1. NUFYP OUTCOMES Atomic structure
Use the atomic number to write the electronic configurations of the first 20 elements in the Periodic Table (H to Ca).
Use the standard notation (e.g. 12C) for any atom to calculate the number of protons, neutrons and electrons in an atom (and so any ion of the atom). Define isotopes as atoms of an element with the same number of protons but different numbers of neutrons. Use data to identify the relative abundances of isotopes.
Use Ar values to calculate the relative molecular (formula) mass, Mr, of a compound.

2. NUFYP OUTCOMES The Periodic Table (IUPAC)
Know and use the relationship between the position of an atom in the Periodic Table (Group and Period) and the electronic configuration
Understand that elements in the same Group have similar chemical properties and that down a metal Group, reactivity increases and down a non-metal Group, reactivity decreases.
Know the physical and chemical properties of the alkali metals (Group 1), the halogens (Group 17) and the noble gases (Group 18) and the position in the Periodic Table of the transition metals (d-block elements) and their common properties (coloured ions, multiple stable ions, use as catalysts).
Calculate the relative atomic mass of an element from its isotopes given their relative isotopic masses and their relative abundances.

3. Bohr's Model of the Atom
Bohr's model:
-electrons orbit the nucleus like planets orbit the sun
-each orbit can hold a specific maximum number of electrons

4. He bombarded a thin gold foil with a beam of fast-moving -particles (+ve charged)
Observation:
most -particles passed through the foil without deflection
very few -particles were scattered or rebounded back

5. Interpretation of the experimental results
- The condensed core is called ‘nucleus’.
- The positively charged particle is called ‘proton’.

6. Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
Representation: X A Z
Symbol of the element
Mass number
Atomic number

7. MASS NUMBER AND ATOMIC NUMBER
Atomic Number (Z) Number of protons in the nucleus of an atom
Mass Number (A) Sum of the protons and neutrons in the nucleus
Mass Number (A)
PROTONS + NEUTRONS Na 23 11
Atomic Number (Z)
PROTONS

8. Cl 35 17 37 e.g. the two isotopes of chlorine are written as:
OR labelled as Cl-35 and Cl-37.

9. Interpret the diagram
Protons are deflected on a curved path towards the negative plate.
Electrons are deflected on a curved path towards the positive plate.
The amount of deflection is exactly the same in the electron beam as the proton beam if the energies are the same - but, of course, it is in the opposite direction.
Neutrons continue in a straight line.

10. Practice Fill in the Chart 27Al3+
Protons
Neutrons
Electrons
Charge
Atomic Number
Mass Number
Symbol A 19 21 B 20 40 C + 11 23 D 6 E 92
235 F 13 G 16 2- H
27Al3+

11. MASS NUMBER AND ATOMIC NUMBER
Protons
Neutrons
Electrons
Charge
Atomic Number
Mass Number
Symbol A 19 21 40
40K B 20
40Ca C 11 12 10 + 23
23Na+ D 6
12C E 92
143
235
235U F 7 13
13C G 16 18 2- 32
32S2- H 14 3+ 27
27Al3+

12. Orbitals Sub-level Number of orbitals in sub-level
Shape (no need to learn)
Maximum number of electrons in sub-level s 1 2 p 3 6 d 5 10 f 7
Even more complicated! 14

13. Arrangement of the Orbitals
Aufbau Principle
Arrangement of the Orbitals
There are 4 main subtypes of orbitals in an atom
s, p, d, and f. And the energy goes s>p>d>f
There is 1 s orbital, 3 p orbitals, 5 d orbitals and 7 F’s
There arrangement follows the periodic table
Or does the periodic table follow their arrangement???
Aufbau Principle
Electrons

14. Aufbau Principle Aufbau Principle
Electrons

15. Electron Configurations of Selected Elements Electron configuration
Aufbau Principle
Ele
You try a orbital filling diagram of the oxygen atom.
Electron Configurations of Selected Elements
Element 1s 2s
2px 2py 2pz 3s
Electron configuration H
1s1 He
1s2 Li
1s22s1 C
1s22s22p2 N
1s22s22p3 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 Na
1s22s22p63s1
Notice: each of the three 2p orbitals has one electron. The remaining electron now pairs with an electron occupying one of the 2p orbitals.

16. Aufbau Principle Orbitals
CHEMISTRY & YOU
Explain why the correct electron configuration of oxygen is 1s22s22p4
and not
1s22s22p33s1.
Answer: it does not obey the Aufbau Principle
Orbitals
Electrons

17. Aufbau Principle of Phosphorous
Sample Problem 5.1
The atomic number of phosphorus is 15. Write the electron configuration of a phosphorus atom. 1s 2s 2p 3p 3s 4s
The electron configuration of phosphorus is 1s22s22p63s23p3.
The superscripts add up to the number of electrons.
Electrons

18. Exceptional Electron Configurations.
Cr 1s22s22p63s23p63d44s2
Cu 1s22s22p63s23p63d94s2
The correct electron configurations are as follows:
Cr 1s22s22p63s23p63d54s1
Cu 1s22s22p63s23p63d104s1
These arrangements give chromium a half-filled d sublevel and copper a filled d sublevel- This is because filled and half filled orbitals are more stable
Electrons
Orbitals

19. Ionization Energy Ionization Energy What is ionisation energy?
Definitions
Film Clip
Ionization Energy
Electrons

20. WHAT IS IONISATION ENERGY?
Ionization Energy
WHAT IS IONISATION ENERGY? -
Attraction between the nucleus and an electron
Ionisation Energy is a measure of the amount of energy needed to remove electrons from atoms.
The greater the pull of the nucleus, the harder it will be to pull an electron away from an atom.
FIRST IONISATION ENERGY - Definition
The energy required to remove ONE MOLE of electrons from each atom in ONE MOLE of gaseous atoms to form ONE MOLE of gaseous positive ions.
e.g. Na(g) Na+(g) + e-
Al(g) Al+(g) + e-
Make sure you write in the (g)
Electrons
Ionization Energy

21. Successive Ionisation Energies
Ionization Energy
Successive Ionisation Energies
What does this mean to you?
Mg(g)  Mg+(g) + e- 1st I.E. =+738 kJ.mol-1
Mg+(g)  Mg2+(g) + e- 2nd I.E.= kJ.mol-1
Mg2+(g)  Mg3+(g) + e- 3rd I.E.= kJ.mol-1
Ionization Energy
Electrons

22. Ionization Energy
1st ionisation energy. What is happening across the table? AND Down the noble gases
Li 520
Na 496
K 419
Rb 408
Cs 376

23. Ionization Energy - Groups
Ionization energies decrease going down a Group, Why?
1) Moving down a group the principle energy level increases
distance of electron from the pull of the positive nucleus
2) AND here are more electrons providing inner shell shielding of the
nucleus. Therefore the valence electrons are less strongly held
Na (IE = 496 KJ/Mol) Rb (IE = 403 KJ/mol)
Ionization Energy

24. 1st ionisation energy (across period)
Orbitals
1st ionisation energy (across period)
Atomic Radius DECREASES, why?
Increased nuclear charge (i.e. more protons)
Thus a stronger attraction from nucleus to electron in outer shell
Atoms get smaller
Orbitals
Electrons
© AS Jun-2015

25. Orbitals Orbitals General Trends
Ionization energy increases as each successive
electron is removed
2) Ionization energy DECREASES going DOWN a GROUP
3) Ionization energy INCREASES as you go ACROSS a
PERIOD
4) The noble gases have the highest Ionization energies
5) The alkali metals have the lowest Ionization energies
Orbitals

26. Ions Orbitals https://www.youtube.com/watch?v=vp9mfW7dqE0
IONS FORM by:
ALWAYS removing the highest level of n.
Or the highest energy electrons are lost when an ion is formed. AND the energy is as follows:
least s<p<d<f (most)
Orbitals
Electrons
© AS Jun-2015

27. ELECTRONIC CONFIGURATION OF IONS
3 Things:
Positive ions (cations) are formed by removing electrons from atoms
Negative ions (anions) are formed by adding electrons to atoms
Electrons are removed first from the highest occupied orbitals
SODIUM Na 1s2 2s2 2p6 3s electron removed from the 3s orbital
Na+ 1s2 2s2 2p6
CHLORINE Cl 1s2 2s2 2p6 3s2 3p electron added to the 3p orbital
Cl¯ 1s2 2s2 2p6 3s2 3p6
Anions are always larger than their parent atom.
Cations are always smaller than their parent atom.
Ions

28. ELECTRONIC CONFIGURATION OF IONS FIRST ROW TRANSITION METALS
Despite being of lower energy and being filled first, electrons in the 4s orbital are removed before any electrons in the 3d orbitals.
TITANIUM Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2
Ti+ 1s2 2s2 2p6 3s2 3p6 4s1 3d2
Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2
Ti3+ 1s2 2s2 2p6 3s2 3p6 3d1
Ti4+ 1s2 2s2 2p6 3s2 3p6
Ions

29. Chemical Periodicity? What?

30. Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
Elements in the same group have similar chemical properties.

31. Periodicity When elements are organized in order of  atomic mass, and
grouped by similarities of chemical properties,
a certain “pattern” or periodicity
of properties becomes evident
Draft for first version of
Mendeleev's periodic table
(17 February 1869).

32. Review of the Periodic Table
Group IA – alkali metals

33. Group IA – alkali metals
H s1
Li s22s1
Na11 - 1s22s22p63s1
K s22s22p63s23p64s1
Rb37 - 1s22s22p63s23p63d104s24p65s1
Cs55 - 1s22s22p63s23p63d104s24p64d105s25p66s1
Fr s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p67s1

34. Review of the Periodic Table
Group IIA – alkaline earth metals

35. Group IIA – alkaline earth metals
Be s2 2s2
Mg12 - 1s22s22p63s2
Ca20 - 1s22s22p63s23p64s2
Rb38 - 1s22s22p63s23p63d104s24p65s2
Cs55 - 1s22s22p63s23p63d104s24p64d105s25p66s2
Fr s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p67s2

36. Review of the Periodic Table
Group VIII – Noble Gases

37. Group VIII – Noble Gases
He s2
Ne10 - 1s22s22p6
Ar18 - 1s22s22p63s23p6
Kr36 - 1s22s22p63s23p63d104s24p6
Cs54 - 1s22s22p63s23p63d104s24p64d105s25p6
Rn s22s22p63s23p63d104s24p64d104f145s25p65 d106s26p6

38. Group trends As you go down a group, first IE decreases because...
The electron is further away.
More shielding.
The electron in the outermost energy level experiences more inter-electron repulsion (shielding).

39. Group trends As we go down a group...H
As we go down a group...
each atom has another energy level,
so the atoms get bigger. Li Na K Rb

40. Periodic trends
All the atoms in the same period have the same energy level.
Same shielding.
But, increasing nuclear charge
So IE generally increases from left to right.
Exceptions at full and 1/2 full orbitals.

41. Periodic Trends As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

42. The Periodic Table Always remove electrons from the highest n value?
Na is: 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6
Same configuration as neon (isoelectronic with NEON.
Ions have noble gas configurations (not transition metals)

43. Size of Isoelectronic ions
Iso- means the same
Iso electronic ions have the same # of electrons
Al3+ Mg2+ Na1+ Ne F1- O2- and N3-
all have 10 electrons
all have the configuration: 1s22s22p6

44. Size of Isoelectronic ions
Positive ions that have more protons would be smaller.
N3-
O2-
F1- Ne
Na1+
Al3+
Mg2+

45. ISOTOPES & AVERAGE ATOMIC OR MOLAR MASS

46. MASS SPECTRA
R.A.M.
Consider neon and its 3 isotopes Ne 21Ne Ne.
We use these values to determine the
average molar mass of an element or
Relative atomic mass, called RAM
m/z
90.92
0.26
8.82
Abundance / %
Calculate the average relative atomic mass of neon using the above information.
Out of every 100 atoms are 20Ne , are 21Ne and are 22Ne
Average = (90.92 x 20) + (0.26 x 21) + (8.82 x 22) =
100
Relative atomic mass =

47. Calculate the relative atomic mass of the following
– give your answers to 3 significant figures
Bromine : 79 Br 50% , 81 Br 50%
Copper : Cu 69% , 65 Cu 31%
Zirconium : Zr 51.5% , Zr 11.2%, Zr 17.1%,
94 Zr 17.4%, Zr 2.8%
Lead : 204 Pb 1.5% , Pb 23.6%, Pb 22.6%, Pb 52.3%
Neon : Ne 90.9% , Ne 0.2%, Ne 8.9%
63.6
91.3
207.
20.2

48. Can we calculate the % composition if we know the isotopes ?
Naturally occurring potassium consists of potassium-39 and potassium-41.
Calculate the percentage of each isotope present if the average is 39.1.
Assume there are x nuclei of 39K in every 100; so there will be (100-x) of 41K
so x (100-x) =
100
therefore x x = 3910
thus x =
and x = 95
ANSWER There will be 95% 39K and
5% 41K