UNIT 3: Atoms: the building blocks of matter

1. UNIT 3: Atoms: the building blocks of matter INSTRUCTIONS FOR COPYING NOTES Copy all notes in 2-column (Cornell) style You don’t need to copy tables, diagrams, or pictures, UNLESS INSTRUCTED TO. You DO need to copy all informational slides (“words”) Leave room for a 3-sentence summary at the bottom of each page. All notes are copied into SPIRAL* *If you are an AVID student, you may keep notes in binder

2. Essential Question #1 How do I explainatomic theory and describe the changes in the atomic model over time and why those changes were necessitated by experimental evidence? (copy the next slide, then write “ See Structure of Atom Foldable for more info.” Continue with notes as usual.)

3. Early Ideas About Matter • Greek philosophers (300 BC) proposed matter was made of 4 elements: earth, air, fire, water. • Democritus coined the word “atom” meaning “cannot be broken.” • Atom seen as a solid sphere

4. Essential Question #2 • What is atomic theory? (Can I describe the structure of atoms in terms of protons, neutrons and electrons, and differentiate among these particles in terms of their mass, electrical charges and locations within the atom? • SC.912.P.8.4

5. Subatomic Particles

6. DON’T COPY

7. Charges in an Atom • The + charge on a proton is equal to the - charge on an electron. • Atoms are N______ (have no overall charge) • Therefore, the # of protons = # electrons in an atom.

8. Atomic number • determines the identity of the atom. • tells us the # of protons in the atom. • also tells us the # of electrons (b/c an atom is neutral in charge.) • Ex: atomic number of carbon, C = 6 • Question: how many protons? How many electrons? How many neutrons?... Slide 2.2

9. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

10. Mass Number Mass number is the number of protons and neutrons in the N_____ of an isotope: Mass # = p+ + n0 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

11. Isotopes • (Def) Atoms of the same element w/different #s of neutrons. • The number of neutrons can vary from atom to atom in an element. • In order to know how many neutrons in an atom you must be told. • The mass number tells you how much mass the atom has. • Since p+ and n0 are the heavy parts, • mass # =p+ + n0.

12. Writing atomic symbols for isotopes B Symbol for element 11 Mass # 5 Atomic #

13. Naming Isotopes • The atom in the prior slide can be called “boron-11” Mass # Name of element

14. QUESTION: If the mass number of a carbon atom is 14, • How many protons? • How many electrons? • How many neutrons? • LET’S PRACTICE! • Whiteboard • Marker • Paper towel

15. Practice Problem #1 • If an element has an atomic number of 34 and a mass number of 78, what is the: • number of protons • number of neutrons • number of electrons • complete symbol

16. Practice Problem #2 • If an element has 78 electrons and 117 neutrons what is the • Atomic number • Mass number • number of protons • complete symbol

17. Atomic Mass Units • Atoms are weighed in a.m.u. • 1 a.m.u. is based on the mass of a Carbon-12 atom. • it has 6 p+ and 6 n0, • 1 a.m.u = 1/12 the mass of a carbon-12 atom.

18. Atomic Mass • (definition) Weighted average of all the isotopes of an element. See p 68 of text. calculating atomic mass • Located below element symbol on periodic table.

19. Isotopesare atoms of the same element having different masses, due to varying numbers of neutrons.

20. Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons.

21. To calculate the average: • Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. • If not told otherwise, the mass of the isotope is expressed in atomic mass units (amu)

22. Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon = 12.011

23. Practice Problem #1 • Copper has the following isotopes ISOTOPE MASS # % ABUNDANCE • Copper-63 63 69.15 • Copper-65 65 30.85 • Calculate the atomic mass (average mass)

24. Understanding Info on P.Table(Leave room for a practice problem at the end of your notes) Finding Atomic Mass on Periodic Table

25. Calculating Atomic Mass • Leave room to copy an example

27. STOP HERE!!!Expanding our model of the atom Ch 5 – Chem IH Ch2.2 & 7 – Chem I

28. Light: Electromagnetic Spectrum • Energy can travel in waves. • There are high energy and low energy waves. • The ones we can see are called “the visible spectrum.” ROY G BIV • Red is the low energy end: violet is the high energy end.

29. Properties of Waves • 1. Wavelength: distance between crests of a wave. Ex: radio waves = 102 m

30. Properties, cont. • 2. Frequency: number of wave cycles to pass a point per unit time.

31. Electrons in Atoms Energy of Electrons • Why electrons don’t crash into the nucleus: they have enough energy to keep them away. • Why e-s (usually) don’t fly off of atoms: they have enough attraction to the nucleus to keep them in “orbit.” (Kind of like planets in orbit around the sun.)

32. Energy of Electrons (cont.) (Don’t write this!) DISCUSS WITH YOUR NEIGHBOR: • You are an electron. If you have a lot of energy, will you stay close to the nucleus or will you move further from it? Answer: you may still stay in “orbit” but you will be able to move further away from the nucleus.

33. Bohr’s Model of Atom • Neils Bohr studied w/Rutherford • His model is also called the planetary model • He discovered that e-s could only exist at certain distances from the nucleus. (Energy Levels)

34. Niels Bohr • "The opposite of a correct statement is a false statement. But the opposite of a profound truth may well be another profound truth." Neils Bohr

35. Bohr’s Model of Atom • See p 75 of text: electron energy levels are like rungs of a ladder. • Ladder • To climb to a higher level, you can’t put your foot at any level, • you must place it on a rung • Electron energy levels • e-s must move to higher or lower e.l.’s in specific intervals

36. Bohr Model of the Atom • Interactive Bohr Model

37. Electrons in Energy Levels • Atoms are arranged in energy levels (e.l.’s), at different distances from nucleus • Close to nucleus = low energy • Far from nucleus = high energy • e-s in highest occupied level are “valence e-s” • Only so many e-’s can fit in energy levels • e-s fill lower e.l.’s before being located in higher e.l.’s* (* There are exceptions we will learn later!)

38. Electrons in Energy Levels • Only so many e-’s can fit in energy levels Energy Level # of electrons 1st 2 2nd 8 3rd 18* 4th 32*

39. KEY CONCEPT!!! • VALENCE ELECTRONS DETERMINE HOW ELEMENTS BEHAVE!!!

40. Drawing Bohr Models Let’s practice drawing some atoms/ions In your teams, pick up enough of the following for your team: 1 white board per person 1 marker per person 1 paper towel per team (Please save a tree & share!)

41. Electron Cloud Model of Atom • Electrons aren’t in perfect orbits. • Energy levels are regions of space in which an e- is likely to be found most of the time. • The area in which they move is like a cloud, an area of space surrounding the nucleus.

42. Drawing Bohr Models • Show # of protons and neutrons in the nucleus • Draw e.l.’s and show each electron in the proper e.l. • Ex: Bohr Model of BORON-11

43. Practice • Hydrogen-2 (Practice together) • Helium-4 • Lithium-6 • Beryllium-8 • Carbon-12 • Magnesium-24

44. Lewis-Dot Diagrams Have 2 parts • Chemical symbol of element • Valence e-s, represented by dots • Are placed in one of four locations • Above • Below • Right • left • Are not paired unless there is 1 e- in each location. • Ex: Oxygen

45. Practice Lewis Dot Diagrams TEACHER DEMONSTRATION • Hydrogen • Helium • Lithium STUDENT PRACTICE • Beryllium • Boron • Carbon

46. PRACTICE WORKSHEET • Bohr Models • Lewis dot diagrams