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“Chemical Reactions”

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Evıdences for chemıcal Evıdences for chemıcal reactıonsreactıons A gas is released.

An insoluble substance is produced. A permanent color change is observed. Changes in energy.

Heat, light, sound, and electrical. Exothermic reaction: releases heat,light

and sound. Endothermic reaction: absorbs heat and

light. pH change Odor given off

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Evıdence of a chemıcal reactıon Evıdence of a chemıcal reactıon usıng your senses:usıng your senses:

Sight- Change in color, formation of solid, formation of gasses (bubbles), light emission

Hearing- reaction leads to a popping noise, fizzes

Smell- pungent odor, change in smell

Touch- heat absorption and emission

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All chemical reactions… have two parts:

1. Reactants = the substances you start with

2. Products = the substances you end up with

The reactants will turn into the products.

Reactants Products

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Reactants

Products

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In a chemical reaction Atoms aren’t created or destroyed (according

to the Law of Conservation of Mass) A reaction can be described several ways:

#1. In a sentence every item is a word Copper reacts with chlorine to form copper (II)

chloride.

#2. In a word equation some symbols used

Copper + chlorine copper (II) chloride

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Symbols in equations? the arrow (→) separates the reactants

from the products (arrow points to products)

–Read as: “reacts to form” or yields The plus sign = “and” (s) after the formula = solid: Fe(s)

(g) after the formula = gas: CO2(g)

(l) after the formula = liquid: H2O(l)

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Symbols used in equations(aq) after the formula = dissolved in

water, an aqueous solution: NaCl(aq) is a salt water solution

used after a product indicates a gas has been produced: H2↑

used after a product indicates a solid has been produced: PbI2↓

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Symbols used in equations■ double arrow indicates a

reversible reaction (more later)■ shows that

heat is supplied to the reaction■ is used to indicate a

catalyst is supplied (in this case, platinum is the catalyst)

heat ,

Pt

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What is a catalyst? A substance that speeds up a

reaction, without being changed or used up by the reaction.

Enzymes are biological or protein catalysts in your body.

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The Skeleton EquationUses formulas and symbols to

describe a reaction

–but doesn’t indicate how many; this means they are NOT balanced

All chemical equations are a description of the reaction.

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Write a skeleton equation for:1. Solid iron (III) sulfide reacts with

gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

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Now, read these equations:

Fe(s) + O2(g) Fe2O3(s)

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

NO2(g) N2(g) + O2(g)

Pt

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Types of Reactions There are probably millions of reactions. We can’t remember them all, but luckily they

will fall into several categories. We will learn: a) the 5 major types. We will be able to: b) predict the products. For some, we will be able to: c) predict

whether or not they will happen at all.How? We recognize them by their reactants

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#1 - Combination Reactions Combine = put together 2 substances combine to make one

compound (also called “synthesis”) Ca + O2 CaO

SO3 + H2O H2SO4

We can predict the products, especially if the reactants are two elements.

Mg + N2 Mg3N2 (symbols, charges, cross)

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Complete: Ca + Cl2 Fe + O2 (assume iron (II) oxide is the product)

Al + O2 Remember that the first step is to write

the correct formulas – you can still change the subscripts at this point, but not later while balancing!

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#1 – Combination ReactionsAdditional Important Notes:

a) Some nonmetal oxides react with water to produce an acid:

SO2 + H2O →→ H2SO3

b) Some metallic oxides react with water to produce a base:

CaO + H2O →→ Ca(OH)2

(This is what happens to make “acid rain”)

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#2 - Decomposition Reactionsdecompose = fall apartone reactant breaks apart into two

or more elements or compounds.NaCl Na + Cl2

CaCO3 CaO + CO2

Note that energy (heat, sunlight, electricity, etc.) is usually required

electricity

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#2 - Decomposition ReactionsWe can predict the products if it is

a binary compound (which means it is made up of only two elements)

–It breaks apart into the elements:H2OHgO

electricity

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#2 - Decomposition Reactions If the compound has more than

two elements you must be given one of the products

–The other product will be from the missing pieces

NiCO3 CO2 + ___

H2CO3(aq) CO2 + ___

heat

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Combination Decomposition

Reaction of a metal oxide with water produces a metal hydroxide; that is, a strong base. Reaction of a nonmetal oxide with water produces an oxyacid in which the nonmetal is in the same oxidation state as in the oxide you started with.

Both of these are combination reactions, and both can be reversed by heating the products.

Metal hydroxides decompose on heating to give the metal oxide and water.

Oxyacids decompose on heating to give water and the nonmetal oxide in the appropriate oxidation state.

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ExamplesExamples

Na2O + H2O NaOH

MgO + H2O Mg(OH)2

SO2 + H2O H2SO3

Cl2O5 + H2O HClO3

HNO3 ∆ N2O5 + H2O

Fe(OH)3 ∆ Fe2O3 + H2O∆: means heating

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Decomposition ExceptionsDecomposition Exceptions

Carbonates and chlorates are special Carbonates and chlorates are special case decomposition reactions that do case decomposition reactions that do not go to the elements.not go to the elements.

• Carbonates (CO32-) decompose to

carbon dioxide and a metal oxide• Example: CaCO3 CO2 + CaO

• Chlorates (ClO3-) decompose to

oxygen gas and a metal chloride• Example: 2 Al(ClO3)3 2 AlCl3 + 9 O2

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#3 - Single Replacement ReactionsOne element replaces anotherReactants must be an element and a

compound.Products will be a different element

and a different compound.Na + KCl K + NaCl F2 + LiCl LiF + Cl2

(Cations switched)

(Anions switched)

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#3 Single Replacement Reactions Metals will replace other metals (and they

can also replace hydrogen) K + AlN Zn + HCl Think of water as: HOH

–Metals replace the first H, and then combines with the hydroxide (OH).

Na + HOH

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#3 Single Replacement Reactions We can even tell whether or not a single

replacement reaction will happen:–Because some chemicals are more

“active” than others–More active replaces less active

There is a list - called the Activity Series of Metals

Higher on the list replaces those lower.

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The “Activity Series” of Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead HydrogenHydrogen Bismuth Copper Mercury Silver Platinum Gold

1) Metals can replace other metals, provided they are above the metal they are trying to replace (for example, zinc will replace lead)

2) Metals above hydrogen can replace hydrogen in acids.

3) Metals from sodium upward can replace hydrogen in water.

Higher activity

Lower activity

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The “Activity Series” of Halogens

Fluorine Chlorine Bromine Iodine

Halogens can replace other halogens in compounds, provided they are above the halogen they are trying to replace.

2NaCl(s) + F2(g) 2NaF(s) + Cl2(g)

MgCl2(s) + Br2(g) ???No Reaction!

???

Higher Activity

Lower Activity

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#3 Single Replacement Reactions Practice:

Fe + CuSO4

Pb + KCl

Al + HCl

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#4 - Double Replacement Reactions Two things replace each other.

–Reactants must be two ionic compounds, in aqueous solution

NaOH + FeCl3

–The positive ions change place. NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1

= NaOH + FeCl3 Fe(OH)3 + NaCl

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#4 - Double Replacement Reactions Have certain “driving forces”, or reasons

–Will only happen if one of the products:

a) doesn’t dissolve in water and forms a solid (a “precipitate”), or

b) is a gas that bubbles out, or

c) is a molecular compound (which will usually be water).

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Complete:assume all of the following

reactions actually take place:

CaCl2 + NaOH

CuCl2 + K2S

KOH + Fe(NO3)3

(NH4)2SO4 + BaF2

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How to recognize which type?Look at the reactants:

E + E =Combination

C = Decomposition

E + C = Single replacement

C + C = Double replacement

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Practice Examples: H2 + O2

H2O Zn + H2SO4 HgO KBr + Cl2

AgNO3 + NaCl

Mg(OH)2 + H2SO3

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Total ionic equations Once you write the molecular equation

(synthesis, decomposition, etc.), you should check for reactants and products that are soluble or insoluble.

We usually assume the reaction is in water We can use a solubility table to tell us what

compounds dissolve in water. If the compound is soluble (does dissolve in

water), then splits the compound into its component ions

If the compound is insoluble (does NOT dissolve in water), then it remains as a compound

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Solubility tableSolubility table

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Total ionic equationsTotal ionic equations

Molecular Equation:

K2CrO4 + Pb(NO3)2 PbCrO4 + 2 KNO3

Soluble Soluble Insoluble Soluble

Total Ionic Equation:2 K+ + CrO4

-2 + Pb+2 + 2 NO3-

PbCrO4 (s) + 2 K+ + 2 NO3-

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#5 – Combustion Reactions Combustion means “add oxygen” Normally, a compound composed of

only C, H, (and maybe O) is reacted with oxygen – usually called “burning”

If the combustion is complete, the products will be CO2 and H2O.

If the combustion is incomplete, the products will be CO (or possibly just C) and H2O.

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Combustion Reaction Examples:

C4H10 + O2 (assume complete)

C4H10 + O2 (incomplete)

C6H12O6 + O2 (complete)

C8H8 + O2 (incomplete)

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The moleThe mole

mole (mol)- SI Unit for the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

A unit of counting, like the dozen.

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Avogadro’s NumberAvogadro’s Number

6.02 x 1023

1 mole of 12C has a mass of 12 g

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Molar massMolar mass

By definition, these are the mass of 1 mol of a substance (i.e., g/mol)

– The molar mass of an element is the mass number for the element that we find on the periodic table

– The formula weight (in amu’s) will be the same number as the molar mass (in g/mol)

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Using molesUsing moles

Moles provide a bridge from the molecular scale to the real-world scale

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Mole relationshipsMole relationships

One mole of atoms, ions, or molecules contains Avogadro’s number of those particles

One mole of molecules or formula units contains Avogadro’s number times the number of atoms or ions of each element in the compound

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Three basic laws of Three basic laws of matter:matter:

Law of conservation of massLaw of definite proportionsLaw of multiple proportions

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Law of Conservation of Mass- Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

CH4 + 2O2 → 2H2O + CO2

16g + 64g → 36g + 44g

Antoine Lavoisier

stated this about 1785

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Law of Definite ProportionsLaw of Definite Proportions – no matter how much salt you have, it is always 39.34% Na and 60.66% Cl by mass.

Example: Sodium chloride always contains 39.34% Na and 60.66% Cl by mass.

2NaCl → 2Na + Cl2

100g → 39.34g + 60.66g116.88g → ? + ?

Joseph Louis Proust stated this in 1794.

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Law of Multiple ProportionsLaw of Multiple Proportions- Two or more elements can combine to form different compounds in whole-number ratios.

Example

John Dalton proposed this in 1803.

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Balanced Chemical Balanced Chemical EquationsEquations

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Stoichiometry is…Stoichiometry is… Greek for “measuring elements”

Pronounced “stoy kee ahm uh tree” Defined as: calculations of the

quantities in chemical reactions, based on a balanced equation.

There are 4 ways to interpret a balanced chemical equation

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In terms of In terms of ParticlesParticles An Element is made of atoms A Molecular compound (made of

only nonmetals) is made up of molecules (Don’t forget the diatomic elements)

Ionic Compounds (made of a metal and nonmetal parts) are made of formula units

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Example: 2H2 + O2 → 2H2O Two molecules of hydrogen and one molecule of oxygen form two molecules of water.

Another example: 2Al2O3 Al + 3O2

2formula unitsAl2O3 form 4 atoms Al

and3moleculesO2

Now read this: 2Na + 2H2O 2NaOH + H2

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In terms of In terms of MolesMoles The coefficients tell us how

many moles of each substance

2Al2O3 Al + 3O2

2Na + 2H2O 2NaOH + H2

Remember: A balanced equation is a Molar Ratio

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In terms of In terms of MassMass The Law of Conservation of Mass applies We can check mass by using moles.

2H2 + O2 2H2O

2 moles H2

2.02 g H2

1 mole H2

=4.04 g H2

1 mole O2

32.00 g O2

1 mole O2

=32.00 g O2

36.04 g H2 + O2

+

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In terms of Mass (for products)

2H2 + O2 2H2O

2 moles H2O18.02 g H2O1 mole H2O

=36.04 g H2O

36.04 g H2 + O2=36.04 g H2O

The mass of the reactants must equal the mass of the products.

36.04 grams reactant = 36.04 grams product

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In terms of In terms of VolumeVolume

At STP, 1 mol of any gas = 22.4 L

2H2 + O2 2H2O (2 x 22.4 L H2) + (1 x 22.4 L O2) (2 x 22.4 L H2O)

NOTE: mass and atoms are ALWAYS conserved - however, molecules, formula units, moles, and volumes will not necessarily be conserved!

67.2 Liters of reactant ≠ 44.8 Liters of product!

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Mole to Mole conversions

2Al2O3 Al + 3O2

– each time we use 2 moles of Al2O3 we will also make 3 moles of O2

2 moles Al2O3

3 mole O2

or2 moles Al2O3

3 mole O2

These are the two possible conversion factors to use in the solution of the problem.

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Mole to Mole conversions How many moles of O2 are

produced when 3.34 moles of Al2O3 decompose?

2Al2O3 Al + 3O2

3.34 mol Al2O3 2 mol Al2O3

3 mol O2 =5.01 mol O2

If you know the amount of ANY chemical in the reaction,

you can find the amount of ALL the other chemicals!

Conversion factor from balanced equation

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Practice: 2C2H2 + 5 O2 4CO2 + 2 H2O

• If 3.84 moles of C2H2 are burned, how many moles of O2 are needed? (9.6 mol)

•How many moles of C2H2 are needed to produce 8.95 mole of H2O? (8.95 mol)

•If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed? (4.94 mol)

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Steps to Calculate Stoichiometric Problems

1. Correctly balance the equation.

2. Convert the given amount into moles.

3. Set up mole ratios.

4. Use mole ratios to calculate moles of desired chemical.

5. Convert moles back into final unit.

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Mass-Mass Problem:

6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?

4Al + 3O2 2Al2O3

=6.50 g Al

? g Al2O3

1 mol Al

26.98 g Al 4 mol Al

2 mol Al2O3

1 mol Al2O3

101.96 g Al2O3

(6.50 x 1 x 2 x 101.96) ÷ (26.98 x 4 x 1) = 12.3 g Al2O3

are formed

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Another example:Another example: If 10.1 g of Fe are added to a

solution of Copper (II) Sulfate, how many grams of solid copper would form?

2Fe + 3CuSO4 Fe2(SO4)3 + 3Cu

Answer = 17.2 g CuAnswer = 17.2 g Cu

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“Limiting” Reagent If you are given one dozen loaves of

bread, a gallon of mustard, and three pieces of salami, how many salami sandwiches can you make?

The limiting reagent is the reactant you run out of first.

The excess reagent is the one you have left over.

The limiting reagent determines how much product you can make

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Limiting Reagents - Combustion

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How do you find out which is limited?How do you find out which is limited?

The chemical that makes the least amount of product is the “limiting reagent”.

You can recognize limiting reagent problems because they will give you 2 amounts of chemical

Do two stoichiometry problems, one for each reagent you are given.

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If 10.6 g of copper reacts with 3.83 g sulfur, how many grams of the product (copper (I) sulfide) will be formed? 2Cu + S Cu2S

10.6 g Cu 63.55g Cu 1 mol Cu

2 mol Cu 1 mol Cu2S 1 mol Cu2S

159.16 g Cu2S

= 13.3 g Cu2S

3.83 g S 32.06g S 1 mol S

1 mol S 1 mol Cu2S

1 mol Cu2S

159.16 g Cu2S

= 19.0 g Cu2S

= 13.3 g Cu2S

Cu is the Limiting

Reagent, since it

produced less product.

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Another example:Another example: If 10.3 g of aluminum are reacted with

51.7 g of CuSO4 how much copper (grams) will be produced?

2Al + 3CuSO4 → 3Cu + Al2(SO4)3

the CuSO4 is limited, so Cu = 20.6 g How much excess reagent will remain?

Excess = 4.47 gramsExcess = 4.47 grams

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What is Yield?What is Yield? Yield is the amount of product made in a

chemical reaction. There are three types:

1. Actual yield- what you actually get in the lab when the chemicals are mixed

2. Theoretical yield- what the balanced equation tells should be made

3. 3. Percent yieldPercent yield = Actual Theoretical

x 100

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Example: 6.78 g of copper is produced when

3.92 g of Al are reacted with excess copper (II) sulfate.

2Al + 3 CuSO4 Al2(SO4)3 + 3Cu

What is the actual yield?

What is the theoretical yield?

What is the percent yield?

= 6.78 g Cu= 6.78 g Cu

= 13.8 g Cu= 13.8 g Cu

= 49.1 %= 49.1 %

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Details on YieldDetails on Yield Percent yield tells us how “efficient” a

reaction is. Percent yield can not be bigger than

100 %. Theoretical yield will always be larger

than actual yield!– Why? Due to impure reactants; competing

side reactions; loss of product in filtering or transferring between containers; measuring