2.1 Flashcards
Niels Bohr Atomic Model
Nuclear model - atom consists of a nucleus made up of protons and neutrons with electrons occupying fixed energy levels (shells) around it
Why do we use relative mass?
Because the charges and masses of the atoms are incredibly minuscule - therefore it makes it more efficient and easier for calculations
Mass of subatomic particles?
P+ - 1
E- - 1/1836
N - 1
Neutron is slightly heavier than a proton bu this factor is negligible
Strong nuclear force?
Neutrons can be thought of as providing the four to overcome the electrostatic forces of repulsion between protons within the nucleus - therefore most atoms contain the same number of (or more) neutrons than protons and as the nucleus gets larger more and more neutrons are added
Importance of atomic number?
The number of protons is what identifies the element- every atom of the same element has the same number of protons and different elements contain atoms with different number of protons ; THIS IS SHOWN IN PERIODIC TABLE WHICH LISTS ELEMENTS IN ORDER OF THE NUMBER OF PROTONS IN NUCLEUS
Isotopes?
Every atom of an element has the same number of protons but different forms of an atom of an element can consist of varying numbers of neutrons (different atomic masses) - MOST ELEMENTS ARE MADE UP OF A MIXTURE OF ISOTOPES
Mass number
Number of protons + number of neutrons - A
Atomic number
Number of protons - Z
Isotopes and chemical reactions
HAVE SAME CHEMICAL PROPERTIES BECAUSE NUMBER OF ELECTRONS DOES NOT DIFFER - number of neutrons does not affect chemical reactions BUT does affect boiling point, melting point and density (physical properties) - higher mass = higher BP, MP, density
Deuterium vs Water
Deuterium heavy water (2 as mass number - 1 neutron) ; therefore it is denser and has higher MP making it freeze earlier - D2O vs H2O - called heavy water because of denser nature
Cations
Positively charged - lost electrons
Anions
Negatively charged - gained electrons
Ions
Charged atoms that have the same number of protons but have either lost or gained electrons to become positively or negatively charged
Why can we not add the relative masses of the subatomic particles to find the relative mass of an isotope?
This is because the strong nuclear force holding together the protons and neutrons comes at the expense of the loss of a fraction of their mass - MASS DEFECT
Standard isotope
Carbon-12 isotope - atomic mass unit is used (because in kg the weight is too minuscule and therefore difficult) ; mass of carbon-12 = 12 atomic mass units and the standard mass for atomic mass in 1u (1/12 * 12 = 1)
1u = mass of proton
Relative isotopic mass
The mass of an isotope relative to 1/12th the mass of a carbon-12 atom ; has no units because it is a ratio of the two masses - in most cases we can assume the mass number (A) = relative isotopic mass
Relative atomic mass
Ar is the weighted mean mass of an Tom of an element compared to 1/12th the mass of an atom of carbon-12
Takes into account :
Percentage abundance of each isotope
Relative isotopic mass of each isotope
Mass Spectrometry
Used to find out percentage abundance and relative isotopic masses
1) Sample is placed in a vaporised state in mass spectrometer
2) They are ionised to form +1 cations
3) Ions are accelerated and the heavier ions move more slowly and are thus more difficult do deflect than lighter ions ; thus this separates the isotopes
4) Ions are detected on a mass spectrum with mass to charge ratio on x-axis and percentage abundance on y-axis (each ion reaching the detector adds to the signal)
Mass to charge ratio
m/z => with 1+ positive charge this is equal to relative isotopic mass (X-axis) - mass spectrometer records the accurate m/z Dario for each isotope so that accurate values of relative isotopic mass can be measured
Relative atomic mass
Percentage abundance * mass number / 100
How can there be inaccuracies in recording relative atomic masses?
DEPENDING ON WHERE SAMPLE OF ELEMENT ORIGINATES
