Presentation on theme: "Chapter 13 “Electrons in Atoms”"— Presentation transcript:
1Chapter 13 “Electrons in Atoms” Mr. Daniel Olympic High SchoolCredits: Stephen L. CottonCharles Page High School
2Section 13.3 Physics and the Quantum Mechanical Model OBJECTIVES:Describe the relationship between the wavelength and frequencyDistinguish between quantum mechanics and classical mechanics. of light.Identify the source of atomic emission spectra.Explain how the frequencies of emitted light are related to changes in electron energies.
3Light Visible light is a type of electromagnetic radiation. Electromagnetic radiation is a form of energy and includes many types: gamma rays, x-rays, radio waves, visible light…Speed of light (c) = 3.00 x 108 m/sAll electromagnetic radiation travels at this same rate when measured in a vacuum
4Parts of a WaveCrestWavelengthAmplitudeNodeTrough
5- Page 373“R O Y G B I V”Frequency IncreasesWavelength Shorter
6Long Wavelength Low Frequency Low ENERGY Short Wavelength Wavelength TableShortWavelengthHigh FrequencyHigh ENERGY
7Wavelength and Frequency: Are inversely relatedAs one increases the other decreases.Different frequencies of visible light are different colors.There is a wide variety of frequenciesSpectrum: A whole range of electromagnetic wavelengths. (e.g. the visible light spectrum)
9So what is Energy? Not all quanta (plural) are the same size. All energy is quantizedA quantum is a “packet” of energy.Not all quanta (plural) are the same size.(eggs are not all the same size either, but all are eggs)
10So what is Light Energy?Light is a form of energy.Therefore, light must be quantizedA quantum of light energy is called a photon.Einstein determined that light is not only a wave, but is also a particle!He demonstrated it in an experiment that showed the photoelectric effect
12Experiment demonstrates the particle nature of light. Photoelectric EffectExperiment demonstrates the particle nature of light.
13So what is Light Energy? (con’t) Therefore, light has what is called wave-particle duality. It has characteristics of both waves and particles.
14Wave-Particle Duality (again) J.J. Thomson won the Nobel prize for describing the electron as a particle.His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.The electron is a particle!The electron is an energy wave!
15Confused? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.”Physicist Sir Arthur EddingtonThe Nature of the Physical World1934
16The Physics of the Very Small Quantum mechanics explains how very small particles behaveQuantum mechanics is an explanation for subatomic particles and atoms as wavesClassical mechanics describes the motions of bodies much larger than atoms
17Section 13.1 Models of the Atom OBJECTIVES:Identify the inadequacies in the Rutherford atomic model.Identify the new proposal in the Bohr model of the atom.Describe the energies and positions of electrons according to the quantum mechanical model.Describe how the shapes of orbitals related to different sublevels differ.
18Ernest Rutherford’s Model Discovered dense positive piece at the center of the atom- “nucleus”Electrons would surround and move around the nucleusAtom is mostly empty spaceIt did not explain the chemical properties of the elements – a better description of the electron behavior was needed
19Niels Bohr’s ModelWhy don’t the electrons fall into the nucleus?He agreed with Rutherford that electrons move around the nucleus, But:In specific circular paths, or orbits (like planets around the sun), at specific energy levels.An amount of fixed energy separates one electron energy level from another.
20The Bohr Model of the Atom I pictured the electrons orbiting the nucleus much like planets orbiting the sun.However, electrons are found in specific energy levels around the nucleus, and can jump from one level to another.Niels Bohr
21Bohr’s Model Electrons occupy specific energy levels analogous to the rungs of a ladderThe electron cannot exist between energy levels, just like you can’t stand between rungs on a ladderA quantum of energy is the amount of energy required to move an electron from one energy level to another (plural: quanta)Since the energy of an atom is never “in between” there must be a quantum leap in energy.
22Changing the energyLet’s look at a hydrogen atom, with only one electron, and in the first energy level.
23Changing the energyHeat, electricity, or light can move the electron up to different energy levels. The electron is now said to be in an “excited state”
24Changing the energyThe electron is unstable at the higher energy level and as it falls back to the ground state, it gives the energy back in the form of light
25Changing the energy They fall down in specific steps Each step has a different energy (quantum) and results in a different color of light.
27Ultraviolet Visible Infrared The further electrons fall, the more energy is released and the higher the frequency of light emitted.This is a simplified explanation!Remember, the orbitals also have different sublevels within the principle energy levels
28Atomic SpectraWhite light is made up of all the colors of the visible spectrum.Passing it through a prism separates it.
29But not all light is white.. By heating a gas with electricity we can get it to give off colors.Passing this light through a prism does something different.
30Atomic SpectrumEach element gives off its own characteristic colors of light.The colors can be used to identify the atom.This is how we know what elements stars are made of.
31This is called a bright line spectrum Unique to each element, like fingerprints!Very useful for identifying elements
32Explanation of Atomic Spectra ground state - the lowest energy level of the electron.In summary. When an electron at ground state receives a quantum of energy it jumps directly to a higher energy level. The electron is unstable and immediately drops to a lower energy level. As it drops it gives off the same amount of
33The Quantum Mechanical Model Problems with Bohr’s theory :It was only successful for H- no other elements followed his predictions.It introduced the quantum idea artificially.
34Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.”You can find out where the electron is, but not its energyOR…You can know how much energy it has, but not where it is!Werner Heisenberg
35Heisenberg Uncertainty Principle It is impossible to know exactly the location and velocity of a particle simultaneously.The better we know one, the less we know the other.Measuring one property, changes the other.
36After Before Photon wavelength changes Photon Moving Electron Electron velocity changes
37In 1926, Erwin Schrodinger derived an equation that described the energy and probable position of the electrons in an atom
38Schrodinger’s Wave Equation His equation determined the probability of finding a single electron along a single axis (x-axis)Erwin SchrodingerErwin Schrodinger
39The Quantum Mechanical Model Particles which are very small and travel very quickly (like electrons) behave very differently from objects big enough to observe.The quantum mechanical model is a mathematical solution describing how those particles act.
40The Quantum Mechanical Model Describes energy levels for electrons.Electrons move in an unpredictable mannerWe can only determine the probability of finding an electron a certain distance from the nucleus.
41The Quantum Mechanical Model The electrons are probably located inside a blurry “electron cloud”The area where there is the greatest chance of finding an electron.
42Atomic OrbitalsPrincipal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc.Within each energy level, there are sub-levels (like theater seats arranged in sections): letters s, p, d, and fThe complex math of Schrodinger’s equation describes several shapesThese are the atomic orbitals - regions where there is a 90% probability of finding an electron.
43Principal Quantum Number The Principle Quantum Number (n) denotes the shell (energy level) in which the electron is located.The maximum number of electrons that fit into an energy level can be calculated:2n2
44Summary # of orbitals Max. electrons Starts at energy level s 1 2 1 p 362d51037144f
47By Energy Level First Energy Level Has only an s sublevel only 2 electrons1s2Second Energy LevelHas s and p sublevels2 e- in s, 6 e- in p2s22p68 total electrons
48By Energy Level Third energy level Has s, p, and d sublevels 2 e- in s, 6 e- in p, and 10 e- in d3s23p63d1018 total electronsFourth energy levelHas s, p, d, and f sublevels2 e- in s, 6 e- in p, 10 e- in d, and 14 e- in f4s24p64d104f1432 total electrons
49By Energy LevelBeyond the fourth energy level, not all sublevels fill up.You simply run out of electronsSo only the s, p, d and f sublevels are usedBecause the energy levels overlap the orbitals do not fill up in a consistent patternHowever, the lowest energy orbitals fill first.
50Section 13.2 Electron Arrangement in Atoms OBJECTIVES:Describe how to write the electron configuration for an atom.Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.
52Electron Configurations… …are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how:Aufbau principle - electrons enter the lowest energy sublevels first.This becomes complex because of the overlap of orbitals of different energies – follow the diagram!Pauli Exclusion Principle – there are at most 2 electrons per orbital - with opposite spins
53Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers.To show the different direction of spin, a pair in the same orbital is written as:Wolfgang Pauli
54Electron Configurations Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until each orbital has one electron.Let’s write the electron configuration for PhosphorusWe need to account for all 15 electrons in phosphorus
55Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The first two electrons go into the 1s orbitalNotice the opposite direction of the spinsonly 13 more to go...
56Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The next electrons go into the 2s orbitalonly 11 more...
57Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The next electrons go into the 2p orbitalonly 5 more...
58Increasing energy The next electrons go into the 3s orbital 2p3p4p5p6p3d4d5d7p6d4f5fThe next electrons go into the 3s orbitalonly 3 more...
59Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p The last three electrons go into the 3p sublevel.They each go into separate orbitals (Hund’s)3 unpaired electrons= 1s22s22p63s23p3
60Orbitals fill in an order: Lowest energy to higher energy.Adding electrons can change the energy of the orbital. Full sublevels are the most stable arrangement.Half filled sublevels have a lower energy than partially filled sublevels, and are next most stable.
61Write the electron configurations for these elements: Zirconium - 40 electrons[Kr] 5s2 4d2Tantalum - 73 electrons[Xe] 6s2 4f14 5d3Chromium - 24 electrons[Ar] 4s2 3d4 (expected)But this is not what happens with Chromium
62Chromium is actually: [Ar]4s13d5 Why? This gives us two half filled orbitals (the others are all still full)Half full is slightly lower in energy.The same principal applies to copper…
63Copper’s electron configuration Copper has 29 electrons so we expect: [Ar] 4s2 3d9But the actual configuration is: [Ar]4s13d10This change gives one more filled orbital and one that is half filled.Remember these exceptions: Groups ending in d4 and d9
64Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULLCopper steals a 4s electron to FILL its 3d sublevel