Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature


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Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature

  1. 1. <ul><li>Atoms, Ions, & Nomenclature </li></ul>
  2. 2. History of Atomic Theory <ul><li>Democritus 460-370 B.C.: proposes idea of matter being made up of small, indivisible particles (atoms) </li></ul><ul><li>Antoine Lavoisier 1743-1794: Law of Conservation of Mass </li></ul><ul><li>Joseph Proust 1754-1826: Law of Constant Composition (Law of Definite Proportion) </li></ul><ul><li>John Dalton 1766-1844: Law of Multiple Proportion & Dalton’s Atomic Theory </li></ul>
  3. 3. Dalton’s Atomic Theory <ul><li>Each element is made up of tiny, indivisible particles called atoms (indivisible disproved) </li></ul><ul><li>The atoms of a given element are identical (disproved); the atoms of different elements are different in some fundamental way or ways. </li></ul><ul><li>Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. </li></ul><ul><li>Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. </li></ul>
  4. 4. 19 th & 20 th Centuries <ul><li>William Crookes: Cathode Ray Tube; negative particle exist; e - </li></ul><ul><li>J.J. Thomson: Cathode ray deflection; mass/charge ratio; e - </li></ul><ul><li>Robert Millikan: Oil Drop; charge; e - </li></ul><ul><li>Ernest Rutherford: Gold Foil; small, dense nucleus present; + nucleus </li></ul>
  5. 5. <ul><li>James Chadwick: proved the existence of neutrons </li></ul><ul><li>Niels Bohr: proposed the idea that the atom is made up of a nucleus containing p + and n 0 that was being orbited by e - s in orbits (disproved) </li></ul><ul><ul><li>This particle model of the e - and atom was expanded a few years after Bohr’s ideas to include the wave nature of electrons </li></ul></ul>Charge Mass Position Proton +1 1 amu nucleus Neutron 0 1 amu nucleus Electron -1 1/1836 amu Outside nucleus
  6. 6. Periodic Table <ul><li>Isotope Notation </li></ul><ul><li>Atomic #s </li></ul><ul><ul><li>Referred to as “Z” </li></ul></ul><ul><ul><li># of p + </li></ul></ul><ul><ul><li>For neutral atoms, also # of e - </li></ul></ul><ul><ul><li>Mass #s </li></ul></ul><ul><ul><ul><li>Referred to as “A” </li></ul></ul></ul><ul><ul><ul><li># of p + + # of n 0 </li></ul></ul></ul>
  7. 7. <ul><li>Isotopes </li></ul><ul><li>Isotopes: atoms with same # of protons and electrons, but different # of neutrons </li></ul><ul><li>Leads to modification of Dalton’s Atomic Theory </li></ul><ul><ul><li>All atoms of the same element contain the same number of protons and electrons, but may have different numbers of neutrons </li></ul></ul><ul><ul><li>Since it is the electrons in atoms that affect chemical properties of a substance, isotopes of the same element have the same chemical properties </li></ul></ul>
  8. 8. Mass Numbers are not integers <ul><li>Atomic mass of Cl is 35.5 and can be “represented” by the following symbol: 35.5 Cl </li></ul><ul><li>Does not mean 17 p + , 17 e - , and 18.5 n 0 </li></ul><ul><ul><li>Not possible to have a fraction of a neutron </li></ul></ul><ul><ul><li>Non-integer means there is more than 1 isotope of Cl that exists in nature, 35 Cl and 37 Cl </li></ul></ul><ul><ul><li>The Cl isotopes exist naturally in the following abundance: 35 Cl = 75% and 37 Cl = 25% </li></ul></ul><ul><li>Average = Σ (% of each isotope • atomic mass of each isotope </li></ul><ul><li>Atomic Mass 100 </li></ul>
  9. 9. <ul><li>Chlorine Example: </li></ul><ul><li>Average Atomic mass = ((35)(75%)) + ((37)(25)) </li></ul><ul><li>100 </li></ul><ul><li>= 35.5 amu </li></ul>
  10. 10. Radioactivity <ul><li>The spontaneous decay of certain atoms with the evolution of alpha, beta, gamma, and positron particles. The radiation comes from the nucleus (it is a nuclear reaction) </li></ul><ul><li> </li></ul>
  11. 11. Alpha α Beta β Gamma γ Helium nucleus Essentially electons Always products Electromagnetic radiation High energy High Frequency Charge +2 -1 0 Mass 4 1/1840 0 Movement To neg. plate To pos. plate None Penetration Least Stopped by paper Intermediate Stopped by lead or glass Greatest Thick layers of lead or concrete
  12. 12. Types of Radioactive Decay <ul><li>Alpha Emission </li></ul><ul><ul><li>4 2 He or 4 2 α </li></ul></ul><ul><ul><li>Restricted to heavy nuclei </li></ul></ul><ul><ul><li>Both protons and neutrons need to be reduced in order to stabilize the nucleus. </li></ul></ul><ul><ul><li>Example: </li></ul></ul><ul><ul><ul><li>218 84 Po -> 4 2 He + 214 82 Pb </li></ul></ul></ul><ul><ul><ul><li>Both mass number and atomic number change </li></ul></ul></ul>
  13. 13. Radioactive Decay <ul><li>Beta Emission </li></ul><ul><ul><li>0 -1 β </li></ul></ul><ul><ul><li>In order to decrease the number of neutrons, a neutron can be converted into a proton & an electron. </li></ul></ul><ul><ul><li>1 0 n -> 1 1 p + 0 -1 β </li></ul></ul><ul><ul><li>An electron is emitted from the nucleus as a β particle. 14 6 C -> 14 7 N + 0 -1 β </li></ul></ul><ul><ul><ul><li>Mass of the nucleus doesn’t change, only the atomic number. </li></ul></ul></ul><ul><ul><ul><li>β particles are associated with elements above the band of stability. </li></ul></ul></ul><ul><ul><ul><li>They are always found on the product side of the rxn. </li></ul></ul></ul>
  14. 14. Radioactive Decay <ul><li>Positron Emission </li></ul><ul><ul><li>0 +1 β </li></ul></ul><ul><ul><li>Decreases the # of protons by converting into a neutron by emitting a positron </li></ul></ul><ul><ul><li>1 1 p -> 1 0 n + 0 +1 β </li></ul></ul><ul><ul><li>Has the same mass as an electron, but (+) charge. </li></ul></ul><ul><ul><li>Example: </li></ul></ul><ul><ul><li>38 19 K -> 38 18 Ar + 0 +1 β </li></ul></ul>
  15. 15. Radioactive Decay <ul><li>Electron Capture </li></ul><ul><ul><li>0 -1 e </li></ul></ul><ul><ul><li>Too many protons in the nucleus </li></ul></ul><ul><ul><li>Inner orbital electrons are captured by nucleus </li></ul></ul><ul><ul><li>Electron + proton = neutron </li></ul></ul><ul><ul><li>0 -1 e + 1 1 p -> 1 0 n </li></ul></ul><ul><ul><li>Electron capture will always be found on the reactant side. </li></ul></ul><ul><ul><li>Example: </li></ul></ul><ul><ul><li>37 18 Ar + 0 -1 e -> 37 17 Cl </li></ul></ul>
  16. 16. Radioactive Decay <ul><li>Gamma Emission </li></ul><ul><ul><li>0 0 γ </li></ul></ul><ul><ul><li>Highest energy; electromagnetic waves </li></ul></ul><ul><ul><li>Has no charge </li></ul></ul><ul><ul><li>Never see these rays (frequency is too high) </li></ul></ul><ul><ul><li>Most powerful </li></ul></ul>
  17. 17. Half Life <ul><li>The half-life of a radioactive nucleus is the time taken for half the atoms to decay </li></ul><ul><li>It is independent of the initial quantity of atoms. There are 3 methods of determining half-life: </li></ul><ul><ul><li>Graphically </li></ul></ul><ul><ul><li>Use of the expressions (see next slide) </li></ul></ul><ul><ul><li>Use of expression </li></ul></ul><ul><ul><ul><li>Fraction of remaining activity = 1/2 n </li></ul></ul></ul><ul><ul><ul><ul><li>Where n = # of ½ lives </li></ul></ul></ul></ul>
  18. 18. it is easy to determine how much of a sample remains after a whole number of half lives it is more difficult to determine how much remains when a complete half life has not passed in order to do this, you need to apply two different equations ln N t /N o = -kt and k= 0.693/t 1/2 N t is the amount of substance left after a given time N o is the original amount of the substance t is the amount of time that has passed t 1/2 is the half life of the substance
  19. 19. Transmutation of Elements <ul><li>Possible by nuclear reactions to artificially produce elements </li></ul><ul><li>Example </li></ul><ul><ul><li>Alpha Bombardment: 14 7 N + 4 2 α 1 1 H + 17 8 O </li></ul></ul><ul><ul><li>Accelerated heavier nuclei: </li></ul></ul><ul><ul><li>250 98 Cf + 11 5 B 257 103 Lr + 4 1 0 n </li></ul></ul><ul><ul><li> (bombard little one w/ big one) </li></ul></ul>
  20. 20. Mass Deficit (Mass Defect) <ul><li>When atoms are formed by the combination of p + , n 0 , and e - ….the mass of the atom is found to be less than that of the sum of the individual particles </li></ul><ul><li>Contradicts law of conservation of mass </li></ul><ul><li>When particles combine, a small amount of the mass is converted to energy (binding energy) and released to the surroundings </li></ul><ul><ul><ul><ul><ul><li>E = mc 2 </li></ul></ul></ul></ul></ul>
  21. 21. Predicting Stability <ul><li>Stable nuclei tend to have neutron-proton ratios close to 1:1 or atomic numbers below 83 </li></ul><ul><li>Zone of Stability is a ratio of 1 – 1.5 </li></ul><ul><li>Nuclei with higher ratios tend to want to lower the ratio by converting a neutron to a proton and e - </li></ul><ul><li>Electrons are then released as β particles </li></ul>
  22. 22. Nuclear Fission vs. Fusion <ul><li>Fission </li></ul><ul><li>Heavy nuclei capturing neutrons, splitting to form other, smaller nuclei and releasing more neutrons </li></ul><ul><li>Large amounts of energy can be released, leading to a potential chain reaction </li></ul><ul><li>Fusion </li></ul><ul><li>Combination of smaller nuclei into larger ones with the release of energy </li></ul><ul><li>Less easy to perform since they involve the combination of two nuclei that are positively charged and therefore repel one another </li></ul>
  23. 23. Uses of Radioactivity <ul><li>Medicine: 133 I for thyroid and brain imaging 67 Ga for lung function </li></ul><ul><li>Isotopic dating </li></ul><ul><li>Thickness control in engineering </li></ul><ul><li>Leak detection </li></ul><ul><li>Nuclear Fission (power and atomic bomb): Uranium nuclei can be bombarded with neutrons and converted to other nuclei </li></ul>
  24. 24. Molecules <ul><li>Formed when a definite number of atoms are joined together by chemical bonds </li></ul><ul><li>Can consist of atoms of one element or atoms of many different elements, but always in a fixed proportion Molecules can be elements or compounds </li></ul><ul><li>Usually formed between non-metal elements </li></ul><ul><li>Formulas show the number of each type of atom present written as subscripts </li></ul>H 2 , N 2 , O 2 , F 2 , Cl 2 , I 2 , Br 2 elements diatomic H 2 O compound polyatomic NH 3 compound polyatomic
  25. 25. Ions <ul><li>When atoms lose or gain e - s, particles become charged </li></ul><ul><li>Positive ions (cations): number of p + is greater than the number of e - </li></ul><ul><li>Negative ions (anions): number of e - is greater than the number of p + </li></ul><ul><li>Metals form cations; non metals form anions </li></ul><ul><li>Oppositely charged ions form ionic compounds by attracting one another </li></ul>
  26. 26. Na + cation monoatomic Cl - anion monoatomic CO 3 2- anion polyatomic NH 4 + cation polyatomic
  27. 27. Nomenclature of Inorganic Compounds <ul><li>Binary compounds (ionic compounds) </li></ul><ul><ul><li>Formed between 2 elements metal and non metal </li></ul></ul><ul><ul><li>To find the formula, positive and negative charges must be balanced </li></ul></ul><ul><ul><li>To name: the unmodified name of the positive ion is written first followed by the root of the negative ion with the ending modified (-ide) </li></ul></ul><ul><ul><li>Transition metals can carry more than 1 charge, so when writing the name of the compound, parentheses must be shown to indicate the charge </li></ul></ul>
  28. 28. <ul><li>Binary Acids </li></ul><ul><ul><li>For nomenclature purposes, an acid can be defined as a compound that produces hydrogen ions when dissolved in water </li></ul></ul><ul><ul><li>Formed when hydrogen ions combine with monoatomic anions </li></ul></ul><ul><ul><li>To name: use prefix hydro-followed by the other nonmetal name modified to an –ic ending </li></ul></ul><ul><ul><ul><li>Example: HCl = Hydrochloric acid </li></ul></ul></ul>
  29. 29. <ul><li>Polyatomic ions and oxoanions </li></ul><ul><ul><li>Polyatomic ions are those where more than one element is combined to create a species with a charge </li></ul></ul><ul><ul><li>Polyatomic ions where oxygen is combined with another non-metal are called oxoanions </li></ul></ul><ul><ul><li>Certain non-metals (Cl, N, P, and S) form a series of oxoanions containing different numbers of oxygen atoms Hypo(element) ite increasing # of (Element) ite oxygen atoms (Element) ate Per(element) ate </li></ul></ul>
  30. 30. <ul><ul><li>Some oxoanions contain hydrogen and are named accordingly, example: HPO 4 2- , hydrogen phosphate </li></ul></ul><ul><ul><li>Prefix (thio-) means that a sulfur atom has replaced an atom of oxygen in an anion </li></ul></ul><ul><ul><li>To name an ionic compound that contains a polyatomic ion, the unmodified name of the positive ion is written first followed by the unmodified name of the negative ion </li></ul></ul><ul><ul><ul><ul><li>Example: K 2 CO 3 , potassium carbonate </li></ul></ul></ul></ul>
  31. 31. <ul><li>Oxoacids </li></ul><ul><ul><li>Oxoacids are formed when hydrogen ions combine with polyatomic oxoanions </li></ul></ul><ul><ul><li>Gives a combination of hydrogen, oxygen, and another non-metal </li></ul></ul><ul><ul><li>To name: use the name of the oxoanion and replace the (-ite) ending with (-ous) or the (-ate) ending with (-ic). Then add the word acid. </li></ul></ul><ul><ul><ul><li>Example: H 2 SO 4 , hydrogen sulfate becomes Sulfuric acid </li></ul></ul></ul>
  32. 32. HClO ClO - Hypochlorite Hypochlorous acid HClO 2 ClO 2 - Chlorite Chlorous acid HClO 3 ClO 3 - Chlorate Chloric acid HClO 4 ClO 4 - perchlorate Perchloric acid
  33. 33. <ul><li>Binary Compounds of 2 non-metals </li></ul><ul><ul><li>Two non-metals combine, then the compound is molecular </li></ul></ul><ul><ul><li>To name: unmodified name of 1 st element is followed by the root of the 2 nd element with the ending (-ide) </li></ul></ul><ul><ul><li>In order to distinguish between compounds of the same element, use prefixes mono, di, tri, tetra, penta, ….. </li></ul></ul><ul><ul><ul><li>Example: SO 2 , sulfur dioxide </li></ul></ul></ul>
  34. 34. <ul><li>Hydrates </li></ul><ul><ul><li>Ionic formula units with water associated with them </li></ul></ul><ul><ul><li>Water molecules are incorporated into the solid structure of the ions </li></ul></ul><ul><ul><li>To name: use the normal name of the ionic compound followed by the term hydrate with an appropriate prefix to show the number of water molecules per ionic formula. Example: CuSO 4 •5H 2 O , copper (II) sulfate pentahydrate </li></ul></ul><ul><ul><li>Strong heating can generally drive off the water in these salts. With water removed, they become anhydrous </li></ul></ul>
  35. 35. Mass Spectrometer <ul><li>Most direct way to determine the atomic and molecular weights </li></ul><ul><li>What Happens? </li></ul><ul><ul><li>A gas is introduced and bombarded by a stream of high energy electrons ( vaporization ) </li></ul></ul><ul><ul><li>Collisions between the electrons and the atoms or molecules of the gas produce positive ions, mostly with a +1 charge ( ionization ) </li></ul></ul><ul><ul><li>Ions are accelerated toward a (-) charged wire grid ( acceleration ) </li></ul></ul>
  36. 36. <ul><ul><li>After they pass through grid, they come to 2 slits that only allow a narrow beam of ions to pass @ a given time: no magnets </li></ul></ul><ul><ul><li>This beam pass through 2 magnetic poles which deflect the ions into a curved path ( deflection ) •extent of curve depends on mass (high mass = small deflection) •ions are separated by masses </li></ul></ul><ul><ul><li>By changing the strength of the magnetic field or acceleration voltage on grid, ions of varying masses can be selected to enter detector @ end of instrument ( detection ) </li></ul></ul>