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Atomic Theory

Atomic Theory. History of Atomic Theory. The Greeks - Democritus. 450 B.C. Believes there is a fundamental particle that cannot be broken. The particle is invisible. Coins the term “atom”. John Dalton. 1810 Believes atoms are invisible and indivisible.

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Atomic Theory

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  1. Atomic Theory

  2. History of Atomic Theory

  3. The Greeks - Democritus • 450 B.C. • Believes there is a fundamental particle that cannot be broken. • The particle is invisible. • Coins the term “atom”

  4. John Dalton • 1810 • Believes atoms are invisible and indivisible. • All atoms of the same element are alike in every respect –especially mass. • Atoms of different elements are different in every respect – especially mass. • Billiard Ball Theory

  5. J.J. Thomson • 1897 • An atom consists of a positively charged globule. • Negative particles, called electrons, are embedded in the globule. • These electrons move about in the globule. • Jellied salad or raisin bun theory.

  6. Positive mass electrons

  7. Ernest Rutherford • 1908 • Performs his famous scattering experiment

  8. Observations Most alpha particles pass through the foil as if there was nothing there A few veer off as they are slightly deflected A very few bounce back Conclusions Atoms are mainly empty space (over 99.9 %) The mass of the atom is concentrated in a core or nucleus which is positively charged. Electrons, which are negative move about in a large region of space outside the nucleus

  9. Rutherford Model Negative Electron Positive Nucleus

  10. Quantum Mechanical Model OR the Nuclear Atom Modern Atomic Theory 1. Atoms have a very dense core called the nucleus.The nucleus contains two sub-atomic particles: A) Protons which have a positive charge B) Neutrons which do not carry a charge (neutral) 2. The electrons, which carry a negative charge, are in a state of constant motion in a relatively large region of space away from the nucleus. 3. Atoms are mainly empty space. (over 99.9 %) I. General Information

  11. II. Summary of Atomic Particles H or p+ 1 Nucleus +1 n or n 1 Nucleus 0 e Constant motion outside nucleus or e-1 0 -1

  12. III. Atomic Number • This is used to identify the element (name and / or symbol) • This number ALWAYS tells us the how many protons are in the nucleus of atoms of that element. • NO EXCEPTIONS • Examples: All carbon atoms (atomic # of 6) have 6 protons in the nucleus. All atoms with 11 protons are sodium atoms and have an atomic number of 11. ….etc. etc. etc……

  13. IV. Nuclear Charge • This is the total positive charge present on the nucleus. • It is always positive because the protons are the only charged particle in the nucleus. • Therefore, the nuclear charge is equal to the number of protons in the atom or ion. (equal to the atomic number) • Examples: The nuclear charge on a potassium atom (atomic number of 19) is +19. An atom has 5 protons. What nuclear charge does it have? Answer: +5 …..etc. etc. etc…….

  14. V. Mass Number • The mass number is the total mass in grams for a constant and specific number of atoms of each element. • This number is very large (because atoms are very small) and is known as the mole number. • 1.0 mole is equal to 6.02 x 1023 particles. Examples: The mass of one mole of Hydrogen atoms is 1.01 g. Therefore 1.01 g/mol is the mass number for Hydrogen.

  15. VI. Number of Neutrons • Since only protons and neutrons contribute to the mass of an atom, the sum of these two particles must be equal to the mass number. • To calculate the number of neutrons in a given atom: MASS NUMBER – # OF PROTONS = # OF NEUTRONS (round the mass number to the nearest whole number) (REMEMBER: the atomic number and # of p+ is the same) Examples:How many neutrons are in each of the following? • A carbon atom? 12 – 6 = 6 • A zinc atom? 65 – 30 = 35

  16. Sheet 20/2—1Do numbers: 1, 2, 3, 6, 8, 9a,9b

  17. VII. ISOTOPES All atoms of the same element are not identical. Isotopes are 2 or more atoms of the same element (same atomic number and same number of protons) that have different numbers of neutrons (and therefore different mass numbers). Examples: Most hydrogen atoms have a single proton and no neutrons in their nucleus. These would have an atomic number of 1 and a mass number of 1. Some hydrogen atoms have a proton and a neutron. These would still have an atomic number of 1 but would now have a mass number of 2.

  18. A carbon isotope has 8 neutrons. What is its mass number? Answer: 14 An isotope of zinc has a mass number of 62. How many neutrons does it have? Answer: 32 A given isotope has a mass number of 52 and contains 26 neutrons. What element is it? Answer: Iron (Fe)

  19. VIII. WRITING ISOTOPES We have 2 methods to represent the different isotopes of an element. 1.) X(Y) where ‘X’ is the element symbol and ‘Y’ is the mass number of that particular isotope. Examples: H(1), H(2), H(3), C(12),C(14),U(238), Zn(64) etc.etc.etc. 2.) where ‘X’ is the element symbol, ‘A’ is the mass number of the isotope and ‘Z’ is the atomic number Examples:

  20. IX. CALCULATING AVERAGE MASS • An example: A theoretical element • Element J has 4 known isotopes. Their mass distribution in nature is as follows: A) J(40) = 30.0% B) J(43) = 40.0% C) J(44) = 10.0% D) J(50) = 20.0% Calculate the mass number for this element. Take the % of each mass and add the results together to get a total.

  21. 30/100 x 40g = 12.0 g 40/100 x 43g = 17.2 g 10/100 x 44g = 4.4g 20/100 x 50.0g = 10.0g 43.6g

  22. Sheet 20/2—1Do numbers: 5, 15, 16

  23. IMPORTANT****when completing tables and answering questions on atomic theory,ALWAYS USE THE GIVEN DATA FROM THE QUESTION.The values from the periodic table are used if no other information is available.

  24. X. NUMBER OF ELECTRONS • In any neutral species – the number of electrons is equal to the number of protons. (equal to the atomic number) • That is: + = - XI. IONS Species where the number of protons is NOT equal to the number of electrons. That is: + = - These are charged particles.

  25. **All ions are formed by the addition or removal of electrons. NEVER CHANGE THE NUMBER OF PROTONS.** Type I: It is possible for an atom to gain one or more electrons. This results in the formation of a negative ion. Energy is released as each electron is accepted. Examples: F + e- → F- + energy S + 2e-→ S2- + energy N + 3e-→ N3- + energy This kind of reaction is known as a REDUCTION reaction.

  26. Type: II It is also possible to remove one or more electrons from an atom or ion. This will result in the formation of a positive ion. Energy must be supplied in order to remove each electron. Examples: Na + energy → Na+ + e- Mg + energy → Mg2+ + 2e- Al + energy → Al 3+ + 3e- This kind of reaction is known as an OXIDATION reaction

  27. To help remember oxidation and reduction

  28. My name is LEO GER ! LEO the lion says GER Loss Electron Oxidation - Gain Electron Reduction

  29. Sheet 20/2—1Do number: 3, 4, 7, 9c, 9d, 10, 13, 14, 19

  30. The Periodic Table

  31. 1. The periodic table is arranged in horizontal rows and vertical columns. 2. Each horizontal row is called a PERIOD or SERIES. 3. Each vertical column is known as a GROUP or FAMILY. Label the following on your blank table: • Families (show numbers & Roman numerals) • Periods (show number of each period) • Alkali metals • Alkali earth metals • Transition elements • Halogens • Inert gases (noble gases)

  32. Stability in ions is represented by the total number of electrons present. The stable numbers are those found in atoms of the inert gases. (2, 10, 18, 36, 54, 86) These are unreactive – that is they are stable. All other atoms will gain (become reduced) or lose (undergo oxidation) in order to achieve one of these numbers. Predicting Ion Formation

  33. Predict the ion that each of the following atoms would form in order to become stable. Write out the oxidation or reduction reaction. • K • Cl • O • P • Ca • Ga • Cs • N • Br

  34. ANSWERS K + energy → K+ + e- Cl + e- → Cl- + energy O + 2e-→ O2- + energy P + 3e- → P3- + energy Ca + energy → Ca2+ + 2e- Ga + energy → Ga3+ + 3e- Cs + energy → Cs+ + e- N + 3e- → N3- + energy Br + e- → Br- + energy

  35. VALENCE ELECTRONS • All of the electrons located in the outermost (highest) energy level are called valence electrons. • For all ‘A’ elements, the group number equals the number of valence electrons. • Eg. C is in group IVA and therefore has 4 valence electrons. • All transition (B family) elements have 2 valence electrons.

  36. General Trends of the Periodic Table

  37. Metallic Properties Metallic properties decrease as we move left to right across the table. Metallic properties increase as we move from top to bottom in the table. The most metallic element is in the bottom – left of the table and the least metallic element is in the top right.

  38. ATOMIC SIZE IN ANY FAMILY As atomic number increases – Atomic size increases This is due to more energy levels holding electrons. Eg. Cs atoms are larger than Na atoms because Na atoms have electrons in 3 energy levels holding electrons and Cs has 6 energy levels holding electrons IN ANY PERIOD As atomic number increases – Atomic size decreases This is due to increased nuclear attraction on electrons in the same energy level. Eg. A Br atom is smaller than a Ca atom. 35 protons pulling on outer electrons vs 20 protons pulling on outer electrons.

  39. IONIZATION ENERGY This is the amount of energy required to remove an electron(s) from an atom or ion. (See formation of a positive ion..) IN ANY FAMILY As atomic number increases – ionization energy decreases This is a function of size. The outermost electrons are further and further from the nucleus and are easier to remove. IN ANY PERIOD As atomic number increases – ionization energy increases This is due to increased nuclear attraction on electrons in the same energy level.

  40. Energy Levels & Writing Electron Configurations

  41. ELECTRON LOCATION • Electrons are located in specific energy levels • These energy levels can be thought of as globe shaped regions of space surrounding the nucleus. • Electrons in the inner level have the lowest amount of energy and energy increases as the levels move outward. • Electrons enter and fill these levels in order and each level has a limit to how many electrons it can hold. • Level 1 holds 2 electrons, level 2 and level 3 hold 8 each.

  42. BACKGROUND • Electrons are located in specific energy levels surrounding the nucleus. These are numbered consecutively from 1 – 7, starting from the nucleus and working outward. • Each energy level has a specific maximum capacity for holding electrons. • Energy levels can be sub-divided into sub-levels (or sub-shells). These are identified with letters: s, p, d, f…etc. • These sub-levels also have maximum capacities. • Each sub-level are, in turn, divided into orbitals that hold 2 electrons each. • Levels and sub-levels fill systematically with electrons starting from the inside (lowest energy) and working outward (highest energy).

  43. That is: level one completely fills before level 2 starts filling…level 2 before level 3….etc. • Note: Not all energy levels contain all sub-levels. • Capacities: s – 2 electrons – 1 orbital p – 6 electrons – 3 orbitals d – 10 electrons – 5 orbitals f – 14 electrons – 7 orbitals

  44. Writing Electron Configurations H: 1s1 He: 1s2 Li: 1s2 2s1 Ca: 1s2 2s2 2p6 3s2 3p6 4s2 Ag: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9 U: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f4

  45. Other Topics • Last electron added (valence electrons) • How to remember the sequence (two methods) • Grouping quantum numbers • Configurations of ions

  46. Short Cut Version(Only use when instructions allow it) Write the symbol for the last complete inert gas and then continue the configuration from the next energy level. Example: Write the abbreviated configuration for Ra. Ra: 86Rn – 7s2 Using this format, write configurations for: Au: Ga: Cf: Sc: 54Xe – 6s2 4f14 5d9 18Ar – 4s2 3d10 4p1 86Rn – 7s2 5f10 18Ar – 4s2 3d1

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