2 Dalton’s Atomic Theory John Dalton ( ) had four theoriesAll elements are composed of submicroscopic indivisible particles called atomsAtoms of the same element are identical. The atoms of anyone element are different from those of any other elementAtoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compoundsChemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction
3 Thomson’s Atomic Model Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model
4 Thomson’s TheoryThomson stated: The atom had negatively charged electrons stuck into a lump of positively charged protons.Thomson never explainedNumber of protons and neutronsThe arrangement of the particles in the atomThe ease with which atoms are stripped of electrons to form ions
5 Rutherford Model Rutherford used the Gold Foil Experiment Rutherford proposed the following:Thomson model was incorrectMost of the mass of the atom and all of its positive charge reside in a very small, extremely dense region, which he called the nucleusMost of the total volume of the atom is empty space in which electrons move around the nucleus
6 Rutherford’s ModelDiscovered dense positive piece at the center of the atomNucleusElectrons moved aroundMostly empty space
7 Bohr ModelBohr changed the Rutherford model and explained how the electrons travel.Bohr explained the following in his model:Electrons travel in definite orbits around the nucleusElectrons are arranged in concentric circular paths or orbitals around the nucleusElectrons don’t fall into the nucleus because electrons in particular path have fixed energy and don’t lose energyHis model was patterned after the motion of the planets around the sun. It is often called the Planetary model.
10 Quantum TheoryBohr explained how electrons were moving via Quantum TheoryKey Terms:Energy Levels- Regions around the nucleus where the electron is likely movingQuantum- Amount of energy required to move an electron from one energy level to the nextQuantum Leap- Abrupt Change
11 } Bohr’s Model Fifth Fourth Increasing energy Third Second First Further away from the nucleus means more energy.There is no “in between” energyEnergy LevelsFifthFourthThirdIncreasing energySecondFirstNucleus
12 Bohr’s Model cont. Energy levels are not equally spaced. Energy levels more closely spaced further from the nucleusHigher energy level occupied by an electron, the more energetic that electron is.Amount of energy gained or lost by an electron is not always the same amount.
13 Bohr Model Cont. The Bohr Model did not account for: Emission spectra of atoms containing more than one electron.So comes along the next model:
14 The Quantum Mechanical Model Energy is quantized. It comes in chunks.A quanta is the amount of energy needed to move from one energy level to another.Since the energy of an atom is never “in between” there must be a quantum leap in energy.Schrodinger derived an equation that described the energy and position of the electrons in an atom
15 The Quantum Mechanical Model Things that are very small behave differently from things big enough to see.The quantum mechanical model is a mathematical solutionIt is not like anything you can see.
16 Schrödinger’s Equation The wave function is a F(x, y, z)Actually F(r,θ,φ)Solutions to the equation are called orbitals.These are not Bohr orbits.Each solution is tied to a certain energyThese are the energy levelsAnimation
18 The Quantum Mechanical Model Has energy levels for electrons.Orbits are not circular.It can only tell us the probability of finding an electron a certain distance from the nucleus.
19 The Quantum Mechanical Model The atom is found inside a blurry “electron cloud”A area where there is a chance of finding an electron.Draw a line at 90 %
20 Atomic OrbitalsThere are the region of space which there is a high probability of finding an electronWithin each energy level the complex math of Schrodinger’s equation describes several shapes.These are called atomic orbitalsQuantum Numbers- numbers that specify the properties of atomic orbitals and their electrons
21 Quantum Numbers There are 4 types of Quantum Numbers Principal – distance from the nucleusAngular Momentum- Orbital ShapeMagnetic- Orbital position with respect to the X, Y, & Z axes.Spin- Has only two values (+1/2 or –1/2) and is needed to specify 1 of 2 positional orientations of an electron
22 Principal Quantum Number Symbolized by the letter N, indicates the main energy levels surrounding the nucleusThere are 7 principal quantum numbersA.K.A. – ShellsValue of N is a whole number ex. 1,2,3 ect..Main Energy Level – N=1; closest to the nucleus or ground stateGround State- state of the lowest energy of the atom.As N increases, the distance from the nucleus increases and the energy increases
23 Angular Momentum Quantum Number Indicates the shape of the orbital.Within each main energy level beyond the first, orbitals with different shapes occupy different regionsA.K.A. – Sublevels or SubshellsThe number of sublevels = Value of the Principal Quantum Number
24 Magnetic Quantum Number Indicates the orientation of a orbital about the nucleusThere are 4 types of orbital orientationS OrbitalP OrbitalD OrbitalF Orbital
25 Spin Quantum NumberHas only two possible values: +1/2 or –1/2. These values indicate two possible states of an electron in an orbitalSpin Quantum # is significant because each single orbital can hold no more than two electrons, which must have opposite spin.
26 S orbitals 1 s orbital for every energy level Spherical shaped Each s orbital can hold 2 electronsCalled the 1s, 2s, 3s, etc.. orbitals.
27 P orbitals Start at the second energy level 3 different directions 3 different shapesEach can hold 2 electrons
32 Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6 d51037144f
33 By Energy Level First Energy Level only s orbital only 2 electrons 1s2 Second Energy Levels and p orbitals are available2 in s, 6 in p2s22p68 total electrons
34 By Energy Level Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d3s23p63d1018 total electronsFourth energy levels,p,d, and f orbitals2 in s, 6 in p, 10 in d, ahd 14 in f4s24p64d104f1432 total electrons
35 By Energy LevelAny more than the fourth and not all the orbitals will fill up.You simply run out of electronsThe orbitals do not fill up in a neat order.The energy levels overlapLowest energy fill first.
36 Question for You How many principal quantum numbers are there? What is the maximum number of electrons that can fill the 3rd energy level?How many orbitals are in the sublevel F?What is the total number of orbitals for the 3rd main energy level?
37 Electron Configuration The way electrons are arranged in atomsThere are three rules which help dictate how electrons are arranged in the atoms.Aufbau Principle- electrons occupy the orbitals of the lowest energy firstHund’s Rule- Orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a second electron. All electrons in a single occupied orbital must have the same spin.
38 Electron Configuration cont. Pauli Exclusion Principle- No two electrons may occupy any given orbital without having opposite spin. No two electrons in the same atom can have the same set of four quantum numbers.Let’s determine electron configuration.Let’s start with Phosphorus.Need to account for all 15 electrons
39 Electron Configurations Distribution of all electrons in an atomConsist ofNumber denoting the energy level
40 Electron Configurations Distribution of all electrons in an atomConsist ofNumber denoting the energy levelLetter denoting the type of orbital
41 Electron Configurations Distribution of all electrons in an atom.Consist ofNumber denoting the energy level.Letter denoting the type of orbital.Superscript denoting the number of electrons in those orbitals.
46 Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the energy of the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order
47 Write these electron configurations Titanium - 22 electrons1s22s22p63s23p64s23d2Vanadium - 23 electrons 1s22s22p63s23p64s23d3Chromium - 24 electrons1s22s22p63s23p64s23d4 is expectedBut this is wrong!!
48 Chromium is actually 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals.Slightly lower in energy.The same principal applies to copper.
49 Copper’s electron configuration Copper has 29 electrons so we expect1s22s22p63s23p64s23d9But the actual configuration is1s22s22p63s23p64s13d10This gives one filled orbital and one half filled orbital.Remember these exceptions
50 Shortcuts for Electron Configuration There are two short handed methods of writing the electron configuration.The 1st method is called the outer-level configuration. That tells you the outer-most configuration for that element.The 2nd method is called the Noble Gas Notation. This tells you the complete notation using Noble Gases.Let’s start with outer-level notation!!!
62 Writing Electron configurations the easy way Yes there is a shorthand
63 Electron Configurations repeat The shape of the periodic table is a representation of this repetition.When we get to the end of the column the outermost energy level is full.This is the basis for our shorthand.
64 The Shorthand Write the symbol of the noble gas before the element. Then the rest of the electrons.Aluminum - full configuration.1s22s22p63s23p1Ne is 1s22s22p6so Al is [Ne] 3s23p1
65 More examples Ge = 1s22s22p63s23p64s23d104p2 Ge = [Ar] 4s23d104p2 Hf=1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2Hf=[Xe]6s24f145d2
66 The Shorthand Again Sn- 50 electrons The noble gas before it is Kr Takes care of 36Next 5s2Then 4d10Finally 5p2[ Kr ]5s24d105p2
67 Quantum NumbersSolving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.Each orbital describes a spatial distribution of electron density.An orbital is described by a set of three quantum numbers.
68 The Wave-like Electron The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.Louis deBroglie
69 The Quantum Mechanical Model A totally new approachDe Broglie said matter could be like a wave.De Broglie said they were like standing waves.The vibrations of a stringed instrument
71 What’s possible?You can only have a standing wave if you have complete waves.There are only certain allowed waves.In the atom there are certain allowed waves called electrons.1925 Erwin Schroedinger described the wave function of the electronMuch math, but what is important are the solutions
72 Schrödinger’s Equation The wave function is a F(x, y, z)Actually F(r,θ,φ)Solutions to the equation are called orbitals.These are not Bohr orbits.Each solution is tied to a certain energyThese are the energy levelsAnimation
73 What does the wave Function mean? nothing.it is not possible to visually map it.The square of the function is the probability of finding an electron near a particular spot.best way to visualize it is by mapping the places where the electron is likely to be found.
80 Quantum Numbers There are many solutions to Schrödinger’s equation Each solution can be described with quantum numbers that describe some aspect of the solution.Principal quantum number (n) size and energy of an orbitalHas integer values >0