Elements, atoms, & the. discovery of atomic structure

Transcription

1 Elements, atoms, & the discovery of atomic structure

2 Chapter 4 EARLY MODELS OF THE ATOM

3 One What is an atom? The smallest particle of an element that can keep the same properties of the element.

4 Democritus (Ancient Greece, 440 B.C.) He stated that atoms are the smallest particles of all matter. Who was the first person to use the word atom? Democritus

5 Atomic Structure- outline Dalton s Atomic Theory Thomson s model vs. Rutherford s model of the atom Discovery of subatomic particles: Protons, neutrons, and electrons Isotopes Quantum mechanical model

6 Early model of the atom John dalton ( ) Dalton s Model was a dense solid sphere; indivisible and unchanged in chemical reactions. Where did he get his ideas?

7 Berzelius s Experiment Law of Definite Proportions

8 Proust s experiments Law of Definite Proportions

9 Dalton s Atomic Theory (pg 103) All matter is made of indivisible atoms; they can be neither created nor destroyed during chemical reactions All atoms of a given element are identical in their physical and chemical properties; they differ from atoms of every other element Atoms of different elements combine in simple whole-number ratios to form compounds (can form more than one compound together) Chemical reactions consist of the combination, separation, or rearrangement of atoms Which of these are no longer valid?

10 Can we see an atom? We can now view individual atoms! Scanning electron microscope (SEM)

11 J.J. Thomson s Experiments Used a cathode ray tube: metal is placed at the positive end (anode) and the negative end (cathode).

12 J.J. Thomson s Experiments Rays produced from the cathode end. The beam bends toward positive plates. A small paddle wheel spins when hit by the cathode rays.

13 Conclusions- Particles were bent by the charged plates particles are charged. Particles set the wheel in motion particles have mass. Particles were the same size no matter what metal was used particles are the same, no matter the element. Atoms are neutral, but are made of negatively charged particles atoms are made of negative and positive charge. What happened to the Dalton model?

14 The Plum Pudding Model electron An early and now obsolete attempt to describe the interior structure of atoms Electrons scattered throughout positively charged matter Electrically neutral sphere of positive charge

15 Rutherford's gold foil experiment Alpha particles (positively charged) bombarded foil of various metals. A fluorescent screen was placed around to detect the particles as they passed through the metal.

16 Animated Tutorial Animation Rutherford's expected vs. actual results Rutherford expected α-particles to pass undeflected through atoms. But, he observed that a small fraction of the α-particles were deflected Evidence that the positively charged part of the atom consisted of a tiny, dense object at the atom's center. He proposed the nuclear model of the atom.

17 Rutherford model- nuclear What is the problem with this model? Charge and mass of atom did not work out!

18 Chadwick- discovers neutrons Act as a kind of glue to hold the nucleus together. Positively charged protons are in a very confined space but shouldn t because they repel each other. Protons and neutrons are all attracted to each other as a result of another force - the strong nuclear force. The neutrons don't contribute any repulsive effects, so having more around can help to hold the nucleus together.

19 Subatomic particles: summary Name Location Charge Mass Discovered Proton Neutron Electron Chemical reactions involve changes Nuclear reactions involve changes

20 Subatomic Particles Mass of nucleus comes from the mass of protons and neutrons (= the nucleus).

21 The nuclear atom

22 How small is an atom? An atom is so small, a single water droplet contains about 5 sextillion(10 21 ) atoms Electrons are on the outside of the atom with very little mass. Most of the mass of the atom is in a central nucleus. Therefore, an atom is mostly empty space You can think of it as being like a marble in the middle of a football stadium. the marble is the nucleus-on the 50-yard line; spectators are the electrons.

23 Size of an atom

24 Atomic Number and Mass Number Atomic Number: number of protons in the nucleus of one atom - number of electrons in a neutral atom Mass Number: total number of protons and neutrons in an atom s nucleus.

25 Atomic Mass Average Atomic Mass: average mass of all known atoms of an element. Unit: amu (atomic mass unit)

26 Atomic Mass

27 Isotopes Naturally occurring isotopes Atoms of the same element that contain different numbers of neutrons. What is the chemical symbol? What is their atomic number? What is the mass number of the atom on the left?

28 Stable vs. Unstable isotopes Radioactive Isotopes: unstable atoms due to a nucleus with too many or too few neutrons No amount of neutrons can hold a nucleus together once it has more than 82 protons. Elements with an atomic number greater than 82 have unstable isotopes. Unstable atoms emit energy in the form of radiation when they break down (decay) Large nucleus (unstable) à nucleus + energy Reaction gives off LOTS of energy (= nuclear energy)

29 Discoveries lead to more about atomic theory 1890 s X-rays given off from anode when cathode operating (light energy) Radioactivity- α, β, γ, rays Quantization of Energy 1900 Max Planck. E = hv 1905 Light as a wave and particle Einstein's Ideas about Light

30 Electromagnetic spectrum

31 Waves If ν 1 = 4s -1 λ=?

32 Electromagnetic Spectrum Speed of light c = speed of light (3.0 x 10 8 m/s) Types of light energy: λ = wavelength ν = frequency E = energy c = λ ν

33 Electromagnetic spectrum

34 Diffraction grating/prism Note: A light bulb is an example of blackbody radiation (continuous spectrum). Most densely packed solids will emit a continuous spectrum when heated to a certain temperature.

35 Absorption or Emission of light The atom can absorb or emit light. Examples of absorption the color of shirt. Photosynthesis Examples of emission Gas discharge tubes Flame tests Neon lights Lamps

36 Excited Electrons and Spectra Line spectrum - can be used to identify an element it is a characteristic property of that element. Examples of practical use: determine the chemical make-up of the stars and plants atmospheres. FIREWORKS! SIMILAR CONCEPT TO OUR FLAME TEST Different metal will burn different colors. -What metallic elements do you think are in these fireworks?

37 Continuous vs. Line Spectrum

38 Hydrogen s line spectrum

39 Another great student Niels Bohr (student of Rutherford) Revised Rutherford s model to include newer discoveries about how an atom could absorb or emit light! Here s his thoughts: Electrons are found in distinct energy levels. This means electrons can t be found in-between these levels. Like Rutherford he proposed e- orbited the nucleus.

40 Bohr Model Electrons absorb energy and move to outer energy levels. When they relax, they give off energy. Your theory is crazy, but it's not crazy enough to be true. Niels Bohr

41 Quantum Theory vs. Classical Theory Quantum Theory

42 Bohr Model 6 5 Energy of photon 4 3 depends on the 2 difference in energy 1 levels Bohr s calculated energies matched the IR, visible, and UV lines for the H atom

43 Bohr Model Each element has a unique bright-line emission spectrum. Atomic Fingerprint Helium Bohr s calculations only worked for hydrogen! L Did not agree with classical physics. L

44 Electrons and energy An electron s P.E. & K.E both change when it relaxes (down orbital/s) or is in an excited state (up orbital/s) EXCITED STATE: Absorbs a photon or quantum of energy elevates to higher energy level GROUND STATE: Electrons in their lowest energy levels

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46 Atomic structure- Bohr model Energy level=n Lowest energy state is closest to nucleus-attracted to the protons When one energy level is filled, electrons are found at higher levels. Each energy level can hold a maximum number of electrons (2n 2 electrons) First shell = two electrons Second shell = eight electrons Third shell = eighteen electrons

47 Quantum Mechanical Model Electrons have properties of waves and light (De Broglie) It is impossible to know both the position and momentum of an electron (Heisenberg) The probability of finding in electron in a certain area around the nucleus. (Schrödinger) Sublevels- defined by energy level/distance from nucleus Orbitals- mathematical function corresponding to a region within atom each with a maximum of 2 e - with opposite spin

48 Quantum Mechanical Model Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of the atom.

49 Where is an electron? Heisenberg Uncertainty Principle It is impossible to know both the position and momentum of an electron.

50 S Orbitals

51 P orbitals

52 D orbitals

53 F orbitals

54 An orbital is a mathematical (3D) graph of the solution to the quantum mechanical wave equation. It defines a region of space that has a high probability of containing up to 2 e-. Movie visual

55 How do concepts of energy levels and orbitals fit together? Each energy level is made of 1 or more sublevels: Each sublevel is made of 1 or more orbitals:

56 Orbitals are filled from lowest to highest energy, in order of the periodic table

57 Electron Configurations Aufbau Principle Electrons fill from lowest energy to highest energy.

58 Electron Configurations Pauli Exclusion Principle Paired electrons must have opposite spins. Each orbital holds 2 electrons.

59 Electron Configurations Hund s Rule Electrons must be unpaired before they are paired in a sublevel. Make sure that everyone gets a helping! WRONG RIGHT

60 Abbreviated Configurations s d (n-1) p f (n-2) by Harcourt Brace & Company

61 Abbreviated Configurations Example - Germanium [Ar] 4s 2 3d 10 4p 2

62 Abbreviated Configurations s d (n-1) p f (n-2) by Harcourt Brace & Company

63 Chapter 6 THE PERIODIC TABLE

64 Periodic Table of Elements Sing-a-long

65 Names and symbols Symb ol Name H Hydrogen He Helium Li Lithium Be Beryllium B Boron C Carbon N Nitrogen O Oxygen F Fluorine Ne Neon Na Sodium Mg Magnesium Al Aluminum Si Silicon Symb ol Name S Sulfur Cl Chlorine Ar Argon K Potassium Ca Calcium Fe Iron Co Cobalt Cu Copper Zn Zinc Ag Silver Sn Tin I Iodine Au Gold Hg Mercury

66 Universe s elements

67 Earth s elements

68 Human Body Elements

69 Diatomic elements Label your PT

70 Metals, nonmetals, metalloids Label your PT

71 Periodic table organization Groups or families = column, similar chemical properties Alkali metals Alkaline earth metals Halogens Noble gases Period = row, chemical and physical trends repeat Other sections Transition metals Metalloids Metals Nonmetals Lanthanide and actinide series (inner transition metals or rare earth)

72 Trends based on # of electrons Groups (columns) Elements in the same group have similar properties; why? They all have the same # of outer electrons= VALENCE ELECTRONS- Use the periodic table note valence electrons Periods (rows) Elements in a period have valence electrons in the same outer energy level. They all have the same # of inner electrons= CORE ELECTRONS- Use the periodic table note energy levels

73 Physical properties of elements Physical state: gas, solid, liquid Label your PT Conductivity: Conductor, semiconductor Physical qualities: