chapter3chemicalformulaeandequations-130319185107-phpapp01

8/12/2019 chapter3chemicalformulaeandequations-130319185107-phpapp01

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Chemical Formulae

And

Equations


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A. Relative Atomic Mass and Relative

Molecular Mass Based on the theory of particles:

particles are very small and discrete. A single atom

is too small and light and cannot be weigheddirectly

Thus, the mass of an atom is obtained bycomparing it with another atom which is taken as a

standard.


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3 types of scale to determine the mass of theparticles

a) Compared with a hydrogen-1 scaleb) Compared with an oxygen-16

c) Compared with carbon-12 (modern comparisonUNTILL TODAY)


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Relative atomic mass, RAM Meaning;

The average mass of one atom of the element when

compared with 1/12 of the mass of an atom ofcarbon-12.

Relative Atomic Mass, RAM

= Average mass of one atom of the element

1/12 x the mass of an atom of carbon-12


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Example:

RAM of magnesium

= 24 = 241/12 x 12

= magnesium is 24 times larger than carbon-12

** THE VALUE OF NUCLEON NUMBER IN THEPERIODIC TABLE OF ELEMENT

= RELATIVE ATOMIC MASS, RAM


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Relative molecular mass, RMM Meaning;

The average mass of one molecule when comparedwith 1/12 of the mass of an atom of carbon-12.

Relative Molecular Mass, RMM

= Average mass of one molecule

1/12 x the mass of an atom of carbon-12

Calculate RMM/RFM by adding up the relativeatomic mass of all the atoms that present in the

molecule/ionic compound


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B. The Mole and the Number of

Particles The number of particles in matter is measured inmole.

Definition:The amount of substance that contains as manyparticles as the number of atoms in exactly 12 g ofcarbon-12

Symbol of mole: mol


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How many atoms are there in 12 g of carbon-12?

= 6.02 1023

The value of 6.02 1023 is called theAvogadroconstant or Avogadro number

Avogadro constant, NAThe number of particles in one mole of a substance


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Point to note:

One mole of any substance contains 6.02 1023 particles

1 mol of atomic substance contains 6.02 1023

atoms 1 mol of molecular substance contains 6.02 1023

molecules

1 mol of ionic substance contains 6.02 1023 ions


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Relationship between the number of moles and thenumber of particles

Number of moles,

(mol)

Number of particles

(atoms, molecules, ions)

NA

NA


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Number ofparticles

Moles NA


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Example 1:A closed glass bottle contains 0.5 mol of oxygen gas, O2(a) How many oxygen molecules, O2 are there in the

bottle?(b) How many oxygen atoms are there in the bottle?

[Avogadro constant: 6.02 1023 mol-1]


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a) The number of oxygen molecules, O2= 0.5 mol 6.02 1023 mol-1

= 3.01 1023

molecules

b) The number of oxygen atoms

= 0.5 mol 6.02 1023 mol-1 2

= 6.02 1023 atoms

Number ofparticles

Moles NA


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Example 2:Find the number of moles of 9.03 1023 molecules in asample containing molecules of carbon dioxide, CO2

[Avogadro constant: 6.02 10

23

mol

-1

]

The number of moles carbon dioxide

= 9.03 1023

6.02 1023 mol-1

= 1.5 mol

Number ofparticles

Moles NA


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C. The Mole and the Mass of

Substances The mass of one mole of any substance is called molar

mass

Units: g mol

-1

The molar mass of substances are numerically equalto relative mass


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ElementRelative

massMass of

1 molMolarmass

Helium 4 4 4 g mol-1

Sodium 23 23 23 g mol-1

Water, H2O 2(1) + 16 = 18 18 18 g mol-1

Ammomia, NH3 14 + 3(1) = 17 17 17 g mol-1


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Relationship between the number of moles and themass of a substance

Number of moles,

(mol)

Mass

(g)

molar mass

molar mass


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Mass

(g)

MolesRAM /RMM /

RFM


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Example 1:What is the mass of

(a) 0.1 mol of magnesium?

(b) 2.408 1023

atoms of magnesium?[Relative atomic mass: Mg=24; Avogadro constant: 6.02 1023 mol-1]


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(a) Molar mass of Mg = 24 g mol-1

Mass of Mg = 0.1 mol 24 g mol-1

= 2.4 g

Mass (g)

Moles

RAM /RMM /

RFM

(b) The number of moles Mg atoms

= 2.408 1023

6.02 1023

mol-1

= 0.4 mol

Mass of Mg atoms

= 0.4 mol 24 g mol-1

= 9.6 g

Number ofparticles

Moles NA


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How many moles of molecules are there in 16 g of sulphurdioxide gas, SO2?

[Relative atomic mass: O=16, S=32]

Example 2:

RMM of SO2= 32 + 2(16) = 64

Molar mass of SO2 = 64 g mol-1

The number of moles= 16 g

64 g mol-1

= 0.25 mol

Mass (g)

Moles

RAM /RMM /

RFM


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D. The Mole and the Volume of Gas The volume occupied by one mole of the gas is called

molar volume

One mole of any gas always has the same volumeunder the same temperature and pressure

The molar volume of any gas is22.4 dm3 at STP or

24 dm3 at room condition


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Relationship between the number of moles and thevolume of gas

Number of moles,

(mol)

Volume of gas

(dm3

)

molar volume

molar volume


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Volume

(dm3)

Moles

22.4 dm3

(STP) /

24 dm3

(roomcondition)


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Example 1:What is the volume of 1.2 mol of ammonia gas, NH3 atSTP?

[Molar volume: 22.4 dm3

mol-1

at STP] Volume(dm3)

Moles

22.4 dm3

(STP) /24 dm3 (RC)

The volume of ammonia gas, NH3

= 1.2 mol 22.4 dm3

mol-1

= 26.88 dm3


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How many moles of ammonia gas, NH3 are present in600 cm3 of the gas measured at room conditions?

[Molar volume: 24 dm3

mol-1

at room condition]

Example 2:

Volume(dm3)

Moles

22.4 dm3

(STP) /24 dm3 (RC)

The number of moles of ammonia gas, NH3

= 600 cm3

1000= 0.6 dm3

= 0.6 dm3

24 dm3 mol-1

= 0.025 mol


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Relationship between the number of moles, number of

particles, mass and the volume of gas

Number of moles,(mol)

Mass(g)

molar volume molar volume

Number ofparticles

Volume of gas(dm3)

NA

NA molar mass

molar mass


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E. Chemical FormulaeA chemical formulae

A representation of a chemical substance using lettersfor atom and subscript numbers to show the numbersof each type of atoms that are present in the substance


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H2

Symbol ofhydrogen atom

Shows that thereare two hydrogen

atom in a hydrogengas, H2molecule


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H2O

Symbol ofhydrogen atom

Shows that thereare two hydrogenatom in a water

molecule

Symbol ofoxygen atom

Shows that thereare one oxygenatom in a water

molecule


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Compound can be represented by two types:

1. Empirical formula

2. Molecular formula


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Empirical formula Meaning

Formula that show the simplest whole number ratio ofatoms of each element in the compound


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ExampleA sample of aluminium oxide contains 1.08 g ofaluminium and 0.96 g of oxygen. What is the empiricalformula of this compound?

[Relative atomic mass: O = 16; Al = 27]


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Element Al O

Mass of element

(g)Number of mole

(mol)

Ratio of molesSimplest ratio


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Experiment question:Describe how you can carry out an experiment todetermine the empirical formula of magnesium oxide.Your description should include

Procedure of experiment

Tabulation of result

Calculation of the results obtained

[Relative atomic mass: O = 16; Mg = 24]


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Procedure:1. Clean (5-15 cm) magnesium ribbon with

sandpaper and coil it

2. Weigh an empty crucible with its lid3. Place the magnesium in the crucible and weigh

again

4. Record the reading

5. Heat the crucible strongly

6. Open and close the lid very quickly


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7. When burning is complete, stop the heating

8. Let the crucible cool and then weigh it again

9. The heating, cooling and weighing process isrepeated until a constant mass is recorded


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Result:Description Mass (g)

Crucible + lid x

Crucible + lid + Mg y

Crucible +lid + MgO z


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Calculation:Mg O

Mass (g) y-x z-yNo. of mole (y-x)/24 (z-y)/16

Ratio 1 1

Empirical formula = MgO


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Discussion1. H2 gas must be flowed through the apparatus toremove all the air

2. H2

gas must be flowed throughout the experiment toprevent the air from outside mixing with the H2 gas

3. H2 gas flowed through the apparatus during cooling toprevent copper being oxidised by air into copper(II)

oxide4. Repeat heating, cooling & weighing process to ensureall the copper(II) oxide changed into copper


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5. This method is to determine empirical formula ofoxide of metals which are less reactive than H2 in thereactivity series

6. Other example: Lead(II) oxide, Iron(II) oxide

7. Anhydrous calcium chloride to dry the H2 gas


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Molecular formula Meaning

Formula that show the actual number of atoms of eachelement that are present in a molecule of thecompound

Molecular formula = (Empirical formula)n


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Example:(CH3)n = 30n [12 + 3(1) ] = 30

15n = 30n = 30/15

= 2

Molecular formula = (CH3)2= C2H6


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Ionic formulaeIonic

compounds

Positive ions

(cation)

Negative ions

(anion)


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Formulae of cations & anionsCation Formula Anion Formula

Sodium ion Na+ Chloride ion Cl-

Potassium ion K+ Bromide ion Br-

Zinc ion Zn2+ Iodide ion I-

Magnesium ion Mg2+ Oxide ion O2-

Calcium ion Ca2+ Hydroxide ion OH-

Aluminium ion Al3+ Sulphate ion SO42-

Iron(II) ion Fe2+ Carbonate ion CO32-

Iron(III) ion Fe3+ Nitrate ion NO3-

Copper(II) ion Cu2+ Phosphate ion PO43-

Ammonium ion NH4+

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