Honors Ch3 and Ch4. Atomic History and the Atom

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1 Honors Ch3 and Ch4 Atomic History and the Atom

2 Ch. 3.1 The Atom is Defined 400 B.C. the Greek philosopher Democritus said that the world was made of two things: Empty space and tiny particles called atoms

3 Ch3.1 What are atoms? Atoms are the smallest part of an element that still has the element s properties.

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5 By 1700 s chemistry was defined by 3 Laws: Law of the Conservation of Mass Law of the definite Proportions Law of Multiple Proportions

6 Law of Conservation of Mass

7 Law of Definite Proportions/Composition Substances contain atoms in the same ratio of mass.

8 Law of Multiple Proportions

9 Early 1800 s John Dalton Came up with first atomic theory that is the basis for today s theory.

10 Early 1800 s John Dalton s Theory-proven 1. Every element is made of tiny, unique particles called atoms These atoms cannot be destroyed, but instead rearrange during a chemical change 2. Atoms of different elements can join to form molecules in constant whole number ratios.

11 John Dalton s Theory-disproved Atoms cannot be broken down into smaller particles. Atoms of the same element are exactly alike in mass

12 Ch JJ Thomson Used a cathoray tube to examine if atoms were made of charged particles.

13 Opposite charges attract

14 1897 JJ Thomson Discovered atoms are made up of particles with negative charges, but little mass. Called them electrons.

15 Thomson Model Electrons Positive charges, not known as protons yet

16 1911-Rutherford Put together a team of physicists to performed the gold foil experiment

17 1 out of 8000 alpha particles were repelled.

18 1913-Rutherford The experiment led to the discovery that atoms are mostly empty space. (expected) It also discovered that atoms contain a positive dense nucleus which contained most of the mass of the atom.

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20 Rutherford Model

21 Inside the Nucleus Particles were discovered. Positive protons (1836 times more massive than electrons) These practices identified the type of atom. Neutral neutrons (1837 times more massive than electrons) Kept the protons from repelling by producing strong nuclear forces

22 Ch 3.3 Parts of an atom Nucleus Proton Neutron Cloud Subatomic Particles Charged Particles Electron

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24 Nucleus center of an atom positively charged makes up 99.9% of the atom s mass

25 Protons Charge (+) Mass is equal to 1 atomic mass unit (u) 1/12 mass of Carbon atom

26 Neutrons Charge (0 net) - neutral Mass is equal to 1amu Determine stability of nucleus

27 Electrons Charge is negative (-) Mass is equal to amu

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29 Atomic Number Identifies # of protons Determines the type of atom because no two elements can have same # of protons.

30 Mass Number Mass of a single atom # of protons # of neutrons Mass #

31 Isotopes Any atoms having the same number of protons but different number of neutrons. thus they have different mass numbers.

32 Springfield Isotopes

33 Isotopes

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35 Average Atomic Mass Atomic mass is the mass of all isotopes of a particular element averaged together Calculating Average Mass is based on the abundance of each isotope

36 Atoms vs Ions All atoms have the same number of protons and electrons. They are neutral. Charges cancel each other out.

37 Atom vs Ions Ions are charged particles. Form when atoms lose or gain electrons. Form in order to have a full outer shell Two Types.

38 Cations Positively charged ions. Form when atoms lose electrons. Form from metal atoms

39 Cations # of protons greater than # of electrons More (+) than (-)

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41 Na Atom Na + Cation

42 Anions Negatively charged ions. Form when atoms gain electrons. Form from nonmetal atoms

43 Anions # of protons less than # of electrons More (-) than (+)

44 Cl atom Cl - Anion

45 Ch4.1 The Duality of Light Led to a New View of the Atom Light has characteristics of both waves and particles All forms of radiation travel at the same maximum speed of 3.00 x10 8 m/s

46 Wave description Different forms of light are defined by their unique wavelengths (ƛ) and frequencies (v) Speed of light (c)= ƛ v

47 Electromagnetic Spectrum Electromagnetic spectrum includes light at all possible frequencies and wavelengths wavelength, frequency

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50 Gamma Rays highest energy and frequency. Nuclear radiation

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54 X-Rays high ionizing energy radiation Used for imaging

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56 Ultraviolet Light (UV) Ionizing energetic radiation Can cause skin cancer when over exposed

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59 Sun Protection (SPF) 7 or less no protection 8 Extra protection but still permits tanning 15 Offers total protection from burning 30 Totally blocks UV

60 ROYGBIV Visible spectrum We see red and orange the best Blue and violet are the hottest colors, and emits the most energy, red the least and coldest White is all the colors combined, black the absence

61 Infrared Light Lower energy radiation Night vision

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65 Microwaves Low energy radiation Used to heat food Also used in telecommunication

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68 Radio Waves Have the Lowest frequencies and highest wavelengths Includes FM, AM, and TVs Radar

69 REALITY TV

70 The Particle Description The Photoelectric effect: Emission of electrons when certain light hits metal Einstein theorized that light can be modeled as a photons (particles of light) Each photon carries an unique quantum of energy

71 Maxwell Planck Theorized that the energy given off by a photon is directly related to the frequency of the radiation emitted. His equation Energy (E) = hv h = planck s constant = x J s

72 Ch s Niels Bohr Suggested that electrons move around nuclei in set paths around the nucleus. (solar system)

73 s Niels Bohr He said each path is a calculated energy level Atom s electrons can jump to different energy levels when absorbing photons

74 Niels Bohr States of Atoms Ground State- An electron s lowest energy state or level Excited State- An electron that is energized will jump to a higher energy state or level. Will last for a short period. (resulting light production)

75 States of Atoms? Ground State Excited state

76 This is how light is produced Ground State Excited state back to ground And Light production

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79 Niels Bohr Bohr noticed that different atoms emitted different radiation when excited.

80 Planck s Flame Test

81 Spectroscopy

82 Emission (Line) Spectrum Atoms and molecules are identified by these spectrums.

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84 What the Heck is Light?

85 Ch4.2 Quantum Theory French Physicist De Broglie Wave-Particle Duality of electrons He proposed that electrons can only exist at certain frequencies, hence the energy they release when excited

86 Today s Theory Werner Heisenberg Uncertainty Principle It is impossible to determine an electron s exact position and speed at the same time. Along with De Brogile, they disproved Bohr s definite orbit assumption

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88 Schrödinger s Work Electrons found in orbitals within different energy levels. a calculated region in an atom where there is a high probability of finding electrons. This led to the Electron Cloud model and Quantum Numbers

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91 Modern Atomic Cloud Model

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93 History Review

94 Quantum # s Principal (n) = number of energy level (1-7) Angular momentum (l) = sublevels: 0,1,2,or 3 Magnetic (m) = the orbital # (-,0,+) Spin (s) = electron spin (1/2, -1/2) Pauli Exclusion Principle No two atoms will have the same configuration or set of quantum # s

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96 Energy levels 1 st level holds 2 e - (s) 2 nd level holds 8 e - (s,p) 3 rd level holds 8 or 18e - (s,p,d) 4 th level holds 18 or 32e - (s, p d, f) Outer (valence) level holds up to 8 e- (s, p)

97 Making Bohr Models

98 Bohr s Model Blue dots represent electrons Rings represent energy Level, NOT orbit (path)

99 ELECTRON CONFIGURATIONS Ch4.2 Electron Energy Levels, SUBLEVELS, and Orbital's

100 Electron Configuration The rows (periods) of the periodic table tell you the energy level There are 4 sublevels s,p,d,f

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102 Electron Configuration Each sublevel contains a different number of orbitals. S =1, p = 3, d= 5, and f = 7 Orbital is represented by: Each orbital holds only two electrons with opposite spin.

103 Writing Atomic Electron Configurations Three ways of writing configurations. 1. orbital box notation. 2. Electron Configuration (spdf) Notation 3. Noble Gas Configuration

104 Orbital Notation Rules: 1. Aufbau Principle- electrons are added one at a time starting at lowest energy level 2. Hund s Rule - When filling orbitals with in the same sublevel, fill one electron into each box before pairing electrons.

105 Orbital Notation ORBITAL BOX NOTATION for He, atomic number = s 1s Arrows depict electron spin

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107 3. Hund s Rule 2P 1.

108 S,p,d,f Notation for H, atomic number = 1 1 s 1 no. of electrons value of n sublevel

109 Phosphorus Group 5A Atomic number = 15 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne] 3s 2 3p 3 3s 2s 3p 2p 1s

110 Aluminum Group 3A Atomic # = 13 1s 2 2s 2 2p 6 3s 2 3p 1 [Ne] 3s 2 3p 1 All Group 3A elements have [core] ns 2 np 1 configurations where n is the period number. 3s 2s 1s 3p 2p

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