3.1.1 Atomic Structure (Physical Chemistry) Flashcards
Isotopes definition
Atoms with the same number of protons and electrons but a different number of neutrons
Why do Isotopes of the same element have the same chemical properties
Isotopes of the same element have the same chemical properties as they have the same electron configuration
Define first ionisation energy
The amount of energy required too remove one mole of electrons from one mole of atoms in the gaseous state to form one mole of 1+ ions
Write an equation for first ionisation energy
X(g) —> X+(g) + e-
Name three key factors that influence ionisation energy
1.Nuclear Charge
-more protons leads to a higher positive charge on the nucleus giving a stronger attraction between the outer electron and nucleus resulting in a higher ionisation energy
2.Distance from the nucleus
-The shorter the distance between the outer electron and nucleus, the stronger the attraction between them leading to a higher ionisation energy
3.Shielding
-The more shells in an atom leads to more shielding resulting in a weaker attraction between the outer electron and nucleus resulting in a lower ionisation energy
State the general trend of first ionisation energies across a period
-The first ionisation energy will increase
-There are more protons in thee nucleus
-Nuclear charge increases
-Shielding remains the same
-Attraction between outer electron and nucleus increases
Describe the first Ionisation Energy trend across period 2
-General increase in ionisation energy across a period
-Same shielding but a greater nuclear charge so a stronger attraction between outer electron and nucleus
-Drop at Boron
-The first electron removed from Be is in a 2s sublevel whereas the first electron removed from B is in a 2p sublevel
-2s sublevel is lower in energy than the 2p sublevel and so more energy is required to remove the electron from Be
-Drop at Oxygen
-First electron in N in the 2p sublevel is unpaired whereas in O it is paired
-less energy is required to remove the electron from O due to electron pair repulsion
Describe the first Ionisation Energy trend across period 3
-General increase in ionisation energy across a period
-Same shielding but a greater nuclear charge so a stronger attraction between outer electron and nucleus
-Drop at Aluminium
-The first electron removed from Mg is in a 3s sublevel whereas the first electron removed from Al is in a 3p sublevel
-3s sublevel is lower in energy than the 3p sublevel and so more energy is required to remove the electron from Mg
-Drop at Sulfur
-First electron in P in the 3p sublevel is unpaired whereas in S it is paired
-less energy is required to remove the electron from S due to electron pair repulsion
Describe the trend in successive ionisation energies and why
-As the series of ionisation energies increases, ionisation energy increases
-The electron is being removed from an increasingly positive ion and so the attraction between outer electron and nucleus is stronger
Why is Li a bigger atom than Be
-Both atoms have the same shielding as they have the same number of shells
-Be has more protons / a higher nuclear charge
-the outer electron in Be is attracted more strongly to the nucleus
-atomic radius decreases
Why is Li a bigger atom than He
-Li has more shielding than He as it has one more electron shell
-The outer electron is more shielded
-It is less strongly attracted to the nucleus
Electron shells and the max number of electrons in each shell
1=2
2=8
3=18
4=32
n=2n squared
Define orbitals
A region within the atom that can hold up to 2 electrons with opposite spins
S and P orbital shape
s- spherical
p- dumbell
State the order in which subshells fill up
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
Electron configurations for chromium and copper and why
Chromium:
1s2 2s2 2p6 3s2 3p6 4s1 3d5
Copper:
1s2 2s2 2p6 3s2 3p6 4s1 3d10
-D block elements are more stable with a full or exactly half full D subshell
Trend in Ionisation energies down a group
-Atoms get bigger
-Atomic radius increases
-Weaker attraction between outer electron and nucleus
-Ionisation energy decreases going down a group
State and Explain the trend in atomic radius going down a group and across a period
Down a group:
-Atomic radius increases
-Number of shells increases
-Shielding increases
Across a period:
-Nuclear charge increases
-Shielding stays the same
-Greater attraction between outer electron and nucleus
Whose model was supported by this experiment?
What was detected at point P?
What was detected at point Q?
Ernest Rutherfords nuclear model - 1911
Alpha scattering experiment:
-He fired He2+ ions at a thin sheet of gold foil
Point P:
-Most arrived at point P, passed through
-Most of the atom was empty space
Point Q:
-Very small amount detected at point Q, deflected back
-Nucleus is small and positively charged
Describe how a TOF mass spectrometer works
- The entire machine is a vaccum preventing particles colliding with air
- Ionised via electrospray Ionisation or electron impact ionisation to form 1+ions
- Accelerate ions towards negatively charged plate, ions with a high m/z ratio will accelerate at lower speeds , all ions have the same kinetic energy
- Ion drift - some pass through a hole in the negatively charged plate to form a beam, ions travel at different speeds
- Detection: Each ion hits the detector, gains an electron generating a current, the size of which is proportional to the number of each type of ion
Electron Impact Ionisation
-The sample is vaporized
-High energy electrons are fired at the sample from an electron gun
-This knocks off one electron = 1+ ion
X(g) —> X+(g) + e-
-can result in fragmentation
-used for substances with a low formula mass
Electrospray Ionisation
-Sample dissolved in a volatile solvent
-Injected through a fine hypodermic needle
-Fine mist
-Tip of needle attached to the positive terminal of a high voltage power supply
-Ionised by gaining a proton
X(g) + H+ –> XH+(g)
-fragmentation rarely occurs
-used for substances with a high molecular mass
Ar calculation formula
Kinetic energy formula
