Unit 3: Solutions/ Acids and Bases

Matter as Solutions, Acids, and Bases

Examples:Pure:

Heterogeneous:

Homogeneous:

Solutions: made by the homogenous mixture of 2 substances (can only see 1 substance)a) Solute: substance that dissolves (normally a solid, but can be a liquid or gas), or the

substance in lesser quantity. b) Solvent: substance that does dissolving (normally a liquid, but can be a gas), the

substance in greater quantity.

When substances are in the same state, the one with greater volume is considered to be the solvent.

Solutions: 1. Allow reactions to happen faster due to the Brownian motion of fluids

-more movement of particles = more likely particles will react in a favorable way

2. Allow reactions to happen that would otherwise not happen – surface area increases, possible entry shape is changed.

Examples of solutions:

Types of solutions:

DATE: NAME: CLASS:

CHAPTER 5Classification of MatterReview

BLM 5.0.1ASSESSMENT

1. Use the terms provided on p. 160 of the text to complete this flow chart on the classification of matter.

2. Classify the following substances using the terms provided on p. 160 of the text.(a) tap water(b) gold(c) sweetened tea

(d) gasoline(e) ocean water(f) seasoning salt

3. Many ionic compounds are soluble in water. Using the solubility table on p. 161 of the text, determine which of the following ionic compounds would dissolve to form an aqueous solution.(a) NaF (s)(b) AgCH3COO(s)(c) (NH4)2S(s)(d) Ba(OH)2(s)(e) Ca(OH)2(s)(f) PbSO4(s)(g) CuCl2(s)(h) CuCl(s)

Anything with mass and

volume

Has a definite composition

Separated by physical means

H2O(ℓ), CH4(g), NaNO3(s)

H2(g), Ne(g), Cu(s)

Air, brass, stainless steel

Orange juice, chocolate chip

cookies

AgCl(s), Na3PO4(s), KOH(s)

CO2(g), CH4(g), O2(g), P2O5(s)

4. List two properties that are different in molecular and ionic compounds.

Dissolving (forming a solution): A physical change (not chemical change-rxn)

‘like’ dissolves ‘like’ ie. Polar molecules will dissolve in a polar substance. Non-polar molecules will not dissolve in a polar substance. (vice-versa)

Water is the most common solvent used to form a solution. A solute is the substance dissolved into the solvent to form a solution. A solution is a mixture of 2 or more things that appears the same throughout (homogeneous mixture). Solutes and solvents may be gases, liquids, or solids, but we are going to focus on solutions that use water as the solvent ie. Aqueous solutions. They are clear and transparent.

Ionic compounds: positive end of water attracts the negative ions and draws them out of the crystal and surrounds them. The same process occurs for the negative end of water and the positive ions.

Ionic crystals dissociate in water

molecular compounds: the intermolecular forces break and the water is attracted to the polar parts of the molecule and forms H-bonds if possible.

Molecular compounds break into individual molecules in water.(if they dissolve)

Any liquid substance that is cloudy (opaque or translucent) has undissolved particles in it and is NOT aqueous. They are heterogeneous!

See sheet 5.1.2-practice with solubility table

Conductivity Electrolyte: An aqueous solution that conducts electricity.Non-electrolyte: An aqueous solution that does NOT conduct electricity.

If a substance ionizes (separates into ions when dissolved) it will be an electrolyte. This happens for any aqueous solutions of ionic compounds and for strong acids and bases. Some substances only partially ionize and therefore are WEAK electrolytes.

Conductivity can be tested with a conductivity apparatus or an ohmmeter.

Place the following terms in the correct column of the chart below:

aqueous ionic compounds, aqueous molecular compounds, low solubility ionic compounds, insoluble molecular compounds, strong acids, weak acids, C8H18(l)

C6H12O6(aq), NaOH(aq), H2SO4(aq), CH3COOH(aq), CaPO4(s), C2H5OH(aq), NaCl(aq),

Electrolytes Non-electrolytes (or weak)

Text questions and 5.1 review pg 175Lab 5A pg 174

Section 5.2 – Solubility

Solubility: the amount of a solute that dissolves in a given quantity of solvent, at a given temperature. (g/mL)

Things that affect the rate or degree of Solubility:1. Nature of solute and solvent – like dissolves like2. Agitation – solute dissolves more rapidly since fresh solvent is brought into

contact with the solute. Only affects the rate not amount of solute that dissolves.3. Temperature – at higher temp kinetic energy of the solvent is higher so more

collisions of solvent molecules with solute 4. Particle size – smaller particles dissolve faster since there is more surface area

available to solvent5. Pressure (partial pressure) – only affects gas in liquids – solubility of gas

increases as partial pressure of the gas above the solution is increased.

Miscible: liquids will dissolve in each other in all proportions (can never reach saturation point) ex. alcohol and waterImmiscible: liquids that are insoluble in one another. Ex. oil and water

Remember: Ionic compounds break apart into positive and negative ions when they are dissolved in solution. KBr(s) in water K+ (aq) and Br- (aq)

See your new solubility chart.

Slightly soluble ionic compounds are solid crystals in solution e.g. LiF(s)

General rules: all ionic compounds are at least partially soluble All compounds containing group 1 ions or ammonium ions are soluble.

(except LiF) All compounds containing nitrates are soluble. Solubility for solutions is dependent on temperature

Temperature:The solubility of solids increase with temperatureThe solubility of liquids does not change appreciably with temperatureThe solubility of gases decreases with temperature

Can you explain why??

What are the problems with the temperature increasing (thermal pollution) in a river or lake?

Pressure:The solubility of solids does not change appreciably with pressureThe solubility of liquids does not change appreciably with pressureThe solubility of gases increases with pressure.

Explain why SCUBA divers need to rise slowly after diving.

Unsaturated solution: a solution that is not saturated and, therefore, can dissolve more solute at that particular temperature.

Saturated solution: a solution that contains the max.amount of dissolved solute at a given temperature in the presence of undissolved solute. It is said to be in a state of equilibrium.

In equilibrium, the forward and the reverse reaction are occurring at the same rate. In a saturated solution, the dissolved solute will crystallize at the same rate as the undissolved solute dissolves

Supersaturated solution: contains more dissolved solute than its solubility at a given temperature.How can you make a supersaturated solution???

http://gizmodo.com/magically-turn-liquid-soda-into-a-delicious-slushie-wit-525655870

Investigation 5B (blm 5.2.2)pg 180 or solution saturation lab

5.2 review pg 183

5.3 Concentration of solutions

Concentration: conc. of a solution is a measure of the amount of solute that is dissolved in a given quantity of solvent.

Dilute solution: contains small amount of solute Concentrated solution: contains a large amount of solute

Ways to express concentration:

1. Percent by mass = 100

Common use: alloys

Example:

Page 186 practice problems

2. Concentration in ppm = 106ppm

Common use: small solute amounts (toxins; CO detectors, HS gas)

3. Concentration in ppb = 109 ppb

Eg. Symptoms of mercury poisoning become apparent after a person has accumulated 20mg of mercury in his or her body.

a) Express this concentration as parts per million for a 60 kg person.

b) Express this concentration as parts per billion for the same person.

c) Express this amount as a percent by mass for the same person.

Page 188 practice problemsSheet 5.3.2

4. Molarity (Molar concentration) : the number of moles of solute dissolved in one liter of solution. Square brackets

(mol/ L)

(M) = n/V= number of moles of solute (mol)/ number of liters of solution (L)

Page 191 practice problemsExample:

Sheet 5.3.3

Molarity is the most common method used for expressing concentration of aqueous solutions. We will use this method for the rest of the unit.

If asked for ion molarity, watch your ratio of compound/ion.

Concentration of solution vs. Concentration of ions (dissociation/ionization)

With strong electrolytes, we said that the solute ionizes in order to dissolve in water. This means that the compound has separated into ions. Different compounds have different ratios of ions within them. A substance such as NaOH has one Na ion and one hydroxide ion. However, Ba(OH)2 would separate to form one barium ion and 2 hydroxide ions. In other words the same conc. of the solutes would result in twice the amount of hydroxide for the barium compound.(write dissociation equations and show mole ratio)

eg. Page 193 practice problems

Sheet 5.3.55.3 review page 196

5.4 Solution Preparation:

Volumetric Techniques must be used when preparing STANDARD solutions (SOLUTIONS OF KNOWN CONCENTRATIONS).

This involves the use of precision equipment like electronic balances and containers called VOLUMETRIC glassware. They come in a variety of shapes and sizes.

Volumetric flask;

volumetric pipette:

A: to prepare from a pure solid:

use the formula C=n/V to find the moles required,

and then the formula n=m/M to find the mass of solid needed to prepare the solution.

Eg. Calculate the moles of SALT required to make 900mL of 0.225 mol/L sodium phosphate solution. (can be used to remove grease from ovens)

What mass is required to prepare the above solution?

What steps would need to be taken?

Step 1 – Using an electronic balance, measure the mass of solute onto weigh paper. Place the paper on the balance. Press

“tare” or “zero” to re-zero the balance (so that you will be massing only the solute).

Add the solute using a scoopula until you reach the desired mass. Transfer the mass to a beaker.

The weigh paper must be rinsed into the beaker to reduce the amount of solute lost during the transfer.

Step 2 – Add distilled water to the solute in the beaker. Stir with a stirring rod to dissolve.

Add approximately ½ of the total volume of water desired for the final solution (ie: if you are making 100 mL of solution, add about 50 mL – this volume does not have to be precise)

At this point, all of the solute must be dissolved. Add more water if necessary or heat the solution to dissolve all of the solute (do not add more than the total volume of the solution and leave some room for the water from rinsing.)

Step 3 – Transfer your solution to a volumetric flask of appropriate size. A volumetric flask is a pear-shaped

glass container with a flat bottom and a long neck. They are used to make up standard solutions because they are carefully calibrated to the precise volume listed.

Ensure that the volumetric flask is clean before using. You may rinse the flask with distilled water. The flask does not need to be dry.

Use a funnel to transfer the solution from the beaker to the volumetric flask.

Step 4 – Rinse the beaker and stirring rod into the funnel with distilled water. Rinse the funnel into the volumetric flask.

Rinse all glassware that came into contact with the solution into the funnel. This will reduce the amount of solute lost during the transfer.

Thoroughly rinse the funnel into the flask.

Be careful not to add too much water – the level MUST remain below the etched line on the neck of the volumetric flask.

Step 5 – Add distilled water until the bottom of the meniscus reaches the etched line on the volumetric flask. Stopper the flask, keep your thumb on the stopper and invert the flask several times to mix the solution.

Add the rest of the water slowly. The narrow neck of the volumetric flask will fill up quickly.

When the flask is almost full, add the water drop by drop until the bottom of the meniscus rests at the etched line. The flask must be on a level surface and the line must be at eye level.

Depending on the size of the volumetric flask used, the volume will be precise to either one or two digits after the decimal point. Check with your teacher as to the precision of your laboratory glassware.

Page 194,195Sheet 5.3.6

Mixed conc review 5.3.7

B: To prepare a dilute solution from a concentrated solution (Dilutions):

When a more dilute solution is being made from a STANDARD solution then how do you know how much of each should be used? The amount of solute (moles) in the solution drawn into a pipette, is the same amount of solute (moles) in the final diluted solution.Therefore:

n=C1V1 before dilution n=C2V2 after dilution

moles is the same so: C1V1(before)= C2V2(after)eg. A 18.0 mol/L concentrated sulfuric acid is used to make 1.00L of 0.100 mol/L solution. How much of the concentrated acid is needed?

**In dilutions, the number of moles of solute stay the same because you are using ALL of the original solute in solution and just adding more solvent (water).

Step 1 – Rinse the pipette with the standard solution using the technique in Step 2. Discard the solution.

A pipette is a thin glass tube that is used to measure a precise volume of solution. There are two types of pipettes: graduated and volumetric.

Pipette A is a graduated pipette. It is calibrated to measure a number of volumes, like a graduated cylinder.

Pipette B is a volumetric pipette. It is calibrated to measure one volume only, like a volumetric flask.

Step 2 – Draw the desired volume of solution into the pipette. Squeeze some air from the pipette

bulb before lightly placing it over the pipette to draw up the solution (draw up more solution than desired without overfilling the pipette and flooding the bulb).

Remove the bulb and quickly place your index finger over the top of the pipette.

By slightly rolling your index finger to the side, release the solution from the pipette until the bottom of the meniscus is on the line etched into the glass. (This must be done while looking at the line at eye level)

You may need to practise to get the hang of this technique.

Step 3 – Release the standard solution into a clean volumetric flask. DO NOT force the last drop from

the pipette using the bulb. The pipette should be vertical and should touch the side of the flask as shown below.

After the solution has visibly stopped flowing, keep the pipette touching the side of the flask for three seconds to ensure the full volume has had time to exit the pipette.

A small volume of standard solution will remain in the pipette. Pipettes are calibrated in the factory based on how much solution they deliver, not on how much solution they contain.

Step 4 – Add distilled water until the bottom of the meniscus reaches the etched line on the volumetric flask. Stopper the flask, keep your thumb on the stopper, and invert several times to mix the solution.

Page 198 practice problemsSheet 5.4.1Lab page 200

** When taking a small sample of solution from a larger volume of solution, the concentration stays the same. You have NOT used ALL of the solute (total moles has changed), but you do have the same amount of solute per volume (conc is the same)

Eg. A 20mL sample is taken from a 0.25mol/L solution of volume 500mL. What is the concentration of the sample?

What is the amount of solute (moles) present in the sample?

5.3 review page 202chapter 5 review/test

Solution Stoichiometry:

Remember that in our gas unit we learned the four steps to stoich:1. Write a balanced chemical equation/reaction.2. Find the moles of a substance in the reaction.3. Convert the moles of the substance from step 2 to moles of the unknown

substance.4. Find the unknown value.

In steps 2 and 4 we are using equations that we have learned to find the moles, or the unknown value. These equations, so far, were;

Gravimetric(solids); n=m/M for solids

ideal gas law(gases); PV=nRT for gases

NOW with solutions we also have the equation C=n/V that can be used to find moles or the unknown value for a solution (liquids).

Eg. What is the minimum volume of 0.250mol/L MgCl2(aq) needed to precipitate all the silver ions in 60mL of 0.30 mol/L AgNO3(aq)?

Pg 282 practicelab

ACIDS AND BASES: Name them IONIC first!

Classical:

Eg.

IUPAC naming for acids:

Pg 209 naming acids chart and questions

DATE: NAME: CLASS:

CHAPTER 6Naming Acids and Bases Review

BLM 6.1.2ASSESSMENT

1. Fill in the following chart:

Name of pure substance

IUPAC name for acid

Classical name for acid

Chemical formula

hydrogen cyanide

HBr(aq)

aqueous hydrogen sulfate

nitrous acid

HSCN(aq)

hydrogen phosphate

aqueous hydrogen chlorite

hydroiodic acid

H3BO3(aq)

2. USING THE TRENDS YOU DISCOVERED IN QUESTION 1, FILL IN THE FOLLOWING NAMING GRID.

Ending of name of pure substance Classical name of acid

-ate

hydro-----ic acid

-ous acid

3. Name or give the formula of the following bases:

(a) NaOH (c) strontium hydroxide

(b) Mg(OH)2 (d) potassium hydroxide

(c) NH3 (e) cesium hydroxide

Empirical Properties of acids and bases (using senses, not theory)

ACIDS BASES Sour taste Bitter taste

Slippery feelConduct electricity Conduct electricityReact with base to neutralize Reacts with acid to neutralizeReacts with active metals to produce gas Doesn’t react with active metalsPh 1-6.9 Ph 7.1-14

Sheet 6.1.4

Arrhenius Theory of acids: is a substance that ionizes to form hydrogen ions, H+ (aq), when placed in water

Ex. HCl H+ (aq) + Cl – (aq)

Arrhenius Theory of bases: is a substance that ionizes to form hydroxide ions, OH- (aq),when placed in water

Ex. NaOH Na + (aq) + OH – (aq)

Ionization/Dissociation: when acidic molecular substances or soluble ionic substances dissolve(dissociate) in water, they ionize or form ions.

Page 213 questions

Limitations of Arrhenius Definition of acids and bases:

His definition did not include substances such as water or hydrogen carbonate ions, who can play both an acid or a base depending on what it is mixed with. Also doesn’t explain the behavior of substances that do not have OH in their formula like ammonia.

Modified Theory of Acids and Bases: demo lab 6b page 214

Acid: reacts with water to produce H3O + (aq) in aqueous solution.

Base: dissociates or reacts with water to produce OH - (aq) in aqueous solution.

Examples:

Limitations; some substances can behave as both! How do you tell whether it is really an acid or base?

Sheet on oxides6.1 review

Strong vs. Weak Acids: according to the theory of acids and bases , all acids produce hydronium ions in water. However not all acids ionize to the same degree.

Strong acids ionize 100% in water, whereas weak acids ionize less than 50% in water.

Dissociation equation for strong acid (single arrow):

Dissociation equation for weak acid (equilibrium arrow):

This means different amounts of hydronium ion are present for strong vs weak acids.

Strong acids: Have high concentrations of hydronium ions, making their pH closer to 1 Have a high conductivity React quicker React with active metals vigorously producing hydrogen gas.

Weak Acids: Have a low concentration of hydronium ions, making their pH closer to 7 Have a low conductivity React more slowly Have a little reaction with active metals.

Compare amount of ionization/strength.

Strong vs. Weak Bases:

Strong base: completely (100%) dissociates into hydroxide (OH-) ions when dissolved in water. All oxides and hydroxides of the alkali metals and alkaline earths – group 1 and 2 are strong bases. (some are not very soluble in water but what does dissolve will dissociate completely.

Examples: (arrow!)

What would you expect with their: Conductivity? pH? Reactivity?

Weak bases: < 50% of the substance dissociates into ions (resulting in hydroxide ions) when combined with water.

Examples: (arrow!)

Pg 222 #5-7Following sheet

CHAPTER 6Properties of Strong and Weak Acids and Bases

BLM 6.2.4ASSESSMENT

1. Fill in the chart of expected results for acids that all have a concentration of 0.1 mol/L.

Acid Chemical formula

pH(slightly less

than 7 or much less

than 7)

Conductivity(high or low)

Reactivity with magnesium

metal (high or low)

hydrochloric acidethanoic acidboric acidhydrofluoric acidsulfuric acid

HClO4(aq)H3PO4(aq)HBr(aq)H2SO3(aq)

2. Two different acidic solutions have a concentration of 0.1 mol/L. Solution A conducts electricity extremely well, while solution B conducts very poorly. Which of the solutions will have a lower pH? Explain using a description of what is happening on a molecular level.

3. You have two basic solutions. One has a concentration of 1.0 mol/L and one has a concentration of 0.1 mol/L. You know one is a strong base and one is a weak base. Can you determine which solution is which based on pH? If so, explain how. If not, explain why not.

Monoprotic acid – will only give one H + ion (PROTON) when in water.

Ex. HCl, CH3COOH, HBr, HNO3

Polyprotic acid contains two or more hydrogen ions. Each hydrogen ionizes less completely than the previous hydrogen ion.

Ex. H2SO4(aq)

H3PO4(aq)

Sulfuric acid, a strong diprotic acid, is a very important industrial acid. It is so important that the mass of sulfuric acid used by a nation is a measure of the strength of a country’s economy. Sulfuric acid is used in all of the products and processes in this diagram.

Polyprotic Bases: reacts with water in two or more steps to generate hydroxide ions:

Examples: a) sodium carbonateFirst dissociate:

Then write base reactions:

Try writing the base reactions for all steps of potassium phosphate:

Dealing with Spills of acids and basesIf acid spills you can either?

a) Dilute it with water

b) Neutralize it

When would you use each method?

Neutralization reactions: When an acid and base react to form water and an ionic salt (ionic compound that does not produce hydrogen or hydroxide ions in water).

Eg:

All neutralization reactions come down to hydronium (or hydrogen ion) and hydroxide ion reacting to form water. Neutralization reactions are used to maintain how acidic or basic a solution should be (e.g. pool water).

Example:

Do flipchart review: acids and bases

HYDROGEN IONS FROM WATER

Hydroxide ion: a negatively charged ion formed when a water molecule loses a hydrogen ion. (OH-)

Hydronium ion: a positively charged ion formed when a water molecule gains a hydrogen ion or when H+ binds to a water molecule . (H3O+)

In pure water or a neutral substance: H3O+= OH- = 1x10-7 mol/L The concentrations of H3O+and OH- must be equal in pure water.(neutral, pH= 7)Kw= [H3O+] x [OH-] = 1.0 x 10-14 (mol/L)2 Ion product constant for water.

This is true for every aqueous solution regardless of the presence of other ions.

Acidic solutions [H3O+] greater then [OH-] [H3O+] greater than 1 x 10-7

Basic/alkaline solutions [H3O+] less than [OH-] [H3O+] less than 1 x 10-7 mol/L

BOTH Hydrogen(hydronium) and hydroxide ions will always be present in aqueous solutions, even in acids or bases. If we know the conc. of one of the ions we can always find the conc. of the other using Kw.

Example

If H3O+= 1 x 10-5 mol/L, is the solution acidic, basic, or neutral? What is the concentration of OH-?

Sheet 6.3.2

The pH concept (power of hydrogen)

The pH of a solution is the negative logarithm of the hydrogen ion concentration.

pH = -log[H+] or –log [H3O+]

Ex 1: [H+] = 1 x 10-7 mol/L of a neutral solution.

pH= -log( 1x 10-7)= -(log 10-7)pH = 7.0 Neutral!

Ex 2: H+ = 1 x 10-10 mol/l

pH: = -log (H+)= -log (1 x 10-10)= -(0.0 + (-10) )= 10.0

pOH = -log[OH-] Fig 18-6 (p 435)

pH + pOH=14

**SI digit rule

Most solutions fall between 0 and 14 on the pH scale.

Hydronium ion conc

Hydroxide ion conc

pH pOH

Acids

Neutral

Bases

A change in pH of one unit is a 10X difference in strength(conc)

Dilutions of acid or base with water decrease the strength of the acid or base but CANNOT change an acidic solution into a basic one! (or vice versa)

Can you find the concentration when give the pH or pOH?Ex 1. If the pH of a solution is 11.638, what is the

a) pOH?b) [H3O+]?c) [OH-]?

Ex 2. What mass of potassium hydroxide is required to make 250mL of solution with a pH of 11.70?

Read pH and acid deposition on page 238‘pH calculations’ worksheet‘more practice’ worsheetSheet 6.3.4Sheet 6.3.8 after dilution

Indicators are substances that change colour at different pH values

Indicators are used to differentiate amoung acidic, basic, and neutral solutions.

Sheet 6.3.6Lab 6D with indicators-6.3.7

6.3.13 putting it all together6.3.14 review

Titrations 8.3???