chapter 2 atomic structure and isotopes

Mass

  • protons have virtually the same mass as a neutron
  • An electron has a negligible mass of 1/1836th the mass of a proton
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Charge

  • A proton has a charge of +1
  • An electron has a charge of -1
  • A neutron has no charge
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Building the atom

  • Nearly all the atoms mass is in the nucleus
  • Atoms contain the same number of protons as neutrons
  • The overall charge is 0
  • Neutrons can be thought of as the glue that holds the nucleus together
  • Atoms contain the same number of or slightly more neutrons as protons
  • As the nucleus gets larger more and more neutrons are needed
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Atomic number

  • The number of protons in an atom defines the element
  • The atomic number is the number of protons in an element
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Isotopes

  • Unlike protons, the number of neutrons in the atoms of an element can change
  • Isotopes are atoms of the same element with different numbers of neutrons
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Isotopes and chemical reactions

  • Different isotopes of the same element will have the same number of electrons
  • Because the number of neutrons has no effect on the reaction, different isotopes will react in the same way
  • There may be small changes in physical properties
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Atomic structure of ions

  • Cations (positive ions) have fewer electrons than protons
  • Anions (negative ions) have more electrons than protons
  • Ions and atoms of the same element will have the same number of protons but a different number of electrons
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Carbon-12

  • Mass defect is the small amount of mass lost due to the strong nuclear force holding together protons and neutrons
  • A standard isotope is needed to base all atomic masses
  • The mass of a carbon-12 isotope is defined as exactly 12 atomic mass units
  • The standard mass for atomic mass is 1u which is equal to 1/12th of an atom of carbon-12
  • On this scale, 1u is equal to the mass of a proton or neutron
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Relative isotopic mass

  • Relative isotopic mass is the weighted mean mass of an atom relative to 1/12th of an atom of carbon-12
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Relative atomic mass

  • Relative atomic mass is the weighted mean mass of an atom relative to 1/12th of an atom of carbon-12
  • The weighted mean mass takes account of the percentage abundance and relative isotopic mass of each isotope
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Determination of relative atomic mass

  • A sample is placed in a mass spectrometer 
  • The sample is vaporised and then ionised to form positive ions 
  • The ions are accelerated
  • Heavier ions move more slowly and are more difficult to deflect than lighter ions
  • Ions of each isotope are separated
  • The ions are detected on a mass spectrum as a mass-to-charge ratio

mass to charge ratio = relative mass of ion / relative charge on ion

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Simple ions

  • Metals lose electrons to form positive ions
  • Non-metals gain electrons to form negative ions
  • Some metals such as some of the transition metals can form several ions with different charges
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Binary compound

  • A compound that contains two elements only
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Polyatomic ions

  • An ion that contains atoms of more than one element
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writing the formulae from ions

  • The overall charge must be zero, so the sum of the positive charges must equal the sum of the negative charges
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Writing the formulae from ions

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Writing the formula

  • The number of each ion is shown in subscript after the ion 
  • The ionic charge is usually omitted in a full formula
  • Brackets are used if there is more than one polyatomic ion
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Elements

  • Elements are shown simply as their symbol unless they exist as a diatomic molecules
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Compounds

  • Covalent compounds do not contain ions
  • For ionic compounds the formula worked out from the ionic charges is used in the equation- this is called the formula unit
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State symbols

  • (g) - gas
  • (l) - liquid
  • (s) - solid
  • (aq) - aqueous
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balancing equations

  • You must not change any chemical formula
  • Balancing numbers go in front of chemical formula
  • The equation is balanced when the same number of atoms are on both sides
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