Unit 5: Chemical Reactions

When chemical reactions occur:

1. Chemical bonds are broken and new bonds form.

2.Energy is produced (exothermic) or absorbed (endothermic).

3.New compounds are formed, or compounds decompose to their elements.

4.The law of conservation of mass is obeyed.

Reactions are represented by chemical equations:

A+B → C+DREACTANTS PRODUCTS

Chemical Equations are the “written” way to represent a

chemical reaction.

Chemical EquationSymbols

, yields

Reading Equations

NaCl (aq) + AgNO3 (aq) NaNO3 (aq) + AgCl (s)

This reads: A water solution of sodium chloride reacts with a water solution of silver nitrate to yield a water solution of sodium nitrate and a precipitate (solid) of silver chloride.

Chemical Reactions

Since matter can not be created or destroyed, chemical reactions must be balanced in terms of mass.

The amount of mass you start with must be equal to the mass of the products.

Reactants → Products

100g total = 100g total

Balancing Steps1. If you are given a word equation,

write the chemical equation with the correct formulas and symbols-don’t forget diatomics

2. Add coefficients to the formulas to make the number of atoms of each element on both sides of the equation the same.

3. You may not add coefficients to the middle of a formula.

Balancing Steps4. You may not change the subscript of

a correctly written formula.

1.

2.

3.

4.

Writing Equations from words:Start with a skeleton equation:

-Write the formulas of the reactants on the left side of the yield sign and products on the right side

-Next add information like physical states, catalysts involved, etc.

-Then you need to balance the equation.

Fe+O2→Fe2O3

Fe(s)+O2(g)→Fe2O3(s)

Types of Reactions

Remember that during a chemical reaction a chemical change occurs…

Indicators of a chemical change include:

1. Energy Change – heat or light is produced, or a decrease in temperature.

– Exothermic – gives off heat, feels hot

– Endothermic – takes in heat, feels cool.

2. Production of a gas (you see bubbles or fizzing)

3. Precipitate – a solid is formed when two liquids are mixed together.

The indicator that a precipitate has formed is that the liquid turns cloudy.

4. Color change (unexpected change)

Combination Reaction (synthesis)• A chemical change where 2 or more substances react to

form a new compound

• only one product

• opposite of a decomposition reaction

A + B → AB

Synthesis

Example: 2H2 + O2 → 2 H2O

Predict the products and write the balanced chemical equation.

1. Potassium reacts with oxygen

2. Lithium oxide reacts with water

3. Carbon dioxide reacts with water

4. Carbon burns

5. Sodium reacts with bromine

Decomposition Reaction

• A chemical change where a single compound breaks down into 2 or more simpler products

• only one reactant

AB → A + B

Decomposition

Example: 2 H2O → 2 H2 + O2

1. Barium hydroxide decomposes when heated

2. Sodium carbonate decomposes when heated

3. Lithium chlorate decomposes when heated

4. Aluminum oxide decomposes during electrolysis

5. Carbonic acid decomposes

6. Calcium chlorate decomposes when heated

Single-Replacement Reaction

• A chemical change in which one element replaces another in a compound

– metal replaces metal (+)

– nonmetal replaces nonmetal (-)

A + BC → B + AC

Single Replacement

Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)

Activity SeriesThe elements at the top are the most reactive.

In a single replacement reaction, a free element will replace anything below itself on the activity series.

For halogens (group 17) the most reactive is fluorine and the least reactive is iodine.

Find this Activity Series on your formula chart.

Other helpful info: Group 1 and 2 metals most active, transition metals less active, and

jewelry metals and H least active

Use the activity series to determine if the following single replacement reactions occur. If they do, write the equation and balance. If the reaction does not occur,

write the reactants, the yield sign, and No Reaction.

1. iron and silver sulfate

2. aluminum + hydrochloric acid

3. potassium + water

4. fluorine + potassium bromide

AB + CD → AD + CB

Double-Replacement Reaction (precipitation rxns)

• elements in two compounds exchange places to make two new compounds.

• generally take place in aqueous solutions

Double-Replacement

These reactions occur between ions in aqueous solutions and

produce at least one of the following: a precipitate, a

gas or water.

Formation of a precipitate- if a product is not soluble, it is a

precipitate. Use the solubility chart/rules.

NaCl(aq) + AgNO3(aq) → NaNO3(s) + 2KNO3(aq)

Solubility rules for double-replacement reactions

Locate the Solubility rules on your formula chart.

Double-Replacement

Formation of a gas-

HCl (aq) + FeS (s) → FeCl2 (aq) + H2S (g)

Formation of water-

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

Formation of product which decomposes-

CaCO3 (s) + HCl (aq) → CaCl2 (aq) + CO2(g) + H2O (l)

*extra rules in notes*

1. potassium iodide + lead (II) nitrate

2. sodium iodide + acetic acid

3. barium nitrate + calcium carbonate

4. sodium carbonate + hydrochloric acid

Combustion Reaction

C4H6 + O2 → CO2 + H2O

When hydrocarbons burn in excess oxygen, the products are always carbon dioxide and water.

If there is too little oxygen, CO is produces which is highly toxic.

1. CH4 + O2

2. CH3OH + O2

3. C3H4 + O2

Identify the type of reaction, predict the products, and balance.

1. __ Na + __ ZnO →

2. __ MgCO3 →

3. __ C4H8 + __ O2 →

4. __ Li3PO4 + __ Ca(NO3)2 →

5. __ Ag + __S →

Redox Reactions and Oxidation #’s

Redox stands for reduction-oxidation reactions.

Electrons move from one atom to another or from one ion to another.

REMEMBER this:

O oxidationI isL losing

R reductionI isG gaining

Losing electrons become more positive

Gaining electrons become more negative

Example:

Which element is oxidized? ___Na__

Which element is reduced? ___Cl__

Assign Oxidation Numbers

0 0 +1 -1

2 Na + Cl2 → 2 NaCl

2 Na + Cl2 → 2 NaCl

1. Complete, balance and identify element oxidized and element reduced.

Al + O2 →

2. Balance and identify element oxidized and element reduced.

Fe + S →

Examples of REDOX Reactions

-Iron rusting.

-Hydrogen peroxide sanitizing wounds.

-Photography development.

-Chlorine bleach whitening laundry.

-Silver tarnishing.

Rusting

Requires oxygen and water to be present.

Rusting is a slow process since water droplets have few ions making them poor electrolytes

CorrosionThe loss of metal resulting from an oxidation-reduction reaction of the metal with substances in the environment.

Corrosion occurs faster with water with abundant ions - ie. seawater or regions where roads are salted.

To prevent corrosion iron can be coated with another metal that is more resistant - like zinc - this is called galvanizing.

Net Ionic Equations

Reactions that are simplified to only show what particles are changing.

What type of reaction is happening…

Pb(NO3)2 (aq) + NaI (aq) →

Pb(NO3)2 (aq) + NaI (aq) → PbI2 (s) + NaNO3 (aq)

Are there any ions that aren’t undergoing a chemical change?

A net ionic equations does not show the ions that don’t change (ions that stay aqueous)

Strong acids completely ionize: HCl, HBr, HI, HNO3, HClO4, H2SO4

Steps for Writing Net Ionic Equations

1. Write the balanced equation with all states labeled. (THIS IS THE CHEMICAL EQUATION –sometimes called MOLECULAR EQUATION)

2. Split any aqueous ionic or strong acids into ions. (THIS IS THE TOTAL IONIC EQUATION)

3. Cancel out any ions that appear on each side of the arrow (called spectator ions). (THIS IS THE NET IONIC EQUATION)

AgNO3 (aq) + NaCl (aq) → NaNO3 (aq) + AgCl (s)

This is the chemical equation – it is what we’ve been doing for weeks.

This is the total ionic equation- it shows ions for aqueous ionic or strong acids.

1. A strip of magnesium is added to a solution of silver nitrate.

2. A solution of hydrogen peroxide is heated.

3. Solutions of sodium hydroxide and hydrochloric acid are mixed.

4. Chlorine gas is bubbled into a solution of sodium bromide.