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National 5 Supported Study

National 5 Supported Study. Atomic Structure and Bonding Week 2. Mandatory Area. Atomic structure and bonding related to properties of materials Nuclide notation. Isotopes and relative atomic mass. Ions. Ionic bonding. Covalent molecular, covalent network and ionic lattices.

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National 5 Supported Study

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  1. National 5 Supported Study Atomic Structure and Bonding Week 2

  2. Mandatory Area • Atomic structure and bonding related to properties of materials • Nuclide notation. Isotopes and relative atomic mass. Ions. Ionic bonding. • Covalent molecular, covalent network and ionic lattices. • Physical properties of chemicals explained through bonding. • Chemical and ionic formulae including group ions.

  3. Structure of the Atom • 1st electron shell holds 2 electrons • 2nd electron shell holds 8 electrons • 3rd Electron shell holds 8 electrons REMEMBER THE NUCLEUS HAS A POSITIVE CHARGE

  4. The Atom

  5. The Atom • Is overall Neutral • Because has the same number of positive protons and negative electrons • Different elements have different number of protons • Recognised as an ATOMIC NUMBER

  6. Atomic Number • Unique for every element • Equals the • Number of protons • Number of electrons (for an atom)

  7. Mass Number • The only two particles which have mass in an atom are • Protons • Neutrons BOTH ARE FOUND IN NUCLEUS

  8. Mass Number • The only two particles which have mass in an atom are • Protons • Neutrons • MASS NUMBER = • PROTONS + NEUTRONS • NO. OF NEUTRONS = • MASS NUMBER – ATOMIC NUMBER BOTH ARE FOUND IN NUCLEUS

  9. Elements • Number of protons = never changes • Is like the elements DNA, will never change • Number of neutrons = affect mass of atom • Some heavier, some lighter atoms • Number of electrons = affects the mood the element in • Neutral • Positive • Negative

  10. Nuclide Notation • Note the numbers are written at LHS of symbol • Atomic Number is at the bottom – can be checked using data book if you forget • Mass Number (the bigger number) at the top.

  11. Nuclide Notation • No. of protons = Atomic Number • = 17 • No. of electrons = Atomic Number • = 17 • No. of neutrons = Mass No. – Atomic No. • = 37- 17 = 20

  12. Isotopes • Are atoms with the • Same atomic number • But, Different mass number • i.e. Different number of neutrons

  13. Isotopes • Most elements exist as a mixture of isotopes

  14. Relative Atomic Mass • This is rarely a whole number • WHY? • The relative atomic mass is an average of all the different Isotopes taking into account the proportion of each present.

  15. Relative Atomic Mass • Relative Atomic Mass = 36.5 • Would expect Average to be = 36 • Closer to 37 = Grater proportion of 37 isotope present

  16. Relative Atomic Mass • Relative Atomic Mass = 36.0 • This tells you that there is exactly 50% of each isotope present as the average lies exactly in the middle.

  17. Ions • Charged particles

  18. Periodic Table • Group 1 – Alkali Metals • Reactive • Group 7 – Halogens • Reactive • Group 8 (0) – Nobel gases • Unreactive

  19. Bonding • Atoms held together by bonds • Bonds are formed so elements can achieve a stable electron arrangement • i.e. a full outer shell like the noble gas elements. REMEMBER GROUP 8 ARE UNREACTIVE AND SO NOT BOND

  20. Covalent Bonding • Non-metals only • Atoms share electrons • Covalent bond held together by • Attraction between the positive nuclei and the negative shared electrons • Molecule = group of atoms held together by COVALENT bonds

  21. Diatomic • 2 elements bonded together • Diatomic Elements • H, N, O, F, Cl, Br, I • Diatomic Compounds • CO, NO

  22. Molecular vs Network • Low Mp and Bp • Because strong covalent bonds not broken • Weak forces BETWEEN molecules broken • High Mp and Bp • Because strong covalent bonds need to be broken and this required a lot of energy. • E.g. Diamond Sand (SiO2) MOLECULAR NETWORK

  23. Covalent • (usually) soluble in non-aqueous solvents • (usually) insoluble in aqueous solvents(water) • NEVER CONDUCT ELECTRICITY!!!!

  24. Bonding Diagrams

  25. Bonding Diagrams – Part 2 Try CH4 NH3 HCl CO2 – very tricky!

  26. Ions • Charged Particle • Metals – lose electrons – form +ve Ions • Non-metals – gain electrons – form –ve Ions MP

  27. Ionic Bonding • Between metals and non-metals • Bond held together by electrostatic attraction between POSITIVE AND NEGATIVE IONS Crystal Lattice Structure of oppositely charged ions

  28. Ionic Compounds • Do not conduct when solid as IONS NOT FREE TO MOVE • Conduct when in solution or molten as IONS ARE FREE TO MOVE • Solid at room temperature • Strong forces of attraction that need to be broken • Soluble is aqueous solvent (i.e. water)

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