1. Atomic Structure
IB Chemistry Topic 2

2. The Atom
Originally thought to be the smallest form of matter due to the fact that the atom can’t be broken down into simpler components by chemical reaction.

3. Dalton’s Atom
All matter is composed of tiny indestructible particles called atoms
Atoms cannot be created or destroyed
Atoms of the same element are alike in every way
Atoms of different elements are different
Atoms can combine together in small numbers to form molecules (compounds)

4. Dalton’s Atom
Dalton’s Law of Constant Composition is seen in the chemical formulas of compounds
Each atom is represented by its element symbol
The number of each type of atom is indicated by a subscript written to the right of the symbol
H2O N2O5 C6H12O Fe2(CO3)3

5. J.J. Thompson
J.J. Thomson ( ) –atoms contain negative particles (electrons)
Proposed the “plum pudding” model of the atom - if a negative charge was present in atoms, there must also be a positive to balance it; negative charges were distributed evenly on a positive sphere (raisins suspended on pudding)

6. Gold Foil Experiment
Ernest Rutherford ( ) aimed positively charged particles at a thin metal foil; most particles went through but some were deflected in different directions

7. (a)If the plum pudding model was correct.
(b)Actual results

8. Structure of the Atom
Protons + neutrons = nucleons  located in the nucleus
Electrons are found in energy levels or shells surrounding the nucleus
Most of the atom is empty space

9. Structure of the Atom

10. Describing the atom
Mass Number (A) = number of protons + number of neutrons in an element X
Charge if it is an ion, zero if not A
n+ / n- Z
Atomic Number (Z) = the number of protons (also electrons in a neutral atom). The atomic number is unchanging, it is the unique identifier of an atom.

11. Label A and Z for the following:
Be Ge O Te Mo Rb

12. Isotopes
Atoms of the same element (same atomic #) with different masses due to a different number of neutrons .
Examples:
1H, 2H, 3H
C-12, C-14
35Cl, 37Cl

13. Deduce the symbol for an isotope given its mass number and atomic number.
1. p = 6 n = 7 2. p = 3 n = 4 3. p = 16 n = 17 e = 18

14. Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. Cl
e=?
p=?
n=?
2. Which subatomic particle occurs in the same amount in both of the following species?
P S 35 17 31 32 2- 15 16

15. Properties of Isotopes
All isotopes of an element have the same chemical properties because chemical properties are determined by the protons and electrons, not the mass.
Isotopes have different physical properties.
Ex: rate of diffusion, mass, density, melting point, boiling point.

16. Radioisotopes
Some isotopes are radioactive (the nuclei of these atoms break down spontaneously and emit radiation).

17. Radioisotopes
Uses:
Nuclear power generation, sterilization of surgical instruments in hospitals, crime detection, finding cracks and stresses in structural materials, food preservation…
Add to your notes: Iodine – 131, Cobalt – 60, Carbon – 14, (brief description and uses)
Shroud of Turin
Pgs

18. The electromagnetic spectrum
𝐸=ℎ𝜐= ℎ𝑐 𝜆
The electromagnetic spectrum
Visible light only makes up a small part of the electromagnetic spectrum.

19. Different types of spectra
A continuous spectrum is produced when white light is passed through a prism and shows all the frequencies.
Line spectra
When white light passes through gases certain absorption will occur
Results in a line spectrum produce with some colors of continuous spectrum missing
Different elements have different line spectra
Therefore elements can be identified by their line spectra similar to products with their barcodes

22. Transitions and EMS Hydrogen spectrum:
The type of radiation given out by an atom is dependent on where an electron falls to from its excited state
Series nf ni
Region of EMS
Lyman 1
2,3,4,5… UV
Balmer 2
3,4,5,6…
Visible and UV
Paschen 3
4,5,6,7… IR
You don’t have to know the series name, but you do need to know which region of the EMS goes with which transitions.

23. Balmer:

24. Heisenberg’s uncertainty principle:
It is impossible to accurately know both the position and the momentum of an electron. The more we know about the position of an electron, the less we know about the momentum, and vice versa.
Schrödinger’s equation:
A very complex math equation whose results are describes by atomic orbitals. The results describe the probability density of the space an electron can occupy.
An atomic orbital is a region in space where there is a high probability of finding an electron.

25. Orbitals Energy shells are divided into 4 levels:
Principal quantum number, n, has integer values (1, 2, 3…) and can hold 2n2 electrons. This is the main energy level.
Sublevels, l: s, p, d, and f. Each of these levels can hold 2 electrons, but the number of orbitals, ml, in each sublevel is different:
Sublevel
# of orbitals
Max e- number s 1 2 p 3 6 d 5 10 f 7 14

26. Orbital Shapes

27. Orbitals Orbitals in sublevel p are labeled px, py, and pz.
**Labels for d and f do not need to be known.
Within each orbital, are two electrons that have spins, ms, with values of +1/2 or -1/2.
Pauli exclusion principle: an orbital can hold two electron and they must have opposite spins.
Hund’s rule: electrons will fill each degenerate orbital (orbitals of equal energy) singly, before occupying them pairs.
ms = magnetic quantum number

28. Orbital Diagrams
When drawing orbital diagrams, we use the arrows-and-boxes method.
For each sublevel, there is one box per orbital.
For s: 1 box For p: 3 boxes For d: 5 boxes For f: 7 boxes

29. Orbital diagrams
Aufbau principle: electrons fill the lowest-energy orbital that is available first. Ex: Sulfur 16 electrons

30. Electron Configuration
Full electron configuration:
Sulfur: 1s12s22p63s23p4
Electron configuration shorthand:
Sulfur: [Ne]3s23p4
Remember this from CP? 