Unit 3: Atomic Structure & The Periodic Table

1. Unit 3: Atomic Structure & The Periodic Table Chapters 13 & 14

2. Energy: The capacity to do work • Potential Energy: stored energy due to position or condition. Chemicals can store energy; thus they have potential energy. • Kinetic Energy: energy in motion.

3. Kinetic Theory • Particles have no attractive or repulsive forces existing between the particles • Particles in gas move rapidly in constant motion. They travel in a straight path. • Total kinetic energy is conserved when particles collide.

4. Models of Atoms Atomic Models: • Chemical properties of atoms, ions, and molecules are related to the arrangement of the electrons within them. • John Dalton: 1st atomic model & considered the atom as a solid indivisible mass.

5. Dalton’s Atomic Theory 1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3. Atoms of different elements can physically mix together or can chemically combine. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of element are never changed into atoms of another element as a result of a chemical reaction.

6. J.J. Thomson:revised Dalton’s model by proposing that electrons were stuck to the outside of the atom. • Ernest Rutherford: proposed the nuclear atom, in which electrons surround a dense nucleus composed of protons and neutrons. • Chadwick: discovered the neutron.

7. Discovery of the nucleus • This theory was discovered by Rutherford who bombarded a sheet ofgold foil with a beam of alpha particles surrounded by a fluorescent screen. They found that most of the particles passed through the foil, while a few were deflected.

8. Niels Bohr: student of Rutherford; proposed that electrons are arranged in concentric circular paths (orbits) around the nucleus. • Erwin Schrodinger: developed the quantum mechanical model

9. Atoms – tiny particles that make up matter Structure of Atoms: • Nucleus – center of the atom • protons – positively charged subatomic particles that is found in the nucleus; dictates the identity of the atom -(Discovered by E. Goldstein using canal rays. Canal rays traveled from the positive metal plate to the negative metal plates)

10. Neutron: subatomic particle with no charge; found in the nucleus - Discovered by James Chadwick • Electron: negatively charged; found outside the nucleus ( electron cloud) - Discovered by J.J. Thomson using a cathode ray - the rays were attracted to a metal plate of positive charge.

12. Atomic Number- the # of protons in the nucleus - the # of protons = the # of electrons • Mass Number – total # of protons & neutrons in an atom - To find the # of neutrons subtract the mass # from the atomic #

13. Example: Nitrogen (147N) Mass number = 14 Atomic number = 7 # of protons = 7 # of electrons = 7 # of neutrons = 14 – 7= 7

14. Isotopes - Has the same # of protons, but different #’s of neutrons • Atomic Mass Unit (AMU) – 1/12 the mass of carbon • Average atomic mass: weighted average of the masses of the isotopes of an element

15. In nature most elements occur as a mixture of two or more isotopes • Each isotope has a fixed mass and a natural percentage of abundance.

16. Average Atomic Mass • Avg atomic mass: what is found on the periodic table • =(mass)(% abundance) + (mass)(% abundance) +…

17. Practice Problem • Assume that element Uus is synthesized and that it has the following stable isotopes: 284Uus (283.4 amu) 34.60% 285Uus (284.7 amu) 21.20% 288Uus (287.8 amu) 44.20% What would the average atomic mass be?

18. Bohr’s Model • Orbits are known as energy levels. • Electrons can move between energy levels. • A quantum of energy is the amount of energy required to move an electron up an energy level. • The higher the energy level the easier the electron can escape.

19. Bohr Models Examples on the board Elements #1-20

20. Atomic Orbitals 1. Principal Quantum Numbers ( ) = 1,2,3,4…. 2. Each principal level contains sublevels * Table 13.1 p. 364 3. Atomic orbitals are regions where electrons can be found. (Letter denotes the orbital)

21. S orbitals are spherical. • P orbitals are dumbbell-shaped. ( exist in three different planes) • D orbitals have clover leaf shapes • F orbitals have complex shapes

22. 4. The number & kinds of atomic orbitals depend on the energy sub level. a. N = 1; 1 sublevel; 1s orbital b. N = 2; 2 sublevels; 2s (1 orbital), 2p ( 3 orbitals) c. N = 3; 3 sublevels; 3s ( 1 orbital), 3p (3 orbitals), 3d (5 orbitals) d. N = 4; 4 sublevels; 4s (1 orbital), 4p (3 orbitals), 4d ( 5 orbitals), 4f (7 orbitals)

23. 2n2 (n = principal quantum number). This equals the maximum # of electrons that the sublevel can hold.

24. Electron Arrangements in Atoms Electron Configurations: 1. Unstable systems tend to lose energy to become stable. 2. Electrons try to form stable arrangements with the nucleus. 3. The way in which electrons are arranged around the nuclei of atoms is called electron configuration.

25. 4. Three rules tell you how to find the electron configuration of atoms. a. Aufbau principle: electrons enter orbitals of lowest energy level first. b. Pauli exclusion principle: and atomic orbital may describe at most two electrons. (arrows show the direction of electron spin) c. Hund’s rule: when electrons occupy orbitals of equal energy, one electron enters each orbital, all of orbitals contain one electron with parallel spins.

26. Electron Configurations of Ions 5. When writing electron configurations for ions you must add or subtract the # of electrons gained or lost to create the ion.

27. Electron Configuration Practice • Elements #1-20

28. PERIODIC TABLE • Periodic Table – an arrangement of elements according to similarities in their properties • There are 92 naturally occurring elements. • Demitri Mendeleev – drew the first periodic table; Russian chemist arranged the first periodic table of elements in 1871. Arranged by atomic mass

29. * The periodic table contains chemical symbol, atomic number, & average atomic mass, physical state of each element, group numbers, and electron configuration. • Moseley: Later arranged the periodic table by atomic number. (Which is the one we use today.)

30. MODERN TABLE • Periods – horizontal rows (7 total) • Groups – vertical columns (has similar physical & chemical properties) • Metals – high electrical conductivity, luster, ductile, & malleable (Group 1 & 2A) - Alkali Metals – Group 1A - Alkaline Earth Metals – Group 2A

31. Transition Metals & Inner Transition Metals – make up Group B (1B – 8B) • Nonmetals – poor conductors, non lustrous - Halogens – 7A - Noble Gases – 0 • Metalloids – elements that border the stair step line • Group # = the outermost electrons

32. Periodic Trends • The elements on the periodic table are arranged periodically so that trends can be recognized…

33. Trend of Ions 1. You can determine the charge of an ion by what group it is in. 1A = +1 5A = -3 2A = +2 6A = -2 3A = +3 7A = -1 4A = +/- 4

34. Trend of Electronegativity • This refers to the ability of an atom to attract the electrons of another atom to it. • Increases across the period ( left – right) • Decreases down the group ( top – bottom)

35. Trend of Electron affinity • Measure of the tendency for atoms to gain electrons. • Increases across the period; this is caused by the filling of the valence shells • Decreases down the group; this is due to the electron entering an orbital far away from the nucleus

36. Trend of Ionization Energy • The exact quantity of energy that it takes to remove the outermost electron from the atom. • Factors affecting Ionization Energy: - nuclear charge - distance from the nucleus

37. Ionization energy increases across the period ( left – right) due to increased nuclear charge • Ionization energy decreases down the group ( top – bottom)

38. Trend of Atomic Radius • Atomic size is determined by how much space the electron takes up. It is also depends on how far its valence electrons are from the nucleus. • The atom will be large if the electron is far from the nucleus - size increases down a group (top – bottom)

39. The atom will be small if the electron is close to the nucleus - size decreases across the period ( left – right) This is due to an increase in nuclear charge pulling them closer… the energy level stays the same

40. Trend of Metallic/Non-Metallic Properties • Metallic properties: elements will form cations as they lose electrons (+ve charge) • Non-Metallic properties: elements form anions as they gain electrons (-ve charge)

41. Trend of Melting / Boiling Points • Melting and Boiling point increase from the right side of the periodic table until it reaches aluminum and silicon • Here, melting point and boiling point then begin to decrease.

42. Trend of Reactivity How likely/vigorously an atom is to react with other substances • Metals: • Period: decreases from left to right • Group: increases down the group The farther left and down you go the easier it is for electrons to be taken away. (Higher Reactivity)

43. Trend of Reactivity • Non-Metals • Period: increases from left to right • Group: decreases down the group The farther right and up you go the higher electronegativity – vigorous exchange of electrons

44. Classification of Elements Elements can be classified into 4 groups based on electrons. 1. Noble gases: outermost s & p sublevels are filled. Belong to group 0. (Also called inert gases.) 2. Representative elements: outermost s or p sublevel is partially filled

45. 3. Transition metals: metallic elements in which the outermost s sublevel and near d sublevel contain electrons. (Group B elements) 4. Inner transition metals: metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. (Lanthanide & Actinide series)

46. Light and Atomic Spectra • Light consists of electromagnetic waves. • Light has a velocity of 3.0 x 10 8 m/s. • Amplitude: is the wave height from origin to crest. • Wavelength (λ): distance between crest. • Frequency (ν): number of wavelength to pass a given point per unit of time. (units = hertz Hz)

47. c=λν • c= speed of light (3.00 x 10 8 m/s) • λ= wavelength • ν= frequency Example: Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s-1). c = 3.00 x 108 m/s Frequency (ν) = 5.10 x 1014 s-1 wavelength (λ) = ??? m

48. Frequency & wavelength are inversely related. • Electromagnetic spectrum: series of waves at different wavelengths (radio waves, radar, microwaves, infrared, visible light, ultraviolet, x-rays, gamma rays, cosmic rays) • Every element emits light when it is excited by the passage of an electric discharge through its gas or vapor.

50. Black & White Light • Black light – All colors absorbed • White Light – All colors reflected • What happens for you to see colors?