1.1 Atomic Structure - The Atom Flashcards
Charge of proton:
+ 1.60 x 10^-19
Charge of neutron:
0 C
Charge of electron:
- 1.60 x 10^-19
Relative charge of proton:
1+
Relative charge of neutron:
0
Relative charge of electron:
1-
Mass of proton:
1.67 x 10^-27
Mass of neutron:
1.67 x 10^-27 kg
Mass of electron:
9.11 x 10^-31
Relative mass of proton:
1
Relative mass of neutron:
1
Relative mass of electron:
1/1836
Atomic number:
The number of protons in the nucleus of an atom.
Why does the atomic number also tell you the number of electrons in an atom?
The atom is neutral therefore it must have an equal number of protons and electrons.
Mass number:
The sum of protons and neutrons in an atom.
Number of neutrons:
= mass number - atomic number
What are isotopes?
Atoms of the same element (same atomic number) but with a different mass number as they have the same number of protons but a different number of neutrons.
Chemical properties of isotopes:
Chemical properties of different isotopes are the same as they have the same number and arrangement of electrons.
Physical properties of isotopes:
Physical properties of different isotopes differ because of their masses (e.g. density, rate of diffusion)
Relative atomic mass:
The average mass of an atom of an element compared to the 1/12th of the mass of the carbon-12 isotope.
Relative molecular mass:
The average mass of a molecule of an element compared to the 1/12th of the mass of the carbon-12 isotope
Relative isotopic mass:
The average mass of an isotope of an element compared to the 1/12th of the mass of the carbon-12 isotope
Why are relative masses compared to carbon-12?
Hydrogen was originally chosen to be measured against other elements’ masses but it had a low accuracy of atomic mass measurement so carbon was chosen as it was more practical being a solid.
Ar of an element calculation:
sum (isotopic mass x % abundance) / 100
% abundance of an isotope calculation:
(n x y) + (m x (100 - y) / 100
19th Century, John Dalton model:
Described atoms as solid spheres and said that these spheres made up different elements
1897, J.J. Thomson model:
Discovered the electron which showed that atoms weren’t solid and indivisible. Known as ‘plum pudding’ model
1909, Ernest Rutherford model:
With his students Hans Geiger and Ernest Marsden, he conducted the gold foil experiment - fired alpha particles at a very thin sheet of gold. Discovered the atom to have a very small positively charged nucleus with a ‘cloud of electrons surrounding it; most of the atom is empty space
Explanation of gold leaf experiment:
Plum pudding model suggested most alpha particles would be slightly deflected by the positive ‘pudding’ of the atom but most passed through with a small number being deflected backwards. This disproved the plum pudding model as the atom must be mostly empty space.
Niels Bohr model:
Scientists realised that if the electrons were in a ‘cloud; they would spiral into the nucleus causing it to collapse. Therefore, Bohr proposed a model in which the electrons are arranged in fixed energy shells. When electrons move between shells, EM radiation with fixed energy is emitted of absorbed.
Refined Bohr model:
Scientists observed that not all electrons in the same shell had the same energy leading to the discovery of sub-shells. This is a useful model as its simple and explains observations such as bonding
