Module 2.1 Atoms and Reactions Flashcards
15th C - The Greek atom
Greek philosopher Democritus developed the first idea of the atom
Suggested you could divide matter only a certain number of times - “átomos” meaning “indivisible”
Early 1800s - Dalton’s atomic theory
Atoms are tiny particles that make up elements
Atoms cannot be divided
All atoms of a given element are the same
Atoms of one element are different from those of every other element
Dalton used his own symbols to represent atoms
Dalton developed the first table of atomic masses
1897-1906 - Joseph John Thomson
He discovers electrons
Discovered “cathode rays” had a negative charge, could be deflected by a magnet and an electric field and had a v v small mass
Cathode rays were electrons
The idea that an atom could not be split any further (proposed by the ancient Greeks and Dalton) had been disproved
1897-1906 - Joseph John Thomson
Proposed that atoms were made up of negative electrons moving around a “sea” of positive charge - “plum pudding model”
He thought that the overall negative charge = the overall positive charge - an atom is neutral with no overall charge
1909-1911 Ernest Rutherford
Gold leaf experiment:
Directed a-particles towards a sheet of very thin gold foil. They measured any deflection of the particles.
What the results showed:
Most particles weren’t deflected at all
A small percentage of particles deflected at large angles
Few particles were deflected back towards the source
His new model (1911):
The positive charge of an atom and most of the mass are concentrated in the nucleus, in the centre
Negative electrons orbit the nucleus - just like how planets orbit the sun
The + and - charges must balance
Rutherford had disproved the plum pudding model - the expected result was that the a-particles would pass through undisturbed but the observed results were that a few particles were deflected - indicating a small concentrated positive charge (the nucleus)
1913 - Niels Bohr and Henry Moseley
Bohr thought electrons could only follow certain paths otherwise they would spiral into the nucleus - “planetary atom” - electrons orbited the central nuclear “sun” in “shells”
Bohr’s model helped explain some periodic properties e.g. spectra lines seen in the emission spectra, the energy of electrons at different distances from the nucleus
Henry Moseley discovered the link between x-ray frequencies and an element’s atomic number - however, he couldn’t explain this at the time
1918- Rutherford
Discovers the proton
Could explain Moseley’s finding that an atom’s atomic number was linked to x-ray frequencies
We now know that atomic no. = no. of protons in the atom
1923-1926
Louis de Broglie suggested that particles could have the nature of both a wave and a particle
Erwin Schrödinger suggested that an electron had wave-like properties in an atom. He also introduced the idea of atomic orbitals
1932- James Chadwick
Discovers neutrons
Observed a new type of radiation emitted from some elements
Uncharged particles - approx. same mass as a proton
These were neutrons
Modern day
Now thought that protons and neutrons are made up of even smaller particles called quarks
Current model of the atom
Protons and neutrons in the nucleus
Nucleus = centre of atom
Electrons orbit the nucleus in “shells”
Nucleus = absolutely tiny
Nucleus = extremely dense, accounts for almost all of the atom’s mass
Most of an atom is empty space
Relative mass and charge of the sub-atomic particles
Proton:
Mass = 1
Charge = +1
Neutron:
Mass = 1
Charge = 0
Electron:
Mass = 1/2000
Charge = -1
Isotopes
Different masses
Same no. of protons
Same no. of electrons
Different number of neutrons
Atomic number
Proton number
Mass number
Protons + neutrons
Relative molecular mass, Mr
Add together the relative atomic masses of each atom making up a molecule
Relative formula mass
Add together the relative atomic masses of each atom making up the formula unit
What is mass spectrometry used for?
Identify an unknown compound
Find the relative abundance of each isotope of an element
Determine structural information about molecules
A mass spectrum shows
Relative or percentage abundance on the y-axis and mass-to-charge ratio on the x-axis
How is mass-to-charge ratio in all mass spectra shown as
m/z
m= mass z= the charge of the ion (usually 1)
Atoms of metals in groups 1-13
Lose electrons
Form positive ions with the electron configuration of the previous noble gas in the periodic table
Atoms of non-metals in groups 15-17
Gain electrons
Form negative ions with the electron configuration of the next noble gas in the periodic table
Why do atoms of Be, B, C and Si not usually form ions?
It requires too much energy to transfer the outer shell electrons to form ions
Ammonium
NH4+
