Presentation is loading. Please wait.

Presentation is loading. Please wait.

Atoms, Molecules, and Ions HOMEWORK n Read your textbook pages 74 to 78 and answer q 1 to 6 n Start a timeline in your notebook of the discoveries that.

Similar presentations


Presentation on theme: "Atoms, Molecules, and Ions HOMEWORK n Read your textbook pages 74 to 78 and answer q 1 to 6 n Start a timeline in your notebook of the discoveries that."— Presentation transcript:

1

2 Atoms, Molecules, and Ions

3 HOMEWORK n Read your textbook pages 74 to 78 and answer q 1 to 6 n Start a timeline in your notebook of the discoveries that lead to the present theory of the atom. Include the date and name of the scientist mentioned in the book and the contribution.

4 HISTORY n Democritus 400 BC suggested that matter is made up of indivisible particles called atoms.

5 The Law of Conservation of Mass n Established in 1789 by French Chemist Antoine Lavoisier n States that mass is neither created nor destroyed in any chemical reaction. n Or more simply, the mass of substances produced (products) by a chemical reaction is always equal to the mass of the reacting substances (reactants).

6 LAW OF DEFINITE PROPORTIONS n Every compound is a combination of a fixed proportion by mass between the elements that compose the sample. n Water is always 1 g of H and 8 g of O n CO2 is always 12 g of C and 32 of O

7 LAW OF MULTIPLE PROPORTIONS n When the same elements combine to form different compounds, they do so in mass ratios that can be expressed by small whole numbers. n CO carbon monoxide 12 g C 16 g O n CO2 carbon dioxide 12 g C 32 g O

8 First Atomic Theory Dalton ( ) n All elements are composed of indivisible atoms. n Atom of the same elements are identical. n Atoms of different elements differ in their physical and chemical properties. n Atoms of different elements combine to form compounds in simple whole number ratios. n Chemical reactions occur when atoms are separated, joined or rearranged, but they can not change.

9 Experiments to determine what an atom was n J. J. Thomson (1897)- used Cathode ray tubes

10 Thomson’s Experiment Voltage source +-

11 Thomson’s Experiment Voltage source +-

12 n Passing an electric current makes a beam appear to move from the negative to the positive end. Thomson’s Experiment Voltage source +-

13 Thomson’s Experiment n By adding an electric field

14 Voltage source Thomson’s Experiment n By adding an electric field, he found that the moving pieces were negative + -

15 n OBJECTIVE : WHAT IS INSIDE THE ATOM? n Experiments that lead to find what is inside the atom n Rutherford gold foil experiment – n Millikan

16 Thomsom’s conclusions n Thomson named the cathodic rays electrons and concluded that they must be a part of atoms of all elements.

17 The Atom, circa 1900: n “Plum pudding” model, put forward by Thompson. n Positive sphere of matter with negative electrons imbedded in it.

18 The Gold foil experiment n In 1911 Ernest Rutherford designed an experiment using alpha particles and thin gold foil - “gold foil experiment”. Rutherford was very surprised by his finding and his conclusions led to which became known as the “planetary” model of the atom.

19 Rutherford’s Experiment n Used uranium to produce alpha particles. n Aimed alpha particles at gold foil by drilling hole in lead block. n Since the mass is evenly distributed in gold atoms alpha particles should go straight through. n Used gold foil because it could be made atoms thin.

20 Lead block Uranium Gold Foil Florescent Screen

21 What he got

22 n Review of Rutherford Experiment n Subatomic particles

23 Discovery of the Nucleus Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

24 How he explained it + n Atom is mostly empty n Small dense, positive piece at center. n Alpha particles are deflected by it if they get close enough.

25 +

26 The Nuclear Atom Since some particles were deflected at large angles, Thompson’s model could not be correct.

27 Observations n Most of the alpha particles pass straight through the gold foil. n Some of the alpha particles get deflected by very small amounts. n A very few get deflected greatly. n Even fewer get bounced of the foil and back to the left.

28 Conclusions n The atom is 99.99% empty space. n The nucleus contains a positive charge and most of the mass of the atom. n The nucleus contains a positive charge and most of the mass of the atom. n The nucleus is approximately 100,000 times smaller than the atom.

29 Vocabulary n Deflected: curved or bent downward. n To Scatter: to separate and drive off in various directions; to disperse n To Bounce: To rebound, move up and down repeatedly.

30 Other Subatomic Particles n Protons were discovered by Rutherford in n Neutrons were discovered by James Chadwick in 1932.

31 The Bohr Model n Niels Bohr proposed a model of the atom based on the solar system. n It was wrong but some of his ideas we still used. It was one step in the direction of the last model for the atom. n First let’s take a look at what is inside the atom.

32 Subatomic Particles n Protons and electrons are the only particles that have a charge. n Protons and neutrons have essentially the same mass. They are located inside the nucleus. n The mass of an electron is so small we ignore it.

33 Atomic Number (Z) n The number of protons in the nucleus. It identifies the element. In the neutral atom n # protons = # electrons

34 Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

35 Mass Number (A) * It is not in the periodic table but is the atomic mass rounded off n Protons plus neutrons in the atom n A= # protons + # neutrons n *In the table we can find the Atomic Mass. n Number of Neutrons =Mass number-Atomic number Number of Neutrons Number of Neutrons

36 ISOTOPIC SYMBOL Atomic number (Z) = number of protons in the nucleus. Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Z A X By convention, for element X, we write Z A X

37 Isotopes, Atomic Numbers, and Mass Numbers Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Z A X. Isotopes have the same Z but different A. We find Z on the periodic table.

38 Mass Number The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

39 Mass Number (A) * It is not in the periodic table n Protons plus neutrons in the atom n A= # protons + # neutrons n *In the table we can find the Atomic Mass. n Number of Neutrons =Mass number-Atomic number Number of Neutrons Number of Neutrons

40 Isotopes n atoms with same number of protons and different number of neutrons. Same atomic number different mass number. n SAME ELEMENT with different # of neutrons!

41 Isotopes: n Atoms of the same element with different masses. n Isotopes have different numbers of neutrons C 12 6 C 13 6 C 14 6 C

42 Atomic Mass Unit (amu) is one-twelfth the mass of a carbon-12 atom.

43 Atomic Mass n is the weighted average of the mass of all the isotopes of one element. n It takes into consideration the mass of the different isotopes of the element and their natural abundance

44 Hw questions 1) 3 2) 2 3) 4 4) 4 5) 4 6) 2 7) 2 8) 4 9) 3 10) 4 11) 2 n 11) 2 n 12) 4

45 Temperature review 17) 2 18) 4 19) 1 20) 2 21) 4 22) 3 23) 4 24) 1 25) 1 26) 3 27) 1 28) 3 29) 3

46 DECEMBER 13 n The Bohr Atom n Energy levels – shells n Electron configuration

47 Planetary Model n Neils Bohr 1913 n Electrons are arranged in orbits around the nucleus. n He proposed that electrons in a particular orbit have a fixed energy. The electrons cannot fall into the nucleus.

48 Bohr’s Atom electrons in orbits nucleus

49 Energy level n is the region around the nucleus where the electron is moving. n As the electrons are further away from the nucleus they have more energy. The closest orbits or energy levels to the nucleus have low energy.

50 n He named the orbits with letters beginning with k,l,m,n… n The first energy level could hold 2 electrons. n Second energy level could hold up to 8 electrons

51 Electron Configurations n Indicates how the electrons are located in the atom. n Niels Bohr proposed that electrons are located in energy levels at different distances from the nucleus. n total of electrons=13 n 3 energy levels

52 HELIUM ATOM + N N proton electron neutron Shell What do these particles consist of?

53 ATOMIC STRUCTURE Particle proton neutron electron Charge + ve charge -ve charge No charge 1 1 nil Mass

54 ATOMIC STRUCTURE MASS NUMBER the number of protons and neutrons in an atom the number of protons in an atom He 2 4 number of electrons = number of protons Atomic number

55 ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shella maximum of 2 electrons second shella maximum of 8 electrons third shella maximum of 8 electrons

56 SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.

57 n n A closer look at the Bohr model n n How did Bohr came to the conclusion that the electron moved in orbits? n n TEXTBOOK PAGES 90 TO 94

58 Bohr Model n 1 An electron in a permitted orbit has a specific energy an is in an “allowed” energy state. It will not spiral into the nucleus n 2 Energy is emitted or absorbed by the electron only as the electron changes from one allowed state to other

59 Energy levels or shells n n Electrons with the lowest energy are found in the energy level closest to the nucleus n n Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

60 The Nature of Energy A mystery involved the emission spectra observed from energy emitted by atoms and molecules. When gases at low pressure were placed in a tube and were subjected to high voltage, light of different colors appeared

61 Continuous Spectra Radiation composed of only one wavelength is called monochromatic. Radiation that spans a whole array of different wavelengths is called continuous. White light can be separated into a continuous spectrum of colors. Note that there are no dark spots on the continuous spectrum that would correspond to different lines. Line Spectra and the Bohr Model

62

63 Line Spectra If high voltage is applied to atoms in gas phase at low pressure light is emitted from the gas. If the light is analyzed the spectrum obtained is not continuous. SPECTROSCOPE

64 Line Spectra. When the light from a discharge tube is analyzed only some bright lines appeared.

65 Waves n n Waves carry energy and have cycles. Each cycle of a wave begins at the origin and ends in the origin. The wavelength (  lambda) of a wave is a measurement of how long it is. It is measured in length units. n n The longer the wavelength, the smaller amount of energy that is associated with. n n * Short wavelength, high energy. n n * Long wavelength, low energy.

66 Waves Light consists of electromagnetic waves

67

68 How is Light Energy related to Color? n n White light is made up of all colors of the spectrum (ROYGBIB) n n When passed through a prism a continuous spectrum (rainbow effect) is obtained

69 Nature of light n n Light is a form of electromagnetic wave. Each color of the visible light is associated with a different wavelength. n n R (700 nm) n n O n n Y n n G n n B n n I n n V (400nm)

70 Line Spectra If high voltage is applied to atoms in gas phase at low pressure light is emitted from the gas. If the light is analyzed with an SPECTROSCOPE the spectrum obtained is not continuous.

71 Line Spectra n n This type of spectra is used to identify atoms of different elements and were used to explain the new model of the atom.

72 Bright line spectrum n n When an atom ABSORBS energy, the electrons JUMP (LEAP) to a higher energy level. They are in the EXCITED STATE. n n When the electrons return to the normal level (GROUND STATE), they emit energy as LIGHT. The light emitted produce the bright line spectrum that is characteristic of each element.

73 Electrons Surround the Nucleus n n Electrons surround the nucleus and travel at the speed of light n n They are found in only certain allowed energy levels or orbitals n n Electrons are at the lowest energy level – the ground state.

74 Different Energy States are Possible n n When the electrons in an atom become excited by absorbing energy from the surroundings, they jump to new higher energy levels. n n The excited state is less stable than the lower energy state therefore the electron falls back or returns to the lower energy ground state.

75 Return to Ground State n n When electrons fall back a wave with a specific amount of energy is emitted – called a photon. If we see color then the emission is in the visible range n n We perceive this as unique colors associated with particular elements

76 Ground state configuration vs excited state configuration n n 11 Na This is the ground state. n n If atom is excited some possible electron configuration would be n n or or n n The number of electrons must add to the atomic number but the order is not the one that appears in the periodic table.

77 The wave-mechanical model of the atom

78 The Wave Model n n Today’s atomic model is based on the principles of wave mechanics. n n According to the theory of wave mechanics, electrons do not move about an atom in a definite path, like the planets around the sun.

79 Quantum Model

80 n n The modern model of the atom is called the quantum model or the wave- mechanical model.

81 n n By 1900 scientist were studying energy and waves. n n It was proposed that energy was made of tiny packets called quanta, these packets acted like particles. n n That implied that light could behave as a wave and also as a particle.

82 n n Quantum – singular – package or bundle or energy. If the energy is light is a photon. n n Quanta is the plural of quantum.

83 Duality wave -particle n n Louis De Broglie discovered that the electron also behaves sometimes as a particle and sometimes as a wave

84 Orbitals n n In the wave-mechanical model the electrons are found not in fixed orbits around the nucleus but in ORBITALS. n n ORBITALS are regions in which an electron of a particular amount of energy is most likely to be found.

85 Probability n n The wave mechanical model describes the atom using PROBABILITY of finding electrons at different distances from the nucleus.

86 Quantum theory n n In 1926 Erwin Schrodinger used the quantum theory to write an equation describing the location and energy of the electron in the atom of hydrogen.

87 December 20 Quiz – DO NOW For an atom of C Name of element n n 2. Number of proton n n 3. Number of electron n n 4. Number of neutron n n 5. Electronic configuration n n 6. Isotopic representation n n 7. One possible excited state

88 5 th Solvay Conference of Electrons and Photons

89 The quantum model principles n n Electrons are found in orbitals. n n There are different energy levels. Electrons can gain or lose fixed amounts of energy ( quanta ) to move to different energy levels. n n When the electrons occupy the lowest available orbitals the atom is in the ground state. n n If the atom absorbs a quantum of energy the electrons jump to a higher energy level. This UNSTABLE CONDITION is called EXCITED STATE. n n When the electrons return to the ground state they emit the same amount of energy they absorbed as LIGHT of different colors. n n The emitted light can be analyzed with a spectroscope and the resulting spectrum can be used to identify the atoms.

90 Energy levels n n A space in which electrons are likely to be found. n n Electrons whirl about the nucleus billions of times in one second n n They are not moving around in random patterns. n n Location of electrons depends upon how much energy the electron has.

91 Energy levels or shells n n Electrons with the lowest energy are found in the energy level closest to the nucleus n n Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

92 Modern Atomic Theory n The atom is mostly empty space. n Two regions n Nucleus- protons and neutrons. n Electron cloud- region where you might find an electron.

93 SET 1 & 2 ATOMIC SET ANSWERS n 1 -1 n n n 4 – 2 n 5 – 3 n 6 – 4 n 7 – 4 n n 9 – – n 11 – n 12 – n 13 do not do 35-1 n n 29-3 n 30-DND

94 n n What happens when the number of protons and the number of electrons in an atom are not equal? n n IONS

95 Nuclear Charge n n Is the number of protons!

96 n Atoms are neutral and contain same number of electrons than protons. In a chemical reactions atoms can lose, gain or share electrons. n When an atom loses an electron it becomes a positive ion. When it gains an electron it becomes a negative ion.

97 n The nucleus always remains the same in a chemical reaction. n The number of protons and neutrons never change.

98 Ion n n A charged particle n n When an atom loses or gains electrons it becomes an ION. n n Remember protons can not be touched! n n If an atom loses electron the ion will be positive (more protons than electrons) n n If an atom gains electrons it will became negative (more electrons than protons)

99 n Metals lose electrons and become positive ions (caTions) n Non metals gain electrons and become negative ions ( aNions)

100 IndivisibleElectronNucleusOrbitElectron Cloud Greek X Dalton X Thomson X Rutherford X X Bohr X X X Wave X X X


Download ppt "Atoms, Molecules, and Ions HOMEWORK n Read your textbook pages 74 to 78 and answer q 1 to 6 n Start a timeline in your notebook of the discoveries that."

Similar presentations


Ads by Google