sport drinks)
  • A nonelectrolyte is a chemical that dissolves in water but does not allow it to conduct electricity

SOLUBILITY RULES

1. Salts of ammonium (NH4+) and Group 1 are always soluble.

2. a. All chlorides (Cl-) are soluble except AgCl, Hg2Cl2, and PbCl2 whichare insoluble.

b. All bromides (Br-) are soluble except AgBr, Hg2Br2, HgBr2, andPbBr2 which are insoluble.

c. All iodides (I-) are soluble except AgI, Hg2I2, HgI2, and PbI2 which areinsoluble.

3. Chlorates (ClO3-), nitrates (NO3-), and acetates (CH3COO-) are soluble.

4. Sulfates (SO4-2) are soluble except CaSO4, SrSO4, BaSO4, Hg2SO4, HgSO4,PbSO4, and Ag2SO4 which are insoluble.

5. Phosphates (PO4-3), and carbonates (CO3-2) are insoluble except NH4+ andGroup 1 compounds.

6. All metallic oxides (O-2) are insoluble except NH4+and Group 1 compounds.

7. All metallic hydroxides (OH-) are insoluble except NH4+ and Group 1 andGroup 2 from calcium down.

8. All sulfides (S-2) are insoluble except NH4+and Groups 1 and 2.

Solubility Equilibra (Equilibrium Reactions with Dissolving Salts)

  • Some ionic compounds will dissociate into their respective ions completely when added to water
  • NaCl(s)  Na+(aq) + Cl-(aq)
  • Some dissociate then return back to solid form and then dissociate again (and again…)
  • BaSO4(s)  Ba2+(aq) + SO42-(aq)
  • You can solve for the equilibrium constant for this dissociate reactions as well

BaSO4(s)  Ba2+(aq) + SO42-(aq)

  • The equation would be Keq = [Ba2+][SO42-] / [BaSO4]
  • In the equilibrium expression above, [BaSO4] is constant because it’s a solid. Therefore, you can cross-multiply (or multiply each side by [BaSO4]) to get Keq * [BaSO4]
  • This creates a new constant called the solubility product constant - Ksp

Solubility Product Constant

  • Ksp is very similar to Keq because it deals with the equilibrium of ions that have not completely disassociated
  • To solve any problems that would occur asking for solubility of ionic compounds, simply set the Ksp equal to the cation and anion concentrations found in the right side of the equilibrium reaction
  • If there are coefficients in the balanced reactions, then you will still use them as exponents

Ksp = [cation+] * [anion-]

Net Ionic Equations

  • These equations include only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution
  • Solids dissolve (dissociate in water)
  • Ca(OH)2 Ca+2 + 2OH-

Example:

Show the net ionic equation for zinc nitrate and ammonium sulfide:

Zn(NO3)2 + (NH4)2S  2 NH4NO3 (?) + ZnS (?)

Zn2+ + 2NO3- + 2NH4+ + S2- 2NH4+ + 2NO3- + ZnS(s)

Zn2+ + S2- ZnS

Mixing Solutions into Solutions

  • If the solution does not want to mix, it is known as immiscible
  • Therefore, miscible solutions are those that are soluble in one another
  • Gases
  • Henry’s Law:
  • The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid
  • If the gas is escaping from a liquid, it is known as effervescence
  • CO2 from sodas, the commercials of alka seltzer

Concentration

  • You can add too much solute to the solvent and it will not mix or dissolve any longer
  • This point is known as the saturation point and we are left with a saturated solution
  • Saturation is defined as the maximum amount of material something can hold onto (Think of humidity levels with weather)
  • Super-Saturated Solutions! – Heat a solvent, dissolve the solute, then allow to cool
  • If you can still add solute to the solution, then it is an unsaturated solution
  • Therefore we can generate yet another equation
  • Concentration ([substance]) = amount of solute / amount of solution
  • Percent by Mass = (mass of solute/mass of solution) * 100
  • This will be used to calculate concentration ratios of mixtures and will later be needed to calculate molarity, molality, etc.
  • The term molarity then refers to the number of moles of solute/liters of solution

Molarity

  • Molarity (M) is the number of moles of solute dissolved per liter of solution
  • M is read as molar
  • Molarity = moles of solute/liters of solution
  • The denominator must always be in L

Molality (m)

  • Yet another form of measuring concentrations gives the use of moles expressed in moles per kilogram
  • Moles/kg
  • Moles of solute / kilogram of solvent
  • Diluting Solutions
  • When making solutions from concentrated stock supplies, you can dilute it down to make it a lower concentration
  • Because the total number of moles of solute does not change, only the volume of the solution you can use…
  • M1V1 = M2V2 : Where M1 & V1 represent the molarity and volume of the stock solution and M2 & V2 represent the molarity and volume of the dilute solution

Preparing Solutions

Example:

A 100.5 mL intravenous (IV) solution contains 5.10 g of glucose (C6H12O6). What is the molarity of the solution?

Example:

If you need 1.50 M of glucose but only 100 mL, how would you prepare this solution?

Normally, you would get 1.50 moles of glucose (~270 g) and place it in (< 1 L of water) to make 1000 mL solution.

100 mL x 1 L x270 g C6H12O6 = 27.0 g

1000 mL 1 L solution

Molarity Practice

1. What is the molarity of a solution containing 21.0 g NaCl in 200 mL of solution?

2. How do you prepare 250 mL of a 0.500 M NaOH solution?

Molality Practice

1. Calculate the molality when 75.0 grams of MgCl2 is dissolved in 500.0 g of solvent.

2. 150.0 grams of sucrose (C12H22O11, mol. wt. = 342.3 g/mol) is dissolved in 1.00 L of water. What is the molality?

3.29.5 grams of KI is dissolved in 1.25 kg of solvent. What is the molality?

4. How many grams MgCl2 will be needed to prepare 3000 grams of a 0.8 molal solution?

Dilution Problems

1. Your teacher needs to make a 0.500 M solution of HCl from concentrated 12.0 M HCl. If the volume of the dilute needs to be 500 mL, then how many mL of the concentrated does he need to mix with how much water? (Remember the final total volume is 500 mL).

2. It is necessary to make a 0.500 M solution of HCl from 500.0 mL of a 3.00 M solution of HCl. What is the volume of the new solution?

3. What is the molarity of a solution which has a volume of 1500.0 mL if it was obtained by diluting 250.0 mL of a 6.0 M solution of H2SO4?

4. Your teacher needs to make 500.0 mL of a 3.00 M solution of H2SO4. Concentrated H2SO4 from the chemical company is 18.0 M. How many mL of the concentrated acid is needed to dilute with how much water to make this solution?

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