LIMITATIONS OF RUTHERFORD S ATOMIC MODEL

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1 ELECTRONS IN ATOMS

2 LIMITATIONS OF RUTHERFORD S ATOMIC MODEL Did not explain the chemical properties of atoms For example, it could not explain why metals or compounds of metals give off characteristic colors when heated in a flame It could not explain why objects might change color when heated to higher temperatures

3 T H E H Y D R O G E N AT O M A N D B O H R

4 THE BOHR MODEL In 1913 Niels Bohr developed a new atomic model Experiment: Tested Hydrogen atoms Conclusion: Bohr proposed electrons orbit around the nucleus in fixed energies

5 ENERGY LEVELS Fixed energy levels of an electron are similar to rungs of a ladder Electrons cannot exist between energy levels To move from one energy level to the next, an atom must gain or lose the correct amount of energy A quantum of energy is the amount of energy required to move an electron from one energy level to the next The amount of energy an electron gains or loses is not always the same Higher energy levels are closer together It takes less energy to move from one rung to the next near the top of the ladder The higher the energy level occupied by an electron, the less energy it takes the electron to move from that energy level to the next higher energy level

6 THE HYDROGEN ATOM Excited State (based off electron placement around nucleus) Bohr proposed that an electron moves into an orbit or higher energy level further from the nucleus when an atom absorbs energy Ground State Electron returns here (home) after being excited

7 THE HYDROGEN SPECTRUM

8 T H E Q U A N T U M M E C H A N I C A L M O D E L A N D S C H R O D I N G E R

9 THE QUANTUM MECHANICAL MODEL Starting Point: Both Bohr and Rutherford s model of the atom described the path of a moving electron Experiment: Austrian physicist Erwin Schrodinger used calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom Conclusion: There is a probability that describes how likely it is to find an electron in a particular location around the nucleus of an atom Location is described as an electron cloud that is dense and the probability of finding an electron there is high For each energy level, the Schrodinger equation also leads to a mathematical expression called an atomic orbital Model: Devised from the mathematical solutions to the Schrodinger equation which is the modern description of the electrons

10 T H E Q U A N T U M M E C H A N I C A L M O D E L D E T E R M I N E S T H E A L L OW E D E N E R G I E S A N E L E C T R O N C A N H AV E A N D H OW L I K E LY I T I S TO F I N D T H E E L E C T R O N I N VA R I O U S L O C AT I O N S A R O U N D T H E N U C L E U S The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller.

11 ATOMIC ORBITALS (CLOUDS) An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. The probability of finding an electron within the electron cloud is 90%

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13 Like the Bohr model, the quantum mechanical model restricts the energy of electrons to certain values. Unlike the Bohr model, the quantum mechanical model does not specify an exact path the electron takes around the nucleus.

14 S, P, D, A N D F O R B I TA L S

15 S - O R B I TA L S-orbitals are spherically shaped Smaller atoms have fewer electrons and take up less space Larger atoms have more electrons and take up more space

16 p-orbitals are dumbell shaped. z-axis

17 p-orbitals are dumbell shaped. x-axis

18 p-orbitals are dumbell shaped. y-axis

19 p-orbitals together x, y, & z axes.

20 ATOMIC ORBITALS Different atomic orbitals are denoted by letters. Four of the five d orbitals have the same shape, but different orientations in space.

21 ATOMIC ORBITALS Describe the probability of finding an electron at various locations around the nucleus s orbitals: groups 1 and 2 on the periodic table p orbitals: groups d orbitals: groups 3-12 f orbitals: lanthanide and actinide series

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23 ORBITAL SUBLEVELS Each energy sublevel (subshells) correspond to one or more orbitals of different shapes. The orbitals describe where an electron is likely to be found.

24 Summary of Principal Energy Levels and Sublevels Principal energy level Number of sublevels Type of sublevel Maximum number of electrons n = 1 1 1s (1 orbital) 2 n = 2 2 2s (1 orbital), 2p (3 orbitals) 8 n = 3 3 n = 4 4 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals) 4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals) 18 32

25 SUBLEVELS The Principal Quantum Number, n, always equals the number of sublevels within that principal energy level The number of orbitals in a principal energy level is equal to n 2 A maximum of two electrons can occupy an orbital Therefore, the maximum number of electrons that can occupy a principal energy level is given by the formula 2n 2

26 VA L E N C E E L E C T R O N S

27 VA L E N C E E L E C T R O N S The electrons in the outermost, furthest from the nucleus, electron shell are called valence electrons The number of valence electrons in orbitals s and p (not transition metals) is the same as the group number The number of electron shells with electrons in them is the same as the period number

28 NOBLE GAS STABILITY Noble gases are usually unreactive This is because they have a full valence shell For two atoms to bond, they must gain, lose, or share electrons Metals tend to lose electrons Non-metals tend to gain electrons

29 ELECTRON CONFIGURATION

30 BLOCK TABLE The periodic table shows the different blocks located on the periodic table It also shows the electron configuration order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

31 E L E C T R O N P L A C E M E N T A N D T H E R U L E S T H AT F O L L O W

32 AUFBAU PRINCIPLE Electrons are placed in the lowest energy levels first PAULI EXCLUSION PRINCIPLE Only 2 electrons can be held in an orbital, different than an electron energy shell, and they must have opposite spins HUND S RULE Every orbital within a sublevel gets an electron before any gets paired

33 ELECTRON CONFIGURATIONS The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel The sublevel is written followed by a superscript with the number of electrons in the sublevel If the 2p sublevel contains 2 electrons, it is written: Number of electrons Energy level 2p 2 Energy sublevel

34 ELECTRONS IN SUBLEVELS REVIEW s-orbital 1 orbital, 2 electrons p-orbital 3 orbitals, 6 electrons d-orbital 5 orbitals, 10 electrons

35 WRITING ELECTRON CONFIGURATIONS First, determine how many electrons are in the atom For example, Iron has 26 electrons Arrange the energy sublevels according to increasing energy 1s 2s 2p 3s 3p 4s 3d Fill each sublevel with electrons until you have used all the electrons in the atom Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 The sum of the superscripts equals the atomic number of iron (26)

36 ELECTRON CONFIGURATION PRACTICE Write a ground state electron configuration of a neutral atom: K: Ne:

37 A SHORTCUT! 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 6f 7s 7p 7d 7f Do not exist in normal ground state atoms 6g 6h 7g 7h 7i

38 MORE PRACTICE! Write a ground state electron configuration of a neutral atom using the shortcut: Cl: Rb:

39 N O B L E G A S C O N F I G U R AT I O N

40 NOBLE GAS CONFIGURATION The Noble Gases are: He, Ne, Ar, Kr, Xe, Rn Notice that each noble gas finishes a row, or energy level Noble gas configurations take advantage of this by condensing what you have to write Example: He: 1s 2 Example: C: 1s 2 2s 2 2p 2 Noble Gas Configuration for C: [He] 2s 2 2p 2

41 MORE EXAMPLES The ground state configuration for Arsenic (As) is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Notice, that the part in purple is the same as Argon s configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 The noble gas configuration will state with the gas in the row before it [Ar] 4s 2 3d 10 4p 3

42 NOBLE GAS CORE ELECTRON CONFIGURATIONS Core Electrons: Electrons in [Noble Gas] Valence Electrons: Electrons outside of [Noble Gas] Recall, the electron configuration for Sodium (Na) is: Na: 1s 2 2s 2 2p 6 3s 1 We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas The preceding noble gas with an atomic number less that sodium is neon, Ne. We rewrite the electron configuration: Na: [Ne] 3s 1

43 NOBLE GAS CONFIGURATION PRACTICE Write the noble gas configuration for the following neutral atoms: Cu: Sr:

44 O R B I TA L D I A G R A M S

45 THE AUFBAU PRINCIPLE Each electron occupies the lowest energy orbital All orbitals related to an energy level are of equal energy Example: The three 2p orbitals are the same energy level

46 ORBITAL FILLING DIAGRAM

47 PAULI EXCLUSION PRINCIPLE A maximum of two electrons may occupy a single orbital, but only if the electrons have opposite spins Spin: Electrons have an associated spin, either one way or the other These spins are called spin up and spin down In the example to the right: Box = orbital Arrow = electron

48 HUND S RULE Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals Example: Nitrogen 1s 2 2s 2 2p 3 1s 2 2s 2 2p 3 NOT 1s 2 2s 2 2p 3

49 THE ORDER OF THINGS Electrons fill up the empty orbitals before sharing orbitals

50 EXCEPTIONAL ELECTRON CONFIGURATIONS Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations. Exceptions to the aufbau principle are due to subtle electron-electron interactions in orbitals with very similar energies

51 ORBITAL DIAGRAM PRACTICE Draw the orbital diagram for the following neutral atoms: N: Al:

52 MORE PRACTICE! Ti: Mg: As: