Chapter Two Flashcards
Atoms: structure and mass
Leucippus stated that
If a piece of matter were subdivided over and over again eventually it would not be able to be divided any further
Democritus
- Expanded on Leucippus’s idea and named the final piece of matter ‘atomos’
- He believed different substances had atomos of different shapes and sizes which gave substances different properties
Atomos means
Indivisible
Aristotle
Believed all matter was made of four elements > fire, water, air and earth
John Dalton
- All matter was made up of these indivisible atoms
- Atoms had different weights
- Compounds are formed through the combination of atoms
- Atoms of a given element are identical in size and mass and when they combine to form compounds they do so only in the simplest ratio (some of these points are untrue)
Joseph John Thomson
- Particles flow from the negative electrode to the positive electrode and the stream of particles is repelled by another negative electrode
- Must be negatively charged particles present in all atoms
- Because atoms are electrically neutral he proposed there must also be positive charges to balance out the negative ones
- Plum pudding model
Plum pudding model
- Negatively charged particles were spread throughout the positively charged sphere of the atom like raisins in a plum pudding
- These negatively charged particles became known as electrons
Ernest Rutherford
- Most of the atom must be empty space
- Most of the mass and positive charge of an atom must be located in a tiny region, the nucleus of the atom
- Electrons orbited the nucleus like planets around the sun
- Called the positive subatomic particle proton
- Believed there was a yet undiscovered neutral particle in the nucleus
James Chadwick
Confirmed the neutral particle and called it the neutron
Electrons
- Negatively charged particles in a cloud surrounding the nucleus
- Very little contribution to an atoms overall mass
- Negative particles attract positive particles which keeps the atom bound together
The nucleus
- 99.97% of an atoms mass
- The subatomic particles inside are referred to collectively as nucleons
Proton charge
+1
Proton relative mass
1
Proton mass in kg
1.673x10^-27
Neutron charge
0
Neutron relative mass
1
Neutron mass in kg
1.675x10^-27
Electron charge
-1
Electron relative mass
1/1800
Electron mass in kg
9.109x10^-31
Atomic number
The number of protons in an atoms nucleus, represented by the symbol ‘Z’
Mass number
- Always a whole number
- The number representing the number of protons + neutrons in the nucleus, symbol = ‘A’
Isotopes
- Atoms of the same element with different numbers of neutrons and in turn mass numbers
- Identical chemical properties but different physical properties
What causes an isotope to be unstable and radioactive
- A significant difference between the number of neutrons and protons
- These isotopes undergo radioactive decay, emitting various forms of radiation and turn into more stable, lighter nuclei
What standard are the masses of individual isotopes of each element ccompared to
Carbon-12
Relative atomic mass
The average relative mass of an atom in this naturally occurring mixture of isotopes
Relative abundance
- %
- The proportion of each isotope in the sample of each element
Equation for relative atomic mass
Relative atomic mass = Relative isotopic mass x Relative abundance/100
OR
Relative atomic mass = (Relative abundance x relative isotopic mass) + (relative abundance x relative isotopic mass)/100
Relative abundance equation
Relative atomic mass = ( relative isotopic mass x a) + (relative isotopic mass x (100-a)) / 100
a = lighter isotope abundance
100-a = heavier isotope abundance
Mass spectrometry
- A technique used to measure the mass of atoms of molecules
- Can identify the presence and relative abundance of isotopes in a sample of an element (in context of atomic structure)
- Can measure the mass of molecules which helps determine structure
Relative isotopic mass
The mass of an individual isotope of each element compared to the standard of carbon-12
Mass spectrometer
- Separates the individual isotopes in a sample of the element
- Determines the mass of each isotope relative to the carbon-12 standard
- Calculates the relative abundance of the isotopes in the sample
Mass spectra
- The information from the mass spectrometer is presented graphically and plots the relative abundance of each ion against its mass-charge ratio (m/2)
- For most elements the proportion of its isotope is approximately the same in every sample
Relative isotopic abundance
The % amount of an isotope in a natural element
Mass spectra equation
% = peak height/total peak height x100
Chemical symbol
One or two letters that represent the element (the first is always capitalised)
