1. UNIT 1 – MATTER AND QUALITATIVE ANALYSIS
The Electromagnetic Spectrum, Bohr’s Model, Chemical Bonding

2. OBJECTIVES

3. The electromagnetic spectrum
Electromagnetic Energy: light energy that travels in the form of waves
Frequency: the number of cycles per second
Wavelength: he distance between successive crests or troughs in a wave
Nanometer: 10-9 m; unit nm

4. WAVES

5. LIGHT AND THE ELECTROMAGNETIC SPECTRUM
There are many different kinds of light: X-rays, ultraviolet, infrared, microwaves are all examples of light
Visible spectrum: the region of the electromagnetic spectrum tat the human eye can see
400 to 700 nm
A rainbow contains all the colours of visible and near visible light
Continuous Spectrum: an uninterrupted pattern of colours that is observed when a narrow beam of white light passes through a prism

6. PRISM

7. VISIBLE SPECTRUM

8. Radio waves have the lowest frequency/longest wavelengths and lowest energy – not dangerous
Gamma Rays have the highest frequency/shortest wavelengths and highest energy – very daengerous

9. LINE SPECTRUM
Different types of matter emit different wavelengths of light:
Neon gas emits red light
Sodium ions emit yellow light
Hydrogen gas emits blue/violet light
Each element has a different ‘line spectrum’
Discontinuous spectrum of light that is produced when light the light from the gas is passed through a type of prism called a spectroscope
You see coloured lines rather than a rainbow

11. Continuous vs. Line

12. DIFFERENT ELEMENT, DIFFERENT LINE SPECTRUM: THE ELEMENT ‘FINGERPRINT’

13. QUESTIONS
Describe the difference between radio waves and X-rays using the concepts of wavelength, frequency, and energy
What range of wavelengths of electromagnetic radiation can the human eye detect
White light is composed of many different colours of light. Explain.
Distinguish between a continuous spectrum and a line spectrum

14. The bohr-model of the hydrogen atom
Negatively charged particles orbiting positive nucleus according to Rutherford… what is the problem with that?
(+) and (–) attract… atom should collapse on itself
Also, laws of moving charges means that orbiting electrons should emit energy and eventually run out....
SPIRAL DEATH!!! (whomp whomp… that’s a problem)
MATTER IS VERY STABLE

15. HYDROGEN GAS LINE SPECTRUM

16. Niels bohr To the rescue!!!
Looking at the discrete line spectrum of hydrogen… observed discrete lines, inferred fixed energy levels of electrons
Inferred electrons are restricted to quantized energy levels
Quantized: possessing specific values or amounts (a quantity)
That is, electron’s can only have specific energy and orbit around the nucleus in orbitals at specific distances, the further away the greater the energy
i.e. a ball sitting on a step… it cannot rest between levels, gravity pulls it down to rest on specific steps only… if they want to change levels, either you must add energy to increase the levels (by picking it up and putting it at the next step) or it must be given energy to roll off the step and on to a lower one

17. Explanation Electrons occupy specific energy states (orbitals)
Can only transfer levels by absorbing the EXACT amount of energy needed between them, to go up OR to give off (emit) the energy difference between lower energy levels..
Hence, discrete energy levels
In doing so, it either absorbs or emits a particular amount of light…

18. each atom has slightly different allowable energy differences and different numbers of electrons that can do the transitions, therefore different spectrums for each
Ground state: level the electron wants to be at
Excited state: level with added energy

20. Each energy level can only have a certain number of electrons in it
1st = 2
2nd = 8
3rd = 8

21. example
Draw the Bohr diagrams for:
Neon
Helium
Carbon
Calcium
Oxygen

22. THE NUMBER OF ELECTRONS IN THE VALENCE (OUTER) ORBITAL (SHELL) CORRESPONDS TO THE COLUMN (GROUP) THE ELEMENT IS N ON THE PERIODIC TABLE
Ex: Oxygen (Group 6) has 6 electrons in its valence shell; Sodium (group 1) has 1

23. Homework questions
State the reasons why Rutherford’s model of the atom failed to describe the observed behaviour of matter
Describe Bohr’s model of the atom. How is it similar and how is it different to Rutherford’s Model of the Atom?
What role did Spectroscopy play in helping Bohr come up with his model?
Why do electrons emit light energy when they drop from a higher to lower energy level?
What does it mean by “energy levels are quantized?
Drawing Bohr diagrams

24. FLAME TESTS

25. Four types of qualitative analysis
Thermal Emission Spectroscopy (TES): substances are identified based on amount of heat they emit
Line spectroscope: substances are identified based on their line spectra
Flame tests: substances are identified based on the colours they emit when placed in a flame
Carbonation: the presence of metals is identified based on the emission of carbon dioxide when reacted with a metal

26. Lewis symbols
The inner electrons are really just place holders; outer electrons determine reactivity
Chemists shorten Bohr Diagrams to Lewis Structures
Example: Draw the Lewis Structure For
Ca H
N Ne
Li Cl

27. Polyatomic ions
Ion that is composed of two or more atoms (all of them end in –ate, you will be given a list of these on a test)
Examples:
Sulfate SO42-
Bromate BrO3-
Chlorate ClO3-
Nitrate NO3-
Phosphate PO43-

28. Forming ionic compounds
Occur when metal and non-metal react with each other
Ionic crystals: solid that consist of a large number of cations and anions arranged in repeating 3D patterns
Ex: Show the formation of Sodium Chloride; Calcium Oxide

29. Naming ionic compounds
USE IUPAC (International Union of Pure and Applied Chemistry)
METAL + NON-METAL (IDE)
Example:
CaO
NaCl
Al2O3

30. Writing formulas Use the criss-cross method Example: Sodium Oxide
Rubidium Fluoride
Strontium Nitride

31. practice
Worksheets!!