Relative atomic mass and the mole

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Relative atomic mass and the mole

• Atom are extremely small that they cannot be
  measured
• For example, the mass of one atom of carbon is
  2x10-23 g
• This is not of great practical value in
  experimental work
• Therefore scientists have came up with the idea
  of relative mass to establish practical masses of
  atoms.

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Relative atomic mass

• E. g. of relative mass
• Lets say I have 2 masses of respectively 50g and
  100 g.
• If I decide to use the first mass as my unit of
  measurement, the second mass is 2 units, relative
  to the first one.
• This is exactly what happens in chemistry
• Every atoms mass is measured relative to the
  mass of an atom of carbon, which is taken as 12
  units.

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Relative atomic mass

• Most elements are a mixture of isotopes.
• E.g. Table 4.2, p55
• The relative isotopic mass is the mass of an atom
  of the isotope relative to the mass of an atom of
  12C taken as 12 units.
• Relative isotopic masses can be obtained using an
  instrument called mass spectometer

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• Knowing the relative isotopic mass we can find
  the relative atomic mass as follows
• The relative atomic mass The sum of relative
  isotopic mass x abundance/ 100
• E.g1. Imagine taking 100 atoms from a sample of
  chlorine. There will be 75.80 atoms of 35Cl and
  24.20 atoms of 37Cl. What is the total abundance
  mass of 100 atoms of chlorine?

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• E.g2. Boron has two isotopes. Their relative
  isotopic masses and abundances are as follows
  10B relative isotopic mass 10.013, relative
  abundance 19.91. 11B relative isotopic mass
  11.009, relative abundance 80.09. Calculate the
  relative atomic mass of boron.

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• E.g3. From fig 4.5, p58, calculate the percentage
  abundance of each of the three isotopes.

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• E.g4. copper has 2 isotopes. 63Cu has a relative
  isotopic mass of 62.95 and 65Cu has a relative
  isotopic mass of 64.95. The relative atomic mass
  of copper is 63.54. Calculate the percentage
  abundances of the two isotopes

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• Chemists can not deal with quantities as small as
  individual atoms, ions or molecules.
• So they use a convenient quantity of atoms,
  molecules or ions that they can weigh out and
  work with in labs.
• The convenient quantity is called the mole

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The mole concept

• 1 mole or mol is equals to 6.02 x 1023
• The same way as one dozen equals 12
• The symbol for a mole is n
• 6. 02 x 1023 is called the Avogadros constant
  and has a symbol NA

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The mole concept

• 1 mole of O2 molecules is equals to 6.02 x 1023
  molecules of O2
• 1 mole of H2 atoms is equals to 6.02 x 1023
  atoms of H2
• 1 mole of H2 ions is equals to 6.02 x 1023 ions
  of H2
• Example 4.2a, 4.2b, p60

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Examples
You can see that n NA / N where, N is the
number of particles in a given substance, NA 6.02
x 1023, and n is the number of mols. Exercise
4.2, Q5, Q6, Q7, p61
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• N m / M, where m is the mass of the given
  substance, M is the molar mass and n is the
  number of mols
• Do exercise 4.3, p65

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The mole concept

• 1 molecule of H2O contains 1 atom of oxygen and
  2 atoms of hydrogen
• Therefore, 1 mole of H2 O will contain moles
  of oxygen atoms and moles of hydrogen atoms
• Worked example 4.2c, 4.2d, p61
• Exercise 4.2, Q8, p61

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The mole concept

• Calculate the amount (in mol) of oxygen atoms in
  5 moles of oxygen
• Calculate the amount (In mol) of hydrogen atoms
  in 2.5 mol of water molecule
• The number of hydrogen atoms in 2.5 mol of water
  molecules
• Exercise 4.2, Q8, p61

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Molar mass

• Is a relative mass expressed in grams
• E.g. The molar mass of C is 12g
• Mg is 24g.

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• The molar mass of a compound is the relative
  formula mass expressed in g.
• E.g. NaCl, H2O
• Calculate the molar mass of table sugar
  (C12H22O11)
• Calculate the mass of 2.5 moles of table sugar

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• We can see that n m/M, where n is the number
  of mol, m is the mass of a given substance and M
  is the molar mass of the substance.
• Do Exercise 4.3, Q9, 10, 11

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• amyl ethanoate C7 H14 O2 is the compound that
  gives banana their characteristic odour. It is
  also the compound released when bees sting. Each
  time a bee stings, 10-6 g of amyl ethanoate is
  released. Calculate the number of amyl ethanoate
  molecules released in each bee sting and the
  total number of atoms released in each sting
• Do Exercise 4.3, Q12, 13

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• Calculate the mass of 0.35 mol of magnesium
  nitrate Mg(NO3)2
• Calculate the amount (in mol) of CO2 molecules
  present in 22 g of carbon dioxide
• What is the number of molecules present in this
  mass of CO2