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Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature

Chemistry- JIB Topic 2 Atoms, Ions and Nomenclature






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

      • Atoms, Ions, & Nomenclature
    • 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
    • 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.
    • 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
      • 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
    • 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
      • 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
    • 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
      • Chlorine Example:
      • Average Atomic mass = ((35)(75%)) + ((37)(25))
      • 100
      • = 35.5 amu
    • 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
    • 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
    • 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
    • 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.
    • 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
    • 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
    • Radioactive Decay
      • Gamma Emission
        • 0 0 γ
        • Highest energy; electromagnetic waves
        • Has no charge
        • Never see these rays (frequency is too high)
        • Most powerful
    • 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
    • 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
    • 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)
    • 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
    • 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
    • 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
    • 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
    • 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
    • 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
    • Na + cation monoatomic Cl - anion monoatomic CO 3 2- anion polyatomic NH 4 + cation polyatomic
    • 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
      • 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
      • 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
        • 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
      • 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
    • 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
      • 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
      • 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
    • 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 )
        • 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 )