Atomic Structure and Bonding

1. Atomic Structure and Bonding Unit 2

2. Major Subatomic Particles Atoms are measured in picometers, 10-12 meters Hydrogen atom, 32 pm radius Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble Radius of the nucleus is on the order of 10-15 m Density within the atom is near 1014 g/cm3

3. Elemental Classification Atomic Number (Z) = number of protons (p+) in the nucleus Determines the type of atom Li atoms always have 3 protons in the nucleus, Hg always 80 Mass Number (A) = number of protons + neutrons [Sum of p+ and nº] Electrons have a negligible contribution to overall mass In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

4. E mass number A elemental symbol Z atomic number Nuclear Symbols Every element is given a corresponding symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and atomic number

5. ATOMIC NUMBER AND MASS NUMBER Mass Number He 4 the number of protons and neutrons in an atom 2 Atomic Number the number of protons in an atom Number of electrons = Number of protons in a neutral atom 5

6. 184 W 74 19 F 9 80 Br 35 Find the number of protons number of neutrons number of electrons atomic number mass number

7. Ions • Cation is a positively charged particle. Electrons have been removed from the element to form the + charge. ex: Na has 11 e-, Na+ has 10 e- • Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge. ex: F has 9 e-, F- has 10 e-

8. 1 2 3 H H H 1 1 1 Hydrogen-1 Hydrogen-2 Hydrogen-3 Isotopes Atoms of the same element can have different numbers of neutrons and therefore have different mass numbers The atoms of the same element that differ in the number of neutrons are called isotopes of that element When naming, write the mass number after the name of the element

9. Calculating Averages Average = (% as decimal) x (mass1) + (% as decimal) x (mass2) + (% as decimal) x (mass3) + … Problem: Silver has two naturally occurring isotopes, 107Ag with a mass of 106.90509 u and abundance of 51.84 % ,and 109Ag with a mass of 108.90476 u and abundance of 48.16 % What is the average atomic mass? Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u) = 107.87 amu

10. Average Atomic Masses • If not told otherwise, the mass of the isotope is the mass number in ‘u’ • The average atomic masses are not whole numbers because they are an average mass value • Remember, the atomic masses are the decimal numbers on the periodic table

11. More Practice Calculating Averages • Calculate the atomic mass of copper if copper has two isotopes • 69.1% has a mass of 62.93 amu • The rest (30.9%) has a mass of 64.93 amu • Magnesium has three isotopes • 78.99% magnesium 24 with a mass of 23.9850 amu • 10.00% magnesium 25 with a mass of 24.9858 amu • The rest magnesium 26 with a mass of 25.9826 amu • What is the atomic mass of magnesium?

12. Bohr • Proposed electrons (e-) orbit around the nucleus in circular paths • Said e- in a particular path have a fixed energy (energy levels) • e- can go from any energy level to another by gaining or losing a specific amount of energy = a “quantum of energy” • When e- absorbs a quantum of energy, it goes from it’s ground state (where it’s normally found) to an excited state • The excited state is at a higher energy level

13. Bohr postulated that: • Fixed energy related to the orbit • Electrons cannot exist between orbits • The higher the energy level, the further it is away from the nucleus • An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive) • Think of Noble gases

14. Atomic Line Emission Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of theLINE EMISSION SPECTRA of excited atoms. • Problem is that the model only works for Hydrogen Niels Bohr (1885-1962)

15. Spectrum of White Light

16. Spectrum of Excited Hydrogen Gas

17. Line Emission Spectra of Excited Atoms • Excited atoms emit light of only certain wavelengths • The wavelengths of emitted light depend on the element.

18. Drawback to Bohr • Bohr’s theory did not explain or show the shape or the path traveled by the electrons. • His theory could only explain hydrogen and not the more complex atoms

19. Energy level populations (Science10) • Electrons found per energy level of the atom. • The first energy level holds 2 electrons • The second energy level holds 8 electrons • The third energy level holds 18 electrons

20. Examples for group 1 • Li 2.1 • Na 2.8.1 • K 2.8.8.1

21. 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. • Schrödinger derived an equation that described the energy and position of the electrons in an atom – an ORBITAL

22. Orbits (Bohr) vsOrbitals (Quantum Mechanics) Bohr said electrons travel in an orbit – can predict exact location of electron at any point in time. Schrodinger used mathematics (calculus) to find the region in space where an electron will be found 90% of the time - this region is called an orbital. There are 4 main types of orbitals – s, p, d, and f.

23. Modern View of the Atom The modern view of the atom suggests that the atom is more like a cloud. Atomic orbitals around the nucleus define the places where electrons are most likely to be found. 23

24. sorbitals • 1 s orbital for every energy level 1s 2s 3s • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals

25. p orbitals • Start at the second energy level • 3 different directions • 3 different shapes • Each orbital can hold 2 electrons

26. The d sublevel contains 5 d orbitals • The d sublevel starts in the 3rd energy level • 5 different shapes (orbitals) • Each orbital can hold 2 electrons

27. The f sublevel has 7 f orbitals • The f sublevel starts in the fourth energy level • The f sublevel has seven different shapes (orbitals) • 2 electrons per orbital

28. Electron Configuration • We use e- configuration as a shorthand to show how e- are arranged around a nucleus • Example: Carbon is … 1s2 2s2 2p2

29. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to .

30. Summary

31. Electron Arrangement 1st Rule: The Aufbau Principle • e- fill orbitals of the lowest energy first • We can use the periodic table to help us!

32. The Diagonal Rule

33. Example #1 1 s • Oxygen 2 s 2 p 1s2 2s2 2p4

34. Example #2 1 s • Magnesium 2 s 2 p 3 s 1s2 2s2 2p6 3s2

35. Example #3 1 s • Iron 2 s 3d 2 p 3 s 3 p 4 s 1s2 2s2 2p6 3s2 3p6 4s2 3d6

36. Practice • Boron • Argon • Calcium • Iodine • Sodium • Zinc • Lead

37. Abbreviations • We can abbreviate electron configurations using the Noble Gases • Ex: Sulfur • 1s2 2s2 2p6 3s2 3p4 • [Ne] 3s2 3p4 • Ex: Lead • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2 • [Xe] 6s2 4f14 5d10 6p2

38. 2nd Rule: Pauli Exclusion Principle • Each orbital orientation can hold up to 2 e- • e- must have opposite spins (up/clockwise or down/counter clockwise) • Therefore: • s has up to 2 e- (1 orientation) • p has up to 6 e- (3 orientations) • d has up to 10 e- (5 orientations) • f has up to 14 e- (7 orientations) • We can use the 2nd rule to draw Orbital Diagrams

39. Example #1 • Oxygen: 1s2 2s2 2p4 1s 2s 2p

40. Example #2 • Magnesium: 1s2 2s2 2p6 3s2 1s 2s 2p3s

41. Example #3 • Iron 1s2 2s2 2p6 3s2 3p6 4s2 3d6

42. 3rd Rule: Hund’s Rule • e- will not pair up until each orbital orientation has 1 e- in it • The first e- in a pair will spin up, the second will spin down • Example: Oxygen is 1s2 2s2 2p4 1s 2s 2p

43. Orbital Notation • Orbital Notation shows us visually the arrangement and spin of electrons • Example: Carbon is 1s2 2s2 2p2 1s 2s 2p

44. Energy Level Diagrams • Energy Level Diagrams give us the same information as orbital diagrams, plus they show us the different energy levels of each orbital • Example: Carbon is 1s2 2s2 2p2 2p 2s 1s

45. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

46. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • Phosphorous, 15 e- to place • The first to electrons go into the 1s orbital • Notice the opposite spins • only 13 more

47. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more

48. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more

49. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more

50. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes • 3 unpaired electrons • 1s22s22p63s23p3