Atoms… and other stuff!

1. Atoms… and other stuff! What is an element? Substance with atoms of only one type

3. *History of the Atom

4. *Rutherford… Didn’t split the atom…. ….more like he chipped it By bombarding nitrogen with alpha particles: 14N + α → 17O + proton. This chipped a proton off (lowering the atomic number by one to produce oxygen). Nb: he didn’t know he had protons as no-one knew about these, he detected hydrogen nuclei

6. Starter Clips • Higgs Boson (staff folder – naked scientists) • Size of an atom: • http://www.youtube.com/watch?feature=player_detailpage&v=yQP4UJhNn0I • http://www.knowmia.com/watch/lesson/802 how small IS an atom? • http://www.knowmia.com/watch/lesson/8469 • See links at bottom too. • You cant touch anything (vsauce) • http://www.youtube.com/watch?v=yE8rkG9Dw4s

7. Content Starts Here

8. ASAP Science Periodic Table Song: https://www.youtube.com/watch?v=zUDDiWtFtEM

9. Physical Properties of Metals • Strong • Hard • Dense • High lustre - shinyness • Conduct heat • Conduct electricity • High melting temperature • High boiling temperature • Ductile – can be shaped into a wire • Malleable – beaten into a shape • Solids at room temperature – except Hg (mercury) • Sonorous – makes a sound when hit Malleability & Ductility Clip: http://www.youtube.com/watch?v=c382ziUpbbc

10. * Reasons for the properties of metals Metallic Bonds and Metallic Properties- Metals are made up of closely packed cations rather than neutral atoms- Valence electrons are moving and can drift freely from metal to metal- Metallic bonds consist of the attraction between the free-floating valence electrons for the positively charged metal ions- Metals are good conductors of electrical currents because electrons can flow freely in them- They are also malleable- The reason for both of these is the mobility of the electrons- The sea of valence electrons insulates the metal cations from one another

11. Identifying Metals Describe the salient (significant) physical properties of the following metals. Pick three • Aluminium • Lead • Gold • Mercury • Copper • Titanium

12. Structure of an atom…? How small is this? http://www.youtube.com/watch?feature=player_detailpage&v=yQP4UJhNn0I

14. What do the numbers mean?

15. Interpreting the Periodic Table Mass number Element symbol !! In some periodic tables the mass number and atomic number are switched. Hint: the mass number is ALWAYS the larger number Atomic number Calculating number of subatomic particles -Number of protons: = the atomic number -Number of electrons: = number of protons -Number of neutrons: = mass number – atomic number Example: Carbon -Number of protons = atomic number = 6 -Number of electrons = number of protons = 6 -Number of neutrons = mass number – atomic number = 12 – 6 = 6 Atoms are electrically neutral because they have the same number of protons (+) and electrons (-)

16. Electron Configuration Electrons are arranged into ‘shells’ or ‘levels’ Rules: • Shells are filled up in order from closest to nucleus • Shells are filled completely before starting the next shell Example: Magnesium has 12 electrons the electron arrangement would be: 2,8,2 Why? Fill the 1st shell with 2 electrons, fill the 2nd with 8 and 2 left over to part fill the 3rd shell

17. *Isotopes • Isotopes are variants found in many chemical elements. • They have different numbers of neutrons compared to each other and they differ in weight because of this. • Chemically they all behave alike, but some may be unstable and radioactive • Examples: Hydrogen-1, Hydrogen-2, Carbon-14, Carbon 12, Uranium (many)

18. *Hydrogen Isotopes The extra neutrons increase the mass number Most H in the universe is H-1, some is H-2 and H-3. To get a periodic table mass number we average ALL the hydrogen which is why the mass number on the table is more than 1 (it is 1.00794) p = 1 e = 1 n = 0 p = 1 e = 1 n = 1 p = 1 e = 1 n = 2 (H-1) (H-3) (H-2)

19. Table Salt(aka common salt, halite) Or: NaCl which is Sodium chloride This is dangerous stuff!!!!

20. Sodium metal Chlorine gas

21. Ions & Ionic Bonding

22. Most atoms are  because they want outer shells to be full = more stable Neutral Atoms Sodium atom (Na)Chlorine atom (Cl) 11p & 11e 17p & 17e Overall charge: 0 Overall charge: 0 e arrangement: 2,8,1e arrangement: 2,8,7 He, Ne, Ar... (the inert gases) refuse to react, their outer shells are full and they are  (stable)

23. Forming Ions Ions are charged atoms – they have charge because they have lost or gained electrons Sodium Atom… Chlorine Atom… Loses 1e (to make outer shell full) Gains 1e (to make outer shell full) Result: 11p and 10e Result: 17p and 18e Overall charge +1 Overall charge -1 Result: Sodium ion Result: Chloride ion

24. Rules for Naming Ions Metal ions • Write the atomic name • Then write “ion” • Eg sodium becomes “sodium ion” • Transition metals (middle of table)can form more than one ion type. Roman numerals indicate the charge • Eg copper (II) becomes Cu2+ (also gets a special name: cupric ion) Non metal ions (with one atom type) • Shorten the atomic name • Then write “-ide” suffix (ending) • Eg oxygen becomes “oxide” Non metal ions (with more than one atom type) • More complicated…. • Usually end in –ite, -ate, ium or –ide • -ite, -ate contain oxygen • EgsNO3- nitrate ion, CO32- SO42- (you can work out easy metals and non metals using the periodic table and you will usually be given a table of ions when needed for the polyatomic ions)

25. Ionic Bonding - Ionic Compound • Opposite charges attract. • So, the Na+ ion and the Cl- ion attract • This attraction holds them together in an ionic bond Atoms are  now, outer shells are now full = stable

27. Compound Formulae #1 Example: Calcium bromide Calcium ion = Ca2+ Bromide ion = Br- If you have one Ca2+ and one Br- charge will be +2 + -1= +1 The final compound must be neutral! A second Br-needed to get an overall charge of 0 Ca2+ Br- Br- The formula is CaBr2

28. Compound Formulae #2If all else fails: swap, drop & lose! “swap the numbers over, drop them to the bottom, lose the charges” Remember: don’t write the 1s (Li2O not Li2O1) simplify eg: Mg2O2 becomes MgO

29. Compound Formulae #3Polyatomic ions (polyatomic = many atoms in the ion) Aka radical Example: Magnesium nitrate Sodium ion = Mg2+Nitrate ion = NO3- If you have one Mg2+ and one NO3-charge will be +1 + -1= +1 A second NO3- is needed to get an overall charge of 0 Mg2+ NO3- NO3- In the final formula we must write brackets around the two NO3- to show that there is two of them ie: Mg(NO3)2

30. Valence Electrons “the number of electrons in the outermost shell” These are the electrons that are available to form chemical bonds Mg (electron arrangement 2,8,2 has two valance electrons) Mg loses 2 electrons (to make its outer shell full) Now: 10 electrons Is now: magnesium ion, Mg2+ Arrangement 2,8

32. Equations – Word • Magnesium burning again… Magnesium + oxygen  magnesium oxide Arrow means ‘reacts’ Products on the right Reactants on the left

33. Equations – Chemical or Symbol Magnesium + oxygen  magnesium oxide Becomes: Mg + O2 MgO Must obey the usual rules for ions and compounds. Eg Mg2+ & O2- balance to become MgO

34. Balanced Chemical Equations 1 Can’t change element type (we’re doing chemical reactions, not nuclear so same types of atoms must be on each side of the reaction) AND Can’t get something from nothing (the total numbers of atoms on each side of the reaction must be the same) Mg + O2 MgO Mg + O2 2MgO 2Mg + O2 2MgO 2 Problem as 2 x O on left and 1 x O on right. Can’t make MgO become MgO2 as that’s not magnesium oxide anymore! 3 But can have extras of any whole molecule by putting numbers IN FRONT 4 For example, the red two fixes the oxygen problem (now two on each side, but now have a magnesium imbalance. 5 The green two fixes that and we now have the same number of each type of atom on both sides. The equation is balanced!

35. Try balancing these: The reaction of sodium in water (makes sodium hydroxide and hydrogen gas) Na + H2O -> NaOH+ H2

36. Covalent Bonding

37. Covalent Bonding – H2 Hydrogen Atom 1p & 1e (needs one extra e to get full shell) Two Hydrogen Atoms One electron from each atom is shared by the other, forming one H2 molecule with one covalent bond

38. Covalent Bonding • Non metal atoms tend to gain electrons rather than lose them • They can form bonds by sharing their electrons between them (to get full shells) this forms a covalent bond

39. Other Covalent Molecules… N2 O2 Also: CO2, CH4 (methane)

40. Ionic vs Molecular (covalent) Ionic - Made of metal & non metal elements - Electrons donated Examples Sodium chloride Lithium fluoride Molecular - Made of non metal elements only - Electrons shared Examples Nitrogen dioxide Water ! This is oversimplified somewhat. There is a spectrum between complete donation and complete shared ! Molecule can mean one unit eg one molecule of NaCl (which is NOT molecular!)

41. Physical vs Chemical Changes Chemical • not easily reversed • new product(s) formed • reactants used up • often heat / light / sound / fizzing occurs • electricity may be produced • *a precipitate may form • e.g. wood burning. Physical • easily reversible • no new products, • often just a state change • e.g. ice melting.

42. What Am I Made Of?