Solutions & Solubility

SCH3UChapter 6

Chemistry Literature Connection

• “Water does not resist. Water flows. When you plunge your hand into it, all you feel is a caress. Water is not a solid wall, it will not stop you. But water always goes where it wants to go, and nothing in the end can stand against it. Water is patient. Dripping water wears away a stone. Remember that, my child. Remember you are half water. If you can't go through an obstacle, go around it. Water does.” ― Margaret Atwood, The Penelopiad

Remember: Periodic Trends - Electronegativity

• Electronegativity = a number that describes the ability of an atom to attract electrons•More electronegative = stronger

pull on electrons being shared• Less electronegative = weaker pull

on electrons being shared

Trend: Electronegativity

Incr

easi

ng

Increasing

Difference in Electronegativity

If the electronegativity difference is:• less than 0.4 = bond is non-polar

covalent• is between 0.4 and 1.6 = bond is

polar covalent• is greater than 1.7 = bond is ionic

Types of Bonds

• Non-Polar Covalent = the attractive forces between two atoms that results when electrons are equally shared by the atoms with similar electronegativities

• Polar Covalent = a covalent bond formed between atoms with significantly different electronegativities resulting in unequal sharing of electrons

• Ionic = a bond formed due a large difference in electronegativity between atoms resulting in a complete transfer of electrons

Comparison…

Non-Polar versus Polar Covalent

Structure & Shape of WaterLone PairLone Pair

Structure & Shape of Water

Polarity of Water Molecules

Electronegativity differenceO = 3.5H = 2.13.5 – 2.1 = 1.4

Recall: between 0.2 and 1.6 = bond is polar covalent

Polarity of Water Molecules

Dipole moment =measure the polarity of a chemical bond, occurs whenever there is a separation of positive and negative charges 

Hydrogen Bonding

Ionic Compounds in Water

Ionic Compounds in Water

Polar Molecules in WaterHydrogen Bonding

Polar Molecules (Sugar) in Water

Polar Molecules (Sugar) in Water

Non-Polar Molecules in Water

“like dissolves like”

• Polar substances dissolve in polar solutions•Non-polar substances dissolve in non-

polar solutions• In other words: • like-polarity substances dissolve in each

other

Concentration =

Solution Concentration

𝑞𝑢𝑎𝑛𝑡𝑖𝑡𝑦 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒𝑞𝑢𝑎𝑛𝑡𝑖𝑡𝑦 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Percentage Concentration

• Percentage Weight by Volume

• Percentage Volume by Volume

Very Low Concentrations

• Parts per million (ppm) = units used for very low concentrations

• Used for very dilute solutions• Example:• Toxic substances found in the environment• Chlorine in a swimming pool• Metals in drinking water

Molar Concentrations•Molar Concentrations (C)= the

amount of solute, in moles, dissolved in one liter of solution

C = n/v

•Units = mol/L = M•Molarity of a solution

Solution Concentration• Dilute = having a relatively small quantity of

solute per unit volume of solution

• Concentrated = having a relatively large quantity of solute per unit volume of solution

Making Solutions…

Solution Preparation

• Standard Solution = a solution for which the precise concentration is known

• Used in research laboratories and industrial processes• Used in chemical analysis and precise

control of chemical reactions

Equipment:• Electronic balance precise measurement of

solids• Pipets (pipettes) precise measurement of liquids• Volumetric flask calibrated to contain a precise

volume at a particular temperature, used for precise dilutions and preparation of standard solutions

Preparing a Standard Solution

Accurate Reading of a Volumetric FlaskBend down to see the meniscus

Pipets and Bulb/Pump to transfer small quantities of liquid

Volumetric Pipet

Serological (Blow Out) Pipet

Mohr (Graduated) Pipet

Automatic Dispensers

Micropipets – Dispense µL (microliters)

Preparing a Solution by Dilution

• Dilution = the process of decreasing the concentration of a solution, usually by adding more solvent

• Stock Solution = a solution that is in stock or on the shelf (i.e., available); usually a concentrated solution

Preparation of a Solution of Known Concentration by Diluting a Stock

Solution

Preparation of a Solution of Known Concentration Using a Solid Solute

Calculating the New Concentration of the Diluted SolutionC1 x V1 = C2 x V2

•C1 = initial concentration•V1 = initial volume•C2 = final concentration•V2 = final volume

Sample Problem• Water is added to 0.200L of 2.40mol/L NH3(aq)

cleaning solution, until the final volume is 1.000L. Find the molar concentration of the final, diluted solution.

Solubility• Solubility = a property of a solute;

the concentration of a saturated solution of a solute in a solvent at a specific temperature and pressure

• Saturated solution = a solution containing the maximum quantity of a solute at specific temperature and pressure conditions• No more solute will dissolve, visible solids in solution

Unsaturated vs. Saturated

Unsaturated vs. Saturated• Unsaturated solution = a solution

containing less than maximum quantity of a solute at specific temperature and pressure conditions

• Supersaturated solution = a solution that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances• Make by heatinga solution to dissolve more solute,

then returning the solute to a lower temperature

Supersaturated Solution

Solubility of Solids• Solubility of a substance changes with

temperature• Solids show a higher solubility at higher

temperatures

• Solubility Curve = a graph of solubility and temperature of a solution

Solubility Curve of Solids

Solubility of Gases• Higher solubility at lower temperatures

• Think: Pop• Can of pop from the fridge Gasses dissolvedvs• Can of pop at room temperature Gasses

escape

• Can of pop is also stored under pressure• Increased pressure = increased solubility

Solubility of Gases

Solubility of Liquids•Difficult to generalize about, but:• For polar liquids solubility usually

increases with increase in temperature

• Immiscible = two liquids that form separate layers instead of dissolving•Miscible = liquids that mix in all

proportions and have no maximum concentration

Solubility Categories• High solubility = with a maximum

concentration at SATP (standard ambient temperature and pressure) of greater than or equal to 0.1mol/L• Low solubility = with a maximum

concentration at SATP of less then 0.1mol/L• Insoluble = a substance that has a

negligible solubility at SATP• SATP = standard ambient temperature

and pressure, = exactly 25°C and 100kPa

Using a Solubility Table• Solubility of ionic compounds•Anions paired with particular

cations have either:• High solubility (equal or greater

than 0.1mol/L)OR• Low solubility (less than 0.1mol/L)

Solubility Table

Reactions in Solution• If a chemical reaction results in the formation

of a compound which has low solubility the compound will not be dissolved in the solution and is said to precipitate out of solution

• Precipitate = the solid formed in a chemical reaction or by decreased solubility

Net Ionic Equations• Total Ionic Equation = a chemical equation

that shows all high-solubility ionic compounds in their dissolved form• Spectator = an entity such as an ion, that

does not change or take part in a chemical reaction • Net Ionic Equation = shows only the

reacting species in a chemical reaction and does not include the spectator ions

1.Write the balanced chemical equation with full chemical formulas for all reactants and products.

2.Using a solubility table, rewrite the formulas for all high-solubility ionic compounds as dissociated ions, to show the total ionic equation.

3.Cancel identical amounts of identical entities appearing on both reactant and product sides (spectator ions).

4.Write the net ionic equation, reducing coefficients if necessary.

Writing Net Ionic Equations

Example: Net Ionic EquationSingle Displacement

Zinc metal and aqueous copper (II) sulfate react

1. Balanced Eqn: Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)

2. Total Ionic Eqn: Zn(s) + Cu2+ (aq)

+ SO42-

(aq) Zn2+(aq)

+So42-

(aq) + Cu(s)

3. Net Ionic Eqn: Zn(s) + Cu2+(aq)

Zn2+(aq) + Cu(s)

Chemical Analysis• Quantitative Analysis = the measurement

of the quantity of a substance present • Describes a quantity of matter or degree of change

of matter• Examples: blood alcohol levels, level of toxins in

drinking water• Qualitative Analysis = the identification of

the specific substances present • Describes a quality or change in matter that has no

numerical value expressed• Examples: litmus paper, limewater test, splint tests

for gases

• Most solutions are colourless (solutions of compounds with group 1,2 and 17 ions) but some have distinct colours, particularly monatomic and polyatomic ions of transition metals• See Table 1 & Sample Problem 1 on page 341

Qualitative Analysis: 1. By Colour of Solution of Ionic Compounds

• A flame test involves dipping a clean nichrome or platinum wire into an ionic solid (or solution of the ionic solid) and then holding the wire into a colourless flame. • Depending on the metal ion in the compound,

a different flame colour will be produced• See Table 3 & Sample on page 342

Qualitative Analysis: 2.By Flame Test Colour

• Qualitative Chemical Analysis = the identification of substances present in a sample; may involve several diagnostic tests• The use of precipitation reactions between

on unknown and one known solution. If a certain precipitate forms, this tells the chemist that a certain ion was present in the unknown solution.

Qualitative Analysis: 3.By Sequential Qualitative Chemical Analysis

How to do a Sequential Analysis Involving Solubility

1.Locate the possible cations on the solubility table.2.Determine which anions precipitate the possible

cations3.Plan a sequence of precipitation reactions that uses

anions to precipitate a single cation a t a time4.Use filtration between steps to remove cation

precipitates that might interfere with subsequent additions of anions.

5.Draw a flow chart to assist your testing and communication.

Solution known to contain Pb2+ and/or Sn2+

No Precipitate

No Pb2+ ions were present

Add Na2SO4(aq)

(or CO32-, PO4

3-, SO32-)

White Precipitate No Precipitate

White Precipitate

Solution contained Pb2+ ions,

precipitated as PbCl (s)

Filter to remove solid

PbCl(s)

Add NaCl(aq)

Solution contained Sr2+ ions,

precipitated as SrSO4 (s)

No Sr2+ ions were present

Pb2+(aq) + 2Cl-(aq)

PbCl2(s)

Sr2+(aq) + SO4(aq)

SrSO4(s)

Acids and Bases: Arrhenius Definition•Base = an ionic hydroxide that

dissociates in water to produce hydroxide ions (OH-)

•Acid = a compound that ionizes in water to form hydrogen ions (H+)

Properties of Acids & BasesAcids Bases

Taste Sour Bitter

Feel Do not feel slippery Slippery

pH Less than 7 More than 7

Litmus test Turn litmus paper red Turn litmus paper blue

Ions released into solution

H+ OH-

Reactivity with meals

Corrode metals Do not react with metals

Strong versus Weak Acids• Strong Acid = an acid that ionizes completely (~100%) in water to form aqueous H+ ions

•Weak Acid = an acid that ionizes only partially (less than 50%) in water to form aqueous H+ ions

• Percentage ionization = the percentage of molecules that form ions in solution

pH of a Solution• pH is a way of indicating the

concentration of H+ ions present in a solution• pH = Power of Hydrogen

Calculating Ion Concentration

[H+] = 10-pH

pH = -log [H+]

Type of Solution

pH [H+] Color of litmus

Acidic < 7.00 > 1x10-7 mol/L

red

Neutral = 7.00 = 1x10-7 mol/L

No change

Basic > 7.00 < 1x10-7 mol/L

blue

pH and Ion Concentration

pH Scale

Acid-Base Titration• Titration = a laboratory procedure involving

the carefully measured and controlled adding of a solution from a buret into a measured volume of a sample solution

• Titrant = the solution in the buret during a titration (KNOWN concentration)

• Standard solution = a solution of precisely and accurately known concentration

Titration Apparatus

Titration of a Strong Acid (with a Strong Base)

• In the buret standard solution of base (KNOWN concentration)

• In the flask precise volume of acid (UNKNOWN concentration)• In the flask indicator to detect the end

point

Titration of a Strong Base (with a Strong Acid)

• In the buret standard solution of acid (KNOWN concentration)

• In the flask precise volume of base (UNKNOWN concentration)• In the flask indicator to detect the end

point

Endpoint vs Equivalence Point• Endpoint = the point in a titration at which a

sharp change in a property occurs (e.g. a colour change) • More than neutral

• Equivalence Point (Stochiometric Point) = the point at which the exact amount of the titrant has added which is stoichiometrically equal to the amount of moles of substance (known as analyte) present in the sample• Exactly neutral

The Endpoint

Steps of a Titration1. Place standard solution in buret2. Place a precise volume of a solution of unknown

concentration in a flask3. Add an indicator to the flask4. Record the volume in the buret as your initial reading5. Open the stopcock of the buret and allow the standard

solution to enter the flask, while swirling the flask6. Slow down the flow of standard solution being added to

ensure you don’t surpass the equivalence point by too much

7. Once the end point is reached, record the final volume in the buret

8. Subtract the initial volume from the final volume in the buret to obtain the total volume of standard solution used to neutralize the unknown solution.

Sample Titration Data