Chemistry

1. Chemistry Mr. McCalla “Mack” Rm 230

2. Class Rules • If your not in your seat by the second bell, you’re late…..! • All assignments must be handed in on time. Late submissions will NOT be accepted. • When I speak……YOUR MOUTH STOPS WORKING……..! • Expect nightly homework, a weekly quiz on Wednesdays and exams on Fridays.

3. What do I need for this class…? Out side of the usual supplies (pens, paper),you are required to have the following: • 3-ringed binder. • 2-3 folders for the 3-ringed binder. • A spiraled notebook with 3-ringed holes for the binder. *You will be graded on the contents of your binder so always make sure it is in order with notes, handouts, returned papers, quizzes and exams*

5. What is Chemistry? • Chemistry: the study of matter, which is anything that has mass and volume. • Your desk your sitting on, the air around you, and your body are all made up of matter. • Everything is composed of Atoms • Chemistry deals with the composition of matter and the changes that matter may undergo. • Water - Ice – Vapor

6. Early Studies of Atoms The Greeks were the first to take a closer look at the physical world around them. They concluded that is composed of fire, air, water and earth. A scientist by the name of Robert Boyle identified Gold and Silver as elemental; that is, they are not themselves made of fire, air water and earth causing the Greek concept to fade. Robert Boyle

7. Dalton’s Atomic Theory • The work of Boyle led a scientist by the name of John Dalton to propose his revolutionary theory in the 1700s. • He theorized that the basic unit of matter is a tiny particle called an atom. • Dalton’s theory of the atom can be summarized by the following points: John Dalton

8. Dalton’s Atomic Theory • All elements are composed of indivisible atoms. • All atoms of a given element are identical. • Atoms of different elements are different; that they have different masses. • Compounds are formed by the combination of atoms of different elements

9. Structure of the Atom • J.J. Thompson was one of the first scientist to identify the smaller units that made up an atom. • He used a cathode ray tube and gold foil to show one of these smaller units (electrons) that make up the atom J.J. Thompson • A concept of the atom developed in which these • negatively charged particles are visible like raisins in bread. • This model called the “plum model”; the mass of the rest • of the atom was distributed and positively charged, taking • space not used by the electrons.

10. Because the ray produced by in the tube was deflected by the gold foil a certain way by an electrical or magnetic field, he concluded that the ray was formed by negatively charged particles. • These negatively charged particles are called electrons.

11. Structure of the Atom • EarnestRutherford led a group of scientist that directed alpha particles (positive) at a piece of gold foil. • Most of the particles passed straight through the foil. • To their surprise some of the alpha particles where greatly deflected and some bounced back. Earnest Rutherfod

13. Structure of Atoms • Rutherford concluded that atoms have a dense central core, called the Nucleus, while the remainder of the atom is essential empty space. • Because alpha particles are positively charged and were repelled by the nucleus, the nucleus must be positively charged because like charges repel each other.

14. Modern Atomic Theory • The Bohr Model of the Atom: In the early 20th century the common model of the atom was the Bohr or “planetary model”. • The model showed a center nucleus (protons and neutrons) and rings with orbiting electrons. • The outer most electrons (electrons that are in the last orbital are called valence electrons.

16. Modern Atomic Theory • The Wave-Mechanical Model: Modern advancements in the study of energy aided in the modifying of the atomic model. • This model of the atom pictures the atom as having a dense, positively charged nucleus as proposed to the planetary model . • The major difference between the wave-mechanical model and the Bohr model is found in which manor the electrons are pictures.

17. Modern Atomic Theory • Instead of moving in definite, fixed orbits around a nucleus as suggested in the Bohr model, the wave-mechanical model portrays electrons with distinct amounts of energy moving in areas called orbitals. • Orbital: a region in which an electron of a particular amount of energy is most likely to be located .

19. The Wave Mechanical Model • Energy and matter were viewed as both waves and particles. • Energy was made up of tiny packets called quanta which acted like particles. • This model pictures the atom as having a dense, positively charged nucleus as proposed in the planetary model.

20. Atomic Number • Atomic Number = the number of protons/electrons • Atomic Mass = protons + neutrons. • Atomic Symbol = symbol for the element

25. Isotopes • Isotopes are atoms of the same element that have different number of neutrons, hence have a different mass number

27. Isotope Symbols • Isotopes can be identified by using a symbol the indicates both the element and its mass number. • Thus C-12 represents a carbon atom with the mass number of 12 which represents the sum of the protons and neutrons. • Because the atomic number of C-12 is 6, the number of neutrons will also be 6.

29. Subatomic Particles -24 • Protons have a mass of only 1.67 x 10 • Because the mass of a proton is so small, it is more convenient to use a different scale called atomic mass units.

30. Location of Electrons • Remember that electrons are found in the space of the atom around the nucleus. • Experiments have shown that electrons are in orbitals.

31. Energy Levels • The orbital in a atom form a series of energy levels in which the electrons can be found. • Each electron in an atom has its own distinct amount of energy that corresponds to the energy level that it occupies. • Electron can gain or lose energy and move to different energy levels.

32. How to calculate AMU • How to calculate AMU – Atomic Mass Units • Atomic mass units can be calculated from the mass and the abundance of naturally occurring isotopes. • Carbon has two naturally occurring isotopes.

33. 98.89% C-12 1.108% C-13 • Convert the percentages to decimal numbers, and multiply the mass of each isotopes by its decimal abundance 12 amu x 0.9889 = 11.87 amu 13 amu x 0.01108 = 0.1440

34. 2. Add these masses of isotopes. 11.87 amu + 0.1440 amu = 12.01 amu

35. Naturally occurring europium (Eu) consists of two isotopes was a mass of 151 and 153. Europium-151 has an abundance of 48.03% and Europium-153 has an abundance of 51.97%. What is the atomic mass of europium?

36. Subatomic Particles

37. Masses • Mass of the electron= 1/1836 amu of a proton • 1 amu = 1.7 x 10 -24

38. Isotopes of Hydrogen • Protium: P:1 N:0 Mass Number: 1 amu • Deuterium: P:1 N:1 Mass Number: 2 amu • Tritium: P:1 N:2 Mass Number: 3 amu

40. Orbits in Detail • Electrons are found in the space of the atom around the nucleus. • Orbitals in an atom form a series of energy levels in which electrons may be found • Each electron atom has its own distinct amount of energy that corresponds to the energy level.

41. Orbits in Detail • Electrons can gain or lose energy and move to different energy levels, but they do so in a unique way. • If an electron moves from one energy level to another, it must give off or absorb the energy difference between the two levels

42. Ground and Excited States • When electrons occupy the lowest available orbital, the atom is said to be in the ground state. • When electrons are subject to stimuli such as heat, light or electricity they move to a higher level; excited state.

43. Electron Configuration

44. Electron Configuration

45. Electron Configuration

46. Electron Configuration • Homework: • Write the Electron configuration for the following elements: Mg Ar Se Sc CL Ne