UNIT 6 - Solution ChemistryWatch: Solutions Lesson 1 Solutions and Solubility – This is a “flipped” intro – so what that means is that you watch the video and take notes, for homework. https://www.youtube.com/watch?v=f5a9xPDreRY

Characteristics of solutions:

solute - the material that is being dissolved into something else. Example : when we dissolve sugar into a cup of coffee,

the sugar is the solute.

solvent - the material into which a solute is being dissolved. In the above example, the coffee is a solvent. Solvents can be almost any material or in any state but for the most part, we will be using water as a solvent.

solution - a combination of a solute and a solvent. We can not distinguish the identify a solute and a solvent- it is a HOMOGENEOUS MIXTURE!

A solution can be in any combination of states:

- a solid/liquid combination (dissolving a solid in water such as salt dissolved in water).

- a liquid/liquid combination (table vinegar is a mixture of water and acetic acid).- solid/solid combination (amalgam or any type of alloy).

- Air is an example of gas solution: nitrogen 78.08% oxygen 20.955argon 0.93% Carbon dioxide 0.03%

-solutions having water as the solvent is referred to as AQUEOUS solution (blood, saliva, etc).

There comes a point where no more solute can dissolve into the solution. At this point in time the solution is said to be saturated.

- When solute can still dissolve in a solution, it is said to be unsaturated. - To increase the solubility of the chemical, we can usually heat the solution up. - The result will be a solution that contains more solute than what would normally be dissolved in the solvent. This is called a supersaturated solution.

Solubility of a solute is the maximum amount of solute that can dissolve in a given amount of solution at a given temperature.

- Solubility is temperature dependent. Usually an increase in temperature increases the solubility and increases how fast a solid dissolves in water.

- Solubility is measured in terms of g/L, g/mL, or mol/L (at a given temp.)- Solubility of solids in liquids vary widely:

Does the solubility of sodium nitrate increase or decrease with an increase in temperature?

Which compound has a solubility of 61g / 100g H2O at 70C

Which compound has a solubility of 41g/100g H2O at 50C

How can we increase solubility?

Check for Understanding:

1. In soda, CO2 is the ____ and water is the ____.

2. How would lowering the pressure change the fizziness of a soda?

3. Which term describes a solution in which only a small amount of solute is dissolved?

- There is an upper limit to the solubility of a solid. It important to note that even the least soluble solid will have a few particles dissolve per litre of solution.

- So, we will say that a substance is soluble if 0.1 mol of the substance dissolves per litre.

When there is no apparent limit to the solubility of one substance in another, the components are said to be MISCIBLE.

THIS IS THE END OF YOUR FLIPPED LESSON. SEE YOU IN CLASS TOMORROW

Let’s Make some Rock Candy – Super-Super-saturation https://www.youtube.com/watch?v=HvKJz4M585c

Assignment: Write out the procedure that you followed in order to make your Rock Candy. For each step, explain what is happening and why. It should be put into a table like the one below.Procedure Explanation

Solubility solids / liquids

Solubility of gases

pressure increasesPressure decreasesTemperature increasesTemperature decreases

Electrical Conductivity: Metals (solid and liquid phase): conduct electricity Non-metals (solid and liquid phase): non-conductor of electricity Ionic compound in solid phase: non-conductor

e.g. NaCl(s), Ionic compound in liquid and aqueous phase: conductor Covalent compounds in solid and liquid phase: non-conductor (usually!) Acid and Base: conductor

Electrolytes and Nonelectrolytes An electrolyte is a substance that will conduct electric

current when it is dissolved. It will conduct electric current because it exists as charged particles (ions) when it is dissolved. Good examples: ionic compounds (such as NaCl), acids, and bases.

A nonelectrolyte is a substance that will not conduct electric current when it is dissolved. It will not conduct electric current because it exists as uncharged molecules when it is dissolved. Good example: table sugar (sucrose, C12H22O11).

Note: Ca2+, Na+, and K+ are the primary electrolytes in your body that are responsible for conducting electricity. Your body uses a lot of electricity to function – for instance, every time you move a muscle! The electricity does not travel through wires, though. It travels through your nerves.

Water, a molecular compound, has very few charged particles in it when it is pure. It has a natural concentration of H3O+ and OH- ions of 0.0000001 M! Other than that, it doesn’t contain any ions. Thus, pure water is a poor conductor of electricity.

The electrolyte will allow the bulb to light up when it’s dissolved in solution. The nonelectrolyte will not.

Examples: Which of the following would you expect to form conducting (ionic) solutions?

a) KI b) ICl c) HBr d) CH4 e) LiOH f) CH3OCH3 g) N2O h) HCl

Examples: Which of the following would you expect to conduct electricity?

a) NaCl(s) b) CH3OH(l) c) Cu(s) d) Fe(l) e) K2CrO4(aq) f) C14H10(s) g) CO2(s)

h) KBr(aq) i) KBr(s) j) Ag(s) k) Hg(l) l) Na(s)

Hebden Page 198 #6-8Crash Course Chemistry: https://www.youtube.com/watch?v=AN4KifV12DA&index=7&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr

Molarity Vs Molality

END OF LESSONPolarity of Molecules:

The attractive forces that exist between molecules are known as intermolecular forces. These include ionic interactions, dipole-dipole interactions and dispersion or London dispersion forces. Dipole-dipole interactions and dispersion forces are weaker than thermal energy (2.4 kJ/mole) at room temperature and are referred to as Van der Waals Force.A. Van der Waal’s Forces (intermolecular interaction):

1. Dipole-dipole interactions:

molecules that are dipolar are attracted electrostatically with one another.

Polar molecules have a partial positive and a partial negative portion of the molecule.

They are made of atoms having differences in ELECTRONEGATIVITY.

Polar molecule must also be ASYMETRICAL.

2. London Forces: https://www.youtube.com/watch?v=3t1Jn_jrsQk

are weak attractive forces which arise as the result of temporary dipolar attraction between neighbouring atoms.

London Forces are always present, even in species that have dipolar attraction.

This is a weaker attraction than dipole-dipole forces.

B. Hydrogen Bonding (intermolecular interaction):

Hydrogen bonding occurs when hydrogen is bonded covalently to N, O, or F (high electronegativity). Highly polar.

Hydrogen bonding is stronger than the Van der Waal’s Forces. The presence of H-O- , H-N-, H-F indicate a highly polar

section of the molecule. Asymetrical molecules are POLAR molecules. *****Hydrogen bonds are not actually bonds at all, they are intermolecular attractions !!!

Hebden Page 198 #9-16Electronegativity of elements

Covalent bonding is the sharing of electrons. Unless the two atoms sharing the molecules are identical, the sharing will not be equal

The more electronegative atom will attract the electrons more strongly than the less electronegative atom. This will result in the more electronegative atom having a slight negative charge (-) The less electronegative atoms will be slightly deficient in electrons and thus have a

slight positive charge (+)

This type of covalent bond in which atoms have slight electrical charges are called polar bonds

Eg.

Electronegativity of elements can be judged from their position on the periodic table.

All elements involved in covalent bonds do have quite high electronegativities.

B & Si < P & H < C & S & I < Br < Cl & N < O < F

High Very High Extremely High

The greater the difference in electronegativity of the atoms involved the greater the polarity of the bond.

In many molecules, the polar bonds result in the molecule have a resultant dipolei.e. there is a positive and negative end to the molecule

In some molecules, the effects of the polar bonds cancel out because of the symmetry of the molecule.

For a molecule to be polar It must contain polar bonds Its shape must be such that the centre of positive and negative charges are not in the

same place

Shapes of molecules Take Notes while watching “ VESPR: Theory Introduction” Video by Tylar DeWitt on You Tube. Video is on the wvssearland website. (https://www.youtube.com/watch?v=nxebQZUVvTg)

VSEPR:

No. of “things” around central atom

Bond Paris and Lone Pairs

Bond Angles

What is “inside” s covalent bond?________________________________

Why is there Repulsion?_______________________________________

2 2 bp 180 LINEAR

3 3bp 120 TRIGONAL PLANAR

2bp + 1lp

116 BENT

4 4bp 109.5 TETRAHEDRAL

3bp + 1lp

107 TRIGONAL PYRAMIDAL

2bp + 2lp

105 BENT

https://www.youtube.com/watch?v=keHS-CASZfc

The shapes of molecules are determined by the repulsion between the electron pairs in the valence level.

This is known as the Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) In the most common molecules, filled valence levels contains four pair of electrons.

Some molecules also contain non-bonding or lone pairs of electrons. These lone pairs affect the shape of the molecule.

In ammonia, NH3, there are three pairs of bonded electrons and one lone pair. Without the lone pair, the molecule would look like an equilateral triangle with a

nitrogen atom at its centre. The lone pair that is not involved in bonding repels the bonding electrons so the

molecule has the shape of a trigonal pyramid. In water, there are two pairs of bonded electrons and two lone pairs. The two lone pairs cause the molecule to be bent, not linear as would be expected

without the lone pairs.

Resonance StructuresOften two or more equivalent structures can be drawn for a molecule. These are known as resonance structures.

Resonance refers to the arrangement of valence electrons in molecules or ions for which more than one Lewis structure can be written. The actual molecule that exists is often referred to as a resonance hybrid of these structures.

Eg. O3

Resonance does NOT mean than the molecule flips from one structure to another. The bonds actually have lengths and strengths intermediate between those of single and double bonds or double and triple bonds.

Steps for writing Lewis structures

1. Write the correct arrangement of the atoms using single bonds. Where necessary, apply the following guidelines.

a) Smaller, more electronegative non-metal atoms surround larger, less electronegative non-metal atoms.

b) Oxygen, hydrogen, and/or halogen atoms often surround a central metal or non-metal atom in a symmetrical arrangement.

c) Carbon atoms are usually bonded to each other.d) Oxygen atoms are bonded to each other only in peroxides (or superoxides)e) In most acids, such as H2SO4, and in many other compounds that contain both oxygen and

hydrogen atoms, the hydrogen atoms are all bonded to oxygen atoms.

2. Find the total number of valence electrons. Add together the number of valence electrons contributed by each atom. If the species is an ion, subtract one electron for each unit of positive charge or add one electron for each unit of negative charge.

3. Assign two electrons to each covalent bond.

4. Distribute the remaining electrons so that each atom has the appropriate number on nonbonded electrons. For elements from the second period, other than beryllium and boron, this is the number of electrons needed so that each atom is surrounded by an octet. For elements of the third period and beyond, except aluminum, this is often the number of electrons needed to complete on octet, although extra electrons can also be placed around atoms if these elements when they are the central atoms in compounds. Remember that atoms bonded to a central atom usually obey the octet rule.

5. If there are not enough electrons to go around, change some of the single bonds to multiple bonds. Multiple bonds can be written between carbon, nitrogen, oxygen, sulphur, selenium and phosphorous atoms. (Note that beryllium, boron, and aluminum do not form multiple bonds.

Steps in using VSEPR to predict geometry

1. Write the Lewis structure of the molecule.

2. Determine the number of bonding pairs and only pairs of electrons around the central atom.

3. Determine the ideal geometry. Then, if necessary taking into account the presence of lone pairs, predict the actual shape of the molecule

4. Keep in mind that lone pairs occupy large site (equatorial in molecules derived from AB5 molecules) or, when sites are equal, occupy sites opposite rather than next to each other.

VSEPR WorksheetEND OF LESSON 2

Polar and Non-Polar Solvents:

Polar Molecules: - Dipole is present- Molecule is ASYMETRICAL (not symmetrical).

Hebden p. 204-205 – complete the table below

Solvent Molecular drawing

Polar or Non-Polar

Water

Methanol

Ethanol

Benzene

Ethoxyethane

Acetone

Solvent Molecular drawing

Polar or Non-Polar

Acetic acid

Chloroform

Carbon tetrachloride

Heptane

Liquid ammonia

“Like dissolves like” :

polar solvents dissolve polar solutes and ionic compounds Non-polar solvents dissolve non-polar solvents

EXPLANATION:The “dissolving process” involves three different attractions

Solvent molecule other solvent molecules Solvent - Solvent Solvent molecule solute particles Solvent - Solute Solute particles other solute particles Solute - Solute

Energy must be added to break a bond and is released when bonds form. If Solvent - Solute bond energy is > Solute - Solute, then the nrg from forming the new bond can be used to break the old one. Strong Solvent – Solvent then allows solutes to be mixed within the solvent = dissolving

Explanation ConclusionSolubility of Polar & Ionic Solutes

Solubility of Nonpolar Solutes

A. Polar and ionic molecules are soluble in polar solvents.

ex: Salt (ionic) dissolves in water (polar solvent).Ex: I2(s) is more soluble in water than gasoline.

non-polar molecules are soluble in non-polar solvents.

B. Polar and ionic compounds are generally held together quite tightly.

Polar solvents are strongly attracted to the ions or the polar compounds causing separation of the solute out of its crystal structure and into solution.

Non-polar solvents do not have this attraction, therefore, polar solutes are unable to separate into a non-polar solution.

C. Non-polar solutes are held together by London Forces.

Solvents with enough London Forces to attract the solutes and break off the solid crystal structure are capable of dissolving the solid. Generally speaking, non-polar solvents have high London Forces, and dissolve other non-polar compounds.

Hydrogen bonding holds water molecules together. For example, water molecules undergo hydrogen bonding. When water freezes, the six sided structure of water crystals is due to hydrogen bonding.

Hebden p. 207 # 18-22Read – 207-208 then do #23-27

END OF LESSON 3 – QUIZ #1 UP TO THIS POINT!The nature of Solutions of ions:

The formation of solution depends on the ability of the solute to dissolve in the solvent.

Ionic solid are crystal structure made up of ions. Covalent solids are crystal structure made up of covalent molecules. Crystal lattice is an orderly arrangement of particles that exists within a crystal.

Water: “Universal Solvent”, polar, V-shaped molecule. Its positive end attracts the negative ions and the negative end attracts the positive ions.

Dissociation: reaction that involves the separation of ions in an ionic compound.

Ionization: reaction that involves the breaking up of a neutral molecule into ions.

Read p. 133-148, Zumdahl. Do Exercises: p. 180-181 # 11, 13, 15, 17, 25

Hebden p210 #28-29END OF LESSON 4

Calculating The concentration of Ions in Solution:

concentration - this is the amount of solute that is in a certain amount of solution. Because 'amount of solute' can be interpreted in a number of different ways (i.e. grams, moles, molecules etc.) we will define a value for concentration :

This is also defined as MOLARITY.

The units for molarity are moles/litre, mol/L or M for short.Examples : 1) What is the concentration of a solution containing 0.055 moles of HCl dissolved in water to make 250.0 mL of solution

2) What is the concentration of a solution containing 0.055 grams of HCl dissolved in water to make 250.0 mL of solution

3) How many molecules of NaOH are dissolved in 125.3 mL of a 3.00 M NaOH solution ?

How to make a solution of a certain concentration:In order to make a solution of certain concentration, we do NOT dissolve the amount of

solute needed in the volume of solvent. If we did that, the volume would increase somewhat and we would not have the precise concentration we require. In order to accomplish this, we simply determine the amount of solute needed, dissolve this in about 1/2 the amount of solvent and then fill up to the proper amount of solution. It is best to use a volumetric flask for this purpose.

Example : Describe how you would make 125 mL of a 0.250 M solution of sodium hydroxide.

Hebden p103 #95 - 98Solutions wks #1

END OF LESSON 5 Ionic concentrations

When a solid dissolves in a solvent it can act one of three ways - it can either dissociate into its ions (only if it is an ionic compound);it can remain as a molecule dissolved in water (if it is not an ionic compound);or it can react with the water (we will not consider this at the present time). For the sake of Chemistry 11, we will discuss only what happens when an ionic substance dissolves in a solvent.

First of all, why do some ionic compounds dissolve readily in water while some others appear to not be as readily soluble?

The reasons for this are due to the relative shapes and structures of both the solute and solvent molecules. Sometimes the molecules have a slight charge on one side of the molecule compared to the other. This type of molecule is called POLAR. Water is a good example of this. A NON-POLAR molecule, such as carbon tetrachloride, can also be used as a solvent. The polarity of the molecules is the basis of dissolution, but we will not get into that aspect of chemistry in this course.

When an ionic compound dissolves in water, it breaks into its ions. Example :

NaCl (s) ----> Na+(aq) + Cl-(aq)

CaF2 (s) ----> Ca2+(aq) + 2 F-(aq)

We are going to assume that all of the compound breaks into its ion state. This is not always true, but we will leave that for Chem 12.

If all of the compound ionizes, then we actually have none of the solute molecules left in the solution. By looking at the ratios between the ions and compound (from the dissociation equation) we can determine the concentrations of each individual ion in the solution.(The abbreviation for concentration of an ion is to place the symbol for the ion inside square brackets. For instance, shorthand for 'the concentration of the hydrogen ion' is [H+].)

Examples :Find the concentration of each ion in the following solutions :

a) 12.0 M HCL b) 18.0 M H2SO4 c) 1.15 M Cr2(SO4)3

DilutionsWhen a solution of known concentration has its volume changed, it stands to reason that

its concentration changes also. If the liquid we are adding does not contain any of the original solute OR if a reaction does NOT take place, the amount of dissolved solute does not change - only the volume changes. In fact the concentration will change by a factor equal to the change in volume. In short, the new concentration can be found by using the following formula :

C1V1=C2V2

old volumenew concentration = old concentration X --------------------

new volume

Examples :1) 125 mL of 3.00 M HCl has 35.8 mL of water added to it. What is the new concentration ?

Solution :

2) 15.0 mL of a solution of 1.50 M HCl has some water added to it changing the concentration of the acid to 1.15 M. What volume of water was added ?

Solution :

3) How much 6.00 M NaOH would we have to use to produce 275 mL of 0.500 M NaOH ?

Solution Worksheet #2Mixing together solutions of different concentrationsAs we said last class, if we mix together two solutions of different concentrations together,

the formula we used can NOT be used. In order to find the new concentration, we must find the amount of solute in each solution.

Example : Determine the new concentration of HCl if 225 mL of 1.50 M HCl is mixed with 115 mL of 2.50 M HCl.

Solutions Worksheet #3

END OF LESSON 6 – QUIZ #2 NEXT DAYReactions in Solutions

1. Precipitation Reaction:

Demo: Mix 0.2 M Pb(NO3)2 + 0.2 M KI. What ions are involved in this mixture?

The ions in solution are free to move around. If a reaction occurred, what ion combinations would be possible?

The Chemical reaction:

The above reaction is a Double replacement reaction. The way the equation is written is called CHEMICAL EQUATION.

On the product side of the equation we have a yellow solid, PbI2 (we'll figure out how to determine which is the precipitate later) and a solution of KNO3. Rewriting the equation to show what we really have in solution:

This is called an OVERALL IONIC EQUATION and is not used that often because of a number of reasons - not the least of which is that it takes up too much space !!

If we look carefully at the overall ionic equation, we see that there are 2NO3-(aq) ions and 2K+

(aq) ions on either side of the equation. If we cancel them out, we get

This is called a NET IONIC EQUATION for the reaction shown above. It shows exactly what is happening in the solution. The NO3-(aq) ions and the K+(aq) ions are actually doing nothing in the reaction. They can be thought of as just watching the lead ions and the iodide ions get together. For

this reason, they are referred to as SPECTATOR IONS. Spectator ions are never included in net ionic equations.

Another example : CuSO4 + NaOH ---> - assume that the precipitate is Cu(OH)2

overall ionic equation:

net ionic equation:

Example : if we knew we got a precipitate of Ca3(PO4)2 in a reaction, the net ionic equation would be:

Examples: Write the a) chemical equation, b) overall ionic equation, and c) net ionic equation when the following chemicals are mixed together:

a) NaOH + Fe(NO3)3 (solid formed is Fe(OH)3)

b) CaCl2 + Na2CO3 (solid formed is CaCO3)

c) AgNO3 + CaCl2 (solid formed is AgCl)

d) Na2SO4 + Ba(NO3) (solid formed is BaSO4)

Solutions Worksheet #4

Reactions in Solutions

2. Acid Base Reaction:This should be review for you!!! In short, the chemical properties of solutions of acids and

bases are as follows:

Acids :

1) taste sour2) conduct an electrical current3) cause certain dyes to change colour (e.g. litmus goes red)4) liberates hydrogen when it reacts with certain metals5) loses the above properties when mixed with a base though the resulting solution conducts electricity.

6) INCREASES [H+]

Bases :

1) taste bitter2) conduct an electrical current3) cause certain dyes to change colour (e.g. litmus goes blue, phenolphthalein goes pink)4) feels slippery5) loses the above properties when mixed with a base though the resulting solution conducts electricity.

6) INCREASES [OH-]

In their respective chemical formula, and more importantly when they dissolve in water, acids have an H+ ion while bases have an OH- ion. (This is an incomplete definition of an acid and a base but will suffice for chemistry 11.)

When an acid and a base react, the result is a salt and water. The type of reaction is called a NEUTRALIZATION reaction because the properties of the acid and the base are neutralized by each other. The net ionic reaction for this reaction is as follows :

H+(aq) + OH-(aq) ---> H2O (l) We must remember that in any neutralization reaction, the moles of H+ MUST equal the moles of OH-.

Example : What volume of 6.00 M HCl must be added to 125 mL of 1.59 M NaOH in order to neutralize it ?

Example : What volume of 0.10 M Ca(OH)2 is needed to neutralize 115.2 mL of 0.55 M H3PO4?

Incomplete Neutralizations:

When acids and bases are mixed together and the moles of OH- and the moles of H+ are not equal to one another.

Moles of OH- or moles of H+ are going to be left over in solution.

Example: Calculate the [OH-] and [H+] which results when 50.0 mL of 0.150 M NaOH is added to 50.0 mL of 0.200 M HCl.

Note that the concentration of OH- is not really equal to zero. But we will not deal with this concept this year.

Example: Find the number of moles of HCl that must be added to 40.0 mL of 0.180M NaOH to produce a solution having [OH-]=0.0316M.

Solutions Worksheet #5END OF LESSON #7

Formation of precipitates

We will use the qualitative definition here that SOLUBLE means that we can make at least a 0.1 M solution of the compound at 25oC without producing as solid. How can we predict if a precipitate will form ?

The answer is not easy. In fact we don't have too many rules that always apply. Chemists have done thousands of experiments and have collated the results into pages and pages of data tables.

A table showing the solubility trends will be provided at first, but you will eventually need to know the solubility trends. To use this table, we determine what the product will be in a reaction (remember, the reactions here are double replacement reactions.) From here we must check out both products to see if there is a precipitate formed.

Examples : 1) Identify the precipitate (if any) when the following solutions are mixed:

a) lead (II) nitrate and sodium chloride

b) ammonium hydroxide and copper (I) sulfate

c) potassium phosphate and cesium sulfide

2) Write the net ionic equation for the formation of any precipitate when solutions of the following are mixed. If no precipitate is formed, write NO REACTION.

a) barium sulfide and ammonium carbonate

b) silver nitrate and sodium acetate

c) iron (III) chloride and copper (II) chloride

d) lead (II) acetate and lithium phosphate

e) Calcium nitrate and sodium hydroxide

f) lead (II) nitrate and potassium iodide

g) chromium (III) chloride + sodium hydroxide

3. When the following solutions are mixed together, what precipitate (if any) will form? If a reaction occurs, write the net ionic equation.

a) barium chloride + sodium sulphate:

b) lead (II) nitrate + potassium chloride:

c) silver nitrate + sodium phosphate:

d) sodium hydroxide + ferric nitrate:

e) ferrous sulphate + potassium chloride:

f) aluminum nitrate + barium hydroxide:

g) calcium chloride + sodium sulphate:

h) potassium sulphide + nickel (II) nitrate:

4. Are the following compounds soluble or low in solubility?a) Fe(OH)3 b) Hg2Cl2c) PbSO4 d) BaCrO4

e) sodium sulphate Bf) aluminum hydroxideg) KNO3 h) NaNO3

i) FeSO4 k) CuCl2Solutions Worksheet #6

END OF LESSON 8 – QUIZ #3 & Lab 16D Next Day

Learning Objectives for the Solutions Unit Test

Students should be able to define solute, solvent, solution, saturated, unsaturated and supersaturated and solubility.

Students should be able to explain the super saturation demo and factors that affect solubility as well as be able to read a solubility chart

Students should understand the connections between bond theories and the conductivity of ionic and covalent aqueous solutions

Student should be able to identify polar and non-polar solvents and use this to predict solubility of give solutes Students should be able to use lewis structures in conjunction with VSEPR to predict the shape and bond angles

for molecules with up to 4 central bonds Students should understand what resonance structures are Students should recongize and be able to explain the differences between dissociation and ionization Students must be able to calculate molarity, dilutions, concentration of ions in solutions, and mixing of solutions

as well as stoichiometry involving titrations. Students should be able to write overall and net ionic equations, identify spectator ions and predict the

formation of precipitates using a solubility chart.

Review Worksheet #1Review Worksheet #2