1. 5. QUANTIFYING CHEMISTRY Chapter 5

2.  Atoms are extremely tiny and have a very very tiny mass. How do we measure atoms?  We have a convenient way to compare and measure masses and amounts of particles in elements and compounds.

3. 5.1 Measuring masses using Mass Spectrometry An instrument called a mass spectrometer can identify naturally occurring isotopes of elements. How? A sample of an element is injected as a gas into an ionisation chamber, where the atoms are ionised into positive ions by bombarding electrons The positive ions are accelerated through an electric field and deflected in a magnetic field along different curved paths depending on their mass to charge ratio.

4. The heavier ions are harder to deflect and so travel in a wider curve (provided the charge of the ion is +1). The different ions are picked up by an ion collector.  A graph called a mass spectrum is produced. This consists of a series of peaks.  Each peak represents a particular isotope. (Remember individual isotopes of an element have a different mass)

5. Mass Spectrometer Its Main Features

6. Mass spectrum The mass spectrum provides us with information about: the number of isotopes in a given sample of an element (the number of peaks) the relative isotopic mass of each isotope (the horizontal position on the mass spectrum) the percentage abundance of the isotopes (the peak height). See page 93 for an example.

7. 5.2 Relative Isotopic Mass (RIM) Chemists have decided on a special scale to measure the mass of atoms. The scale is based on the Carbon 12 isotope. This is assigned the mass of 12.00 EXACTLY. All other atoms are measured against this The relative isotopic mass is the mass of a single atom of an isotope relative to carbon 12. It is determined by comparing the mass of ions of the isotope using a mass spectrometer to the value of carbon 12

8. Mass Spectrum Example Consider the mass spectrum for the element neon (page 93). The element neon is made up of two isotopes. By looking at the peak height we see that 90% of the element neon is made up of the isotope with relative mass of 20, and 10% is the isotope with relative mass of 22.

9. 5.3 Relative Atomic Mass Most elements consist of a mixture of isotopes with different abundances. We need to know what the ‘average’ mass of an atom of an element is The RELATIVE ATOMIC MASS (A r ) of an element represents the average mass of one atom, taking into consideration the number of isotopes of the element, their relative isotopic mass (RIM) and their relative abundance. It is the weighted mean of the isotopic masses. See formula on bottom of page 93

10. Example Find the RAM for Neon if it is made up of 90% of an isotope of isotopic mass 20 and 10% of an isotope of isotopic mass 22. So some Ne atoms weigh 20 (90%) and some weigh 22. (10%) But the RAM tells us what the ‘average’ is, which is 20.2

11. 5.5 Relative Molecular Mass (M r ) The Relative Molecular Mass of a compound is the sum of the relative atomic masses of the elements (relative to carbon 12) Example Find the M r for H2O (the RAM for H is 1.0 and for O is 16)

12. f 5.6 Counting Atoms- the mole Chemists need to count very, very large numbers of atoms. The usual counting units such as 10s, 100s, dozens, billions etc isn’t very useful because they are too small!. A new counting scale has been invented. The number of carbon atoms in 12 gram of carbon 12 has been counted. There are (approximately) 6.02 X 10 ^23 atoms This number has been assigned the value of ONE MOLE.

13. UV-VISIBLE (UV-Vis) SPECTROSCOPY

14. Definitions One mole represents 6.o2 X 10^23 particles. 6,02 x 10^23 is called Avogadro’s number, N A A mole is the number of particles in exactly 12 gram of carbon 12. The mole has the symbol n and the unit mol. Example How many mole of atoms in 24 gram of carbon 12? How many atoms is this?

15. A Clever trick Its no accident that 12 grams of carbon 12 was used. See how the numbers match. We can use this to calculate the mass of a mole of a substance easily. We simply find the RAM (or RMM) and add the unit grams/mol.

16. Molar mass (M) This is the mass of one mole of a substance. The unit is grams/mole. Symbol is M Most Important formula!! n = m/ M n =mole m = mass of substance in grams M = molar mass (either RAM or RMM)

17. 5.5 Mole Formulae 1. n = m/ M n =mole m = mass of substance in grams M = molar mass (either RAM or RMM) 1. n = N/ N A n =mole N = number of particles N A = Avogadro’s number (6.02 X 10^23)

18. Questions 1. A gas balloon contains 5.5 mol of helium atoms. How many helium atoms are present? n = N/ N A N = n x N A N = 5,5 x 6.02 x 10^23 = 33.11 x 10^23 = 3.311 x 10^24 2. Determine how many mol of ethanoic acid molecules are present in 16.2 g of ethanoic acid, CH 3 COOH. n = m/ M (M = 12 + 3x1 + 12 + 2x16 + 1 = 60) n = 16.2/60 = 0.27 mol

19. Questions 3. Pure ethanoic acid, CH 3 COOH, can be used to make vinegar when dissolved in water. If 3.5 moles of ethanoic acid was used, what mass was weighed out? n = m/ M (M =60 from Q2) m = n x M = 3.5 x 60 = 210 g 4. Consider 5.2 × 10 24 molecules of methane, a) calculate the amount (in mol) of CH 4.methane molecules n = N/ N A

20. Questions b) amount (in mol) of carbon atoms c) amount (in mol) of hydrogen atoms