Chapter 10 Chemical ReactionsChemical changes and reactions between compounds

Chemical formula revisited:

types of elements and ratio making up compound

empirical formula : simplest whole number ratio of

elements in a compound

-general description of how to assemble

the compound

ionic-exists as separate ions

-empirical formula is the right description

-simple mixture of ions Na+ and Cl- ions

molecular (covalent)-exists as a particular

bonded entity

-empirical formula tells what it’s made of

-does not describe how it is bonded

molecular formula : identifies actual numbers of atoms

in a molecule-discrete units of atoms

-not appropriate for ionic – array of separate ions

-molecules described by molecular formula

-you know a formula is molecular if the numbers

are not the simplest whole number ratio

glucose: C6H12O6 molecular since divisible by 6

empirical – CH2O

Making compounds – cannot count atoms

need to know about masses of compounds

Formula weight (mass) –mass of one formula unit (amu)

sum of atomic weights for ALL atoms in chemical formula

general term for ionic and molecular

Molecular weight (mass) - mass of molecule

add up all atomic weights for each atom in molecule

specific to molecules only (covalent)

use atomic mass since it is an average

over all isotopes – naturally abundant species

Formula weight of NaCl:

Molecular weight of C12H22O12

Also want to be able to describe percent amount of substances in compound

percent composition: what percent by mass is from each element

% composition = mass of the element in compoundtotal formula weight X100%

Find % comp of each element in NaCl:

Ex: find % comp of each element in C8H18

empirical easier C4H9

Try: H3BO3

Ba(NO3)2

CAN USE EMPIRICAL OR MOLECULAR

Identification of substances

Chemical Equations

Describes what it takes for a chemical reaction to occur (A recipe for producing compounds)

Which involves changes in compounds (one to another), states of matter (liquid, solid, gas), and exchanges in energy.

GENERAL CHEMICAL EQUATIONReactants Products

To quantitatively show how to make reaction succeed, you have to describe the numbers of atoms involved

LAW OF CONSERVATION OF MASS

Matter is neither created nor destroyed

The products must balance the reactantsC + O2 CO2

C = 1 C = 1O = 2 O = 2

NEW REACTION: synthesis of carbon monoxideC + O2 CO (unbalanced)

Unequal – mass of Reactants different than for products2C + O2 2CO

Now mass of Reactants is equal with mass of Products

Balance equations by changing coefficientsNOTE: The formula CAN’T be changed

Be careful to take inventory of ATOMS in equation!Ex: CH4 + 2O2 CO2 + 2H2O

C=1 C=1O=4 O=4H=4 H=4

Many times, Balancing equations is a trial & error processEx: Combustion of Gasoline (Octane)2C8H18(g) + 25O2(g) 16CO2(g) + 18H2O(g)

However, you should be familiar with the rules whichdescribe balanced chemical reactions.

1. Number of Atoms of each element conserved in reactants and products2. Cannot change formula of reactants or products 3. Can only change coefficients to balance equation

Hints to help in balancing equations1. Balance compounds with biggest numbers of atoms first2. Treat polyatomics as single units, especially when on both sides of equation3. Fractional coefficients are useful

Can also identify physical states of compounds in reaction(g) – compound is gaseous(l) – compound is liquid(s) – compound is solid(aq) – compound is an aqueous solution (in water)

() - temperature change (energy) (or ) - solid precipitates out (or gas escapes)

Precipitation of calcium bicarbonate from hard water - changes in state

Ca(HCO3)2 (aq) + Na2CO3 (aq)

2NaHCO3 (aq)+ CaCO3

HYDROCARBONS-combinations of hydrogen and carbon-important energy source - combustion

gasoline, acetylene, propane

-react with O2 to form H2O and CO2 plus ENERGY

Example:butane 2C4H10 + 13O2 8CO2 + 10H2O

(lighters) + energyCarbohydrates: hydrocarbon with oxygen plants store sun energy as carbs (photosythesis)

we eat and respire oxygen energy + CO2

GLUCOSE : C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(g)

Enzymes - slowly release energy in plants and animals free burning produces quick energy release and flame

precipitate

Types of Chemical Reactions

Oxidation-Reduction (REDOX) Reaction - GENERAL

-electrons transferred from one atom to another-one element oxidized (loses e-) and the other is reduced (gains e-) <balance of charge-losses cancel increases>

-oxygen often involved, but F, Cl & other nonmetals do the same thing

OXIDIZING AGENT- takes electrons from other substances- takes e- away causing substances to be oxidized

- O2 in food and fuels gains e- to form octet

- Cl & F in bleach kill bacteria - pools

REDUCING AGENT- provides e- to substance being reduced- carbon: gives e- to form octet- REDUCTION OF IRON ORE [iron(III) oxide]

2Fe2O3(s)+ 3C(s) 4Fe(s) + 3CO2

BUT reactions are normally defined in terms of the effect on reactants and products: reaction type 3 types of redox eqs.: combine, decompose, replace

Ion exchange - not redox since no change in chargeions just replaced

Note reverse crossoverTo get charge for name

Reaction Types

Combination (Synthesis) Reactiontwo or more substances combine to form one

X + Y XY lower energyoxidation of metals- rust, burning

burning of Mg metal

2Mg(s) + O2(g) 2MgO(s)

rusting of iron

4Fe(s) + 3O2(g) 2Fe2O3(s)

burning non-metals

C(s) + O2(g) CO2(g)

Decomposition Reactiona compound is broken down into simpler compounds or constituent elements

XY X + Y requires energyHYDROLYSIS - decompose water

with electricity - hydrogenfuel cells

2H2O(l) 2H2(g) + O2(g)

decomposition of mercury(II) oxide

2HgO(s) 2Hg(s) + O2(g)

electricity

Replacement Reaction

One atom or polyatomic ion in a compound is replacedXY + Z XZ + Y or XY + A AY + X (-) part replaced (+) part replaced

chemical activity - tendency of an element to give up electrons

more chemically active if unable to keep electrons MORE ACTIVE METALS GIVE UP ELECTRONS TO LESS ACTIVE

ACTIVITY SERIES - Fig. 10.12shows most active elements on top a metal will replace any metal it is above

has a larger chemical activity

2Al(s) +3CuCl2(aq) 2AlCl3(aq) + 3Cu(s)

Al more chem active - above Cu

- gives e- to Cu+2

- Cu precipitates, Al goes into solution if less chem active, e- would

stay where they are (no reaction)

Alkali and Alkaline metals- very activeoften replace other elements

2Na(s) + H2O(l) 2NaOH + H2

H takes e- from Na

makes Na+ and H2

Ion Exchange Reaction

Each ion of one compound is replaced by the ions of a second compound

Ions mix together to form - a precipitate (insoluble) - a gas - or water

removes ions from the solution of ions

AX + BY AY + BXone of products must leave solution

(precipitate, gas or water)Example:

3Ca(OH)2(aq) + Al2(SO4)3(aq) 3CaSO4(aq) +2Al(OH)3Application:Dissolved ions form insoluble aluminum hydroxide “net” for water treatment- mixture traps suspended impurity

particles Solubility Tables (Appendix B) if the products are soluble --- no reaction if one product is insoluble ---reaction occurs

precipitate or insoluble gasExamples: identify the reaction type

2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(s)

2Ag(s) + S(g) Ag2S(s)

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

2KClO3(s) 2KCl(s) + 3O2

Summary of Chemical Equations

Tells how compounds combine together to form new substances

what atoms are present before and after(inventory)

2H2(g) + O2 2H2O

atomic description:4 atoms H and 2 atoms O form 4 atoms H and 2 atoms O

MOLECULAR description: 2 molecules H2 and 1 molecule O2 forms 2 molecules of H2O

Also describes masses (formula weight):reactant mass = 2x2.016amu + 32 amu = 36 amu

H2 O2

product mass = 2x18amu =36 amu

H2O

LAW OF CONSERVATION OF MASSNOTE: 2g of H combines with 8 g of O

always the same mass ratioremember amu based on weight of C-12

exactly 12 amu!

relationship between mass and number historically misunderstood by Dalton

mis-measured mass of H and O to form H20

he thought water was HOLed to failure of Daltons atomic theory

Gay-Lusaac and Avogadrogases combine in whole number ratio

(at constant temperature and pressure-STP)

2H2(g) + O2 2H2O

2 volumes of H2 for each volume O2

gives one volume of water vapor

LAW OF COMBINING VOLUMESequal volumes of gas at STP have equal numbers of molecules

FIXED formula for water:

(twice as much H2) H2O

SHOWED H and O diatomic (H2&O2)

since two volumes of water produced

New interpretation of chemical equation:coefficients in chemical equation give the volumes of gas

BUT we need to know how to get numbers of molecules, and atoms by working with mass

mass is easier to determine (measure)for high yield reactions (chemical co.

pharmacy, etc)

How do masses and numbers of atoms relate?

Each element has a averaged atomic weight based on the mixtures of isotopes (natural abundance)

each weight a comparison to carbon-12:

Study weight mass relationship for C-12the number of atoms in 12.0 g of

C-12 is 6.02x1023 experimentallycounted

Avogadro’s number - 6.02x1023

SI unit for counting numbers

Defines the mole - an amount of a substance that contains Avogadro’s number

6.02x1023 H2O molecules

- 1 mole of water- 1 mole of O atoms- 2 moles of H atoms- 3 moles total of atoms

All element weights based on a comparison to carbon-12!remember: mass of one mole C-12=12 g numerically equal to the atomic weight

Weight of one mole of any element is equivalent to its atomic weight (in grams)

Define weights based on large amounts of atoms

gram-atomic weight - the mass (in grams) of one mole of an element -numerically equal to atomic wt.

(in grams)

-represents 6.02x1023 atoms - atomic weight since a mixture

of isotopes - one gram-atomic weight of any element contains the same number

of atoms 6.02x1023

gram-formula weight - the mass (in grams) of one mole of the compound

-represents 6.02x1023 formula units -numerically equal to formula wt.

(in grams) - applies to ionic or covalent

gram-molecular weight - the mass (in grams) of one mole of the molecules (covalent only)

-gram-formula weight of a molecular compound

-represents 6.02x1023 molecules -numerically equal to formula wt.

(in grams)Can now go back and forth between number and mass

like conversion factors

Conversions between mass and number

Gram weights can be thought of as conversion factors: grams per mole

Example: NaCl- formula weight = 58.44 amu gram-formula wt.= 58.44 grams

58.44 g for every moleconversion identity 58.44 g = 1 mole NaCl

24.7 g of NaCl is how many moles?

24.7 g =0.423 moles

Example: 0.773 moles of C2H4 has what mass?

gram formula wt.=2x12.01g + 4x1.008g

= 28 g (per mole C2H4)

Convert:

0.773 moles = 21.64 g of C2H4

How many molecules?

0.773 moles

= 4.65x1023 C2H4 molecules

How many C atoms: 2 C atoms per mole

0.773 moles =1.546 moles

1 mole58.44 g

28 g1 mole

6.02x1023 molecules 1 mole

2 C atoms

1 mole C2H4

Suggested problems for Chapter 10:Parallel Exercises - # 1, 2, 3, 5Additional problems:

1. Convert the following to numbers of moles

a) 6.7 g of iron

6.7g = 0.12 moles Fe

b) 45.3 g of Fe2(SO4)3

45.3 g =0.113 moles Fe2(SO4)3

c) 8.4x1024 molecules of C6H12O6

8.4x1024 molecules = 13.95 moles

d) 9.4x1023 Na+ ions of C6H12O6

2. Convert the following into masses• 0.55 moles of silver =59.34 g

• 4.7x1022 molecules of CH4 =1.26 g

• 2 moles of AgCl - Answer: 286.8 g

• 5.6x1024 Cl- ions - Answer: 330.23 g

3. In two moles of BaF2:

a) how many moles of F are there? Answer: four

b) how many Ba+2 ions are there? Answer: two

c) what is the mass of F? Answer: 76 g

d) what is the mass percent of F? (Ch. 9) Answer: 21.7%

1 mole Fe55.8 g

1 mole Fe2(SO4)3

399.9 g

1 mole C6H12O6

6.02x1023 molecules

107.9 g1 mole silver

16.16 g in a mole6.02x1023 molecules