1. Atoms -Basics on atomic structure -Isotopes -Isotopic Mass & Relative Atomic Mass

2. Revision from Year 11 The Bohr model of the atom. It’s not perfect, but it is a simple model.

3. ParticleChargeMass (u) Neutron Proton Electron

4. Mass and charge With your partner, try to describe the distribution of mass in an atom (i.e where is it and why?) Do the same for charge, where is the charge distributed within an atom? – A diagram may help.

6. The Atomic number. Each element has a unique atomic number. This number tells us two things: 1.How many protons the atom has 2.How many electrons the atom has. Why are these numbers the same?

7. Atomic Mass. Try not to confuse the atomic MASS and the atomic NUMBER. The Atomic Mass tells us how much stuff is in the atom. How many PROTONS and NEUTRONS it has. Why might we ignore electrons when we talk about atomic mass?

8. Atomic number & atomic mass Atomic number – Found in the periodic table – Table arranged by increasing atomic number – Tells us the number of protons in the nucleus. Atomic Mass – Found in the periodic table – Tells us the of protons AND neutrons in the nucleus – A larger number than the atomic number.

9. We’re pretty lazy We have a short way of writing the atomic number and the mass number for an element: Which number is atomic number, which is the mass number?

10. Isotopes. Isotopes are elements of the same type but with different mass.

11. Chlorine has two stable isotopes. How many protons and neutrons does each have? Try the others. 35 Cl 131 I 12 C 37 Cl 127 I 13 C

12. Carbon 12. It’s important. 12 C is used as a standard for the measurement of relative mass. Atomic masses are measured using a unit called unified atomic mass unit. This is a tiny number- 1 u is 1.66 x 10 -27 Mass of carbon atom- 12 u So one twelfth of an atom of carbon is 1 u

13. Relative isotopic mass The Relative Isotopic Mass is the mass of an atom of an isotope compared to one twelfth of the mass of an atom of Carbon-12.

14. Relative isotopic mass The Relative Isotopic Mass is the mass of an atom of an isotope compared to one twelfth of the mass of an atom of Carbon-12.

15. Relative atomic mass. A r In a lump of carbon we have 12 C and 13 C. – These weigh a different amount. – We use a ‘weighted mean mass’ to account for this. – This depends on: The percentage abundance of the isotope The relative mass of the isotope.

16. A r Definition The relative atomic mass is the weighted mean mass of an atom of an element compared to one-twelfth of the mass of an atom of carbon-12

17. A r Definition The relative atomic mass is the weighted mean mass of an atom of an element compared to one-twelfth of the mass of an atom of carbon-12

18. ArAr For example: Boron contains 19.77% 10 B and 80.23% 11 B A r (B)= (19.77/100 x 10) + (80.23/100 x 11) = 1.977+8.8253 = 10.8023 u