Chapter 2.1- Atoms and Reactions Flashcards
Early 1800s, Dalton’s atomic theory
Stated that…
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 also used symbols to represent elements and developed the first table of atomic masses
1897-1906 Joseph John Thomson and his discoveries
Thomson discovered that cathode rays (recently discovered) were a stream of particles that…
Had a negative charge
Could be deflected by a magnet and an electric field
Had a very small mass
These were electrons. He disproved the theory suggested by the Greek and Dalton
What was Thomson’s plum pudding model?
Plum pudding model:
Sea of positive charge with negatively charged electrons on the surface
Charges are balanced
1909-11 Rutherford’s gold leaf experiment
Alpha particles directed at a very thin sheet of gold foil
They measured the deflection of these particles
The results…
Most didn’t deflect
Some that did deflect, deflected through large angles
Very few deflected towards the source
Rutherford’s new model
The new model stated that…
The positive charge of an atom and most of its mass are concentrated in a central nucleus
Negatively charged electrons orbit around this nucleus
Most of the atom’s volume would be the space between the nucleus and the electrons
Positive and negative charges must balance
1913, Niels Bohr’s planetary model
Altered Rutherford’s model to allow electrons to follow only certain paths. Otherwise, electrons would spiral into the nucleus
What did Bohr’s model help explain?
Spectral lines seen in emission spectra
Energy of electrons at different distances from the nucleus
What did Henry Moseley discover in 1913?
He discovered a link between X-Ray frequencies and an element’s atomic number
Couldn’t explain it at the time
Fifth century BCE, Greek Atom
Said that a sample of matter could only be divided a certain number of times
Eventually end up with a particle that doesn’t split
What did Rutherford discover in 1918?
The proton
What did the discovery of the proton reveal?
It explained Moseley’s finding that an atom’s atomic number was linked to x-ray frequencies
atomic number = number of protons
What did Louis de Broglie suggest in 1923?
Particles could have the nature of both a wave and particle
What did Schroedinger suggest in 1926?
He suggested that an electron had wave-like properties
Also introduced the idea of atomic orbitals
What did Chadwick discover in 1932?
And how did he discover this?
The neutron
He observed a new type of radiation that was emitted from some elements. This radiation was made up of uncharged particles with approximately the same mass as a proton
What are protons and neutrons thought to be made up of?
They are thought to be made up of even smaller particles called quarks
What are isotopes?
Isotopes are atoms of the same element with different numbers of neutrons and different masses
However, they have the same number of protons and electrons
The current model of the atom
Protons and neutrons are found in the nucleus, which is at the centre of the atom
Electrons orbit the nucleus in shells
The nucleus is tiny compared with the total volume of an atom
Nucleus is however extremely dense and accounts for almost all of the atom’s mass
Most of an atom consists of empty space between the tiny nucleus and electron shells
The proton
Positively charged, 1+
Relative mass of 1
The neutron
No charge
Relative mass of 1
The electron
Negatively charged, 1-
Relative mass of 1/2000
Reactions of isotopes
Different isotopes of the same element react in the same way
This is because…
Isotopes have the same number and arrangement of electrons
Neutrons make no difference to chemical reactivity
Why are ions charged?
Ions are charged because they have unbalanced numbers of protons and electrons, therefore the charges no longer cancel each other out
What is relative isotopic mass?
The mass of an atom of an isotope compared with 1/12th mass of an atom of carbon-12
What is relative atomic mass?
The weighted mean mass of an atom of an element compared with 1/12th mass of an atom of carbon-12
