Understanding the Structure of the Atom: Elements, Subatomic Particles, and the Periodic Table

1. Chapter 3Structure of the Atom Chapter 3 Notes

2. Brainteaser • What makes one element different from another? • In the beginning….there were only the elements which were divided based on their properties.

3. Chemical Symbols • A letter or letter pair that represents the name of an element, similar to an abbreviation. • C = Carbon • Ag = Argentum (latin) = Silver

4. Types of Sub-Atomic Particles • Type Symbol (amu) Location • Proton (+) 1 nucleus • Neutron (n) 1 nucleus • Electron (-) 0 around the nucleus

5. Types of Sub-Atomic Particles • Type Symbol R. Mass Location • Proton (+) 1836 nucleus • Neutron (n) 1839 nucleus • Electron (-) 1 around the Nucleus • Differences in atoms arise from differences in the amounts of each particle in the atom.

6. About The Atom • Mass Number – the number of protons + number of neutrons, very similar to the atomic mass. • Atomic Number – The number of protons in the nucleus of an atom. • Number of Neutrons = mass number – atomic number

9. Practice Problems • How Many Protons/Neutrons/Electrons are present in: • Lithium • Sodium • Beryllium • Gold • What is the atomic number for: • Carbon • Helium • Iron • Copper

10. Changing the Atom • Change the number of Protons • New Atom, new properties • Change the number of Neutrons • New Isotope (could be radioactive) • Change the number of Electrons • New Ion (holds a charge)

11. Changing the Number of Protons New Element, New Properties

12. The Periodic Table • Brainteaser:How would you orderly arrange the pencils and pens of each member of our class?

13. 1800 – Dimitri Mendeleev (Russian) • Organized the periodic table of the elements • Periodic – repeated in a pattern • Chemical properties of the lighter elements repeated in the heavier elements • Left empty spaces based on the properties of elements not yet discovered (Ge) • Organized left to right based on mass

14. Dimitri’s Periodic Table

15. Modern Modifications • 1913 - Mosley reorganized left to right based on atomic number • always increases

16. Groups • Groups – Vertical columns • Similar properties • Same number of electrons in their outer shell

17. Periods • Periods – Horizontal rows • L to R increase in atomic number • Up to down increase in electron shell energy level

18. The Periodic Table • Metals– good conductors, solids at room temperature, shiny. • Nonmetals – mostly gases, brittle, bad conductors • Metalloids– have properties of both metals and non-metals • Hg – Br only liquids on the periodic table

19. The Periodic Table

20. The Families • Alkali Metals – Group 1 • Alkaline Earth Metals – Group 2 • Chalcogens – Group 16 • Halogens – Group 17 • Nobel Gases – Group 18 • Transition Metals – Groups 3-12 • Lanthanides – Top Row Bottom • Actinides – Bottom Row Bottom

21. Periodic Properties • More reactive at the edges of the table • Noble gasses do not react Periodic Trends: increases if… • Electronegativity – electron greedy (Up and Right) • Size – (Down and Left) • Ionization Energy – electron give away (down and left)

22. The Periodic Table • Diatomic Molecules – molecules made of two atoms • Naturally occurring on the periodic table • Br-I-N-Cl-H-O-F • Allotropes – Different forms of the same element

23. Allotropes

24. Changing the Number of Neutrons Isotopes…Radioactivity

25. Isotopes • Isotopes – atoms of the same element that have a different number of neutrons. • Question: Iodine is found in one of two isotopes, I-131 and I-130. How many Neutrons are present in each?

26. Average Atomic Mass • Average atomic mass – the average mass of the mixture of its isotopes (different from the actual atomic mass) • Question: If we have 5 atoms, 4 with the atomic mass of 11 and one with the atomic mass of 10, what is the average atomic mass?

27. Practice Problems • Question: a chlorine atom has 17 protons and 18 neutrons. What is the mass number? What is the atomic number?

28. Radioactivity • Some Isotopes are radioactive, they are too heavy (filled with neutrons) to be stable • They give off radiation to “transmutate” and become a lighter, more stable element

29. Types of Radiation • Alpha – two neutrons and two protons given off (one electron shy of being helium) • Beta – losing a positive charge, a proton converts to a neutron. • Gamma– no mass, but high energy given off by the nucleus • Types of Radiation

30. Calculating Half Life • The amount of radioactive isotope remaining can be calculated: Nt = No x (0.5) (number of half-lives) • Nt = amount of radioisotope remaining • No = original amount of radioisotope • number of half-lives = time ÷ half-life

31. Uses of Radiation • Half Life – Radio Carbon Dating (Archiology) • Radioactive Tracers (Medicine) • Effects of Radiation on the Body • Fission vs. Fusion • Atomic Weapons (Fission and Fusion) • Using Radiation for Energy (Nuclear Power) (Fission)

32. Changing the Number of Electrons IONS

33. Drawing Planetary Models • Find Atomic Number = Number of Electrons • Circle for nucleus including number of protons and neutrons (+ =, n =) • Draw orbits with correct number of electrons in each, filling in to out

34. Drawing Planetary Models

35. Practice Problems • Draw Planetary Models for the Following: • Lithium • Carbon • Iron

36. Lewis Dot Structures • A diagram of an atoms showing the valence electrons • Valence electrons – outer electrons • Where all exciting chemistry happens

37. How to Draw Lewis Dot Structures • Period = Energy level • Skipping the transition metals and counting group 1 as 1, count right to the position of your element. • The number = the electrons in the outer shell

38. How to Draw Lewis Dot Structures • Draw the element symbol • Draw one dot on each of the 4 sides, then double up when necessary (max 8 dots) • Octet Rule – Atoms considered to be stable have 8 electrons in their outermost shell

39. Practice Problems Give the periodic group number and number of valence electrons for each of the following atoms: • C Cl • Ne Si • Se Al Write Lewis Dot Structures for: • K Mg • S Ar

40. Ions • All atoms are naturally neutral (0 charge) • Electrons can be gained and lost to create a net + or – charge • Na → Na+ + e- (Ionic Equation)

41. Ions • Cations – positive ion – formed when one or more electrons are lost from an atom • Na → Na+ + e- • Anion – negative ion – formed when one or more electrons are gained • Cl + e- → Cl-

42. Ions • Ions are formed in repeated predictable patterns found on the periodic table • Gaining or losing electrons to reveal an octet – atom’s goal • Alkalli metals – all (+1) • Halogens – all (-1) • Transition metals – form (+) ions

43. Periodic Ions

44. Ionic Compounds • Ions separate in solution • High melting points • Mostly between metals and nonmetals • Charged ions in solution carry current through solution or when melted – conduct electricity • Saltwater vs. Water

45. Ionic Compounds • Must have one cation and one anion (metal and non-metal) • Total charge must equal zero • Multiply either the positive or negative ion by a factor until sum is zero • Chemical Bond – a force that holds together the atoms in a substance

46. Writing Molecular Formulas • Listing of the types of atoms (chemical symbols) and numbers (subscripts) of each in a molecule • N2O5 • SO3

47. Why atoms bond • Atoms wish to be chemically stable • Octet Rule – an element is stable when it has 8 outer electrons • An atom will seek to have 8 outer electrons by bonding or by gaining or loosing electrons

48. Ionic Compounds

49. Practice Problems • What is the molecular formula for: • Sodium and Chlorine • Magnesium and Chlorine • Show how each compound is formed

50. Answers