Atoms and the Periodic Table

1. Atoms and the Periodic Table Physical Science Chapter 4

2. Section 1: Atomic Structure Objectives • Explain Dalton’s atomic theory, and describe why it was more successful than Democritus’s theory. • State the charge, mass, and location of each part of an atom according to the modern model of the atom. • Compare and contrast Bohr’s model with the modern model of the atom.

3. What is an atom? An atom is the smallest unit of an element that maintains the properties of that element. The word atom is derived from the Greek word meaning “unable to be divided.” Atoms are all around us… in the air you are breathing, in our chairs, even in our food!

4. How were atoms discovered? In the fourth century BC, a Greek philosopher, Democritus, suggested that the universe was made of invisible units called atoms. He thought that the movement of atoms caused the changes in matter that he observed. Democritus tried and tried to prove his theory but was unable to provide evidence needed to convince people that atoms even existed!

5. How were atoms discovered? John Dalton developed his own atomic theory in 1808. Dalton’sAtomicTheory: • Stated that atoms could not be divided into smaller parts • Stated that all atoms of a given element were exactly alike • Stated that different elements could join together to form compounds Dalton’s Theory is considered the foundation for the modern atomic theory.

6. Parts of an Atom • Nucleus: small, dense center of an atom that has a positive electric charge • Made of protons (positive charge) and neutrons (neutral charge) • Protons and neutrons are almost identical in size and in mass • Electrons move around outside of the nucleus in energy levels and orbitals. • Electrons have a negative charge and are smaller than protons or neutrons • Electrons are constantly on the move but do remain in predictableenergylevels

7. Atoms and their Charges • Atoms have an overallneutralcharge • Atoms are made of an equalnumber of protons (positive charge) and electrons (negative charge) • The number of neutrons (neutral charge) canchange and not effect the overallcharge on the atom

8. Models of the Atom • In 1913, NielsBohr suggested that electrons in an atom move in setpaths around the nucleus • The electron’s path determines the electron’s energylevel • Electrons can only move in certain energy levels. They must gain or lose energy to move from one energy level to another.

9. Bohr Model of a Lithium Atom

10. Bohr Model v. Modern Model • In 1925, Bohr’s Model of the atom no longer explained electron behavior. • In the BohrModel, electrons travel around the nucleus the way that planets orbit the sun in our solar system. • In the modernmodel of the atom, electrons are believed to behave more like waves on a vibratingstring.

11. Electron Energy Levels • Within the atom, electrons with different amounts of energy exist in different energy levels. • The number of filled energy levels depends on the numberofelectrons in the atom • First Energy Level: Holds 2 Electrons • Second Energy Level: Holds 8 Electrons • Third Energy Level: Holds 18 Electrons • Fourth Energy Level: Holds 32 Electrons

12. Electron Orbitals • ElectronOrbitals: the regions in an atom where electrons are likely to be found • Four different kinds of orbitals: • s-orbital • p-orbital • d-orbital • f-orbital

13. Valence Electrons • ValenceElectrons: electrons that are found in the outermost energy level • Valence electrons are involved in chemicalbonding • Every atom has between one and eight valence electrons • Noblegases have eight valence electrons and, therefore, are stable

14. n n n n n n n n n n n n n n n n n n n n n n n n n n n n n + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- H He O Cl N Na Let’s Practice “Drawing Atoms” Let’s Practice “Drawing Atoms”

15. Section 2 Objectives: A Guided Tour of the Periodic Table • Relate the organization of the periodic table to the arrangement of electrons within an atom. • Explain why some atoms gain or lose electrons to form ions. • Determine how many protons, neutrons, and electrons an atom has, given its’ symbol, atomic number, and mass number • Describe how the abundance of isotopes affects an element’s average atomic mass.

16. Organization of the Periodic Table • The periodic table groups similar elements together • Makes it easier to predict the properties of an element based on placement • The order is based on the number of protons an atom of the element has in its’ nucleus • The periodic law states that the repeating chemical and physical properties of elements change periodically with the atomic numbers of the elements

17. Periodic Table Arrangement • Horizontal rows in the periodic table are called periods. • The number of protons and the number of electrons increase as you move left-to-right across a period. • You can determine how the atom’s electrons are arranged as you move across a period.

18. Periodic Table Arrangement • Vertical columns in the periodic table are called groups. • Atoms of the same group have the same number of valence (outermost) electrons and react in a similar way. • Atoms of the same group also have similar chemical properties.

19. Forming Ions • An ion is an atom or a group of atoms that have a net electrical charge. • Neutral atoms form ions by gaining or losing electrons. • By using the groups in the periodic table, you can predict whether an atom will gain or lose an electron!

20. Example: Li atom • Lithium is found in Group 1. • This means that there is ONE valence electron in the outer shell. Li Li+ Li loses one electron Lithium Atom: One valence electron Lithium Ion: More positive now that electron is lost

21. Example: F atom • Fluorine is found in Group 17. • This means that there are SEVEN valence electrons in the outer shell. F- F F gains one electron Fluorine Atom: Seven valence electrons Fluoride Ion: More negative now that electron is gained

22. Periodic Table Information Atomic Number Element Symbol Element Name Atomic Mass Number Atomic Number: How many protons are in the nucleus of the atom Atomic Mass Number: How many protons and neutrons are in the nucleus of the atom

23. Let’s Do Some Math! • How many protons does the atom have? • How many electrons does the atom have? • How many neutrons does the atom have? 20 Ca 40.078

24. Parts of the Nucleus

25. Isotopes & Avg. Atomic Mass • An isotope is an atom that has the same number of protons as other atoms of the same element but have a different number of neutrons. • The average atomic mass for an element is a weighted average that includes all the element’s isotopes. • Average atomic mass = amu (units)

26. Isotopes of Chlorine

27. The AtomConcept Map

28. Section 3 Objectives: Families of Elements • Locate alkali metals, alkaline-earth metals, and transition metals in the periodic table. • Locate semiconductors, halogens, and noble gases in the periodic table. • Relate an element’s chemical properties to the electron arrangement of its atoms.

29. Elements: 3 Basic Groups • Metals: • Shiny solids that can be stretched and shaped • Goodconductors of heat & electricity • Nonmetals: • May be solids, liquids, or gases • Dull and brittle • Poorconductors of heat & electricity • Semiconductors or Metalloids: • Under certain conditions, some nonmetals can conduct heat & electricity

30. Elements: 3 Basic Groups

31. Metal-Type Classification • Alkali Metals: Group 1 • Reacts easily because of ONE valence electron • Usually not found in nature alone because of reactivity—found in compounds • Examples: • Na • K • Li Sodium Potassium Lithium

32. Metal-Type Classification • Alkaline-Earth Metals: Group 2 • Have TWO valence electrons • Less reactive than alkali metals • Form important compounds in living things • Shells and bone (calcium compounds) • Enzymes (magnesium compounds) • Examples: • Mg • Ca • Ba Magnesium Calcium Barium

33. Metal-Type Classification • Transition Metals: Groups 3-12 • Have different numbers of valence electrons • Can form both cations and anions • Cations: positively charged ions (lose e-) • Anions: negatively charged ions (gain e-) • Properties of transition metals change across a period • Examples: • Au • Ag • Hg Gold Silver Mercury

34. Nonmetals • Found on the right side of the periodic table (includes H) • Generally have 4 or more valence electrons • Form anions easily (gain e-) • Nonmetal compounds are plentiful on earth • Carbon is a common nonmetal and is found in all living things

35. Types of Nonmetals • Halogens: Group 17 • Have SEVEN valence electrons • Form –1 anions easily (gain 1 electron) • Highly reactive elements • Examples: • F • Cl • Br Fluorine Chlorine Bromine

36. Types of Nonmetals • Noble Gases: Group 18 • Have EIGHT valence electrons • Do NOT react with other elements • Also known as inert gases • Examples: • Ne • He • Kr Neon Helium Krypton

37. Synthetic Elements • Synthetic elements are man-made. • Elements with atomic numbers greater than 92 are man-made. • Most are radioactive • Radioactive: decaying into other elements • Found in the last two periods at the bottom of the periodic table

38. Section 4 ObjectivesUsing Moles to Count Atoms • Explain the relationship between a mole of a substance and Avogadro’s Number. • Find the molar mass of an element by using the periodic table. • Solve problems converting the amount of an element in moles to its mass in grams and vice versa.

39. Bell Ringer: Chapter 4 Section 4

40. What is a mole anyway?? • It’s not really a furry creature but a way to count in chemistry! • Just like: 12 eggs = 1 dozen; One mole of a substance = 6.022 x 1023 atoms!

41. Avogadro’s Constant • Chemists use large amounts of very small particles and needed a way to count them. • Avogadro’s Constant, named after Italian scientist, Amedeo Avogadro, defines the number of atoms in ONE mole of a substance. • One mole = 6.022 x 1023 atoms! • A mole is abbreviated : 1 mol

42. Why such a weird number? • Since carbon-12 is found in abundance on earth, • The mole is defined as the number of atoms in 12.00 grams of carbon-12. • There are 6.022 x 1023 atoms of carbon-12 in 12.00 grams. • There is a relationship between atoms and grams!

43. Moles and Grams are Related • The mass in grams of 1 mol of a substance is called its molar mass. • The molar mass (g) is the same as the average atomic mass (amu) listed on the periodic table. • We can use conversion factors to convert: • To moles from grams • To grams from moles

44. Gumball Example p. 132 • Given : 10 gumballs has a mass of 21.4 g • Question: What is the mass of 50 gumballs? • How to solve: Conversion Factor: 21.4 g 10 Gumballs Given: 50 Gumballs 107 g = X Arrange so that Gumballs Cancel

45. Practice These! • What is the mass of exactly 150 gumballs? • If you want 50 eggs, how many dozens must you buy? How many extra eggs will you have leftover? • If a football player is tackled 1.7 ft short of the end zone, how many more yards does the team need to get a touchdown?

46. Given: 5.50 mol Fe Conversion Factor: 55.85 grams 1 mol Fe 307 g Fe = X Moles to Mass Problem • Determine the mass in grams of 5.50 mol of iron. Molar Mass From Periodic Table

47. Given: 352 grams Fe Conversion Factor: 1 mol Fe 55.85 grams Fe 6.30 mol Fe = X Mass to Moles Problem Determine the amount of Iron present in 352 grams. Molar Mass From Periodic Table

48. Graphic Organizer 1 mol Molar Mass X = Number Of Moles Mass in Grams Molar Mass 1 Mol = X

49. Practice Problems • Moles to Mass • P. 133 (# 1-4); P. 134 (# 6) • Mass to Moles • P. 134 (# 7-9)