Chapter 3

1. Chapter 3 Atomic Structure and Electron Configuration

2. Atomic theory • Atomic theory states that all matter is composed of atoms. • The above theory was proposed by a Greek philosopher.

3. The 3 laws in chemistry • Law of definite proportions states that a given compound contains the same elements in exactly the same proportions by mass, regardless of the size of the sample or the source of the compound. • Ex. If a sample of ethylene glycol is found to have the formula C2H6 O2 , then the law of definite proportion tells you that all other samples will have the same molecular formula.

4. Law Of Conservation of mass • This law states that the mass of the products of a reaction equals the mass of the reactants. This law applies when two or more elements combine to produce a compound when the compound decomposes or when the atoms in a compound are rearranged.

5. Law of conservation of mass

6. Law of multiple proportion • This law applies to different compounds formed from the same two elements. • Ex there are 3 different compounds formed from the elements nitrogen and oxygen. All 3 are gases, but each has its own physical and chemical property.

7. Law of multiple proportions

8. Dalton’s atomic theory 1. All matter is made of indivisible and indestructible atoms. 2. All atoms of given element are identical in their chemical and physical property.

9. 3. Atoms of different elements differ in their chemical and physical properties. • 4. Atoms of different elements combine in simple whole-number ratios to form compounds. • 5. Chemical reactions consist of the combination, separation, or rearrangement of atoms.

10. Atomic mass • The mass of an atom in atomic mass units (amu). • Carbon-12 is the standard for the atomic mass scale. 1amu=1/12 th the mass of carbon-12. • Units such as milligrams, grams, or kilograms created extremely small numbers when measuring the masses of atoms or molecules.

11. What is mole? • A unit that serves as a bridge between the indivisible atom and the macroscopic world of materials and objects. It is a collection of Avogadro’s number of particles. • 1 mole= 6.022 x 1023 particles.

12. Internal structure of atoms • Subatomic particles are protons, neutrons, and electrons. • J.J Thompson’s theory led to the discovery of electrons. He wanted to study the flow of electric current. • He proposed that the electrons of an atom were embedded in a positively charged ball of matter. His picture of an atom was named the “plum pudding” model of an atom because it resembled plum pudding a desert consisting of a ball of sweet cake and pieces of fruit embedded in it.

13. Rutherford’s discovery of the nucleus: Rutherford's atom resembled a tiny solar system with the positively charged nucleus always at the center and the electrons revolving around the nucleus. • Interpreting Rutherford's Gold Foil Experiment • Alpha particles are positively charged partices produced by some nuclear disintegration

14. Protons, neutrons, and electrons • Protons are the subatomic particles that were present in the nucleus and having a positive charge.

15. Atomic number and mass number • The number of protons in the nucleus is atomic number. • The total number of particles in the nucleus that is the number of protons and neutrons is the mass number. • Ex. 115B

16. Isotopes and radioisotopes • Atoms of the same atom having different neutrons.

17. Radioisotopes • Atoms having unstable nuclear configuration.

18. Electrons and Light • Electrons are the most important parts of an atom. Electrons occupy most of an atom’s volume and determine virtually most of its chemistry. Our knowledge of electrons in atoms come from studying the light it emits.

19. Electromagnetic Spectrum • The EMR is composed of radiation with a broad range of spectrum. The visible spectrum is a tiny portion of the EM spectrum. Light is an electro magnetic wave.

20. Light and all other electromagnetic radiation are considered as moving waves. • The frequency (number of waves that pass a fixed place in a given amount of time ) and wavelength of a wave are inversely related. As frequency increases wavelength decreases.

21. Frequency= speed/wavelength • Frequency x wavelength = speed. • Excited electrons emit light. • Each electron in an atom is in a state of lowest possible energy a ground state. If an electron acquires additional energy then it is in excited state. • The electron will quickly fall back to the ground state and when it does the excess energy is released as light.

22. When electrons are excited they absorb energy and move to a higher energy level (gold arrow). When they emit light, they move to a lower energy level. Violet light (violet arrow) is produced when electrons move from level n=6 to n=2 Bohr’s Atomic Structure for hydrogen atom

23. Blue light (blue arrow) is produced when electrons move from level n=5 to level n=2. Green light (green arrow) is produced when electrons move from level n= 4 to level n= 2. Red light (red arrow) is produced when electrons move from level n= 3 to level n=2. This series of lines which is the visible hydrogen line spectrum is called the Balmer series. The energies of these emissions just happen to be in the visible range so we can see the colors.

24. Homework • Page 105 • Term review all.

25. Bohr’s Atomic Structure • n is also called as the quantum number. • Bohr’s atomic model suggested electrons in terms of their energy state. He postulated that electrons did not radiate energy while in orbit around the nucleus. • The present day quantum model suggests that electrons have both the properties of particles and waves.

26. Quantum Theory • This theory called as quantum theory suggests each electron in an atom is assigned 3 quantum numbers n, l, m. • Ex. If you have been in a concert your ticket specifies your seat by a series of numbers and letters. Like specifies seat number 20 in row K in section 3 of the south set of stands.

27. Figure showing the different energy levels around the nucleus.

28. In quantum theory electrons are located in orbitals ( a region in an atom where there is a high probability of finding one or more electrons) .The orbital is designated by a particular set of values of the quantum numbers n, l, and m.

29. Rules for assigning quantum numbers • The principle quantum number n can take the values 1,2,3, 4 and so on. N values larger than 7 are not encountered. The larger the value of n the farther the orbital is from the nucleus and higher its energy is. • The l quantum number can take any whole number value from 0 to n-1. ex. If n=3 l can have values 0,1, or 2. The m quantum number can take whole number values depending on the value of l. ex. If l=1 m can take values -1,0,1

30. Relation of n and l. • n=principle quantum number in which the level the electron is located. • L=sublevel within that energy level.

31. m= orbital within sublevel the electron is located.

32. Pauli’s Exclusion principles • This principle states that not more than two electrons can occupy a single orbital.

33. Electrons are assigned a spin quantum number ms = -1/2 and ms = +1/2. There are only two possible values for the spin quantum number. Two electrons in the same orbital spin in opposite directions

34. Aufbau’s principle • The electrons in an atom will occupy the lowest available orbital.

35. Hund’s Rule • Orbitals of the same n and l quantum numbers are each occupied by one electron before any pairing occurs.

36. Home work • Page 107 • #21 to 27. • Page 109 • Test prep all