T1 P1+2: atomic structure

• 1.1 Describe how the Dalton model of an atom has changed over time

  because of the discovery of subatomic particles

  JOHN DALTON: 1803

  • All​ ​substances​ ​are​ ​made​ ​of​ ​atoms.​ ​Atoms​ ​are​ ​small​ ​particles​ ​that​ ​cannot be​ ​created,​ ​divided,​ ​or​ ​destroyed.​ ​
  • Atoms​ ​of​ ​the​ ​same​ ​element​ ​are​ ​exactly​ ​alike,​ ​and​ ​atoms​ ​of​ ​different elements​ ​are​ ​different.​ ​
  • Atoms​ ​join​ ​with​ ​other​ ​atoms​ ​to​ ​make​ ​new​ ​substances.​ ​
  • Much​ ​of​ ​Dalton’s​ ​theory​ ​was​ ​correct,​ ​but​ ​some​ ​of​ ​it​ ​was​ ​later​ ​proven​ ​incorrect
• 1.1 Describe how the Dalton model of an atom has changed over time because

  of the discovery of subatomic particles

  J.J. Thomson (1897):

  • Used a cathode-ray tube to conduct an experiment
  • This discovery identified an error in Dalton's atomic theory, Atoms can be divided into smaller parts. because the beam moved away from negatively charged plate and towards positively charged, so he knew atoms must have a negative charge
  • Thomson proposed a model of an atom called the "plum-pudding" model, in which negative electrons are scattered throughout soft blobs of positively charged material.
• 1.1 Describe how the Dalton model of an atom has changed over time because

  of the discovery of subatomic particles

  Ernest Rutherford (1909):

  • shot beams of positively charged particles into a sheet of gold foil
  • Most of the particles continued in a straight line. some of the particles were deflected to the sides, few bounced straight back.
  • Rutherford developed new model: most of the atom's mass is found in a region in the center called the nucleus.
  • In Rutherford's model: atom is mostly empty space, and electrons travel in random paths around the nucleus.
• 1.2 Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells
• 1.3 Recall the relative charge and relative mass of protons, neutrons and electrons

  proton: mass 1 charge +1

  electron: mass 1/1836 charge -1

  neutron: mass 1 charge 0
• 1.4 Explain why atoms contain equal number of protons and electrons

  • atoms are neutral and the charges of protons are +1, electrons are -1
  • amount of protons= amount of electrons, so charges cancel
• 1.5 Describe the nucleus of an atom as very small compared to the overall size of the atom
• • 1.6 Recall that most of the mass of an atom is concentrated in the nucleus
• 1.7 Recall the meaning of the term mass number of an atom

  Mass (nucleon) Number = number of protons + neutrons
• 1.8 Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element
• 1.9 Describe isotopes as different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei
• 1.10 Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number

  • Atomic (proton) Number = number of protons (= number of electrons if it's an atom, because atoms are neutral)
  • therefore, you can calculate number of neutrons by doing mass number - atomic number
• 1.11 Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

  •  isotopes have the same number of protons but different numbers of neutrons, still atoms of the same element, but have different atomic masses
  • relative atomic mass: calculated using abundance of different isotopes, because its an average, can lead to the ram not being a whole number (atomic+mass number will always be whole numbers)
  • mass of atoms are small, compare the masses to each other. carbon atom having a mass number 12, relative atomic mass is 12.
  • Ar or R.A.M.
• 1.12 Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

  R.A.M.

  = (mass of isotope-A x % of isotope-A) + (mass of isotope-B x % of isotope-B) divided by 100