AP Chapter 2

Chemical Equations and Reactions

Formulas

• H2O

• Ba(C2H3O2)2

Empirical and Molecular Formulas

Substance Molecular Formula

Empirical Formula

Water H2O H2O

Benzene C6H6 CH

Acetylene C2H2 CH

Glucose C6H12O6 CH2O

Atomic mass unit

• The atomic mass unit (u), is a unit of mass used to express atomic and molecular masses.

• It is the approximate mass of a hydrogen atom, a proton, or a neutron.

• By definition the atomic mass unit is equal to one-twelfth of the mass of a carbon-12 atom.

The Atom

• The nucleus is very small, dense, and positively charged.

• Electrons surround the nucleus.

• Most of the atom is empty space

Subatomic ParticlesPARTICLE SYMBOL CHARGE MASS

(u)LOCATION

electron e- -1 0 orbit nucleus

proton p+ +1 1 inside nucleus

neutron n0 0 1 inside nucleus

Atomic Number (Z)

• The number of protons in the nucleus of an atom.

• The identifying characteristic of an element.

Mass Number

• The sum of the protons and neutrons in the nucleus of an atom.

Isotopes

Isotopes

• This is the symbol for carbon-12.

• Atomic number is 6.

• Mass number is 12.

Isotopes

• This is the symbol for carbon-12.

• Atomic number is 6.

• Mass number is 12.• Write the symbols for

carbon-13 and carbon-14.

Isotopes

What is the average mass of a carbon atom?

What is the average mass of a carbon atom?

12.01

Atomic Mass

• The atomic mass of carbon is 12.01u.

Atomic Mass

• The atomic mass of carbon is 12.01u.

• Atomic mass is the average mass of all the isotopes of an atom. It takes into account the different isotopes of an element and their relative abundance.

• How many electrons, protons and neutrons are in an atom of actinium with a mass number of 221?

• How many electrons, protons and neutrons are in an atom of actinium with a mass number of 221?

• 89p+

• 89e-

• 132n0

• How many electrons, protons and neutrons are in an atom of rhodium-105?

• How many electrons, protons and neutrons are in an atom of rhodium-105?

• 45p+

• 45e-

• 60n0

Families of the Periodic Table

The Noble Gases

• Elements in group 18• All are gases.• VERY non-reactive.• Have a full outer

energy level.

The Octet Rule

• The octet rule states that an element's outer energy level is full and most stable when it contains eight electrons.

• This stability is the reason that the noble gases are so non-reactive.

Exception to the Octet Rule

• The first energy level can only hold two electrons and so elements such as Hydrogen and Helium that only have one energy level follow a “duet rule”.

Ion Vocabulary• An ion is an atom or group of atoms that have a

charge. • A monatomic ion is an atom with a charge. • The charge on the atom is called an oxidation

number.• A polyatomic ion is a group of atoms with a

charge.• A cation is a positive ion.• An anion is a negative ion.

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.

anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons Cl-

17 protons18 electrons

2.5

Ionic

• When an element that easily loses electrons (a metal) reacts with an element that easily gains electrons (a nonmetal), one or more electrons are transferred.

• This creates two ions which are held together by an ionic bond.

• A compound that contains ions is called an ionic compound.

a

A formula unit is the empirical formula of an ionic compound. It is the lowest whole number ratio of ions represented in an ionic compound. Examples include ionic NaCl and K2O.Ionic compounds do not exist as individual molecules; a formula unit thus indicates the lowest reduced ratio of ions in the compound.

Formula Unit

Covalent

• When atoms share electrons the bond created is said to be covalent. Covalent bonds often form between nonmetal atoms.

• These covalently bonded atoms act as single units called molecules.

• A compound made up of molecules is a covalent compound.

Polyatomic IonsH2PO4

- Dihydrogen Phosphate C2H3O2

- Acetate HSO3

- Hydrogen Sulfite (Bisulfite) HSO4

- Hydrogen Sulfate (Bisulfate) HCO3

- Hydrogen Carbonate (Bicarbonate) NO2

- Nitrite NO3

- Nitrate CN- Cyanide OH- Hydroxide MnO4

- Permanganate ClO- Hypochlorite ClO2

- Chlorite ClO3

- Chlorate ClO4

- Perchlorate HPO4

2- Hydrogen Phosphate C2O4

2- Oxalate SO3

2- Sulfite SO4

2- Sulfate CO3

2- Carbonate CrO4

2- Chromate Cr2O7

2- Dichromate SiO3

2- Silicate PO3

3- Phosphite PO4

3- Phosphate NH4

+ Ammonium Hg2

2+ Mercury(I)

Note that these are charges and not oxidation numbers.

Types of monatomic ions and the rules for naming them

• The periodic table is useful in naming the monatomic ions.

Monatomic cations with one oxidation number

• The cations from the periodic table which have a single oxidation number are as follows: Group 1 (+1), Group 2 (+2), Ag+, Cd2+, Zn2+, and Al3+.

• These types of ions are named by using the name of the element followed by the word ion.

• Na+ sodium ion• Ba2+ barium ion• Zn2+ zinc ion

• We can use the roman numeral from the periodic table to identify the oxidation number for these ions.

Monatomic cations with multiple oxidation numbers

• All other cations that are not listed in the previous category (“cations with one oxidation number) are considered to have the possibility of multiple oxidation numbers.

• These type of ions are named by using the name of the element followed by a Roman numeral to indicate the oxidation number.

• Cu2+ copper (II)• Pb4+ lead (IV)• Mn7+ manganese (VII)

Monatomic anions

• All anions from the periodic table are named by changing the ending of the element’s name to –ide.

• F- fluoride ion

• O2- oxide ion

• N3- nitride ion

• Count back from the noble gases starting at zero to determine the oxidation number.

Name these monatomic ions

• Rb+

• P3-

• Fe3+

• Br‾

• Mn4+

• Cd2+

Write the formula for these monatomic ions.

• Sulfide

• Lead (II)

• Barium ion

• Chromium (IV)

• Aluminum ion

• Carbide

Formulas of Ionic Compounds:

• Formulas for ionic compounds can be written by the following steps:

Formulas of Ionic Compounds:

• Formulas for ionic compounds can be written by the following steps:

• (1) Write the formula for the cation and anion (Don’t forget to include the charge of each ion).

Formulas of Ionic Compounds:

• Formulas for ionic compounds can be written by the following steps:

• (1) Write the formula for the cation and anion (Don’t forget to include the charge of each ion).

• (2) Decide how many cations and anions are needed so that the sum of their charges balances out to be zero.

Formulas of Ionic Compounds:• Formulas for ionic compounds can be written by

the following steps:• (1) Write the formula for the cation and anion (Don’t

forget to include the charge of each ion).• (2) Decide how many cations and anions are needed

so that the sum of their charges balances out to be zero.

• (3) Write the formula of the compound by writing the number of cations followed by the number of anions which you used in step #2. Remember not to include the charges of the ions since now they balance out to be neutral. (*Note when using more than one polyatomic ion the polyatomic ion must be written in parentheses).

Write the formula for barium chloride

Write the formula for iron (II) oxide

Write the formula for calcium phosphate

Write the formula for ammonium carbonate

Write the formulas for the following Write the formulas for the following compoundscompounds

(a) cobalt (II) chloride

(b) lithium sulfate

(c) ammonium dichromate

(d) aluminum oxide

(e) boron (III) phosphide

(f) Chromium (V) nitrate

Naming Ionic Compounds

• When naming ionic compounds the following steps are followed:

Naming Ionic Compounds

• When naming ionic compounds the following steps are followed:

(1) Separate the compound into its positive and negative parts (Note that the positive part of a compound will be only the first element with the exception of ammonium which is NH4

+)

Naming Ionic Compounds

• When naming ionic compounds the following steps are followed:

(1) Separate the compound into its positive and negative parts (Note that the positive part of a compound will be only the first element with the exception of ammonium which is NH4

+)(2) Write the name of the cation followed by

the name of the anion.

Write the name of ZnO and determine the oxidation numbers of the elements within this compound.

Write the name of CuO and determine the oxidation numbers of the elements within this compound.

Write the name of MnCO3 and determine the oxidation numbers of the elements within this compound.

Write the name of Fe2(SO4)3 and determine the oxidation numbers of the elements within this compound.

Name the following compounds and determine the oxidation numbers of each element.

(a) SrCl2

(b) Cr(OH)2

(c) KClO4

(d) NH4MnO4

(e) CuP

Binary Molecular Compounds

• Binary molecular compounds are composed of two different nonmetals – examples: CO, SO2, N2H4, P4Cl10

• These compounds are named by using a prefix to indicate the number of atoms of each element present.

• The prefix mono- is often omitted especially when the first element would have the prefix mono-

CO• (example: CO is named

carbon monoxide, not monocarbon monoxide).

• NF3 is nitrogen trifluoride

• N2O4 is dinitrogen tetraoxide

• P4S10 is tetraphosphorous decasulfide

Name the following compounds: NF3 N2O4 P4S10

• dichlorine heptaoxide is Cl2O7

• carbon hexasulfide is CS6

• octaphosphorous pentaoxide is P8O5

Write formulas for the following compounds: • dichlorine heptaoxide• carbon hexasulfide • octaphosphorous pentaoxide

Acids

• Acids are compounds that give off hydrogen ions, (H+) when dissolved in water. When a compound has hydrogen as its cation the substance is generally an acid – Examples: HCl, H2SO4, H3PO3

• The rules for naming acids are based on the anion portion of the acid formula.

Rules for Naming Acids

• The names of acids are based on the ending of the anion name.– Examples: HCl, H2SO4, H3PO3

• Cl‾ = chloride

• SO42‾ = sulfate

• PO33‾ = phophite

Rules for Naming Acids

Name the acids

• HNO2

• HCN

• H3PO4

Write formulas for the following acids

• chromic acid

• hydroiodic acid

• chlorous acid

Balanced Chemical Equation

• A chemical equation is a written representation of a chemical reaction.

2Na + 2H2O → H2 + 2NaOH• Reactants• Products• Coefficients

– You should be able to balance equations using coefficients.

Symbols Used in EquationsSymbols Used in Equations

2Na(s) + 2H2O(l) → H2(g) + 2NaOH(aq)

• Solid

• Liquid

• Gas

• Aqueous solution

• Solid (cr) or (s)

• Precipitate (↓)

• “Heated”

• Escaping gas () • Catalyst H2SO4

• A word may be written above an arrow to indicate something is necessary for the reaction to occur.

Other Symbols Used in EquationsOther Symbols Used in Equations

electricity

∆

Net Ionic Equations

• Solutions of sodium chloride and silver nitrate are mixed.

• Step 1: Change the word equation into a chemical equation by writing the formulas for the reactants.

NaCl + AgNO3

NaCl + AgNO3• Step 2: Classify each reactant as a particular type of substance.

Classification of Substances• Acids: compounds with formulas that begin with H. (Examples: HCl, H2SO4).• Bases: compounds that end with OH. (Examples: NaOH, Ba(OH)2)• Metal Oxides: binary compounds of a metal and oxygen. (Examples: CaO,

Na2O).• Nonmetal Oxides: binary compounds of a nonmetal and oxygen. (Examples:

SO2, P4O10).• Salts: Ionic compounds other than bases and metal oxides. (Examples: NaCl,

Mg3(PO4)2, NH4NO3).• Other Compounds: All compounds not classified as one of the five types

above. (Examples: CH4, NH3).

• NaCl and AgNO3 are salts

Salt

NaCl

Salts

• A salt is an ionic compound other than a base or oxide.

K2Cr2O7 CuSO4

NaCl + AgNO3

(salt + salt)

• Step 3: Based on your classification of the substances determine the type of reaction.

Types of Net Ionic Equations 1. Double Replacement Reactions: These reactions start with two compounds and produce two different compounds. Such reactions can be

expected when the two reactants are some combination of acid, base, or salt. The products can be predicted by exchanging the positive parts of the two reactants.

If carbonic acid, H2CO3 is produced as a product it should be written as H2O and CO2. If ammonium hydroxide, NH4OH is produced as a product it should be written as NH3 and H2O.

2. Single Replacement Reactions: The reactants are an element and a compound and the products are a different element and compound. A metallic

element will replace the positive part of a compound or a nonmetallic element will replace the negative part of a compound. 3. Synthesis (Combination) Reactions: (a) Two elements combine to form a binary compound. (b) A metal oxide and water combine to form a base.

(c) A nonmetal oxide and water combine to form an acid. (d) a metal oxide and a nonmetal oxide combine to form a salt. In these reactions it is necessary to know the charges of certain ions in order to predict the formulas of your products. You should determine these ion charges by using their charges within the reacting substances. If this is impossible use your prior experience or the periodic table to make a prediction.

4. Decomposition Reactions: These reactions begin with a single compound and decompose into more that one product. In general they are simply

the reverse of the synthesis reactions listed in #3 above. There are however a few other common decomposition reactions that you should learn: (a) Hydrogen peroxide, H2O2 will decompose into water, H2O and oxygen, O2. (b) Potassium chlorate, KClO3 will decompose into potassium chloride, KCl and oxygen O2.

5. Complex Ion Formation (Coordination Compounds): These reactions involve the reaction of two compounds to form a complex ion

(coordination compound). One of the reacting compounds serves as a source of metal ions and the other reacting compound serves as a source of ligands. A complex ion (coordination compound) is a combination of metal ions and ligands. Ligands are normally ammonia (NH3), hydroxide (OH-), or cyanide (CN-). There needs to be a large supply of ligand for a complex ion to form. This is normally indicated within a reaction by words such as concentrated and/or excess. To form a complex ion take the metal ion and add a number of ligands which equal twice the metal ions oxidation number. For example if the metal ion has an oxidation number of 2+ you should add four of the ligands to it. Then simply add up the charges within the complex ion and determine the final charge.

6. Combustion Reactions: Generally involve a hydrocarbon and oxygen, and if so will produce CO2 and H2O. Some form of the word “burn”

usually identifies combustion reactions. The most common hydrocarbons are alkanes and alcohols. Some examples are listed below: Alkane Formula Alcohol Formula

Methane CH4 Methanol CH3OH Ethane C2H6 Ethanol C2H5OH

Propane C3H8 Propanol C3H7OH Butane C4H10 Butanol C4H9OH Pentane C5H12 Pentanol C5H11OH Hexane C6H14 Hexanol C6H13OH

NaCl + AgNO3

(salt + salt)

• Step 3: Based on your classification of the substances determine the type of reaction.

• This is a double replacement reaction.

NaCl + AgNO3 →

• Step 4: Predict the products of the reaction based on the reaction type.

NaCl + AgNO3 → NaNO3 + AgCl

NaCl + AgNO3 → NaNO3 + AgCl

• Step 5: Use solubility rules if necessary.

• Solutions of sodium chloride and silver nitrate are mixed.

Solubility Rules

• Soluble (strong) acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

• Soluble (strong) bases: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

• Soluble salts: All salts of lithium, sodium, potassium, and ammonium cations. All salts of nitrate and acetate anions. All chloride, bromide, and iodide salts except silver, lead and mercury (I). All sulfates except silver, lead, mercury (I), calcium, strontium and barium.

NaCl + AgNO3 → NaNO3 + AgCl

• Step 5: Use solubility rules if necessary.

Na+ + Cl- + Ag+ + NO3- → Na+ + NO3

- + AgCl

Na+ + Cl- + Ag+ + NO3- → Na+ + NO3

- + AgCl

• Step 6: Eliminate all spectator ions.– A spectator ion appears as both a reactant

and a product in a chemical equation.

Na+ + Cl- + Ag+ + NO3- → Na+ + NO3

- + AgCl

• Step 7: Write the final net ionic equation

Ag+ + Cl- → AgCl

Ag+ + Cl- → AgCl

The AgCl is a precipitate• A precipitate is a insoluble solid

formed when solutions are mixed.

• Precipitates are normally formed by reacting two salts or by changing the temperature to affect the solubility of a compound within a solution.

PrecipitateSolutions of lead nitrate and potassium iodide are mixed.

What is the yellow precipitate?

Solid calcium phosphate is added to excess hydrochloric acid.

Ca3(PO4)2 + H+ → H3PO4 + Ca2+

Equal volumes of 0.1M sulfuric acid and 0.1M sodium hydroxide are mixed.

H+ + OH- → H2O

Solid barium carbonate is added to an excess of dilute nitric acid.

BaCO3 + H+ → Ba2+ + H2CO3

Types of Net Ionic Equations

1. Double Replacement Reactions: These reactions start with two compounds and produce two different compounds. Such reactions can be expected when the two reactants are some combination of acid, base, or salt. The products can be predicted by exchanging the positive parts of the two reactants.

• If carbonic acid, H2CO3 is produced as a product it should be written as H2O and CO2. If ammonium hydroxide, NH4OH is produced as a product it should be written as NH3 and H2O.

Solid barium carbonate is added to an excess of dilute nitric acid.

BaCO3 + H+ → Ba2+ + H2CO3

BaCO3 + H+ → Ba2+ + H2O + CO2

Solid ammonium carbonate is added to a saturated solution of barium hydroxide.

(NH4)2CO3 + Ba2+ + OH- → NH4OH + BaCO3

(NH4)2CO3 + Ba2+ + OH- → NH3 + H2O + BaCO3

Single Replacement Reactions:

2. The reactants are an element and a compound and the products are a different element and compound. A metallic element will replace the positive part of a compound or a nonmetallic element will replace the negative part of a compound.

Teacher Example: Calcium metal is added to dilute nitric acid.

Aluminum metal is added to a solution of copper (II) chloride.Al + Cu2+ → Al3+ + Cu

Liquid bromine is added to a solution of potassium iodide.

Br2 + I- → Br- + I2

Solid calcium is added to warm water.

Ca + HOH → H2 + Ca2+ + OH-

Synthesis (Combination) Reactions:

3. (a) Two elements combine to form a binary compound. (b) A metal oxide and water combine to form a base. (c) A nonmetal oxide and water combine to form an acid. (d) a metal oxide and a nonmetal oxide combine to form a salt.

• In these reactions it is necessary to know the charges of certain ions in order to predict the formulas of your products. You should determine these ion charges by using their charges within the reacting substances. If this is impossible use your prior experience or the periodic table to make a prediction.

Teacher Example: Magnesium metal is heated strongly

in nitrogen gas.

Teacher Example: Calcium oxide is added to water.

Teacher Example: Dinitrogen trioxide gas is bubbled

through water.

Calcium metal is heated strongly in nitrogen gas.

Ca + N2 → Ca3N2

Teacher Example:Excess chlorine gas is passed over

hot iron filings.

Cl2 + Fe → FeCl3

A piece of lithium metal is dropped into a container of nitrogen gas.

Li + N2 → Li3N

Solid barium oxide is added to distilled water.

BaO + HOH → Ba2+ + OH-

Lithium oxide powder is added to excess water.

Li2O + H2O → Li+ + OH-

Solid dinitrogen pentoxide is added to water.

N2O5 + H2O → H+ + NO3-

Phosphorus (V) oxide powder is sprinkled over distilled water.

P2O5 + HOH → H3PO4

Metal oxide + Nonmetal oxide

Solid calcium oxide is exposed to a stream of carbon dioxide gas.

CaO + CO2 → CaCO3

Solid calcium oxide is heated in the presence of sulfur trioxide gas.

CaO + SO3 → CaSO4

Decomposition Reactions:

4. These reactions begin with a single compound and decompose into more than one product. In general they are simply the reverse of the synthesis reactions listed in #3 above. There are however a few other common decomposition reactions that you should learn: (a) Hydrogen peroxide, H2O2 will decompose into water, H2O and oxygen, O2.

(b) Potassium chlorate, KClO3 will decompose into potassium chloride, KCl and oxygen O2.

Decomposition of hydrogen peroxide.

H2O2 → H2O + O2

Decomposition of potassium chlorate.

KClO3 → KCl + O2

Here are some decompositionreactions from previous AP Tests• Solid calcium sulfite is heated in a

vacuum.

• A solution of hydrogen peroxide is exposed to an iron catalyst.

• Solid potassium chlorate is heated in the presence of a manganese dioxide catalyst.

Complex Ion Formation (Coordination Compounds):

5. These reactions involve the reaction of two compounds to form a complex ion (coordination compound). One of the reacting compounds serves as a source of metal ions and the other reacting compound serves as a source of ligands. A complex ion (coordination compound) is a combination of metal ions and ligands. Ligands are normally ammonia (NH3), hydroxide (OH-), or cyanide (CN-). There needs to be a large supply of ligand for a complex ion to form. This is normally indicated within a reaction by words such as concentrated and/or excess. To form a complex ion take the metal ion and add a number of ligands which equal twice the metal ions oxidation number. For example if the metal ion has an oxidation number of 2+ you should add four of the ligands to it. Then simply add up the charges within the complex ion and determine the final charge.

Teacher Example: Excess concentrated ammonia is added

to a solution of nickel (II) sulfate.

An excess of ammonia gas is bubbled through a solution saturated with silver chloride.

NH3 + AgCl → Ag(NH3)2+ + Cl-

What happens to the AgCl in this reaction?

Excess concentrated potassium hydroxide solution is added to a precipitate of zinc hydroxide.

OH- + Zn(OH)2 → Zn(OH)42-

Excess sodium cyanide solution is added to a solution of silver nitrate.

CN- + Ag+ → Ag(CN)2 -

Concentrated (15M) ammonia solution is added in excess to a solution of copper (II) nitrate.

NH3 + Cu2+ → Cu(NH3)42+

Excess concentrated sodium hydroxide solution is added to solid aluminum hydroxide.

OH- + Al(OH)3 → Al(OH)63-

Combustion Reactions:6. Generally involve a hydrocarbon and oxygen, and if so will produce CO2 and H2O. Some form of

the word “burn” usually identifies combustion reactions. The most common hydrocarbons are alkanes and alcohols. Some examples are listed below:

Alkane Formula Alcohol Formula Methane CH4 Methanol CH3OH Ethane C2H6 Ethanol C2H5OH

Propane C3H8 Propanol C3H7OH Butane C4H10 Butanol C4H9OH Pentane C5H12 Pentanol C5H11OH Hexane C6H14 Hexanol C6H13OH

Combustion

• Hexane is burned in excess oxygen.

• Propanol is burned completely in air.

Hydrates

CoCl2 ∙ 6H2O CoCl2