Unit 4 Atomic Theory, Electron Arrangement and Periodic Table

1. Unit 4 Atomic Theory, Electron Arrangement and Periodic Table Topics from Chapters 3-5

2. Periodic Table Terminology 1. Periods: Horizontal row of elements in the periodic table. 2. Group (family): A vertical column of elements in the periodic table. (Have the same number of valence electrons)

4. Atomic Theory Continues… • We have come so far in our understanding of atoms. Centuries of researching and countless scientists devoting their lives to create the understanding of the atom today (textbook concepts). • However it is NOT over. The more we understand about atoms (how they work, their make up, etc…) the greater our ability to advance science and technology in all aspects of our lives (i.e. medicine).

5. Atomic Theory Reading and Graphic Organizer

6. Atoms: The Building Blocks of Matter • Basic Laws (1790s) • Law of Conservation of Mass • Mass is neither created nor destroyed during chemical reactions or physical changes. • Law of Definite Proportions • A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. • Law of Multiple Proportions • If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combine with a certain mass of the first element is ALWAYS a ratio of small WHOLE numbers.

7. Atomic Theory - Video • John Dalton (1808) • English Schoolteacher • Proposed an explanation of the 3 basic laws (as mentioned previously) • Dalton’s Atomic Theory (5 Statements) • All matter is composed of extremely small particles called atoms. • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in simple whole number ratios to form compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

8. Modern Atomic Theory • Dalton’s Theory has been modified to fit new findings. (Science is NOT static) • Example: Today we know that atoms are divisible into even smaller particles • 1. Subatomic Particles (protons, neutrons, and electrons) • 2. Protons and Neutrons are made up of quarks • Important concepts • All matter is composed of atoms • Atoms of any one element differ in properties from atoms of another element remain unchanged.

9. Electron Discovery • JJ Thomson (1897) • Cathode-ray tubes • 1. Electric current was passed through various gases at low pressures. • 2. The glow was caused by a stream of particles – cathode ray • 3. The ray was deflected by a magnetic field and away from a negatively charged object. • Concluded that all cathode rays are composed of identical negatively charged particles (electrons) • Plum pudding model • 1. Negative electrons spread throughout the positive charge of the rest of the atoms. • 2. Plums= electrons ; pudding=positive charge

10. Electron Discovery Cont… • Robert Millikan (1909) • Measured the charge of the electron. (Charge to mass ratio) • Mass of electron = 9.109 x 10-31kg • Or about 1/2000 the mass of a hydrogen atom • Inferences from findings… • Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons. • Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

11. Nuclear Atom • Ernest Rutherford (1911) • Bombarded a thin piece of gold foil with fast-moving alpha particles (positively charged particles) • Concluded that deflected alpha particles must have experienced some powerful force within the atom • The source of this force must occupy a very small amount of space • Force must be caused by a very densely packed bundle of matter with a positive electric charge (Nucleus) • Volume of nucleus was very small compared with the total volume of an atom. • If the nucleus were the size of a pea, then the size of the atom would be about the size of a football field.

12. Unit 4 The Arrangement of Electrons Ch. 4

13. What Do You Think?Electron Placement Describe the location of electrons in a single lithium (Li) atom.

14. Check for Understanding • What is the difference between the excited and ground state? • List the different types of electromagnetic radiation found in the electromagnetic spectrum. • How are frequency, wavelength, and energy related?

15. The Development of a New Atomic Model • Ernest Rutherford • Model incomplete did not describe why the electrons where not pulled to the center of the atom (positively charged nucleus), opposite charges attract. • Early Twentieth Century – New/Revised Atomic Model • Absorption and Emission of light by matter • Relationship between light and an atom’s electrons • Light behaves like a particle and a wave. (Particle-wave duality) • Particle-Wave Duality 1 • Particle-Wave Duality 2

16. Electromagnetic Radiation • A form of energy that exhibits wavelike behavior as it travels through space. • Gamma Rays, X-rays, Ultraviolet, Infrared, Microwaves, Radio waves. • Electromagnetic Spectrum • All types of light waves move at a constant speed (3.0 x 108 m/s – “the speed of light”) through a vacuum (relatively the same for air) • Speed = Wavelength x Frequency c = λν • Wavelength – distance between corresponding points on adjacent waves (meters). • Frequency – number of waves that pass a given point in a specific time.

18. Photoelectric Effect • Emission of electrons from a metal when light shines on the metal • Electromagnetic radiation (light) strikes the surface of the metal ejecting electrons from the metal and causing an electric current, if the frequency was below a certain minimum.

19. The Particle Description of Light • Quantum • The minimum quantity of energy that can be lost or gained by an atom. • Photon • A particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

20. Check for Understanding • What is the difference between the excited and ground state? • List the different types of electromagnetic radiation found in the electromagnetic spectrum. • How are frequency, wavelength, and energy related?

21. The Hydrogen-Atom Line-Emission Spectrum • Ground State • Lowest energy state of an atom • Excited State • An atom has higher potential energy than it has in its ground state • When returning to ground state, the atom gives off the energy it gained in the form of electromagnetic radiation (light) • Video • Line-Emission Spectrum • When a narrow beam of emitted light is shined through a prism, it then is separated into four specific color(s) of the visible spectrum.

22. Bohr Model of the Hydrogen Atom • Niels Bohr • Proposed a hydrogen-atom model that linked the atom’s electron to photon emission • The electron can circle the nucleus only in allowed paths called ORBITS • Absorption • Energy must be added to an atom in order to move an electron form a lower energy level to a higher energy level. • Emission • When the electron falls to a lower energy level • A photon is emitted (light)

23. The Quantum Model of the Atom • The Heisenberg Uncertainty Principle • States that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. • One of the fundamental principles of our present understanding of light and matter.

24. The Quantum Model of the Atom • The Schrodinger Wave Equation • Hypothesis that electrons have a dual wave-particle nature • Helped lay the foundation for Modern Quantum Theory • Describes mathematically the wave properties of electrons and other very small particles. • Gives only the probability of finding electrons at a given place around the nucleus • Electrons do not travel in neat orbits, now in regions called orbitals • 3-D region around the nucleus that indicates the probable location of an electron • Different shapes and sizes

25. Flame Test Lab • See hand-out. • Demo • 15 minutes to complete

26. Emission Spectra • Bulbs (H, He, Ne, Ar, and N) and Emission Spectra. “The relationship between electrons and energy is...... The emission spectra is produced by......”

27. Electron Orbital -Electrons orbit the nucleus in orbital clouds. -Electrons with different amounts of energy exist in different energy levels.

28. The Electron Cloud Model

29. Electrons in each energy level • Each energy level can hold a limited number of electrons. • The lowest energy level is the smallest and the closest to the nucleus.

30. Electron Orbital Continued Energy Level # of Electrons 1 2 2 8 3 18 4 32

31. Draw the Atoms Lithium Atom: 3 Protons 3 Neutrons 3 Electrons Aluminum Atom: 13 Protons 13 Neutrons 13 Electrons

32. Valence Electrons • The electrons in the outermost energy level are called valence electrons. • See Periodic Table

33. - Complete in YOUR notes! • Determine the number of protons, neutrons, electrons AND draw the Bohr models of the following atoms: • Lithium • Calcium • Phosphorus • Neon

34. ......Thoughts..... Review/Update One’s Understanding… • Where are lithium’s electrons located in the atom? Turn to your elbow partner and share your ideas, be prepared to explain to the class.

35. POGIL : Electron Configurations • Clean Sheet of Paper / Notes • Label “Electron Configuration POGIL” • Work with partner – same face value on card. • Expectations: • Complete solutions on your note sheet. (Model) • Talk through it with your partner. • Stay on-task. • Be prepared to share out: • Pages 1 & 2: 5 Minutes • Pages 3-5 : 10 Minutes • Page 6: 5 Minutes

36. Electron Configuration • The arrangement of electrons in an atom. • Arrangements in lowest possible energies. • Ground-state electron configurations.

37. Aufbau Principle • An electron occupies the lowest-energy orbital that can receive it. • Writing electron configurations of atoms by successively filling subshells with electrons in a specific order (the building-up order). • You obtain electron configurations by filling the subshells in the following order: • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p • Think back to POGIL

38. 1s 1s 2s 2s 2p 2p Hund’s Rule • Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. • All electrons in singly occupied orbitals must have the same spin state. Ex: Carbon 1s22s22p2 Ex: Oxygen 1s22s22p4

39. Pauli’s Exclusion Principle • No two electrons in the same atom can have the same set of four quantum numbers. • An orbital can hold at most two electrons, and then only if the electrons have opposite spins. • Electron spin has only two possible directions it can spin on it’s axis. • ms= +1/2 or ms= -1/2

40. 1s 2s Writing Electron Configurations with Pauli’s Principle Instead of writing electron configurations using numbers and letters, one must draw an orbital diagram. -You still label the subshells (s, p, f, d) -Write electrons as arrows. Example: Lithium--1s22s1

41. 1. 2. 1s 1s 1s 2s 2s 2s 2p 2p 2p Practice with Pauli’s Principle • According to Pauli’s Principle are the following configurations possible? If no, explain why. 3.

42. Writing Electron Configurations • Determine the number of electrons using atomic number from the periodic table. • Follow the order listed by Aufbau (Diagonal Rule) to fill out the electron configuration. • The superscripts should add up to equal the number of electrons. Ex: Lithium’s Electron Configuration -Lithium has 3 electrons. 1s22s1

43. More Examples 1. Magnesium’s Electron Configuration -model 2. Carbon’s Electron Configuration- model 3. Iron’s Electron Configuration-partner 4. Bromine’s Electron Configuration-partner 5. Uranium’s Electron Configuration-partner

44. Electron Configuration and the Periodic Table

46. For electron configuration and the s,p,d, and f blocks we will use the following: • Lanthanium (La) and Actinium (Ac) are part of the f block. • (La - 4f1 and Ac 5f1) • Lutetium (Lu) and Lawrencium (Lr) are part of the d block. • (Lu - 5d1 and Lr -6d1) • There are discrepancies from source to source. However, we will use this way as it supports our diagonal rule and the stability of an atom.

47. Making It Easier • Using Noble Gas Notation, electron configurations can be shortened. • Noble gases are in the last column of the periodic table (He-Rn) • For any element’s electron configuration, find the noble gas that comes beforeit. • Start with this gas and then finish filling in the subshells. Lithium: [He]2s1 Phosphorus: [Ne]3s23p3

48. Whiteboard Practice Write the electron configuration for the following: • Be • P • Cd • Ce • Br • Gd

49. Check for Understanding Complete individually, turn over when finished. Write the electron configurations AND orbital notation for the following elements using the diagonal rule for order of filling. • Sulfur (atomic number 16) • Silver (atomic number 47) • Europium (atomic number 63)

50. Atomic Orbitals (Subshells) Atomic Orbital: The most likely place one will find an electron. S orbitals: Spherical in shape. Single Pockets. Two electrons are associated with an s orbital. P orbitals: Dumbbell shape. Groups of three. Six electrons fill the p orbitals.