Presentation is loading. Please wait.

Presentation is loading. Please wait.

Chapter 13 Electrons in Atoms. Section 13.1 Models of the Atom OBJECTIVES: l Summarize the development of atomic theory.

Similar presentations


Presentation on theme: "Chapter 13 Electrons in Atoms. Section 13.1 Models of the Atom OBJECTIVES: l Summarize the development of atomic theory."— Presentation transcript:

1 Chapter 13 Electrons in Atoms

2 Section 13.1 Models of the Atom OBJECTIVES: l Summarize the development of atomic theory.

3 Models of the Atom OBJECTIVES: l Explain the significance of quantized energies of electrons as they relate to the quantum mechanical model of the atom.

4 Greek Idea l Democritus and Leucippus l Matter is made up of solid indivisible particles l John Dalton - one type of atom for each element

5 J. J. Thomson’s Model l Discovered electrons l Atoms were made of positive stuff l Negative electron floating around l “Plum-Pudding” model

6 Ernest Rutherford’s Model l Discovered dense positive piece at the center of the atom- nucleus l Electrons would surround it l Mostly empty space l “Nuclear model”

7 Niels Bohr’s Model l He had a question: Why don’t the electrons fall into the nucleus? l Move like planets around the sun. l In circular orbits at different levels. l Amounts of energy separate one level from another. l “Planetary model”

8 Bohr’s planetary model l Energy level of an electron l analogous to the rungs of a ladder l electron cannot exist between energy levels, just like you can’t stand between rungs on ladder l Quantum of energy required to move to the next highest level

9 The Quantum Mechanical Model l Energy is quantized. It comes in chunks. l A quanta is the amount of energy needed to move from one energy level to another. l Since the energy of an atom is never “in between” there must be a quantum leap in energy. l Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

10 l Things that are very small behave differently from things big enough to see. l The quantum mechanical model is a mathematical solution l It is not like anything you can see. The Quantum Mechanical Model

11 l Has energy levels for electrons. l Orbits are not circular. l It can only tell us the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model

12 l The atom is found inside a blurry “electron cloud” l A area where there is a chance of finding an electron. l Draw a line at 90 % l Think of fan blades The Quantum Mechanical Model

13 Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron. l Within each energy level, the complex math of Schroedinger’s equation describes several shapes. l These are called atomic orbitals - regions where there is a high probability of finding an electron.

14 Summary s p d f # of shapes Max electrons Starts at energy level

15 By Energy Level l First Energy Level l only s orbital l only 2 electrons l 1s 2 l Second Energy Level l s and p orbitals are available l 2 in s, 6 in p l 2s 2 2p 6 l 8 total electrons

16 By Energy Level l Third energy level l s, p, and d orbitals l 2 in s, 6 in p, and 10 in d l 3s 2 3p 6 3d 10 l 18 total electrons l Fourth energy level l s,p,d, and f orbitals l 2 in s, 6 in p, 10 in d, and 14 in f l 4s 2 4p 6 4d 10 4f 14 l 32 total electrons

17 By Energy Level l Any more than the fourth and not all the orbitals will fill up. l You simply run out of electrons l The orbitals do not fill up in a neat order. l The energy levels overlap l Lowest energy fill first.

18 Electron Arrangement in Atoms l OBJECTIVES: Apply the aufbau principle Pauli’s exclusion principle Hund’s rule For electron configurations.

19 Section 13.2 Electron Arrangement in Atoms l OBJECTIVES: - Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.

20 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Aufbau diagram

21 Electron Configurations l The way electrons are arranged in atoms. l Aufbau principle- electrons enter the lowest energy first. l This causes difficulties because of the overlap of orbitals of different energies. l Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

22 Electron Configuration l Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to. l Let’s determine the electron configuration for Phosphorus l Need to account for 15 electrons

23 l The first two electrons go into the 1s orbital l Notice the opposite spins l only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

24 l The next electrons go into the 2s orbital l only 11 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

25 The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

26 The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

27 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3

28 The easy way to remember 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons

29 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 4 electrons

30 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 12 electrons

31 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 20 electrons

32 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 38 electrons

33 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 56 electrons

34 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 88 electrons

35 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p electrons

36 Exceptional Electron Configurations

37 Orbitals fill in order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. l Half filled orbitals have a lower energy. l Makes them more stable. l Changes the filling order

38 Write these electron configurations l Titanium - 22 electrons - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 l Vanadium - 23 electrons - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 l Chromium - 24 electrons - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 expected - But this is wrong!!

39 Chromium is actually: l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 l Why? l This gives us two half filled orbitals. l Slightly lower in energy. l The same principal applies to copper.

40 Copper’s electron configuration l Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 l But the actual configuration is: l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 l This gives one filled orbital and one half filled orbital. l Remember these exceptions: d 4, d 9


Download ppt "Chapter 13 Electrons in Atoms. Section 13.1 Models of the Atom OBJECTIVES: l Summarize the development of atomic theory."

Similar presentations


Ads by Google