chem unit 4 review

UNIT 4: CHEMICAL REACTIONS 💥

Collision Theory and Kinetics

Effective collisions occur when the reactant species collide with enough kinetic energy and in the correct orientation to break the original reactant bonds and form new bonds in the product species.

Factors that affect reaction rate (kinetics):

• The surface area of a solid reactant (greater surface area = faster rate)

• Concentration or pressure of a gaseous reactant (greater concentration or pressure = faster rate)

• Temperature (higher temperature = faster rate)

• The nature of the reactants (states of matter)

• The presence of a catalyst (a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change)

Chemical Equations 📝

2Na(s)+2H2O(l)→2NaOH(aq)+H2(g)2Na(s)+2H2​O(l)→2NaOH(aq)+H2​(g)

This chemical equation shows that the effective collision of two Na atoms with two water molecules produces two sodium hydroxide formula units and one hydrogen molecule.

General notes about chemical equations:

• Chemical formulas represent each species (type of chemical) involved in the reaction.

• The subscript in a chemical formula tells you how many of that atom are present in one unit (molecule or formula unit) of that substance.

• Reactants are written on the left of the arrow (which means "yields" or "chemical change"), and products are written on the right. We do not use "=" in chemical equations.

• If the energy absorbed is larger than the energy released, we refer to the reaction as endothermic, which is noted in the chemical reaction using the Greek symbol uppercase delta (ΔΔ) written on top of the arrow.

• If the energy absorbed is less than the energy released, we refer to the reaction as exothermic.

• The physical states: gas (g), liquid (l), solid (s), or aqueous (aq) follow the chemical formula to identify the phases of all substances in the reaction. Liquid does NOT mean aqueous; aqueous substances are mixtures, while liquids are pure substances.

• Chemical equations are always balanced to respect the Law of Conservation of Mass.

To respect the Law of Conservation of Mass:

Write in stoichiometric coefficients (large numbers to the left of each chemical formula) to ensure the number of each type of atom is conserved. The coefficients and subscripts of the same chemical formula multiply. If there is no visible coefficient, the coefficient is 1.

Balancing Chemical Equations

Step-by-Step Guide

1. Count the total number of each type of atom on each side of the arrow.

   • Keep polyatomic ions as single units when possible.

     • For example, Ca(NO3)2Ca(NO3​)2​ means you have two NO3−NO3−​ units.

2. Balance the quantity of atoms (or polyatomic ions) on each side of the arrow by adding coefficients.

   • Multiply the new coefficient and re-count the total quantity.

   • For instance, to balance the HHs in Na(s)+H2O(l)→NaOH(aq)+H2(g)Na(s)+H2​O(l)→NaOH(aq)+H2​(g), we start by placing a coefficient in front of water and re-counting:

   Na(s)+2H2O(l)→NaOH(aq)+H2(g)Na(s)+2H2​O(l)→NaOH(aq)+H2​(g) Reactant Side:

   • Na=1Na=1

   • H=4H=4

   • O=2O=2 Product Side:

   • Na=1Na=1

   • H=3H=3

   • O=1O=1 Now, to balance the OOs we place a coefficient of 2 in front of NaOHNaOH and re-count. Na(s)+2H2O(l)→2NaOH(aq)+H2(g)Na(s)+2H2​O(l)→2NaOH(aq)+H2​(g)Reactant Side:

   • Na=1Na=1

   • H=4H=4

   • O=2O=2 Product Side:

   • Na=2Na=2

   • H=4H=4

   • O=2O=2

3. Continue adding coefficients until all species are balanced. 2Na(s)+2H2O(l)→2NaOH(aq)+H2(g)2Na(s)+2H2​O(l)→2NaOH(aq)+H2​(g) Reactant Side:

   • Na=2Na=2

   • H=4H=4

   • O=2O=2 Product Side:

   • Na=2Na=2

   • H=4H=4

   • O=2O=2

4. Check to make sure that all coefficients are in their lowest whole number ratio to one another by dividing by the lowest common denominator.

Tips and Tricks for Success

• NEVER EVER EVER change any of the subscripts. Only write in coefficients!

• If there is no coefficient before a formula, it means 1. Writing in the 1 is optional.

• Try to balance HH second to last, and OO last if they are NOT part of a polyatomic ion.

• Keep practicing! The more you practice the faster you'll get at it.

Reaction Types ⚗

Chemists organize chemical reactions by the reactants and products that each type forms.

Synthesis/Addition Reactions

• Two or more substances react to form a single product. A+B→ABA+B→AB

• Examples:

  • 6Li+N2→2Li3N6Li+N2​→2Li3​N

  • C+O2→CO2C+O2​→CO2​

  • PCl3+Cl2→PCl5PCl3​+Cl2​→PCl5​

Decomposition Reactions

• A single reactant is broken down into two or more products. AB→A+BAB→A+B

• Examples:

  • 2NaCl→2Na+Cl22NaCl→2Na+Cl2​

  • 2H2O→2H2+O22H2​O→2H2​+O2​

  • 2Al2O3→4Al+3O22Al2​O3​→4Al+3O2​

Single-Replacement Reactions

• Atoms of an element replace the atoms of a second element in a compound; usually in an aqueous solution (aqaq).

  • Metals replace metals.

  • Nonmetals replace nonmetals. A+BX→AX+BA+BX→AX+B

• Examples:

  • 3Mg+2FeCl3→3MgCl2+2Fe3Mg+2FeCl3​→3MgCl2​+2Fe

  • Cl2+2KI→2KCl+I2Cl2​+2KI→2KCl+I2​

Double-Replacement Reactions

• The exchange of positive ions between two compounds; usually between two compounds in aqueous solution resulting in the formation of two NEW compounds. AB+CD→AD+CBAB+CD→AD+CB

• Examples:

  • AgNO3+LiBr→AgBr+LiNO3AgNO3​+LiBr→AgBr+LiNO3​

  • 2HCl+K2SO3→SO2+H2O+2KCl2HCl+K2​SO3​→SO2​+H2​O+2KCl

  • NaOH+NH4Cl→NaCl+NH3+H2ONaOH+NH4​Cl→NaCl+NH3​+H2​O

  • LiOH+HBr→H2O+LiBrLiOH+HBr→H2​O+LiBr

Combustion Reactions

• An element or a compound reacts with oxygen (O2O2​), often producing energy in the form of heat or light.

• Examples:

  • 2Mg+O2→2MgO2Mg+O2​→2MgO

  • CH4+2O2→CO2+2H2OCH4​+2O2​→CO2​+2H2​O

Special Case: Balancing Hydrocarbon Combustion Equations

• The term combustion is applied to any burning process that adds oxygen to an element or compound.

• Hydrocarbons are covalent compounds that contain, at a minimum, carbon and hydrogen, and often contain oxygen, nitrogen, halogens, and/or sulfur.

• In a complete hydrocarbon combustion, the reaction will always produce carbon dioxide gas and water vapor.

• If there is insufficient oxygen, incomplete combustion results and produces either carbon or carbon monoxide as a product. (Carbon monoxide poisoning is the reason why house fires can be fatal.)

• Balancing this type of reaction can be tricky because when you count the total number of each type of atom, you frequently get an odd number on one side and an even number on the other side as you count. For example, if you follow the set of steps above and balance the CC and HH first (because OO is last), you get:

  C2H6+O2→2CO2+3H2OC2​H6​+O2​→2CO2​+3H2​O

  • # OOs on left =2=2

  • # OOs on right =7=7

  You’re lacking OOs on the reactant side, so a coefficient is needed for the O2O2​ but what? Use 3.53.5 (since 2×3.5=72×3.5=7):

  C2H6+3.5O2→2CO2+3H2OC2​H6​+3.5O2​→2CO2​+3H2​O

  • # OOs on left =7=7

  • # OOs on right =7=7

• Since we do not use fractions or decimals as coefficients, we must multiply ALL coefficients by the denominator (22 in this case) to cancel out the fraction:

  2C2H6+7O2→4CO2+6H2O2C2​H6​+7O2​→4CO2​+6H2​O

  and now the equation is balanced.

• Not all combustion equations will need this trick, but it does come in handy!

Redox Reactions ⚡️

• Recall that there are five types of chemical reactions: synthesis, decomposition, single replacement, double replacement, and combustion.

• Many of these types involve the direct transfer of electrons from one reactant to another.

• These are called redox reactions, where redox means reduction-oxidation.

• In fact, you could also classify all chemical reactions as being either redox or not redox.

• Water can be involved in redox reactions because the water facilitates electron movement, or because it participates directly in the reaction.

Assigning Oxidation Numbers

• To understand how redox reactions work, we first need to learn how to identify oxidation numbers.

• If changes in oxidation number occur during a chemical reaction it means that the reaction is classified as redox.

• Oxidation numbers are used as a bookkeeping tool to determine where the electrons are located in an ion, compound, or molecule. They do not have any actual scientific meaning.

Rules for Determining Oxidation Numbers

Another Example: Assign all the oxidation numbers in the compound, calcium permanganate.

1. Write the chemical formula first. Determine the oxidation numbers for elements on the left and right and then do algebra to determine the oxidation number for the element in the middle.

2. Ca(MnO4)2Ca(MnO4​)2​ is a compound whose oxidation numbers all add up to 00.

3. CaCa is a Group 2A metal, so its oxidation number is +2+2.

4. OO is oxygen, and this compound is not a peroxide, so its oxidation number is −2−2.

5. MnMn is unknown. Do algebra, so x=x= the oxidation number of MnMn.

   +2+2(x+(−2×4))=0+2+2(x+(−2×4))=0 2+2(x−8)=02+2(x−8)=0 2+2x−16=02+2x−16=0, so x=+7x=+7; thus, the oxidation number of MnMn in this compound =+7=+7.

• In some cases, the oxidation number of an element in a compound or ion can be zero, or a fraction. It’s okay! (Sometimes different atoms of the same element in a compound can have different oxidation numbers.) Another Example: Assign all the oxidation numbers in the compound Fe3O4Fe3​O4​.

1. Fe3O4Fe3​O4​ is a compound whose oxidation numbers all add up to 00.

2. OO in a compound is always −2−2, and this is obviously not a peroxide.

3. FeFe is unknown. Set x=x= the oxidation number of FeFe in Fe3O4Fe3​O4​.

   3x+(−2×4)=03x+(−2×4)=0 3x−8=03x−8=0, so x=x= oxidation number of Fe=+83Fe=+38​

Reduction and Oxidation

• At first glance this term seems contradictory to its meaning; however, this chemical change is called reduction because the charge of the reactant reduces, i.e., its charge gets smaller.

• Here’s an example, the reduction of chlorine gas, where e−e− stands for electron. The electrons are also reactants, so expect to see electrons on the left of the arrow:

  Cl2+2e−→2Cl−Cl2​+2e−→2Cl−

• In the particle drawing above, each chlorine atom has gained one electron, which makes the radius larger (due to electron-electron repulsion).

• The chlorine gas, Cl2Cl2​, is an element, so its oxidation number is 00.

• The oxidation number of the chloride ion, Cl−Cl−, is −1−1.

• You need two electrons to oxidize the Cl2Cl2​ molecule because each chlorine atom in the Cl2Cl2​ molecule will gain one of those electrons.

• Since the oxidation number of chlorine reduced from 00 to −1−1, this is a reduction reaction.

• (When you decide on an oxidation number, record it per atom or ion, not all together. Don’t count the coefficient.)

• The charge of the reactant will increase since the reactant is losing negatively charged electrons.

• Here’s an example, the oxidation of sodium metal. The electron is a product, so it will be located on the right-hand side of the arrow.

• Sodium metal has an oxidation number of 00, and sodium ion has an oxidation number of +1+1.

• In the particle drawing below, the NaNa loses one electron, gaining a positive charge in the process.

• Its radius gets smaller since the effective nuclear charge increases.

  Na→1e−+Na+Na→1e−+Na+

• Reduction and oxidation reactions often happen in conjunction with one another, meaning that one reactant will give up electrons (and get oxidized) while the other reactant will gain those electrons (and get reduced).

• Since reduction and oxidation make up two halves of a redox reaction, we refer to them as reduction or oxidation half-reactions.

• One easy way to tell which reactant is which, is to remember the acronyms LEOthe lion says GER (loss-electrons-oxidation, gain-electrons-reduction) or OIL RIG (oxidation is loss, reduction is gain).

Summary

Redox Reaction Example:

2Na(s)+Cl2(g)→2NaCl(s)2Na(s)+Cl2​(g)→2NaCl(s)

• Oxidation numbers: 00 00 +1+1 −1−1

• Reduction

• Oxidation Reduction half-reaction: Cl2+2e−→2Cl−Cl2​+2e−→2Cl− Oxidation half-reaction: Na→1e−+Na+Na→1e−+Na+

• Cl2Cl2​ is reduced to Cl−Cl−

• NaNa is oxidized to Na+Na+

• Cl2Cl2​ gains electrons

• NaNa loses electrons

• The charge of ClCl reduces from 00 to −1−1

• The charge of NaNa increases from 00 to +1+1

Predicting Reactions 🔮

• Reaction prediction is the art of writing balanced chemical equations starting from only the reactants.

• In other words, we will give you a description of the reactants, and the conditions of the change, but we will not tell you the reaction type nor the products.

• You must be able to use patterns (which we will tell you about, in this lesson) to infer what the products will be.

• You must also be very familiar with your periodic table, the states at room temperature of the monatomic and diatomic elements, and the charges of elements on the periodic table.

• Also, knowledge of the Common Polyatomic Ions list is extremely helpful.

Prerequisite Knowledge

• Let’s start with basic knowledge of elements on the periodic table.

• Most of the elements are monatomic and solid at room temperature (25∘C25∘C or 298K298K).

  • The Seven Twins are diatomic: H2(g)H2​(g), N2(g)N2​(g), O2(g)O2​(g), F2(g)F2​(g), Cl2(g)Cl2​(g), Br2(l)Br2​(l), I2(s)I2​(s). Note their states.

  • The noble gases (Group 18) are all, well, nonmetals and gases: He(g)He(g), Ne(g)Ne(g), Ar(g)Ar(g), etc.

  • Two elements are liquids at room temperature (25∘C25∘C or 298K298K): Hg(l)Hg(l) and Br2(l)Br2​(l).

  • The stairstep line divides metals from nonmetals, and elements along the stairstep line are metalloids.

  • The metalloids (all solid at room temperature, 25∘C25∘C or 298K298K) are: BB, SiSi, GeGe, AsAs, SbSb, and TeTe.

  • All elements under and to the left of the metalloids are metals and solid at room temperature (25∘C25∘C or 298K298K) except for H2(g)H2​(g) (a gaseous nonmetal) and Hg(l)Hg(l) (a liquid metal).

  • Everything else is a nonmetal.

Predicting and Writing Synthesis and Decomposition Reactions

• Synthesis (otherwise known as addition) is a reaction type where two or more pure substances react to form a single product.

• The oxidation numbers for an element change as the chemical reaction proceeds from left to right.

• Decomposition is a single compound breaking down to form all elements, an element and a compound, or all compounds.

• Here are the most common patterns associated with synthesis reactions.

• Know that decomposition is the reverse of synthesis.

• Decomposition normally requires a very high heat, so we will reference a Bunsen burner, etc.

• You can also assume that all products will be solid unless the reactants are aqueous.

• If the reactants are aqueous (when mixed) the products would also be aqueous.

• None of the reactants or products will have charges.

Common Patterns

• To help you conceptualize reaction prediction, here are two fully worked out examples.

Another Example: Calcium oxide pellets react with carbon dioxide gas.

1. First, write the chemical formulas for the two reactants. Realize that calcium oxide is a metal oxide, and carbon dioxide is a nonmetal oxide. This fits pattern 3).

2. Calcium has a charge of +2+2 and oxide is −2−2, so its formula is CaOCaO. Pellets mean CaOCaO is solid.

3. Carbon dioxide is CO2(g)CO2​(g). Write an arrow after the two formulas.

   CaO(s)+CO2(g)→CaO(s)+CO2​(g)→

4. Pattern 3) shows that the two compounds combine together in a straightforward manner. Put the OOs at the end. All products are ionic, therefore solid.

   CaO(s)+CO2(g)→CaCO3(s)CaO(s)+CO2​(g)→CaCO3​(s)

5. Balance the equation; however, you won’t have to do anything in this case. All coefficients are 1. This is the answer! Another Example: Carbonic acid decomposes at room temperature.

6. First, write the chemical formula for carbonic acid. Remember your naming convention for acids -ic acid came from -ate, and is a polar covalent compound that begins with a hydrogen.

7. To determine the number of HHs, you must look at the charge of the carbonate, carbonate is a polyatomic ion and can be found on the polyatomic ions chart if you’re not sure of its formula. All acids are aqueous solutions.

8. Then, realize that a temperature is being mentioned and there is no other reactant. This reaction type must be decomposition and fits pattern 4) in reverse. Write an arrow after the formula.

   H2CO3(aq)→H2​CO3​(aq)→

9. Pattern 4) says we will form a nonmetal oxide and water. The nonmetal oxide is CO2CO2​, because the only other nonmetals in the formula are hydrogen and oxygen, which will form water.

10. Carbon dioxide is a well-known gas at room temperature, and we stated in pattern 3) we would use common nonmetal oxides SO2SO2​ and CO2CO2​. Water is a pure substance, liquid, at room temperature.

    H2CO3(aq)→H2O(l)+CO2(g)H2​CO3​(aq)→H2​O(l)+CO2​(g)

11. Double-check that the equation is balanced; then you’re finished! All coefficients are 1. This is the answer!

Predicting and Writing Combustion Reactions

• We have seen some combustion reactions in the previous section.

• Combustion is an exothermic redox reaction involving the addition of oxygen gas.

• Any synthesis involving oxygen gas is also a combustion reaction. See the example in pattern 2).

• We will focus here on the combustion of hydrocarbons.

• Hydrocarbons are organic compounds containing mostly carbon and hydrogen and sometimes oxygen or nitrogen.

• If there is a complete combustion which you can assume the only products will be carbon dioxide gas and water vapor.

• Since we have not covered formula writing for hydrocarbons, we will give you all hydrocarbon formulas.

• Balancing these equations will be a bit more challenging, but overall, this problem type is easy to predict. Example: Liquid benzene, C6H6C6​H6​, burns in oxygen gas.

• Balanced Chemical Equation: 2C6H6(l)+15O2(g)→12CO2(g)+6H2O(g)2C6​H6​(l)+15O2​(g)→12CO2​(g)+6H2​O(g)

Single Replacement Reactions ⇆

The Activity Series

• When we say active we mean reactive, or quick/liable to react easily and quickly with other elements.

• We use this activity series to help predict and write equations for single replacement reactions because it is based upon thousands of experiments and research.

• There is an activity series for metals and an activity series for nonmetals (halogens).

Halogens

• We will discuss the halogens first because their activity series is actually very intuitive.

• It is the Group 17A group, and the most active nonmetals are at the top.

• In other words, the most active halogen is fluorine, and the least active halogen is astatine.

• This is because fluorine has the highest electronegativity and the effective nuclear charge is very high (due to its small atomic size), so it attracts electrons easily.

• Astatine, the least active halogen, has the lowest electronegativity due to its large atomic radius, shielding effect and low effective nuclear charge.

• A single replacement reaction involving two halogens will occur (i.e., you can write an equation for it) if the less active halogen can be replaced by a more active halogen. Here is an example of a single replacement reaction that will occur:

2NaI(aq)+F2(g)→2NaI(aq)+F2​(g)→

• Reading from reactants to products, you should be able to see that fluorine is poised to replace the iodine in NaINaI, because fluorine is more active and at the top of Group 17. A single replacement reaction that will not occur is:

2NaCl(aq)+Br2(l)→2NaCl(aq)+Br2​(l)→

• This is because bromine, which is lower in Group 17 than chlorine, is less active.

• The bromine will not replace the chlorine.

• The Na+Na+ will stay with chlorine because chlorine is more active.

Metals

• Here is an activity series for metals.

  Element
  Lithium
  Potassium
  Barium
  Calcium
  Sodium
  Magnesium
  Aluminum
  Manganese
  Zinc
  Chromium
  Iron
  Cobalt
  Nickel
  Tin
  Lead
  Hydrogen
  Copper
  Silver
  Gold

• It works just like the one for halogens in that the metals closer to the top are more active than the metals closer to the bottom.

• If you take a good look at your periodic table, you should notice that Group 1 metals are at the very top (lithium, potassium), then Group 2 metals, then the transition metals from the left of the d-block to the right of the d-block.

• This goes in general order of chemical reactivity.

• Recall that Group 1 and Group 2 metals are the most reactive metals due to their low electronegativity and low first ionization energy.

• They will not tend to attract electrons and require little energy to release their 1 or 2 valence electrons.

• Here are examples of how this activity series works. Example: K(s)+Zn(NO3)2(aq)→???K(s)+Zn(NO3​)2​(aq)→???

• This reaction should occur because potassium is higher in the activity series and more active than zinc.

• Potassium will replace zinc. Example: Co(s)+NiCl2(aq)→???Co(s)+NiCl2​(aq)→???

• This reaction should occur because cobalt is a more active metal than nickel.

• Cobalt will replace nickel. Example: Ag(s)+CaBr2(aq)→???Ag(s)+CaBr2​(aq)→???

• This reaction will not occur because silver is less active than calcium.

• It will not replace CaCa in the compound CaBr2CaBr2​.

Predicting Steps

1. To predict and write single replacement reaction equations, you will first need to determine if the reaction will occur.

2. If it wont, simply write NRNR (no reaction) - you are finished, and there is no answer to write down.

3. If it does, replace the less active metal/nonmetal with the more active metal/nonmetal, writing the formulas correctly to reflect their charge.

Complete vs Net Equations

• The full, balanced chemical equation is also called a complete equation or a molecular equation.

• Aqueous ionic species dissociate in water to form ions, so we must also take these ions into account.

• Sometimes chemists are only interested in the species that are undergoing change, i.e., are having their oxidation numbers change.

• Any ions that don’t undergo change are called spectator ions, and balanced equations written without these spectator ions are called net ionic equations.

• This does not mean the spectator ions are absent!

• It means the spectator ions are present but not participating in the reaction.

Here are some examples and practice problems. Note that particle diagrams are coming back, since all ions in solution will be hydrated by water molecules.

Example: Predict and write the balanced molecular equation and net ionic equation for the reaction between chlorine gas and aqueous sodium iodide.

Cl2(g)+NaI(aq)→Cl2​(g)+NaI(aq)→ (reaction will happen)

Cl2(g)+NaI(aq)→NaCl(aq)+I2(s)Cl2​(g)+NaI(aq)→NaCl(aq)+I2​(s)

Balanced molecular equation:

Cl2(g)+2NaI(aq)→2NaCl(aq)+I2(s)Cl2​(g)+2NaI(aq)→2NaCl(aq)+I2​(s)

dissociate all aqueous species:

Cl2(g)+2Na+(aq)+2I−(aq)→2Na+(aq)+2Cl−(aq)+I2(s)Cl2​(g)+2Na+(aq)+2I−(aq)→2Na+(aq)+2Cl−(aq)+I2​(s)

cancel out spectator ions:

Cl2(g)+2Na+(aq)+2I−(aq)→2Na+(aq)+2Cl−(aq)+I2(s)Cl2​(g)+2Na+(aq)+2I−(aq)→2Na+(aq)+2Cl−(aq)+I2​(s)

Balanced net ionic equation:

Cl2(g)+2I−(aq)→2Cl−(aq)+I2(s)Cl2​(g)+2I−(aq)→2Cl−(aq)+I2​(s)

Example: Predict and write the balanced molecular equation and net ionic equation for the reaction between potassium metal and aqueous aluminum nitrate.

K(s)+Al(NO3)3(aq)→K(s)+Al(NO3​)3​(aq)→ (reaction will happen)

K(s)+Al(NO3)3(aq)→Al(s)+KNO3(aq)K(s)+Al(NO3​)3​(aq)→Al(s)+KNO3​(aq)

Balanced molecular equation:

3K(s)+Al(NO3)3(aq)→Al(s)+3KNO3(aq)3K(s)+Al(NO3​)3​(aq)→Al(s)+3KNO3​(aq)

dissociate all aqueous species:

3K(s)+Al3+(aq)+3NO3−(aq)→Al(s)+3K+(aq)+3NO3−(aq)3K(s)+Al3+(aq)+3NO3−​(aq)→Al(s)+3K+(aq)+3NO3−​(aq)

cancel out spectator ions:

3K(s)+Al3+(aq)+3NO3−(aq)→Al(s)+3K+(aq)+3NO3−(aq)3K(s)+Al3+(aq)+3NO3−​(aq)→Al(s)+3K+(aq)+3NO3−​(aq)

Balanced net ionic equation:

3K(s)+Al3+(aq)→Al(s)+3K+(aq)3K(s)+Al3+(aq)→Al(s)+3K+(aq)

Reaction of Active Metal with Water

• Before we leave single replacement reactions, there is one unique reaction pattern that is worth mentioning, and that is the reaction of an active metal with water.

• A metal ion## Active Metals and Water

Pattern 7: Active Metal + Water 💧

Example: Sodium metal is added to pure distilled water in a beaker.

• Balanced Chemical Equation:

  2Na(s)+2HOH(l)→2NaOH(aq)+H2(g)2Na(s)+2HOH(l)→2NaOH(aq)+H2​(g)

• Balanced Net Ionic Chemical Equation:

  2Na(s)+2HOH(l)→2Na+(aq)+2OH−(aq)+H2(g)2Na(s)+2HOH(l)→2Na+(aq)+2OH−(aq)+H2​(g)

Double Replacement Reactions and Solubility Rules

Solubility Rules 🤔

• Ionic compounds can be pulled apart by forming ion-dipole forces with the negative and positive poles of water molecules.

• Ionic compounds vary in their ability to dissolve in water.

• If a particular ionic compound is very soluble, it has a high solubility in water.

• An insoluble ionic compound has a very low to little solubility in water.

• Solubility rules determine which ionic compounds are water-soluble.

• Water-soluble compounds will dissociate and form aqueous ions, which enable the resulting solution to be electrolytic (conduct electricity well).

• Strong acids such as HClHCl, HIHI, H2SO4H2​SO4​, HNO3HNO3​, HClO3HClO3​ and HClO4HClO4​ are soluble and strong electrolytes.

Solubility Rules 1-6 📜

Rule #1 takes priority over rules #2-6.

1. All alkali metal (Group IA) and NH4+NH4+​ (ammonium) salts are soluble.

2. All Cl−Cl−, Br−Br−, I−I− salts are soluble, *except for Ag+Ag+, Hg2+2Hg2+2​, Pb+2Pb+2 salts. (Hg2+2Hg2+2​is mercury (I) ion.)

3. All F−F− salts are soluble, *except for Group IIA salts.

4. All NO3−NO3−​, ClO3−ClO3−​, ClO4−ClO4−​, and CH3COO−/C2H3O2−CH3​COO−/C2​H3​O2−​ salts are soluble.

5. All SO4−2SO4−2​ salts are soluble, *except for Ca+2Ca+2, Sr+2Sr+2, Ba+2Ba+2, Ag+Ag+, Hg2+2Hg2+2​, Pb+2Pb+2salts.

6. All CO3−2CO3−2​, PO4−3PO4−3​, C2O4−2C2​O4−2​, CrO4−2CrO4−2​, OH−OH− salts and and all other salts are INSOLUBLE *except for hydroxide salts of Ca+2Ca+2, Sr+2Sr+2, Ba+2Ba+2.

Soluble vs. Insoluble Species 🧪

• Soluble species in water form electrolytes.

• Ionic bonds break, and ions are hydrated by water.

• Free charges in water permit the flow of electric current.

• Dissociation equations represent this process.

Writing Dissociation Equations 📝

1. Positive ions separate from negative ions.

2. Indicate the number of cations and anions as coefficients.

3. Polyatomic ions remain polyatomic.

Example 1: Sodium nitrate (NaNO3NaNO3​)

• All nitrates are soluble in water (Solubility Rule #4).

• NaNO3(aq)→Na+(aq)+NO3−(aq)NaNO3​(aq)→Na+(aq)+NO3−​(aq)

Example 2: Iron (III) sulfate (Fe2(SO4)3Fe2​(SO4​)3​)

• All sulfates are soluble in water (Solubility Rule #5).

• One formula unit will separate into two Fe+3Fe+3 ions and three SO4−2SO4−2​ ions.

• Fe2(SO4)3(aq)→2Fe+3(aq)+3SO4−2(aq)Fe2​(SO4​)3​(aq)→2Fe+3(aq)+3SO4−2​(aq)

• Incorrect: Fe2(SO4)3(aq)→Fe2+3(aq)+(SO4−2)3(aq)Fe2​(SO4​)3​(aq)→Fe2+3​(aq)+(SO4−2​)3​(aq)

Insoluble Species

• Cannot form ions, so they are nonelectrolytes.

Precipitates and Spectator Ions 🪨

• Precipitates form when two soluble salts react in solution to form one or more insoluble products in double replacement reactions.

• Spectator ions do not participate in the formation of the precipitate.

• Spectator ions are present but do not participate in the formation of the precipitate.

Types of Double Replacement Reactions

• Precipitate-forming

• Gas formation

• Water formation

Predicting and Writing Precipitation Reaction Equations

1. Write the correct formulas of the reactants, along with their states.

2. Write the correct formulas of the products based on the charges of the ions.

3. Balance the equation.

4. Consult the solubility rules and assign the correct state symbol to the products.

5. Write the total ionic equation. (Aqueous compounds break up into individual cations and anions. Do not dissociate the precipitate!)

6. Eliminate spectator ions from the total ionic equation.

7. Write the net ionic equation using the lowest common coefficients.

Example: Reaction of potassium chloride and lead (II) nitrate solutions.

1. KCl(aq)+Pb(NO3)2(aq)KCl(aq)+Pb(NO3​)2​(aq)

2. KCl(aq)+Pb(NO3)2(aq)→KNO3+PbCl2KCl(aq)+Pb(NO3​)2​(aq)→KNO3​+PbCl2​

3. 2KCl(aq)+Pb(NO3)2(aq)→2KNO3+PbCl22KCl(aq)+Pb(NO3​)2​(aq)→2KNO3​+PbCl2​

4. 2KCl(aq)+Pb(NO3)2(aq)→2KNO3(aq)+PbCl2(s)2KCl(aq)+Pb(NO3​)2​(aq)→2KNO3​(aq)+PbCl2​(s)

5. 2K+(aq)+2Cl−(aq)+Pb+2(aq)+2NO3−(aq)→2K+(aq)+2NO3−(aq)+PbCl2(s)2K+(aq)+2Cl−(aq)+Pb+2(aq)+2NO3−​(aq)→2K+(aq)+2NO3−​(aq)+PbCl2​(s)

6. 2K+(aq)+2Cl−(aq)+Pb+2(aq)+2NO3−(aq)→2K+(aq)+2NO3−(aq)+PbCl2(s)2K+(aq)+2Cl−(aq)+Pb+2(aq)+2NO3−​(aq)→2K+(aq)+2NO3−​(aq)+PbCl2​(s)

7. Pb+2(aq)+2Cl−(aq)→PbCl2(s)Pb+2(aq)+2Cl−(aq)→PbCl2​(s)

Predicting and Writing Neutralization Reactions

• The other product in a neutralization reaction is a salt.

• A salt is formed when the hydrogen ion of an acid is partly or wholly replaced by a metal. Salts are always ionic compounds.

Example:

HCl(aq)+NaOH(aq)→NaCl(aq)+HOH(l)HCl(aq)+NaOH(aq)→NaCl(aq)+HOH(l)

NaOH(aq)+HC2H3O2(aq)→NaC2H3O2(aq)+HOH(l)NaOH(aq)+HC2​H3​O2​(aq)→NaC2​H3​O2​(aq)+HOH(l)

• For complete neutralization to occur, every H+H+ must combine with one OH−OH−to form one H2OH2​O molecule.

• If the number of moles of reactants are equal, acid-base neutralization always has the net ionic equation:

  H+(aq)+OH−(aq)→HOH(l)H+(aq)+OH−(aq)→HOH(l)

Balanced Neutralization Equations

1. Write the correct formulas for each reactant.

2. Break apart the reactants into ions and swap partners.

3. Recombine the reactants and products balancing the charges with subscripts.

4. Balance the equation using coefficients.

Example: Reaction between nitric acid and barium hydroxide.

1. HNO3(aq)+Ba(OH)2(aq)HNO3​(aq)+Ba(OH)2​(aq)

2. H+(aq)+NO3−(aq)+Ba+2(aq)+OH−(aq)→H+(aq)+OH−(aq)+Ba+2(aq)+NO3−(aq)H+(aq)+NO3−​(aq)+Ba+2(aq)+OH−(aq)→H+(aq)+OH−(aq)+Ba+2(aq)+NO3−​(aq)

3. HNO3(aq)+Ba(OH)2(aq)→HOH(l)+Ba(NO3)2(aq)HNO3​(aq)+Ba(OH)2​(aq)→HOH(l)+Ba(NO3​)2​(aq)

4. 2HNO3(aq)+Ba(OH)2(aq)→2HOH(l)+Ba(NO3)2(aq)2HNO3​(aq)+Ba(OH)2​(aq)→2HOH(l)+Ba(NO3​)2​(aq)