Properties of atoms and materials Flashcards
Describe Thomson’s model of the atom.
The plum pudding model of the atom was an important step in the development of atomic theory such as the discovery of electrons and the notion that the atom is a non-inert and divisible mass. Its appearance consists of a sea of uniform positive charge with electrons distributed throughout.
What determines the ability of atoms to form chemical bonds?
The arrangement of electrons within an atom and the stability of the valence electron shell.
Describe Rutherford’s gold leaf experiment and explain how the findings lead to a change in the atomic strucuture from Thomson’s model.
The gold leaf experiment observed how alpha particles affect matter. In 1911, Rutherford discovered that shooting alpha particles into a sheet of gold surrounded by a sheet of ZnS made most of the particles go straight through as thought it was empty space or a small proportion of the particles were deflected and scattered at measurable angles. He concluded that there was a small fraction of the total volume that contains most of the mass of the atom. The gold foil experiment demonstrated that the atom has a tiny heavy nucleus with electrons moving around it in circular orbits.
Describe the Bohr model of the atom and explain why it is different to Rutherford’s model.
He put forward a model in which electrons were only allowed to have specific energies. These specific energies are the energy levels of the electron. Electrons in energy levels would not radiate energy. Electrons in the lowest energy levels could only accept electrons. Electrons in the higher energy levels could either accept or release energy, but only in discrete quanta, by moving between levels. The Bohr model of the atom proposes that the atom comprises a nucleus surrounded by electrons. The electrons revolve in shells located at specific distances from the nucleus. They are bound to the nucleus in specific energy levels.
Explain the trend of Atomic radii as you go across a group and down a period.
It decreases as you go across a period due to the nuclear charge and core charge increasing by 1 each time. The greater the charge the greater the attraction for the electrons pulling them in more slightly. It increases down the group. Size is determined by the outermost electrons. In going down a group, the outermost electrons are assigned to shells with higher energy levels which are further from the nucleus.
Describe the strucutre of the atom.
A positively-charged nucleus surround by electrons in distinct energy levels, held together by electrostatic forces of attraction.
Explain the trend of Valencies as you go across a group and down a period.
Remains the same as you go down a group as all elements in the group have the same number of electrons on their outermost shell. It will increase across a period because the number of electrons on the outermost she’ll increase by 1 each time.
Explain the trend of Electronegativity as you go across a group and down a period.
It is the ability of an atom to attract electrons. It increases across a period because the core charge is increasing and the atom is getting smaller, so the incoming electron will have a greater attraction. It decreases down a group because the atomic radius is increasing and the core charge remains the same, the incoming electrons will not have such a great attraction.
Explain the trend of First ionisation energies as you go across a group and down a period.
First ionisation energy decreases down a group and increases across a period. The outer electron is further from the nucleus as it is located in a higher energy electrons. As the number of protons in the nucleus increase, the number of electrons shielding the outer electron from the nucleus increase. This means that the election requires a lot less energy to remove.
Explain the trend of successive ionisation energies as you go down a group and across a period.
Successive ionisation energies increase both down a group and across a period. The ion that is produced each time has a progressively larger positive charge. The attraction between the electron and the nucleus is greater and would require a lot more energy to overcome this attraction.
Explain how flame tests and atomic absorption spectroscopy are used to identify elements.
- Atomic absorption spectrometry uses the absorption of light by electrons in the atom to measure how much of an element is present in a sample of substance.
- Flame tests use the colour of the flame of a burning substance to determine what elements are contained within the substance.
Define isotope.
Atoms of the same element that always have the same number of protons, they may have different numbers of neutrons.
Why do isotopes have similar chemical properties but different physical properties?
The isotopes of an element are virtually identical in their chemical reactions. This is because they have the same number of protons and the same number of electrons.
Define relative atomic mass (atomic weight) of an atom.
The average mass of an element that takes into account the isotope masses and the relative abundance on Earth. It is measure against carbon-12.
Describe what happens in a mass spectrometer.
Gaseous atoms are bombarded by electrons from an electron gun and are ionised. A sufficient amount of energy is given to form ions of +1 charge. The ions are charge so they can be accelerated through an electric field. The charged particles will be deflected by a magnetic or electric field. They are detected by electric or photographic methods.
