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

OBJECTIVES

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

WAVES

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

PRISM

VISIBLE SPECTRUM

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

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

Continuous vs. Line

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

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

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

HYDROGEN GAS LINE SPECTRUM

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

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…

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

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

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

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

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

FLAME TESTS https://www.youtube.com/watch?v=d8hpUtRnsYc

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

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

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-

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

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

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

practice Worksheets!!