1. Unit 5 - Atomic Theory and Chemical Bonding Emission of Energy by Atoms (pg 284) Energy Levels of Hydrogen (pgs 285-287) Hydrogen Orbitals (pgs 289-292) Electron Arrangements (pgs 295-298) Electron Configuration (pgs 299-303) Types of Chemical Bonds (pgs 317-319) Electronegativity (pgs 319-321) Stable Electron Configurations (pgs 323-326) Lewis Structures (pgs 328-332)

2. Unit 5 - Atomic Theory and Bonding Upon completion of this unit, you should be able to do the following: 1. Write the electron configuration for an atom or ion using the periodic table. 2.Be able to draw an orbital notation diagram for any atom or ion. 3.Determine the valence electrons of an atom. 4.Predict the types of bonds formed between two atoms and describe the properties of each. 5.Use electronegativity to predict the percent ionic character of bonds and the polarity of molecules. 6.Draw Lewis dot structures to represent how atoms share or transfer valence electrons to become more stable. 7.Explain how multiple bonds can form between the same two atoms. 2

3. Atomic Theory The concept of atoms explains many important observations, such as why compounds always have the same composition (a specific compound always contains the same types and numbers of atoms) and how chemical reactions occur (they involve a rearrangement of atoms). We learned to picture the atom as a positively charged nucleus composed of protons and neutrons at its center and electrons moving around the nucleus in a space very large compared to the nucleus. In this unit, we develop a picture of the electron arrangements in atoms. 3

4. Emission of Energy by Atoms 4 When compounds are heated, they emit a color characteristic of the cation. Li +, for example, emits a red flame when heated. Na + emits a yellow flame, Cu 2+ a green flame. The colors of the flames result from atoms releasing energy in the form of visible light of specific wavelengths, or colors. The heat from the flame causes the atom to absorb energy. The atom becomes excited. Some of the excess energy is released as light. The atom moves to a lower energy state as it emits a photon of light.

5. Emission of Energy by Atoms 5 When atoms receive energy, they become excited. They can release the energy by emitting light. The emitted energy is carried away by a photon. The energy of the photon corresponds exactly to the energy change of the emitting atom. High energy photons correspond to short wavelength light. Low energy photons correspond to long wavelength light. The photons of red light have less energy than the photons of blue light because red light has a longer wavelength than blue light.

6. Energy Levels of Hydrogen 6 When we study the photons of visible light emitted, we see only certain colors. Only certain types of photons are produced. Because only certain photons are emitted, only certain energy changes are occurring. So, hydrogen atoms must have certain discrete energy levels. We say the energy levels of hydrogen are quantized, that is, only certain values are allowed. Energy levels of all atoms are quantized.

7. Energy Levels of Hydrogen 7

8. The Hydrogen Orbitals 8 The probability map is called an orbital. The orbital shown in Figure 10.20 is called the 1s orbital and describes the ground (lowest) state of energy for hydrogen.

9. The Hydrogen Orbitals 9 Hydrogen has discrete energy levels. They are called principal energy levels and labeled with an integer. Each principal energy level has sublevels.

10. The Hydrogen Orbitals 10 Principal level 2 has 2 sublevels. They are called 2s and 2p. Principal level 3 has 3 sublevels called 3s, 3p and 3d. Principal level 4 has 4 sublevels called 4s, 4p, 4d and 4f.

11. The Hydrogen Orbitals 11 The principal levels describe size and shape. The s orbital is spherical. Level 1 is smaller than level 2, which is smaller than level 3.

12. The Hydrogen Orbitals 12 The three 2p orbitals are lobed, not spherical. They are oriented along the x, y or z axis.

13. 13

14. Electron Arrangements 14 An atom has as many electrons as it does protons, so all atoms beyond hydrogen have more than one electron. Each electron appears to spin like a top on its axis. It can only spin in one direction. We represent spin with an arrow, ↑ or ↓. Electrons in the same orbital must have opposite spins. This leads to the Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons and those two electrons must have opposite spins.

15. Electron Arrangements 15 Hydrogen has an atomic number of 1 (Z =1) and therefore a single electron to have a net charge of zero. To show its electron configuration, we write the principal energy level followed by the sublevel, 1s. The number of electron in the orbital is placed as a superscript, 1s 1. The electron configuration can also be shown using an orbital diagram, or box diagram, as below.

16. Electron Configuration 16 Hydrogen (Z=1)1s 1 Helium (Z=2)1s 2 Lithium (Z=3) 1s 2 2s 1 Berylium (Z=4) Boron (Z=5) 1s 2 2s 2 2p 1 Carbon (Z=6) Nitrogen (Z=7) Oxygen (Z=8) Fluorine (Z=9) Neon (Z=10) 1s 2 2s 2 2p 6 The orbital diagram for nitrogen is below.

17. Electron Configuration 17 Sodium (Z=11) 1s 2 2s 2 2p 6 3s 1 or [Ne] 3s 1 Magnesium (Z=12) Aluminum (Z=13) Silicon (Z=14) Phosphorous (Z=15) Sulfur (Z=16) Chlorine (Z=17) Argon (Z=18)

18. Electron Configuration 18 Valence electrons are the electrons in the outermost (highest) principal energy level of an atom. These are the electrons involved in bonding of atoms to each other. Also note that the atoms of elements in the same group have the same number of electrons in a given type of orbital, except that the orbitals are in different principal energy levels. Elements with the same valence electron arrangement show very similar chemical behavior.

19. Electron Configuration 19 The order of filling orbitals changes for Z=19. Experiments show that the chemical properties of potassium are very similar to lithium and sodium. We predict that the 4s orbital will fill before the 3d orbital. This means that principal energy level 4 begins to fill before level 3 is full. Potassium (Z=19)[Ar] 3s 2 3p 6 3d 1 Potassium (Z=19)[Ar] 3s 2 3p 6 4s 1 Calcium (Z=20)[Ar] 3s 2 3p 6 4s 2 Scandium (Z=21)[Ar] 3s 2 3p 6 4s 2 3d 1

20. Electron Configuration 20

21. Electron Configuration 21

22. Electron Configuration 22 Practice problems: Determine the electron configuration for: 1.C (Z= 6) 2.Al (Z=13) 3.Cl (Z=17)

23. Types of Chemical Bonds A bond is a force that holds two or more atoms together and makes them function as a unit. In water, the fundamental unit is the H-O-H molecule, which is held together by the two O-H bonds.

24. Types of Chemical Bonds Ionic compounds are formed when an atom that loses an electron relatively easily reacts with an atom that accepts an electron. This occurs when a metal reacts with a non-metal. The resulting bonds are called ionic bonds. In an ionic bond, electrons are transferred.

25. Types of Chemical Bonds Consider diatomic hydrogen H – H. When two hydrogen atoms are brought close together, the electrons are equally attracted to both nuclei. When two similar atoms form a bond, the electrons are equally attracted to the nuclei of the two atoms. This is called a covalent bond. In a covalent bond, the electrons are shared by nuclei.

26. Types of Chemical Bonds Ionic bonding and covalent bonding are extremes. Between the extremes are cases where atoms are not so different that electrons are transferred, but different enough that unequal sharing of the electrons results. These bonds are called polar covalent bonds. In HF, the fluorine atom has a stronger attraction for the shared electrons than the hydrogen atom does.

27. Types of Chemical Bonds 27

28. Electronegativity The unequal sharing of electrons between two atoms is described by a property called electronegativity, the relative ability of an atom in a molecule to attract shared electrons to itself. The higher the atoms electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond. Fluorine has the highest electronegativity value at 4.0. Cesium and Francium have the lowest electronegativity value at 0.7

29. 29 Electronegativity

30. The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bonds. If the atoms have similar electronegativities, the electrons are shared almost equally and the bond shows little polarity. If the atoms have very different electronegativities, a very polar bond is formed.

31. Electronegativity In extreme cases, one or more electrons are actually transferred and ions and an ionic bond are formed. Consider NaCl, for example. When a Group 1 metal reacts with a Group 17 element, ions are formed and an ionic substance results.

32. Electronegativity Page 321, example 11.1 Using the electronegativity values given in Figure 11.3, arrange the following bonds in order of increasing polarity: H-H, O-H, Cl-H, S-H, F-H

33. Stable Electron Configurations 33 Representative metals form ions by losing enough electrons to attain the configuration of the previous noble gas that occurs before the metal. For example, sodium will lose one electron to attain the configuration of neon. Nonmetals form ions by gaining enough electrons to attain the configuration of the next noble gas. For example, chlorine will add one electron to attain the configuration of argon.

34. Stable Electron Configurations 34

35. Stable Electron Configurations 35 In observing millions of stable compounds, chemists have observed that in almost all stable compounds, all of the atoms have achieved a noble gas configuration.

36. Stable Electron Configurations 36 When a non-metal and a Group 1, 2 or 3 metal react to form a binary ionic bond, the ions form so that the non-metal completes the valence-electron configuration of the next noble gas and the metal empties the valence orbitals to achieve the configuration of the previous noble gas.

37. Stable Electron Configurations 37 When two non-metals react to form a covalent bond, they share electrons in a way that completes the valence-electron configuration of both atoms.

38. Lewis Structures 38 Bonding involves just the valence electrons of atoms. (See slide 18.) The Lewis structure is a representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule. Dots are used to represent the valence electrons. Examples: KBr

39. Lewis Structures 39 Duet rule – hydrogen Octet rule – elements are surrounded by 8 electrons. Electrons that are shared form bonds.

40. Lewis Structures 40 Practice problems: Draw the Lewis structure for: 1.HF 2.NH 3 3.CH 4