Academic Chemistry

Academic Chemistry UNIT 10

Solutions

Class Period: Test Date:

1 Chemistry Calendar

Monday Tuesday Wednesday Thursday Friday FEBRUARY 20 21 22 23 24 Notes #1: Water Writing in the Notes #2: Notes #2: Covalent STUDENT (p. 3-4) Content Areas: Covalent (p. 5-8) (p. 5-8) HOLIDAY Science (TELPAS) Learner.org: Water Video Topic: Water

27 28 29 MARCH 1 2 Notes #3: Notes #4: Notes #5: Notes #6: Lab: Dilutions Solutions (p. 9-11) Solvation (p. 12- Solubility Curves Concentrations by 15) (p. 16-17) Molarity Calculations (p. 18-19)

5 6 7 8 9 Notes #7: Test Review: Test Review Sheet UNIT 10 TEST Dilutions Calculations (p. 23-25) Calculations (p. Review (p. 22) 20-21)

2 NOTES #1: Unique Role of Water 2 9 1 Describe the unique role of water in chemical and biological systems (Supporting Standard) 5 SCI:CHEM.10A 0 Importance of Water to Life  Living cells are 70% - 95% water.  The surface of the earth is approximately 70.9% water.  Water is the only substance on earth that naturally occurs in all three states – solid , liquid , and gas .  Life is dependent upon water and its unique properties .  Most of the reactions that take place in living organisms involve water, such as photosynthesis and cellular respiration.

Background: What is water?  Water is a pure substance, a compound , of hydrogen and oxygen chemically bonded to form a single substance.  Electrolysis breaks down water into hydrogen and oxygen.  Physical Properties: - Boiling point: 100 o C - Melting point: 0 o C - Density: 1 g/mL

- Specific heat: 1 cal/g o C, which is also equal to 4.184 J/g o C - Heat of vaporization: 540 cal/g

What Makes Water so Special?  Most liquids, when frozen, become denser and sink. Solid water – ice – floats.  Something to think about: What impact do you think that has on sea life? What do you think would happen to aquatic organisms if ice didn’t float?

 Why does ice float?

3  Water molecules are polar . The water molecule has a bent shape giving it partial charges on its ends. The hydrogen atom is less electronegative than the oxygen atom so it loses the electron tug of war and becomes partially positive, and the oxygen pulls the shared electron closer to it, becoming partially negative.  Hydrogen bonding: - Definition: an attractive force between a hydrogen atom of one molecule and an

electronegative atom (such as nitrogen, oxygen or fluorine) of a second molecule.

- The hydrogen bonding of water gives it some unique properties: 1. High Surface Tension: Caused by uneven attraction between water and air (water

molecules will clump together)

2. High Specific Heat: absorbs a lot of heat. Requires a lot of energy to break bond.

3. Low Vapor Pressure: pressure of vapor above a liquid, caused by escape of surface

molecules

4. Heat of Vaporization: Amount of energy needed to convert liquid to gas

5. High boiling point: Takes more heat to disrupt the attractions between water molecules

 Water as a Universal Solvent - Solution = A homogeneous mixture or two or more substances. - Solvent = Dissolving agent of a solution. - Solute = Substance dissolved in a solution. - Aqueous Solution = A solution in which water is the solvent

4  A Rule of Thumb: Like dissolves Like o Water dissolves ionic compounds (salt) o Water dissolves polar covalent compounds (sugar) o Water does not dissolve nonpolar compounds (oil)

Polar vs. Nonpolar Covalent Molecules  Drawing Lewis Structures can help determine whether a molecular compound is polar or nonpolar, and whether it will dissolve in water .  Lewis Structure: D iagrams that show the bonding between atoms of a molecule and the lone pairs

of electrons that may exist in the molecule.

5 NOTES #2: Covalent Bonds and Lewis Structures

Like Dissolves Like  Most ionic compounds will dissolve in water to form an aqueous solution , but when examining covalent molecules, only polar molecules – those with partial charges – will dissolve in water.  This goes back to our rule, like dissolves like : Water is polar , and as a polar substance, it will dissolve polar molecules and ionic compounds.  RECALL: Ionic bonds are between a cation (metal or ammonium) and an anion (nonmetal or polyatomic)

What are Covalent Bonds?  Bond between two nonmetals where electrons are shared to fill their outer shells  Remember, all atoms want 8 valence electrons ( octet rule)  What do nonmetals do among themselves? - Neither one will give away an electron - So they share their valence electrons - This is a covalent bond

There are 2 kinds of covalent molecules:  Between 2 or more different nonmetals.

- Example : H2O, C6H12O6  Diatomic: molecules w/ same kind of atoms

- Example : O2, H2, Cl2 - There are seven diatomic elements:

▸ Br2 , I2 ,N2 , Cl2 , H2, O2 , F2 (Br-I-N-Cl-H-O-F)

Drawing Lewis Structures  Directions: 1. Start by writing the atomic symbol. 2. Determine the number of VALENCE electrons. 3. Draw dots for each of the valence electrons on each of the four sides of the symbol. 4. Put one dot on each side before you start to pair the electrons 6 EXAMPLES:

INDEPENDENT PRACTICE: Draw Lewis Dot Structures for each of the following elements: 1. Oxygen 5. Lithium

2. Magnesium 6. Chlorine

3. Cesium 7. Aluminum

4. Krypton 8. Phosphorus

From Lewis Dot Structures to Covalent Bonds  Fluorine has seven valence electrons  A second atom of fluorine also has seven  How do these two atoms become stable? By sharing electrons to get a full octet

 When atoms share a pair of electrons , this is called a covalent bond.  A single covalent bond is shown as a dash in a Lewis Diagram (a.k.a. - structural formula) and represents two electrons being shared between the atoms it links together.

 Draw Lewis Dot Structure for Fluorine and the structural formula for a molecule of F2:

 A tip to remember when drawing structural formulas (Lewis Structures): SNP 2 it! - Add up all the valence electrons Present on the atoms involved (call this value P) - Count the total number of electrons Needed to give all atoms full octets (call this value N) - To find the electrons Shared, (call this S), subtract . ▸ This gives us, S = N – P (SNAP!)

7 - Divide by 2 because there are two electrons in a single covalent bond. - Your answer tells you how many bonds - draw them. - Draw in the rest of the valence electrons to fill all octets (except for H, which is filled w/ 2)

8 GUIDED PRACTICE #1:

1. H2O 3. N2

2. CO2 4. NH3

Polar or Nonpolar? 1. The general principle in predicting molecular polarity is the comparison of similar regions of the molecule. - If all similar regions are not the same , the chemical species is polar. A general idea of the polarity direction may be obtained from electronegativity values , the unequal concentration of lone pairs on the ends of a molecule, etc. - If all similar regions are the same , the chemical species is nonpolar.  It boils down to symmetry . If the molecule looks like a copy of itself from top to bottom and left to right, then it’s going to be nonpolar .

GUIDED PRACTICE #2:  Take a look at the Lewis Structures you drew in the GUIDED PRACTICE #1 portion above and identify those substances as either polar or nonpolar. THEN place them in the table on the next page.

1. Polar 3. Nonpolar

2. Nonpolar 4. Polar

9 ASSESSMENT: Determining Molecular Geometry

Bonded Pairs Lone Pair(s) Molecular Geometry Bond Angle(s) Example

N2 2 0 180 CO2

H2O 2 2 <109.5 OF2

SO3 3 0 120 - (NO3)

3 1 <109.5 NH3

4 0 109.5 CH4

INDEPENDENT PRACTICE: Draw Lewis Structures for the following, label as polar or nonpolar, and then add them as examples in the table above.

1. CH4 3. Oxygen difluoride

2. SO3 4. Nitrate

10 NOTES #3: Solutions 2 9 5 Distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions (Readiness Standard) 5 SCI:CHEM.10E 0 RECALL:  What is a solution? A homogeneous mixture of two or more substances

 Solute? Substance dissolved in a solution

 Solvent? The dissolving agent of a solution

Solubility  Definition: Maximum amount solute that can be dissolved in a given amount of solvent at a specific

temperature and pressure.

o Remember, Like Dissolves Like! o Nerdy Riddle of the Day: Why do white bears dissolve in water?

 Soluble – Substances able to be dissolved

 Insoluble – Substances unable to be dissolved

Types of Solutions  Aqueous Solutions: Homogeneous mixtures in which water is the solvent

- Liquids can dissolve in other liquids to form solution. Such liquids are defined as miscible .

o Example: alcohol and water

- If two liquids do not dissolve in one another, but rather separate upon standing, those liquids are defined as immiscible .

o Example: oil and water, or water and the “lava” in a lava lamp

11  Electrolytes – Compounds that conduct electricity . All ionic compounds form electrolytes when dissolved in water.  Nonelectrolytes – Do not conduct an electric current in aqueous solution or in molten form. Many molecular compounds are nonelectrolytes.  Alloys – Solids that dissolve in other solids

- Examples: brass, steel, etc Solutions vs. Things That Just Look Like Them…  Is milk a solution? What about muddy water, dust particles floating or suspended in the

air, or just the different gases – N2 , O2 , CO2 , H2O(g), Ar, etc. – that make up the pure, unpolluted air we breathe?

 So just what makes a solution…a solution?

o Solutions are always transparent , meaning that light passes through with no scattering from tiny solute particles – a solution may have a "color" but it will still be transparent.

o The solute component of a solution can be atoms, ions, or molecules , and will measure between 0.1 to 2 nanometers in diameter.

o Solutions are homogeneous and do not settle out . o Solutions cannot be filtered , but can be separated using the process of distillation.  Back to Basics: Mixtures

o Homogeneous Mixture: Mixture that has the same uniform appearance and

composition throughout. May exist as gas, liquid or solid, depending on state of solvent.

o Heterogeneous Mixture: Mixture that is not uniform in composition

o There are 3 special types of heterogeneous mixtures that look just like homogeneous mixtures – in fact, they look very similar to the solutions we’re discussing…but they’re not! Example: Example: Muddy Muddy 1. Suspensions: Heterogeneous mixtures with particles that have diameters greater than water water 1000 nm. The particles of a suspension will separate upon standing.

12 Example: 2. Colloids: Heterogeneous mixture containing particles that are intermediate Example: Milk, fog, Milk, fog, etc. in size between those found in suspensions and true solutions that remain evenly etc. distributed without settling out due to Brownian Motion. ▸ Brownian Motion is the chaotic movement of colloidal particles caused by collisions between the particles of the dispersing medium with the dispersed colloidal particles. These collisions prevent the colloidal particles from settling. Example: Example: Mayonnais Mayonnais 3. Emulsions: Heterogeneous, colloidal dispersions of liquid particles in a liquid medium. e e  How to tell the difference between a solution, suspension, colloid, and emulsion:

o Suspensions will eventually separate ; whereas colloids and solutions will not. o Colloids can be distinguished from solutions using the Tyndall effect.

▸ A beam of light passing through a true solution, such as pure unpolluted air/water, is not visible. Light passing through a colloidal dispersion, such as smoky or foggy air, will be reflected by the larger particles and the light beam will be visible.

13 ASSESSMENT: Solutions Directions: Place the number for each description (1-9) in the correct column(s) of the chart. Descriptions may be used more than once if applicable. (Hint: 2 statements will be used in more than one column of the table)

1. Can be distinguished from solutions by the appearance of a beam of light when light is passed through such a mixture 2. These must be transparent 3. Particles are ions or molecules 4. Particles will settle out upon standing 5. Is a homogeneous mixture 6. An emulsion is an example of this in which liquid particles are dispersed in a liquid medium 7. Is a heterogeneous mixture 8. Salt water is an example of this 9. Particles are evenly distributed and will not settle out

SOLUTIONS SUSPENSIONS COLLOIDS 2 4 1 3 7 6 5 7 8 9 9

Directions: For questions 10 – 13, identify which of the following will dissolve in water to form electrolytes that conduct an electrical current. Indicate your answer by writing: “Electrolyte” or “Nonelectrolyte”

10. NaCl 12. H2

Electrolyte Nonelectrolyte

11. CH4 13. KNO3 Nonelectrolyte Electrolyte

Directions: For questions 14 & 15, read and answer the questions and provide an explanation when asked. 14. Oil and vinegar are the two main components of any vinaigrette salad dressing. Why does a bottle of vinaigrette have to be shaken before being used? Explain through mention of polarity. They do not dissolve in one another because oil is nonpolar, while vinegar is a polar substance. Therefore to mix them before use, the bottle should be shaken. 15. Are oil and vinegar miscible or immiscible?

Oil and vinegar are immiscible.

14 NOTES #4: Solvation To Dissolve or Not to Dissolve, That is the Question  Solvation, also called dissolution , is the process of attraction of molecules of a solvent with molecules or ions of a solute .  Forces between solute-solute molecules and solvent-solvent molecules must be broken before new solute-solvent attractive forces can form.

o Dissolution will occur if the solute-solvent forces of attraction are stronger than the attractive forces keeping the solute together and those forces keeping the solvent together.

o Dissolving will not occur if particles of a solute are more attracted to other solute

particles than to the solvent . The same is true for solvent particles.

o Generally, if all three of the intermolecular forces of attraction are roughly equal, the substances will be soluble in each other .  Dissolution depends on the relative strength of three intermolecular attractive forces.

o Three Types of Intermolecular Forces: 1. Ion-Dipole Attractive Forces: . This occurs between ionic compounds and a polar solvent (such as water) . A tug-of-war occurs for positive and negative ions between the other ions in an ionic crystal and the surrounding water (or other polar solvent) molecules. ▸ If the water “pulls” harder on the solute, the ionic compound will dissolve

▸ If the “pull” from the ionic crystal is stronger, the ionic compound will not

dissolve. - Example: Sodium Chloride in the solvation process. Because NaCl dissolves, that means that the water won the tug-of-war.

15 2. Dipole-Dipole Attractive Forces: . Occurs between two polar molecules ▸ A polar molecule contains a bond in between ionic and covalent. - Neither atom is “strong” enough to pull the electrons away from the other to form positive and negative ions like in ionic bonds .

- The electrons also aren't shared completely equally either as they are in a typical covalent bond. ▸ The bonds in these molecules are said to be polar, because they have partial

positively and negatively charged ends , or poles, and are said to have a dipole movement. . Hydrogen bonding is a particularly strong type of dipole-

dipole force.

3. Induced Dipole-Induced Dipole Attractive Forces:

. Also called London Dispersion forces . Occur between nonpolar molecules and result from temporary charge imbalances.

16 ▸ The temporary charges exist when the nucleus of one atom attracts electrons from the neighboring atom. At the same time, the electrons in one particle repel the electrons in the neighbor and create a short lived charge imbalance.

Factors Affecting the Rate of Dissolution – 2 9 2 Investigate factors that influence rates of dissolution such as temperature, agitation, and surface area (Readiness Standard) 5 SCI:CHEM.10F 0 1. Agitation or Stirring :  Increases rate of dissolution, meaning the solute dissolves faster . 2. Particle size :  Smaller particle size increases surfaces area and dissolves faster . 3. Temperature :  Heating increases collisions between particles and causes the solute to dissolve faster . 4. Pressure (gases only):  Increasing pressure increases collisions of gases, meaning that a gaseous

solute will dissolve faster . S 1 = S 2  Henry’s Law explains the relationship between pressure and solubility. Formula: P 1 P 2

Concentration  This is one of the important properties of a solution. Concentration is a measure of solute dissolved in the solvent of a solution. - Concentrated – Solution containing a large amount of solute dissolved in solvent - Dilute – Solution containing a little amount of solute dissolved in solvent.  Example: Vinegar is a dilute acetic acid solution o Vinegar is used in cooking or combined with oil in vinaigrette salad dressings. 17 o Concentrated acetic acid will kill you if ingested. o The only difference between such solutions is the concentration of the solute.  Measures of Solute Concentration: - Once you have identified the solute and solvent in a solution, you are ready to determine its concentration. Concentration may be expressed several different ways. We will use only two: 1. Molarity – number of moles of solute per liter of solution, abbreviated, M.  Formula: M = mole solute / Liters solution 2. Percent by Mass – This is the mass of the solute divided by the mass of the solution (mass of solute plus mass of solvent), multiplied by 100.  Formula: Mass solute .X 100 Mass solution o Ex: Find the % composition by mass of a 100 g NaCl solution containing 20 g NaCl. o Solution: 20 g NaCl / 100 g solution x 100 = 20% NaCl solution

18 ASSESSMENT: Solvation Directions: For questions 1 – 4, identify which of the two solutes in each scenario will dissolve faster in water to form aqueous solution AND explain your choice. 1. Regular granulated table salt or rock salt? Granulated table salt will dissolve faster because the particle size is smaller, and thus more overall surface area of the salt is in contact with the water. 2. Carbon dioxide gas under a pressure of 1 atm or under a pressure of 3 atm? At a pressure of 3 atm because at higher pressures, gasses dissolve faster.

3. Sugar in iced tea or hot tea? Hot tea because solids will dissolve faster in water at higher temperatures. 4. Koolaid in a pitcher being stirred or in a pitcher that is not stirred? Being stirred because agitation (stirring) increases the rate of dissolving.

Directions: For questions 5 – 9,match each term to its definition. a) Covalent bond d) Ionic bond b) London Dispersion Force e) Ion-Dipole Force c) Dipole-Dipole _____d 5. Force holding a cation and an anion together within a compound because of opposite charges _____b 6. Force that occurs between two nonpolar molecules because of temporary charge imbalances _____c 7. Force that occurs between two polar molecules because of the partial charges on their ends (poles) _____a 8. Force holding a nonmetal to other nonmetals within a molecule because of the sharing of electrons _____e 9. Force that occurs between an ionic compound and a polar solvent because of the attraction of opposite charges.

Directions: Use Henry’s Law to answer the following two questions. 10. If the solubility of a gas in water is 0.77 g/L at 3.5 atm of pressure, what is its solubility (in g/L) at 1.2 atm of pressure? (Assume temperature is constant 25°C) .77 g/L = S 2 3.5 atm 1.2 atm S2 = .264 = .26 g/L

11. If the solubility of a gas in water is 1.77 g/L at 3.7 atm of pressure, at what pressure would the solubility of the gas be 0.82 g/L? (Assume temperature is constant 25°C) 1.77 g/L = 0.82 g/L 3.7 atm P2 P2 = 1.714 = 1.7 atm

Directions: Calculate concentration in percent by mass:

12. Find the % composition by mass of a 327 g KNO3 solution containing 34 g KNO3. 34 g x 100 = 10.4% or 10% 327 g

19 NOTES #5: Solubility Curves 1 3 11 Organize, analyze, evaluate, make inferences, and predict trends from data such as solubility curves 5 SCI:CHEM.2H 0 2 9 5 Distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions (Readiness Standard) 5 SCI:CHEM.10E 0 Solubility Curves  Graphs that show the maximum amount of solute dissolved in 100 grams of water at a specific temperature.  Solubility: The amount of solute that dissolves per unit of solvent at a specific temperature and

pressure to produce a saturated solution.

Concentration of Solutions: Saturated, Unsaturated, or Supersaturated?  Saturated: A solution containing the maximum amount of solute per given amount of solvent at a constant temperature and pressure.

 Unsaturated: A solution that contains less solute than a saturated solution at a given temperature and pressure.

 Supersaturated: A solution that contains more solute than it can theoretically hold at a given temperature.

GUIDED PRACTICE Directions: Use the solubility graph to the right to answer the following.

1. How many grams of KNO3 can be dissolved in

100g of H2O at 50°C?

80 g KNO3

2. If 70g of KNO3 is dissolved in 100g of water at 50°C, is the solution unsaturated, saturated, or supersaturated? unsaturated

3. If 105g of KNO3 is dissolved in 100g of water at 50°C, is the solution unsaturated, saturated, or supersaturated? saturated

4. Indicate whether the following are saturated, unsaturated, or supersaturated solutions:

a. 30g of NaCl in 100g of H2O at 80°C. unsaturated

b. 120g of KNO3 in 100g of H2O at 60°C. supersaturated c. 92 g of NaNO3 in 100g of H2O at 10°C. 20 supersaturated INDEPENDENT PRACTICE: Solubility Curves

Directions: Use the solubility graph below (or from the previous page) to answer the following. 5. Which substances are gases? How do you know?

NH3 and CE2(SO4)3it’s a molecular compound and it slopes down 6. At 30°C, 90 g of sodium nitrate is dissolved in 100g of water. Is this solution saturated, unsaturated, or supersaturated? unsaturated

7. Which salt is least soluble in water at 20°C?

8. Which salt is most soluble at 10°C? KI

9. Which gas is most soluble at 50°C?

10. Which salt is least soluble in water at 50°C?

11. You have 300 grams of water at 10°C.

You dissolve as much NaNO3 as you can. If you heat water to 30°C, how much more solute can be added and made to dissolve? (96*3) – (79*3) = 51 g

12. At 40°C, how much potassium nitrate can be dissolved in 300 g of water? 62 g * 3 = 186 g

13. How many grams of KCl can be dissolved in 200g of water at 80°C? 50g*2 = 100g

14. A saturated solution of potassium chlorate is formed from 100 g of water, if the saturated solution is cooled from 80°C to 50°C, how many grams of precipitate form? 40 g – 20g = 20 g

15. Extrapolate the following values:

a. Solubility of NH3 at 110°C

About 4g per 100 g H2O

b. Temperature of KCl at 65g/100g H2O About 125°C

21 NOTES #6: Concentrations by Molarity 2 9 6 Calculate the concentration of solutions in units of molarity (Supporting Standard) 5 SCI:CHEM.10C 0

Variables: M = molarity

FORMULA: M = n solute / V solution (in mol/L) n = moles

V = volume (in liters) NOTES:  Liters must be used as the unit for volume because molarity is measured in moles per liter, abbreviated, M.  Conversion Factor: 1000 mL = 1 L

GUIDED PRACTICE

1. How many grams of sodium sulfate (Na2SO4) are required to prepare a 250 mL solution whose concentration is 0.683M? 0.683 M * 0.250 L = n 0.683 M = n . n = 0.17075 mol 0.250 L n = 0.171 mol.

2. What is the molarity of an 85 mL ethanol (C2H5OH) solution containing 1.77 grams of ethanol? C: 2*12.01g = 24.02 g M = 0.452 mol/L M = 0.0384 mol . H: 6*1.01 = 6.06 g M = 0.45 M. 0.085 L O: 1*16.00 = 16.00 g 46.08 g / mol 1.77 g x 1 mol / 46.08 g = 0.0384 mol

INDEPENDENT PRACTICE 3. Calculate the mass of NaOH in grams required to prepare a 500 mL solution with a concentration of

2.8M. 2.8 M * 0.5 L = n Na: 22.99g n = 1.4 mol Cl: 35.45g 2.8 M = n . 58.44g / mol (molar mass) 0.5 L 1.4 mol * 58.44g / mol = 81.816 = 82 g.

4. How many grams of HCl are needed to make a 200 mL of 18.0M hydrochloric acid in aqueous solution? 18.0 M * 0.2 L = n H: 1*1.01g n . 18.0 M = n = 3.6 mol Cl: 1*35.45g 0.2 L 36.46g / mol (molar mass) 3.6 mol * 36.46g / mol = 131.256g m = 100 g.

5. How many moles of ammonium nitrate are in 335 mL of 0.425M NH4NO3? 0.425 M * 0.335 L = n 0.425 M = n . n = 0.142375 mol 0.335 L n = 0.142 mol.

22 6. Calculate the volume in mL of a solution requiring provided it is a .30M solution using 0.85 g of acetic

acid (CH3COOH). C: 2*12.01g V * 0.30 M = 0.0142 mol H: 4*1.01g 0.0142 mol . 0.30 M = V = 0.0142 mol / 0.30 M O: 2*16.00g V V = 0.047175 L 60.06g / mol (molar mass) 0.85 g * (1 mol / 60.06g) = 0.0142 mol V = 0.047 L = 47 mL

7. What is the volume of a 0.125M NiCl2 solution containing 3.25 g of NiCl2? Ni: 1*58.69g V * 0.125 M = 0.0251 mol Cl: 2*35.45g 0.0251 mol . 0.125 M = V = 0.0251 mol / 0.125 M 129.59g / mol (molar mass) V V = 0.200633 L 3.25 g *(1 mol /129.59g) = 0.0251 mol V = 0.201 L

8. Calculate the molarity of a solution containing 15.4 g of sucrose (C12H22O11) in 74 mL of solution? C: 12*12.01g M = 0.607899 mol/L H: 22*1.01g 0.04498 mol . M = M = 0.61 M. O: 11*16.00 0.074 L 342.34g / mol (molar mass) 15.4 g *(1 mol /342.34g) = 0.04498mol

9. Determine the molarity if 7.82 grams of naphthalene (C10H8) were dissolved in 85.2 mL of benzene solution. C: 10*12.01g M = 0.71606 mol/L H: 8 * 1.01g 0.06101mol . M = M = 0.716 M. 128.18g / mol (molar mass) 0.0852 L 7.82 g *(1 mol /128.18g) = 0.06101mol

10. A quantity of 5.25 g of NaOH is dissolved in a sufficient amount of water to make up exactly 1 liter of solution. What is the molarity of the solution? Na: 1*22.99g M = 0.13125mol/L O: 1*16.00g 0.13125mol . M = M = 0.131 M. H: 1*1.01g 1 L 3 sig figs b/c it’s exactly 1 L 40.00 g / mol (molar mass) 5.25 g *(1 mol /40.00g) = 0.13125mol

23 NOTES #7: Dilutions 2 9 7 Use molarity to calculate the dilutions of solutions (Supporting Standard) 5 SCI:CHEM.10D 0

Variables: M = Molarity (in mol/L)

FORMULA: M 1 V 1 = M 2 V 2 V = Volume

1 & 2 represent before & after diluting NOTES:  Volume units MUST match so that they cancel when used in the above equation.  Conversion Factor: 1000 mL = 1 L

RECALL: Define the following –  Dilute:

 Concentrated:

GUIDED PRACTICE 1. If a 12.5 M solution was diluted to 3 L that had a new molarity of .5 M solution, what volume of the original solution was added? M1 = 12.5 M (12.5 M)(V1) =(0.5 M)(3 L) V1 = 0.12 L V1 = ? 1 (12.5 M )(V 1) = 1.5 M*L V = 0.1 L. M2 = 0.5 M 12.5 M 12.5 M V2 = 3 L

2. If 80 ml of an 8M solution was diluted into a 1 Liter container, what would the new molarity be? M1 = 8 M (8 M)(80 mL) =(M2)(1000 mL) V1 = 0.64 L V1 = 80 mL 1 640 M *mL = (M2)(1000 mL) V = 0.6 L. M2 = ? 1000 mL 1000 mL V2 = 1000 mL

INDEPENDENT PRACTICE 3. If a 14 M solution was diluted to a .5 M solution that had a volume of 200 ml, how many ml of the original solution had to be added? M1 = 14 M (14 M)(V ) =(0.5 M)(200 mL) 1 V1 = 7.14286 mL V1 = ? V1 = 7 mL. 14 M (V1) = (0.5 M)(200 mL) M2 = 0.5 M 14 M 14 M V2 = 200 mL

4. If a 14.5 M solution was diluted to 4 L of a 2 M solution, how much of the original solution was added?

M1 = 14.5 M (14.5 M)(V1) =(2 M)(4 L) V1 = 0.5517 L V1 = ? 1 (14.5 M)(V1) = (2 M)(4 L) V = 0.6 L. M2 = 2 M (14.5M) (14.5M) V2 = 4 L 24 5. If 50 ml of an 8M solution was diluted into a 1 Liter container, what would the new molarity be?

M1 = 8 M (8 M)(50 mL) =(M2)(1000 mL) V1 = 0.4 M V1 = 50 mL 1 (8 M)(50 mL) = (M2)(1000 mL) V = 0.4 M. M2 = ? M (1000 mL) (1000 mL) V2 = 1000 mL

6. If a 12M solution was diluted to a 1.25 M solution that had a volume of 2 L, how many liters of the original solution were added? M1 = 12 M (12 M)(V1) =(1.25 M)(2 L) V1 = 0.2083 L V1 = ? L 1 (12 M)(V1) = (1.25 M)(2 L) V = 0.2 L. M2 = 1.25 M (12 M) (12 M) V2 = 2 L

7. If 150 ml of a 4 M solution were added to a solution of 5 L, what would the new molarity be?

M1 = 4 M (4 M)(150 mL) =(M2)(5000 mL) V1 = 0.12 M V1 = 150 mL 1 (4 M)(150 mL) = (M2)(5000 mL) V = 0.1 M. M2 = ? M (5000 mL) (5000 mL) V2 = 5000 mL

8. If you were making a 0.5 M solution of hydrochloric acid in a 3 liter bottle, how many liters of 6 M HCl would need to be added? M1 = 0.5 M (0.5 M)(3 L) =(6 M)(V2) V1 = 0.25 L V1 = 3 L 1 (0.5 M)(3 L) = (6 M)(V2) V = 0.3 L. M2 = 6 M (6 M) (6 M) V2 = ? L

9. If you add .75 L of a stock solution to make 2.25 L of a 1.25 M solution, what was the concentration of the original solution? M1 = ? M (M1)(0.75 L) =(1.25 M)(2.25 L) V1 = 3.75 M V1 = 0.75 L 1 (M1)(0.75 L) = (1.25 M)(2.25 L) V = 3.8 M. M2 = 1.25 M (0.75 L) (0.75 L) V2 = 2.25 L

10. If you were making a 2 M solution of sulfuric acid and you added .25 L of 18 M sulfuric acid, how many liters of the solution were produced?

M1 = 2 M (2 M)(V1) =(18 M)(0.25 L) V1 = 2.25 L V1 = ? 1 (2 M)(V1) = (18 M)(0.25 L) V = 2 L. M2 = 18 M (2 M) (2 M) V2 = 0.25 L

25 Academic Chemistry Unit 10 Test Review Solutions TEST on

Directions: Solve the following problems using the appropriate formula. n – # moles S – Solubility FORMULAS: M = nsolute / Vsolution (in mol/L) VARIABLES: M – Molarity P – Pressure M1V1 = M2V2 S 1 = S 2 V – Volume m – Mass P1 P2

% by mass = msolute/msolution

1. Determine how many grams of potassium dichromate (K2Cr2O7) would be needed to make 250 mL of a 0.10M solution.

2. If 220mL of a 12.0M HCl solution is diluted to 1000mL, what will be the concentration of the new solution?

3. How many grams of methanol (CH3OH) are needed to make 10 L of aqueous 4.0M methanol solution?

4. If a gas has a solubility of .55 g/L at 2.0 atm, how much will dissolve at 12.6 atm?

5. Find the % composition by mass of a 152 g NaNO3 solution containing 34 g NaNO3.

26 6. What volume of 16.0M sulfuric acid must be diluted in order to prepare 1.5L of 0.10M H2SO4 solution?

7. A gas has a solubility of 0.66 g/L at 10,0 atm of pressure. What is the pressure of a sample that contains 1.5 g/L of gas?

8. If enough water is added to a 122mL solution of 6.0M HCl to increase the volume to 244mL, what is the concentration of the new, dilute solution?

9. Carbon Dioxide contains a solubility of 0.878 g/L at 1 atm. What would the solubility of the gas be at 0.2 atm?

10. How many grams of HCl are needed to make a 200 mL of 18.0M hydrochloric acid in aqueous solution?

11. Describe how hydrogen bonds affect the surface tension of water.

12. List factors that will increase the rate of solvation.

27 13. What is the difference between a concentrated and diluted solution?

14. Why is water called the universal solvent?

15. At 60oC how much HCl can be dissolved in 100 grams of water?

o 16. At 50 C how much SO2 can be dissolved in 100 grams of water?

17. Which substance is least soluble at 20oC?

18. Which substance is most soluble at 20oC?

19. How would you describe a solution that o contained 20 grams of NH4Cl at 20 C? (saturated, supersaturated, unsaturated)

20. What is the difference between covalent and ionic bonds?

21. Identify the following as covalent or ionic: a. KCl

b. NF3

c. BeF2

d. PCl5

28 22. CO2 a. Shape:

b. Polar/Nonpolar:

23. CCl4 a. Shape:

b. Polar/Nonpolar:

24. H2O a. Shape:

b. Polar/Nonpolar:

25. KCl a. Shape:

b. Polar/Nonpolar:

26. PCl3

a. Shape:

b. Polar/Nonpolar:

29 Test Review: Solutions Terms and Definitions

DOWN 2 Liquids that do not dissolve in one another 3 Measure of concentration measured in moles of solute per kilogram of solvent 4 Solution containing more solute than can theoretically ACROSS dissolve at a given temperature 1 Maximum amount solute that can be dissolved in a given amount of 5 Mixture that has the same uniform appearance and solvent at a specific temperature and pressure composition throughout 6 Substances able to be dissolved 8 A liquid that is a homogenous mixture of two or more 7 Two liquids that dissolve in each other substances 9 Mixture that is not uniform in composition 10 Compounds that conduct electric current in aqueous 12 Dissolving agent of a solution solution 14 Substance dissolved in a solution 11 Solution containing a large amount of solute dissolved in 15 Solids that dissolve in other solids solvent 18 Homogeneous mixtures in which water is the solvent 13 Substances unable to be dissolved 20 Solution containing maximum amount of solute 16 Solution process 22 Physical blend of 2 or more substances whose compositions vary 17 Bond between two nonmetals where electrons are shared 23 Measure of concentration measured in moles of solute per liter of to fill their outer shells solution 19 Solution containing less than the maximum amount of 24 Says that solubility increases as pressure of gas increases dissolved solute

21 Solution containing a little amount of solute dissolved in solvent 30