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Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature
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Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature

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  • 1.
    • Atoms, Ions, & Nomenclature
  • 2. History of Atomic Theory
    • Democritus 460-370 B.C.: proposes idea of matter being made up of small, indivisible particles (atoms)
    • Antoine Lavoisier 1743-1794: Law of Conservation of Mass
    • Joseph Proust 1754-1826: Law of Constant Composition (Law of Definite Proportion)
    • John Dalton 1766-1844: Law of Multiple Proportion & Dalton’s Atomic Theory
  • 3. Dalton’s Atomic Theory
    • Each element is made up of tiny, indivisible particles called atoms (indivisible disproved)
    • The atoms of a given element are identical (disproved); the atoms of different elements are different in some fundamental way or ways.
    • Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms.
    • 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.
  • 4. 19 th & 20 th Centuries
    • William Crookes: Cathode Ray Tube; negative particle exist; e -
    • J.J. Thomson: Cathode ray deflection; mass/charge ratio; e -
    • Robert Millikan: Oil Drop; charge; e -
    • Ernest Rutherford: Gold Foil; small, dense nucleus present; + nucleus
  • 5.
    • James Chadwick: proved the existence of neutrons
    • 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)
      • This particle model of the e - and atom was expanded a few years after Bohr’s ideas to include the wave nature of electrons
    Charge Mass Position Proton +1 1 amu nucleus Neutron 0 1 amu nucleus Electron -1 1/1836 amu Outside nucleus
  • 6. Periodic Table
    • Isotope Notation
    • Atomic #s
      • Referred to as “Z”
      • # of p +
      • For neutral atoms, also # of e -
      • Mass #s
        • Referred to as “A”
        • # of p + + # of n 0
  • 7.
    • Isotopes
    • Isotopes: atoms with same # of protons and electrons, but different # of neutrons
    • Leads to modification of Dalton’s Atomic Theory
      • All atoms of the same element contain the same number of protons and electrons, but may have different numbers of neutrons
      • Since it is the electrons in atoms that affect chemical properties of a substance, isotopes of the same element have the same chemical properties
  • 8. Mass Numbers are not integers
    • Atomic mass of Cl is 35.5 and can be “represented” by the following symbol: 35.5 Cl
    • Does not mean 17 p + , 17 e - , and 18.5 n 0
      • Not possible to have a fraction of a neutron
      • Non-integer means there is more than 1 isotope of Cl that exists in nature, 35 Cl and 37 Cl
      • The Cl isotopes exist naturally in the following abundance: 35 Cl = 75% and 37 Cl = 25%
    • Average = Σ (% of each isotope • atomic mass of each isotope
    • Atomic Mass 100
  • 9.
    • Chlorine Example:
    • Average Atomic mass = ((35)(75%)) + ((37)(25))
    • 100
    • = 35.5 amu
  • 10. Radioactivity
    • 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)
    • http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/animations_center.html
  • 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. Types of Radioactive Decay
    • Alpha Emission
      • 4 2 He or 4 2 α
      • Restricted to heavy nuclei
      • Both protons and neutrons need to be reduced in order to stabilize the nucleus.
      • Example:
        • 218 84 Po -> 4 2 He + 214 82 Pb
        • Both mass number and atomic number change
  • 13. Radioactive Decay
    • Beta Emission
      • 0 -1 β
      • In order to decrease the number of neutrons, a neutron can be converted into a proton & an electron.
      • 1 0 n -> 1 1 p + 0 -1 β
      • An electron is emitted from the nucleus as a β particle. 14 6 C -> 14 7 N + 0 -1 β
        • Mass of the nucleus doesn’t change, only the atomic number.
        • β particles are associated with elements above the band of stability.
        • They are always found on the product side of the rxn.
  • 14. Radioactive Decay
    • Positron Emission
      • 0 +1 β
      • Decreases the # of protons by converting into a neutron by emitting a positron
      • 1 1 p -> 1 0 n + 0 +1 β
      • Has the same mass as an electron, but (+) charge.
      • Example:
      • 38 19 K -> 38 18 Ar + 0 +1 β
  • 15. Radioactive Decay
    • Electron Capture
      • 0 -1 e
      • Too many protons in the nucleus
      • Inner orbital electrons are captured by nucleus
      • Electron + proton = neutron
      • 0 -1 e + 1 1 p -> 1 0 n
      • Electron capture will always be found on the reactant side.
      • Example:
      • 37 18 Ar + 0 -1 e -> 37 17 Cl
  • 16. Radioactive Decay
    • Gamma Emission
      • 0 0 γ
      • Highest energy; electromagnetic waves
      • Has no charge
      • Never see these rays (frequency is too high)
      • Most powerful
  • 17. Half Life
    • The half-life of a radioactive nucleus is the time taken for half the atoms to decay
    • It is independent of the initial quantity of atoms. There are 3 methods of determining half-life:
      • Graphically
      • Use of the expressions (see next slide)
      • Use of expression
        • Fraction of remaining activity = 1/2 n
          • Where n = # of ½ lives
  • 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. Transmutation of Elements
    • Possible by nuclear reactions to artificially produce elements
    • Example
      • Alpha Bombardment: 14 7 N + 4 2 α 1 1 H + 17 8 O
      • Accelerated heavier nuclei:
      • 250 98 Cf + 11 5 B 257 103 Lr + 4 1 0 n
      • (bombard little one w/ big one)
  • 20. Mass Deficit (Mass Defect)
    • 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
    • Contradicts law of conservation of mass
    • When particles combine, a small amount of the mass is converted to energy (binding energy) and released to the surroundings
            • E = mc 2
  • 21. Predicting Stability
    • Stable nuclei tend to have neutron-proton ratios close to 1:1 or atomic numbers below 83
    • Zone of Stability is a ratio of 1 – 1.5
    • Nuclei with higher ratios tend to want to lower the ratio by converting a neutron to a proton and e -
    • Electrons are then released as β particles
  • 22. Nuclear Fission vs. Fusion
    • Fission
    • Heavy nuclei capturing neutrons, splitting to form other, smaller nuclei and releasing more neutrons
    • Large amounts of energy can be released, leading to a potential chain reaction
    • Fusion
    • Combination of smaller nuclei into larger ones with the release of energy
    • Less easy to perform since they involve the combination of two nuclei that are positively charged and therefore repel one another
  • 23. Uses of Radioactivity
    • Medicine: 133 I for thyroid and brain imaging 67 Ga for lung function
    • Isotopic dating
    • Thickness control in engineering
    • Leak detection
    • Nuclear Fission (power and atomic bomb): Uranium nuclei can be bombarded with neutrons and converted to other nuclei
  • 24. Molecules
    • Formed when a definite number of atoms are joined together by chemical bonds
    • 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
    • Usually formed between non-metal elements
    • Formulas show the number of each type of atom present written as subscripts
    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. Ions
    • When atoms lose or gain e - s, particles become charged
    • Positive ions (cations): number of p + is greater than the number of e -
    • Negative ions (anions): number of e - is greater than the number of p +
    • Metals form cations; non metals form anions
    • Oppositely charged ions form ionic compounds by attracting one another
  • 26. Na + cation monoatomic Cl - anion monoatomic CO 3 2- anion polyatomic NH 4 + cation polyatomic
  • 27. Nomenclature of Inorganic Compounds
    • Binary compounds (ionic compounds)
      • Formed between 2 elements metal and non metal
      • To find the formula, positive and negative charges must be balanced
      • 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)
      • Transition metals can carry more than 1 charge, so when writing the name of the compound, parentheses must be shown to indicate the charge
  • 28.
    • Binary Acids
      • For nomenclature purposes, an acid can be defined as a compound that produces hydrogen ions when dissolved in water
      • Formed when hydrogen ions combine with monoatomic anions
      • To name: use prefix hydro-followed by the other nonmetal name modified to an –ic ending
        • Example: HCl = Hydrochloric acid
  • 29.
    • Polyatomic ions and oxoanions
      • Polyatomic ions are those where more than one element is combined to create a species with a charge
      • Polyatomic ions where oxygen is combined with another non-metal are called oxoanions
      • 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
  • 30.
      • Some oxoanions contain hydrogen and are named accordingly, example: HPO 4 2- , hydrogen phosphate
      • Prefix (thio-) means that a sulfur atom has replaced an atom of oxygen in an anion
      • 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
          • Example: K 2 CO 3 , potassium carbonate
  • 31.
    • Oxoacids
      • Oxoacids are formed when hydrogen ions combine with polyatomic oxoanions
      • Gives a combination of hydrogen, oxygen, and another non-metal
      • 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.
        • Example: H 2 SO 4 , hydrogen sulfate becomes Sulfuric acid
  • 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.
    • Binary Compounds of 2 non-metals
      • Two non-metals combine, then the compound is molecular
      • To name: unmodified name of 1 st element is followed by the root of the 2 nd element with the ending (-ide)
      • In order to distinguish between compounds of the same element, use prefixes mono, di, tri, tetra, penta, …..
        • Example: SO 2 , sulfur dioxide
  • 34.
    • Hydrates
      • Ionic formula units with water associated with them
      • Water molecules are incorporated into the solid structure of the ions
      • 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
      • Strong heating can generally drive off the water in these salts. With water removed, they become anhydrous
  • 35. Mass Spectrometer
    • Most direct way to determine the atomic and molecular weights
    • What Happens?
      • A gas is introduced and bombarded by a stream of high energy electrons ( vaporization )
      • Collisions between the electrons and the atoms or molecules of the gas produce positive ions, mostly with a +1 charge ( ionization )
      • Ions are accelerated toward a (-) charged wire grid ( acceleration )
  • 36.
      • 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
      • 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
      • 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 )