Atomic Structure Flashcards
Define Nucleus
Most of the mass of an atom and very small. Contains protons and neutrons
Define electrons
Orbit the nucleus in shells and take most of the space of an atom.
Proton relative charge
+1
Proton relative mass
1
Neutron relative charge
0
Neutron relative mass
1
Electron relative charge
-1
Electron relative mass
1/1840
Mass number
number of protons and neutrons in the nucleus
Atomic number
Protons in the nucleus
Define ion
Different number of electrons and protons
Negatively charged ions
Gained an electron
Positively charged ions
Lost an electron to form a full shell
Define isotopes
Elements with the same proton number but a different number of neutrons
Define relative atomic mass
the weighted mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12
Define relative molecular mass
mean mass of a molecule, compared to 1/12th of the mass of an atom of carbon-12
Relative isotopic mass
the mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon-12
What axis is the abundance on?
y
What is abundance shown as
%
What is m/z
the mass of an isotope divided by charge. As most have just a + charge this is the same as the isotopic mass.
How to find relative atomic mass (mass spectra)
Relative atomic mass = (abundance A X m/z A) + (Abundance B X m/z B)
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Total abundance
Stages of predicting mass spectra
1) Write the percentages as decimals
2) Create a table showing the isotope combinations in a molecule of whatever your dealing with. Multiply the decimal from abundances of each isotope to get the relative abundance for each molecule.
3) Any molecules which are the same add the abundances up
4) . Divide all the relative abundances worked out before by the smallest value. This will give you a whole number ratio which can be used to predict your spectra.
Mass spectra- molecules
What do peaks show?
What is the significance of the M+
Fragments of the original molecules
The M+ peaks is the last peak or the molecular ion peak, which is the same as the relative molecular mass of the molecule
Mass spectra- molecules
What is the m/z meaning
m/z is just the mass of a fragment divided by charge. As most have just a +1 charge this is the same as the fragment mass.
Electron configuration- subshells
Electron shells are split into 4 subshells
S Subshell
Orbital number + how many electrons can they hold?
s – has 1 orbital can hold 2 electrons
P subshell
Orbital number + how many electrons can they hold?
p – has 3 orbitals can hold 2 × 3 = 6 electrons
d subshell
Orbital number + how many electrons can they hold?
d – has 5 orbitals can hold 2 × 5 = 10 electrons
f subshell
How many orbitals + how many electrons can they hold?
f – has 7 orbitals can hold 2 × 7 = 14 electrons
S orbital shape
The s orbital is spherical and the 2 electrons can move anywhere within this sphere.
P orbital shape
There are 3 p orbitals in the shape of dumbbells each can hold up to 2 electrons can move anywhere within this shape.
What is Hund’s Rule
‘all orbitals of equal energy will be singly occupied before any are doubly occupied’, meaning electrons will only pair up only when other degenerate orbitals aren’t empty
What is Pauli’s Exclusion Principle?
‘no two electrons can have the same quantum numbers/ be in identical quantum states’. Hence two electrons in the same orbitals must have opposite spins
With ions what is the rule for removing/adding electrons?
With ions you just add or remove electrons from the highest energy level first.
What is different about Chromium and Copper’s electron configuration?
An electron from the 4s orbital moves into the 3d orbital to create a more stable half full or full 3d sub-shell respectively.
Which groups are the s block elements?
1/2
Which groups are d block elements?
transition elements
Which groups are known as the p block elements?
3-8
Where are the f block elements located?
The block at the bottom
What does the electromagnetic spectrum show?
The electromagnetic spectrum shows types of radiation at different frequencies.
What is line spectra ?
a way of identifying elements and is evidence for quantum shells
Stages of mass spectrometry
Vaporisation Ionisation Acceleration Deflection Detection
Briefly describe vaporisation
The sample is converted to gaseous state so it can move through the machine
Briefly describe ionisation
An electron gun fires fast-moving electrons at the particles to knock an electron away and form 1+ ions
Briefly describe acceleration (mass spectrometry)
The ions are all accelerated so they have the same kinetic energy
Briefly describe deflection
An electromagnet is used to deflect the ions by creating a magnetic field: the more massive the ions are, the less they are deflected, and the higher the charge on the ions, the more they are deflected
Briefly describe detection
The ions that have made it through the whole machine are detected at the end and a mass spectrum is made electronically
Why is the vacuum pump an important feature?
It is important that the ions produced in the ionisation chamber are able to move through the machine without bumping into air molecules; hence a vacuum is created
Why is the electron gun an important feature?
The particles have to be ionised or else they would not be able to be repelled by the field created by the electromagnet as they would be uncharged
What is changed during mass spectrometry?
The intensity of the magnetic field- to allow for the masses of ions to be calculated
Define first ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions (measured in kJmol-1)
Define successive ionisation energies
The energies required to remove each mole of electrons in turn from one mole of gaseous atoms/ions (measured in kJmol-1)
How does the proton number affect ionisation energy?
The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons. Having more protons increases the ionisation energy
What is the importance of the number of electrons between the outer electrons and the nucleus increasing?
The number of electrons increase so therefore the outer electrons feel less attraction towards the nuclear charge. This is caused by repulsion from electrons in inner shells which lessens this attraction; shielding is the name of this effect.
More greatly shielded electrons decreases the ionisation energy
What is the importance of the distance between outer electrons and the nucleus?
Outer electrons in shells close to the nucleus are much more strongly attracted to it those that reside in much further away shells, and so therefore a larger distance between outer electrons and the nucleus decreases ionisation energy
Explain why there is an increase in first ionisation energies of period 2 elements
Proton number increases, so there is a stronger attraction between the nucleus and outer electrons and so more energy is required to remove it
Trends in first ionisation energy down a group
general decrease
Increase nuclear charge …
Increased nuclear charge + Higher energy quantum shells + More shielding = General decrease
What do big jumps in ionisation energy show?
Big jumps in ionisation energy show where an electron is being removed from a new shell and hence looking at an element’s successive ionisation energies can show which group it is in.
Why is there a drop in between groups 2 and 3
In group 2- the outer electrons are in s orbital, while group 3 outer electrons are in the p orbital
P orbital is higher in energy- further distance, shielded meaning an decrease in ionisation energy
Why is there a drop between group 5 and 6
group 5- singly occupied
group 6- 4th electron pairs with the 1st. Electron-electron repulsion which makes the p orbital easier to remove (unstable configuration)
Identical nucleus distance, identical shielding, decrease in ionisation energy
Define orbital
Orbitals are 3D regions within an atom where electrons have a high probability of being found. Electrons must have opposite spins in order to exist within the same orbital.
Explain how electron configuration determines the chemical properties of an element
Electronic configuration determines the likelihood that an element would donate/accept an electron / multiple electrons.
This therefore determines how an element may choose to bond with other elements. The nature of this bonding determines an element’s physical characteristics as well as its chemical characteristics.
Define periodicity
Periodicity is a regularly repeating pattern of atomic, chemical, and physical properties as atomic number increases.
Trends in electronegativity
Electronegativity increases across a period due to an increase in effective nuclear charge.
This decreases the atomic radius, and the combination of the greater nuclear charge and smaller distance between electrons from other atoms, increases the atoms tendency to attract bonding electrons, hence increasing its electronegativity.
What is atomic radius
the farthest point form the nucleus where an electron can be found.
Trends in atomic radius across a period
Atomic radius decreases across a period due to an increase in effective nuclear charge.
This decreases the atomic radius because it attracts the electrons more strongly to the nucleus, hence pulling them closer towards it, decreasing the distance between them.
