1. Chemistry (14 - 16) The Atom (Basic Structure - 2) © SSER Ltd.

2. The electron levels fill as we move across and down the Periodic Table. Electronic Structure

3. In Group 1 of the Periodic Table, all of the elements have one electron in their outer shell. This similarity in electron structure explains the fact that the elements in Group 1 have similar chemical properties. Group 1 - Electronic Structure Li 3 2,1 Na 11 2,8,1 K 19 2,8,8,1 Rb 37 2,8,18,8,1 Cs 55 2,8,18,18,8,1

4. Group 1 - Electronic Structure Group 1 elements all have similar chemical properties, but the presence of an increasing number of complete electron shells inside the outer level does affect the properties. Li 3 2,1 Na 11 2,8,1 K 19 2,8,8,1 Rb 37 2,8,18,8,1 Cs 55 2,8,18,18,8,1 In most of their chemical reactions, the atoms release this electron and become positively charged ions.

5. Group 1 - Electronic Structure Lithium reacts with cold water on contact and slowly dissolves to form a solution of lithium hydroxide. Caesium reacts explosively on contact with cold water to form a solution of caesium hydroxide. As we descend the group, the metals become more reactive. Li 3 2,1 Na 11 2,8,1 K 19 2,8,8,1 Rb 37 2,8,18,8,1 Cs 55 2,8,18,18,8,1

6. In Group 7 of the Periodic Table (Halogens) all of the elements have an incomplete outer shell - one electron is needed to complete it. This similarity in electron structure gives the elements in Group 7 similar chemical properties. Group 7 - Electronic Structure F 9 Cl 17 Br 35 I 53 2,7 2,8,7 2,8,18,7 2,8,18,18,7

7. Group 7 elements all form diatomic molecules, e.g. F 2, Cl 2. Group 7 - Electronic Structure F 9 Cl 17 Br 35 I 53 2,7 2,8,7 2,8,18,7 2,8,18,18,7 As the group is descended, the elements become less reactive.

8. The elements in Group 7 all form diatomic molecules, e.g. F 2, Cl 2. They are all colourful non-metals.

9. In Group 0 of the Periodic Table, all of the elements have full outer shells of electrons. This similarity in electron structure explains gives the elements in Group 0 similar chemical properties. The elements are unreactive and monatomic (single atoms). Group 0 - Electronic Structure He 2 Ne 10 Ar 18 Kr 36 Xe 54 2 2,8 2,8,8 2,8,18,8 2,8,18,18,8

10. In Group 0 the elements all have similar chemical properties, but the presence of an increasing number of complete electron shells does affect their properties. As the group is descended, the gases become heavier and ‘more’ reactive. Helium and Xenon have no known natural compounds. Group 0 - Electronic Structure He 2 Ne 10 Ar 18 Kr 36 Xe 54 2 2,8 2,8,8 2,8,18,8 2,8,18,18,8

11. A Space Filling Model of Xenon Tetrafluoride Although Xenon has no known natural compounds, it can be induced to react with Fluorine under extreme conditions.

12. Electron Levels (Shells) Drag the electronic configurations and names into the correct places for the elements below…

13. Electron Levels (Shells) Drag the electronic configurations and names into the correct places for the elements below…

14. Li 3 B 5 N 7 C 6 O 8 F 9 Ne 10 Be 4 If we examine the elements in the first short period of the Periodic Table we can see a pattern in their atomic number... 1st Short Period - Electronic Structure Across the Periodic Table (left to right), the atomic number increases by one each time. As the atomic number is also the same as the number of electrons in an atom – then the number of electrons also increases with the increase in atomic number.

15. The numbers below each atom are the electron configurations of the atoms, e.g. the electron configuration of a carbon atom is 2,4. This means that a carbon atom has two electrons in the first level and four electrons in the second level. 1st Short Period - Electronic Structure

16. The electron configuration of a sulfur atom is 2,8,6. This means that a sulfur atom has:...two electrons in the first level......eight electrons in the second level......six electrons in the third level. 2nd Short Period - Electronic Structure

17. Electronic Structure - Properties The elements change gradually in both physical and chemical properties across the period from left to right. The elements on the left are metallic elements, and the elements to the right are non-metallic.

18. Electronic Structure - Properties Sodium is a reactive metal Chlorine is a reactive gas Argon is an unreactive gas

19. Electronic Structure and Chemical Properties The electronic structures dictate that elements with only a few electrons in their outer levels (e.g. Na, Mg, Al) can most easily lose electrons, gaining an inert gas configuration and thus form positive ions.

20. Electronic Structure and Chemical Properties Elements with nearly complete outer shells (e.g. Cl) can most easily gain electrons, gaining an inert gas configuration and thus form negative ions. The inert gases (e.g. Ar) have a complete outer electron shell and so are unreactive.

21. Elements differ from one another in that their atoms contain different numbers of the sub-atomic particles. The number of protons (or the atomic number) of an atom, is the factor which decides the identity (element) of an atom. All Lithium atoms have 3 protons. All Carbon atoms have 6 protons. All Uranium atoms have 92 protons. What Determines The Element?

22. For each element in turn, calculate the missing entries... AtomSymbolAtomic No.Mass No.ProtonsNeutronsElectrons HydrogenH11 Helium24 Carbon612 Chlorine1737 Calcium20 Copper6429 Zinc3065 Tin50 Tungsten74184 Radon86136 10 1 He 2 22 C 6 6 6 Cl 172017 Ca 4020 Cu 29 35 29 Zn 30 35 30 Sn 119 50 69 50 W 74 110 74 Rn 222 86 Calculate the Particles in an Atom

23. Atomic Structure and the Periodic Table Li 3 B 5 N 7 C 6 O 8 F 9 Ne 10 Be 4 Cl 17 If we examine the elements in the two short periods of the periodic table we can see a pattern in their atomic number... 11 Na 12 Mg 14 Si 13 Al 16 S 15 P 18 Ar Elements are arranged in the Periodic Table in order of the number of protons present in the nucleus (atomic number). Across the periodic table (left to right) the atomic number increases by one each time. The number of electrons is the same as the atomic number.

24. The mass of an atom is mainly in the nucleus. Protons and neutrons have an equal mass of one atomic mass unit. Electrons have negligible mass. Therefore, the mass of an atom is equal to the sum of the number of protons and neutrons. This number is called the mass number of the atom. However, some atoms of carbon have different numbers of neutrons. Atoms of the same element with different numbers of neutrons are called isotopes of that element. Isotopes For example, carbon atoms each have six protons and six neutrons, and therefore the atomic mass is twelve atomic mass units. The mass number of the most common carbon atom is 12. Therefore, isotopes of the same element have their own mass number.

25. 12 6 C Mass Number Atomic Number The Carbon Isotopes 14 6 C The carbon 12 isotope has 6 protons, 6 neutrons and 6 electrons The carbon 14 isotope has 6 protons, 8 neutrons and 6 electrons = Neutron= Proton= Electron

26. For each isotope, calculate the missing entries... Isotope SymbolAtomic No. Mass No.Protons Neutrons Electrons Hydrogen H1 1 Hydrogen (Deuterium) H111 Hydrogen (Tritium) H13 Carbon 12 6 12 Carbon 14 C6 14 101 C 686 Calculating the Atomic Particles in Isotopes 666 121 12

27. Summary 1.Atoms are tiny and in constant motion. 2.Atoms are composed of a nucleus (contains protons, neutrons) and orbiting electrons. 3.The nucleus contains most of the mass of the atom. 4.The number of particles an atom contains determines the identity of the atom. 5.The atomic number of an atom is equal to the number of protons in its nucleus. 6.The mass number of an atom is equal to the number of protons plus neutrons in its nucleus. 7.Atoms have the same number of protons as electrons, so their overall charge is neutral.

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