CHAPTER 5 The wave nature of light Electromagnetic radiation: is a form of energy that exhibits wavelike behavior as it travels through space. Examples: Visible light X – rays Gamma rays =m4t7gTmBK3g

Electromagnetic spectrum

Transverse wave

Short wavelength Long Wavelength

c = λν All electromagentic waves travel at the speed of light C = 2.99 x 10 8 m/s (3.0 x 10 8 m/s) (sound only travels at 340 m/s)

c = λν All electromagnetic waves travel at the speed of light C = 2.99 x 10 8 m/s (3.0 x 10 8 m/s) (6.7 x 10 8 miles per hour) (Sound only travels at 340 m/s) Wavelength is in meters Frequency is in hz (1/sec)

c = λν 1. What is the frequency of green light, which has a wavelength of 4.90 x m? 2. An X-ray has a wavelength of 1.15 x m. What is its frequency? 3. Z100 broadcasts at a frequency of MHz. What is the wavelength of the broadcast? (HINT: UNITS)

c = λν 4. A compound emits blue light at a wavelength of 4.50 x cm. What is the frequency of the light? 5. What is the frequency of a wave if its wavelength is 3.6 nm?

6. Calculate the wavelength of an electromagnetic wave that has a frequency of 1.50 x Hz. (c = λν) 7. Is the wavelength longer or shorter than red light? c = λν

E = hv Energy of a photon (J) Plank’s Constant x Js Frequency (Hz) Energy and frequency have a direct relationship High energy = more violet light

E = hv h = x Js An atom releases blue light with a frequency of 6.8 x Hz. What is the energy of the photon emitted?

1.Calculate the wavelength of an electromagnetic wave that has a frequency of 1.50 x Hz. (c = λν) 2.Calculate the energy of the same wave c = λν E = hv

Review: Relationships: Direct or inverse? Wavelength and frequency Frequency and energy Wavelength and energy

Emission spectrum Incomplete spectra that are characteristic of a substance (like atom fingerprints)

Bohr’s model with the hydrogen atom

Bohr Model Electrons are in ‘shells’ Each electron ‘shell’ is quantized (the larger the radius, the larger the energy) Principal Energy levels (n) 1  7

Ground state vs. Excited State Electrons gain energy and jump to higher energy levels = Excited State The excited state is very unstable and the electrons jump back to their previous energy levels releasing a photon (a bundle of energy : LIGHT ENERGY!)

Filling order of the first 20 elements: 2, 8, 8, 2

Filling order: For the first 20 elements: 2,8,8,2 For elements 21 – 57: 2,8,18,18,32 For elements : 2,8,18,32,32

Valence e- for representative elements

Organization of the periodic table: Columns = groups Rows = periods A = representative B = transition Inner transition

Organization of the periodic table Rrepresentative elements families: chemical and physical properties are similar 1: Alkali metals –soft metals, very reactive 2: Alkaline metals – reactive metals 15: Pnictogens – Nitrogen column 16: Chalcogens – Oxygen column 17: Halogens – very reactive nonmetals Fluorine: Very reactive and poisonous Chlorine: Yellow – green gas Bromine: Redish-brown liquid 18: Noble gases – very UNreactive gases

Quantum Mechanical Model – ‘electron cloud’ Based on the quantum theory, which says matter has wave-like properties According to quantum theory, it’s impossible to know the exact position of an electron (Heisenberg’s Uncertainty Principle) LOE9eY4

Quantum Mechanical model – ‘electron cloud’ (cont) The model uses complex shapes called orbitals (volumes of space in which there is likely to be an electron) w.youtube.com/watc h?v=cPDp tc0wUYI

Principal energy level (n) Sublevels ________ ________

Sublevels ___S___ ___P__ __D___ __F__ # orbitals _____ _____ _____ _____ Orbital = the most probable location for an electron (90%)

Shapes of s, p, and d- Orbitals s orbital p orbitals d orbitals

Each orbital can hold _____ electrons # of electrons all together: _____ _____ _____ _____

Ground State - Electron configuration “Address” for an atoms electrons: nL e Aufbau principle: electrons occupy the lowest energy orbit available in each principal energy level first Every principal energy level adds 1 new sublevel Total number of electrons possible on each sublevel = 2n 2

s-block1st Period 1s 1 1 st column of s-block Periodic Patterns Example - Hydrogen Courtesy Christy Johannesson

Energy levels Overlap 4s  3d  4p 6s  4f  5d  6p

Valence Electrons : The valence electrons are the electrons in the outermost principal energy level. (Last energy level on the bohr model) They are always ‘s’ or ‘s and p’ orbital electrons There can be no more than 8 valence electrons Other electrons are called ‘kernel’ electrons Carbon = 1s 2 2s 2 2p 2 Kernel/core Valence

s-block [Ar]4s 2 2 nd column Shorthand / Noble gas notation Example -Calcium Courtesy Christy Johannesson 4 th Period Noble gas that precedes element

Orbital diagrams Uses boxes and arrows to represent electrons Aufbau principle: electrons occupy the lowest energy orbit available in each principal energy level first

Orbital diagrams Uses boxes and arrows to represent electrons Pauli Exclusion principle: each electron in an orbital must have a different spin (one up, one down) Hund’s rule: Electrons repel each other so they must be distributed: single electrons occupy each orbital before electron pairs can be made