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Unit 02 Atomic Structure. Just How Small is an Atom? You don’t need to write. A speck 0.1 mm in diameter (about half the size of a period at the end of.

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Presentation on theme: "Unit 02 Atomic Structure. Just How Small is an Atom? You don’t need to write. A speck 0.1 mm in diameter (about half the size of a period at the end of."— Presentation transcript:

1 Unit 02 Atomic Structure

2 Just How Small is an Atom? You don’t need to write. A speck 0.1 mm in diameter (about half the size of a period at the end of the sentence) requires one million atoms. It would require a million atoms, edge to edge, to match the thickness of a page of paper.

3 Can you see an atom? Technically, you cannot "see" anything smaller than the shortest wavelength of light that you can see it with. But there are ways to "visualize" it, like Atomic Force Microscopy. But these are all just measurements converted to computer images, and are not in any real sense "seeing" the atom. You can't see atoms in any normal sense of using an optical microscope. You don't get an optical image, but it does allow you to map out an image of the atoms of a molecule. To do this you use a metallic tip which interacts with the atoms you want to image. As you move the tip over the atoms, you pass a current, called a tunneling current, between the tip and the atom. This current is extremely sensitive to the distance between the atom and the tip.

4 - REMEMBER FROM: Elements, Mixtures, and Compounds - Element - a pure substance made up of one type of atom. - organized on periodic table - each element has a unique protons number of protons…its atomic number

5 Atomic Structure: Atoms contain three subatomic particles… 1. Protons… 2. Neutrons… 3. Electrons… These are located in NUCLEUS! Electrons surround the nucleus in orbitals positive charge neutral charge negative charge

6 Atomic Structure An atom is considered electrically neutral. Electrically neutral means the number of protons (+) = the number of electrons (-) 4 red protons = 4 blue electrons

7 Properties of Subatomic Particles ParticleSymbol Relative electrical charge Relative massActual mass (g) Electrone-e- 1-1/ × Protonp+p × Neutronn0n × Protons and Neutrons have the same mass.

8 A. Discovery of the Atom Ernest Rutherford discovered the nucleus by shooting alpha particles (have a positive charge) at a very thin piece of gold foil. He predicted that the particles would go right through the foil at some small angle.

9 Rutherford’s Gold Foil Experiment

10 some particles (1/8000) bounced back from the foil this meant there must be a “powerful force” in the foil to hit particle back Predicted ResultsActual Results Rutherford’s Gold Foil Experiment

11 Discovery of the Atom Purpose: The students will find the shape of different items and relate this to the early scientist that made discoveries about the shape and size of the atom. Procedure: 1. Title the left side of your spiral Discovery of the Atom. 2. For each item you will write the letters then draw your predicted shape of the item. 3. Then you will write 1 sentence describing why they think your prediction is the shape of the item. A: Item in brown bag – Use your hands to feel the shape of the item. B: Item in clay – Using the toothpicks provided find the shape of the object enclosed in the modeling clay. C: Black box – Maneuver the black box with a marble inside to discover the shape of the object enclosed.

12 B. Models of the Atom

13 “Plum pudding” atom negatively charged e - stuck into a lump of positively charged material – similar to chocolate chip cookies J.J. Thomson

14 “The Blow Pop” Ernest Rutherford In Rutherford’s gold foil experiment he discovered electrons surround a dense positive nucleus

15 Bohr Model electrons are arranged in fixed orbits around the nucleus. ex. Orbits gum

16 Quantum Mechanical Model Quantum mechanics was developed by Erwin Schrodinger Estimates the probability of finding an e - in a certain position Electrons are found in an “electron cloud”

17 I. Nuclear Symbols B 11 5

18 A. Mass Number mass # = protons + neutrons  always a whole number  NOT on the Periodic Table! © Addison-Wesley Publishing Company, Inc.

19 B. Isotopes Atoms of the same element with different mass numbers. (different number of neutrons) Mass # Atomic #  Nuclear symbol:  Hyphen notation: carbon-12

20 B. Isotopes

21 You must know how to find: A.# of protons = atomic number B.mass # = # of n 0 + # of p + (atomic #) – What’s in the nucleus of the atom C.# of electrons = # of protons (in a neutral atom) Boron 5 B atomic number (average) atomic mass (Not the same as the mass #) C. Nuclear Symbols

22 How to write a Nuclear Symbol B 11 5 Mass Number = p + + n 0 Atomic Number = p + Element Symbol -3 Charge if ion

23 C. Nuclear Symbols Chlorine-37 – atomic #: – mass #: – # of protons: – # of electrons: – # of neutrons:

24 Mg Nuclear Symbol Examples Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons Cl Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons

25 D. Relative Atomic Mass 12 C atom = × g  1 p= amu 1 n = amu 1 e - = amu © Addison-Wesley Publishing Company, Inc.  atomic mass unit (amu)  1 amu= 1 / 12 the mass of a 12 C atom

26 E. Average Atomic Mass weighted average of all naturally occuring isotopes on the Periodic Table round to 2 decimal places Avg. Atomic Mass

27 Avg. Atomic Mass E. Average Atomic Mass EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O amu

28 Avg. Atomic Mass E. Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine amu

29 Example: A sample of cesium is 75% 133 Cs, 20% 132 Cs and 5% 134 Cs. What is the average atomic mass? Answer:.75 x 133 = x 132 = x 134 = = average atomic mass

30 II. The Periodic Table Periodic Law Periodic Law – properties of elements can be predicted by their position on the periodic table

31 A. History of the Periodic Table Dmitri Mendeleev (1871) – Developed the first periodic table – It was arranged by atomic mass because atomic number had not been discovered – He was able to predict properties of elements

32 A. History of the Periodic Table Henry Moseley (1913) - developed the modern periodic table - arranged in order of increasing atomic number

33 B. Organization of the Periodic Table Period – horizontal rows numbered – Energies of outermost electrons are similar

34 Groups/ Families – vertical columns – have similar chemical & physical properties B. Organization of the Periodic Table

35 Group 1 (IA) – Alkali Metals

36 Group 2 (IIA) – Alkaline Earth Metals

37 Group 18 (VIIIA) – Noble Gases

38 Group 17 (VIIA) - Halogens

39 Left of stair step On the stair step Right of stair step METALS NONMETALS Metals- Nonmetals- Metalloids- B. Organization of the Periodic Table METALLOIDS

40 TRANSITION “Group B” INNER TRANSITION REPRESENTATIVE “Group A” Group A- Representative Group B - Transition B. Organization of the Periodic Table

41 The Extended Periodic Table


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