1. Atomic Theory

2. Democritus The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) His atomic model was called the “solid sphere” model His atomic model was called the “solid sphere” model He believed that atoms were indivisible and indestructible He believed that atoms were indivisible and indestructible He named the smallest piece of matter “atomos,” meaning “not to be cut.” He named the smallest piece of matter “atomos,” meaning “not to be cut.”

3. Democritus Based on his theory: Atoms were infinite in number, always moving and capable of joining together Atoms were infinite in number, always moving and capable of joining together Atoms were small, hard particles that were all made of the same material but were different shapes and sizes. Atoms were small, hard particles that were all made of the same material but were different shapes and sizes. Problems with the theory: Problems with the theory: did not explain chemical behavior did not explain chemical behavior was not based on the scientific method – but just philosophy was not based on the scientific method – but just philosophy

4. John Dalton 1766 - 1844 1766 - 1844 All matter is made of atoms. All matter is made of atoms. Atoms of an element are identical. Atoms of an element are identical. Each element has different atoms. Each element has different atoms. Atoms of different elements combine in constant ratios to form compounds. Atoms of different elements combine in constant ratios to form compounds. Atoms are rearranged in reactions. Atoms are rearranged in reactions. Was called the “billiard ball” or “solid sphere” model Was called the “billiard ball” or “solid sphere” model

5. John Dalton Improved on Democritus’s theory Improved on Democritus’s theory Dalton’s theory was based on scientific experimentation Dalton’s theory was based on scientific experimentation His ideas account for the law of conservation of mass - atoms are neither created nor destroyed His ideas account for the law of conservation of mass - atoms are neither created nor destroyed Agree with the law of constant composition - elements combine in fixed ratios Agree with the law of constant composition - elements combine in fixed ratios

6. John Dalton Problems with Dalton’s Theory: Problems with Dalton’s Theory: Discoveries were made that showed that atoms are divisible into smaller subatomic particles: Discoveries were made that showed that atoms are divisible into smaller subatomic particles: Electrons, protons, and neutrons Electrons, protons, and neutrons There are many other particles There are many other particles

7. J.J. Thomson Called the “plum pudding” model Called the “plum pudding” model Provided the first hint that an atom is made of even smaller particles. Provided the first hint that an atom is made of even smaller particles.

8. Thomson Model He proposed that atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding. He proposed that atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding.

9. Thomson Model Thomson studied the passage of an electric current through a gas. As the current passed through the gas, it gave off rays of negatively charged particles.

10. J. J. Thomson Improvements: Improvements: Discovered particles smaller than the atom exist, therefore, the atom was divisible Discovered particles smaller than the atom exist, therefore, the atom was divisible He also concluded that there must be positively and negatively charged particles within the atom. He also concluded that there must be positively and negatively charged particles within the atom. Discovered the electron Discovered the electron Problems: Problems: Could not discover the proton Could not discover the proton

11. Ernest Rutherford’s Gold Foil Experiment - 1911 Positively charged alpha particle were fired at a thin sheet of gold foil. Some of the positively charged particles bounced away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges.

12. Rutherford’s Findings a) The nucleus is small b) The nucleus is dense c) The nucleus is positively charged d) The electrons move around the nucleus like planets orbit the sun. e) Called the nuclear model or the planetary model Problems: Based on classical physics, later abandoned because of the discoveries in quantum physics. Conclusions/Improvements:

13. Bohr Model In 1913, the Danish scientist Niels Bohr proposed an improvement to Rutherford’s model. His model is sometimes referred to as the Rutherford-Bohr model. In 1913, the Danish scientist Niels Bohr proposed an improvement to Rutherford’s model. His model is sometimes referred to as the Rutherford-Bohr model.

14. Bohr’s Model Electrons orbit the nucleus in orbits that have a set size and energy. Electrons orbit the nucleus in orbits that have a set size and energy. The energy of the orbit is related to its size. The lowest energy is found in the smallest orbit. The energy of the orbit is related to its size. The lowest energy is found in the smallest orbit. Radiation is absorbed or emitted when an electron moves from one orbit to another. Radiation is absorbed or emitted when an electron moves from one orbit to another. Problems: Problems: We cannot predict the exact location and orbit of electrons. We cannot predict the exact location and orbit of electrons. Predictions fail to work on larger atoms. Predictions fail to work on larger atoms.

15. James Chadwick  1932 – confirmed the existence of the neutron – a particle with no charge, but a mass nearly equal to a proton  He was trying to identify the extra mass of an atomic nucleus by firing alpha particles into a beryllium target and allowing the resulting radiation to interact with paraffin wax.  The interactions between the radiation and the hydrogen in the wax led to the discovery of the neutron.

16. James Chadwick Improvements: Improvements: With the discovery of the neutron, the atomic model seemed more complete than ever. With the discovery of the neutron, the atomic model seemed more complete than ever. The overall charges remained the same The overall charges remained the same Now there no longer seemed to be a discrepancy between the atomic mass and the atomic number. Now there no longer seemed to be a discrepancy between the atomic mass and the atomic number.

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

18. Electron Cloud Depending on their energy they are locked into a certain area in the cloud. Depending on their energy they are locked into a certain area in the cloud. Electrons with the lowest energy are found in the energy level closest to the nucleus Electrons with the lowest energy are found in the energy level closest to the nucleus Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus. Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

19. Atomic Structure

20. An atom consists of: nucleus – contains protons and neutrons nucleus – contains protons and neutrons electrons in space around the nucleus. electrons in space around the nucleus. The atom is mostly empty space The atom is mostly empty space The Atom Nucleus Electron cloud

21. ATOMIC COMPOSITION Protons (p + ) Protons (p + ) + electrical charge + electrical charge mass = 1.672623 x 10 -24 g mass = 1.672623 x 10 -24 g relative mass = 1.007 atomic mass units (amu) but we can round to 1 relative mass = 1.007 atomic mass units (amu) but we can round to 1 Electrons (e - ) Electrons (e - ) negative electrical charge negative electrical charge relative mass = 0.0005 amu but we can round to 0 relative mass = 0.0005 amu but we can round to 0 Neutrons (n o ) Neutrons (n o ) no electrical charge no electrical charge mass = 1.009 amu but we can round to 1 mass = 1.009 amu but we can round to 1

22. Atomic Number, Z All atoms of the same element have the same number of protons in the nucleus, Z 13 Al 26.981 Atomic number Atom symbol AVERAGE Atomic Mass

23. Atomic Number Atoms are composed of identical protons, neutrons, and electrons Atoms are composed of identical protons, neutrons, and electrons How then are atoms of one element different from another element? How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS Elements are different because they contain different numbers of PROTONS The “atomic number” of an element is the number of protons in the nucleus The “atomic number” of an element is the number of protons in the nucleus # protons in an atom = # electrons # protons in an atom = # electrons

24. Atomic Number Element # of protons Atomic # (Z) Carbon66 Phosphorus1515 Gold7979

25. Mass Number, A A carbon atom with 6 protons and 6 neutrons is the mass number = 12 atomic mass units A carbon atom with 6 protons and 6 neutrons is the mass number = 12 atomic mass units Mass Number (A) = # protons + # neutrons Mass Number (A) = # protons + # neutrons NOT on the periodic table…(Round the AVERAGE atomic mass on the table) NOT on the periodic table…(Round the AVERAGE atomic mass on the table) A boron atom can have A = 5 p + 5 n = 10 amu A boron atom can have A = 5 p + 5 n = 10 amu

26. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p + + n 0 Nuclide p+p+p+p+ n0n0n0n0 e-e-e-e- Mass # Oxygen - 10 -3342 - 31 - 3115 8 8 18 Arsenic 7533 75 Phosphorus 15 31 16

27. Complete Symbols/ Nuclear Notation Contain the symbol of the element, the mass number and the atomic number. Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number Subscript → Superscript →

28. Symbols n Find each of these: a) number of protons b) number of neutrons c) number of electrons d) Atomic number e) Mass Number Br 80 35

29. Symbols n If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons b) number of neutrons c) number of electrons d) complete symbol

30. Symbols n If an element has 91 protons and 140 neutrons what is the a) Atomic number b) Mass number c) number of electrons d) complete symbol

31. Symbols n If an element has 78 electrons and 117 neutrons what is the a) Atomic number b) Mass number c) number of protons d) complete symbol

32. Isotopes

33. Isotopes Dalton was wrong about all elements of the same type being identical Dalton was wrong about all elements of the same type being identical Atoms of the same element can have different numbers of neutrons. Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. Thus, different mass numbers. These are called isotopes. These are called isotopes.

34. Isotopes Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.

35. Naming Isotopes We can also put the mass number after the name of the element: We can also put the mass number after the name of the element: carbon-12 carbon-12 carbon-14 carbon-14 uranium-235 uranium-235

36. Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) (protium)110 Hydrogen-2(deuterium)111 Hydrogen-3(tritium)112

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

38. Isotopes & Their Uses Bone scans with radioactive technetium- 99.

39. Isotopes & Their Uses The tritium content of ground water is used to discover the source of the water, for example, in municipal water or the source of the steam from a volcano.

40. Atomic Mass  How heavy is an atom of oxygen?  It depends, because there are different kinds of oxygen atoms.  We are more concerned with the average atomic mass.  This is based on the abundance (percentage) of each variety of that element in nature.  We don’t use grams for this mass because the numbers would be too small.

41. Measuring Atomic Mass Instead of grams, the unit we use is the Atomic Mass Unit (amu) Instead of grams, the unit we use is the Atomic Mass Unit (amu) It is defined as one-twelfth the mass of a carbon-12 atom. It is defined as one-twelfth the mass of a carbon-12 atom. Carbon-12 chosen because of its isotope purity. Carbon-12 chosen because of its isotope purity. Each isotope has its own atomic mass, thus we determine the average from percent abundance. Each isotope has its own atomic mass, thus we determine the average from percent abundance.

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

43. Atomic Masses IsotopeSymbol Composition of the nucleus % in nature Carbon-12 12 C 6 protons 6 neutrons 98.89% Carbon-13 13 C 6 protons 7 neutrons 1.11% Carbon-14 14 C 6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon = 12.011

44. Average Atomic Mass weighted average of all isotopes weighted average of all isotopes on the Periodic Table on the Periodic Table round to 2 decimal places round to 2 decimal places Avg. Atomic Mass

45. Avg. Atomic Mass D. 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. 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. 16.00 amu

46. Avg. Atomic Mass D. Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine- 35 and 2 are chlorine-37. EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine- 35 and 2 are chlorine-37. 35.40 amu