Atomic Structure Flashcards
Subatomic particles
All elements are made of atoms. - particles in atoms: protons, electrons and neutrons
- electrons found surrounding nucleus in orbitals
Charges:
- protons = +1
- neutrons = 0
- electrons = -1
Mass
- protons = 1
- neutrons 1
- electrons = 0
Describe the arrangement of subatomic particles
- protons and neutrons are in the nucleus - nucleons
- held together by strong nuclear force
- electrons surround the nucleus in orbitals
- held in the atom by electrostatic forces between protons and electrons - nuclear force is much stronger than electrostatic forces
Explain what john dalton stated about the atom
- atoms were solid spheres that make up elements, different spheres for each elements
What did JJ Thompson discover
- plum pudding model, showing atoms contain electrons
What did Ernest Rutherford discover?
- conducted the alpha scattering gold foil experiment to produce nuclear model.
- fired positively charged alpha particles at very thin sheet of gold
- ppm suggested: most particles slightly deflected by positive ‘pudding.’
- actually —> most straight through (most atom is empty space) + some deflected (positive cenbtre)
What did Bohr discover
- adaptive nuclear model, putting electrons in shells/orbitals in fixed energy
- realised electrons in a ‘cloud around the nucleus spiral down into the nucleus, causing it to collapse
- when electrons move to a new shell, electromagnetic radiation is emitted or absorbed.
Mass number (A)
- total number of protons + neutrons in the nucleus. Nucleons are responsible for almost all the mass as electrons have tiny mass
Atomic number (Z)
- number of protons in the nucleus, equals the number of electrons in the atom meaning atoms are electrically neutral
- number of electrons in outer shell determines the chemical properties of the element.
- atoms of same element = same atomic number
relative atomic mass (Ar)
- average mass of an atom of an element, taking into account its naturally occurring isotopes, relative to 1/12th the mass of an atom of carbon-12.
- Ar = (isotope mass number x % abundance) + (isotope mass number x % abundance) / sum of % isotope abundance
Define relative isotopic mass
Mass of an atom of an isotope relative to 1/12th the mass of an atom of carbon-12. Always a whole number
Define relative molecular mass (Mr)
- average mass of a molecule in relation to 1/12 the mass of a carbon 12 atom
Mr = sum of Ar in the formula for the molecule
Define the term ‘isotopes’
An isotope is an atom of the same element with the same number of protons (atomic number) and a different number of neutrons (so diff mass number)
What are some characteristics of isotopes
- same chemical properties as they have the same electron configuration
- slightly different physical properties as they depend on the mass of the atom. E.g - density
- isotopes can be radioactive if they are unstable (extra protons/neutrons in nucleus creates extra energy)
—> emit radiation as they decay and the rate of decay measured by half-life.
Time of flight mass spectrometer
- powerful instrumental technique that’s useful for accurate determination of relative isotopic masses + relative abundances.
- work out relative atomic masses to identify elements and relative molecular masses
Describe the ionisation stage of TOFMS
- where each atom becomes an ion and can be done one of two ways depending on its mass:
- electron impact ionisation
- electrospray ionisation
Electron impact ionisation
- for compounds with low molecular mass.
- sample is vaporised then high energy electrons are fired at sample using an electron gun, one electron is knocked off each atom forming a 1+ ion.
X (g) —> (X)+ (g) + e-
Electrospray ionisation
- for compounds with high relative molecular mass
- sample dissolved in volatile solvent + injected into ionisation chamber through a hypodermic needle which has high voltage as its positively charged
—> particles gain a proton and become 1+ ions as a fine mist. Solvent evaporates leaving 1+ ions
X(g) + H+ —> (XH)+ (g )
Describe the stage of acceleration
- ions are accelerated using an electric field
- the positively charged ions accelerate towards a negatively charged plate
- all ions have same kinetic energy = same kinetic energy but their velocity than heavier ions
Describe the ion drift stage
- ions pass through a hole in the plate into the flight tube where they enter a region with no electric field so they drift through towards detector. —> ions with diff masses have diff time of flight
Describe the detection stage of TOFMS
- detector is a negatively charged plate and a current is produced when the ions pick up electron from detector, cause current to slow
- current is proportional to the abundance
Describe the data analysis stage
- signal from the detector is passed to a computer which generates a mass spectrum
Mass spectra of elements
- shows mass:charge ratio which is same as mass number as most have 1+ charge and relative % abundance of each isotope.
How to identify elements - mass spectra
- each isotope produces a line on the spectra, as they all have different masses.
- always put a + after the isotope of elements when identifying which reach detector first
Identifying molecular mass —> main peak on spectra at the Mr, highest, furthest right, most abundant
What does higher M/Z ratios isotopes indicate?
- isotopes of carbon and hydrogen.
- in organic compounds - often a peak at Mr, +1 due to C-13 and H-2 isotopes
What does single isotope peaks indicate ?
- often peaks at the mass of the isotopes, mononuclear ions forming monatomic fragments
What does lower M/Z ratios - fragmentation indicate?
- molecular ion argument into smaller particles, smaller ion - free radical
- if organic, CH covalent bonds break, CH3 often at 15 -> electron impact
Describe some characteristics of electron configurations
- electrons have fixed energy
- they move around the nucleus in shells/energy levels
- the further away from nucleus the shell is, the higher its energy
- sub shells have orbitals in which there are 2 electrons which spin in opposite directions
Describe aufbau’s principle
- always fill the lowest energy subshells first
- 1s,2s,2p,3s,3p,4s,3d,4p,4d
- 4s subshell has lower energy than 3d subshell, so fills up first
Describe hund’s rule
- electrons fill orbitals singularly before they start sharing
Explain chromium’s exception
- expect to be 4s2, 3s4 but, it donates of its 4s electrons to the 3d subshell so the 3d subshell is more stable as its half full.
- configuration actually: 1s2, 2s2, 2p6, 3p6, 4s1, 3d5
Explains copper’s exception
- expect it to be 4s2, 3d9 but, it donates one of its 4s electrons to the 3d subshell so 3d subshell more stable as it’s now full. Its configuration is actually: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
Describe the ‘ first ionisation energy.’
- the energy needed to remove 1 electron from each atom of an element in one mole of gaseous atoms to form one mole of gaseous ions a 1+ charge
X (g) —> X- (g) + e-
Describe the second ionisation energy
- the energy needed to remove a second electron from each ion in a mole of gaseous 1+ ions, forming a mole of gaseous 2+ ions
X+ (g) —> X2+ (g) + e-
How does atomic radius affect ionisation energy
- the smaller the atomic radius, the higher the first ionisation energy
- due to the outer electron more strongly attracted to the nucleus.
How does nuclear charge affect ionisation energy
- higher nuclear charge, the higher the first ionisation energy
- because a higher nuclear charge means a smaller atomic radius
—> as more protons to attract the outer electrons so more attracted to the nucleus
How does shielding affect ionisation energy
- more shielding, the lower the first ionisation energy
- inner shells repel the outer electrons making them easier to lose.
Describe the ionisation energy across a period
- increases —> higher nuclear charge, smaller atomic radius
- outer electrons more attracted to the nucleus
- shielding is constant
Describe the trend in ionisation energy down a group
- decreases —> larger atomic radius
- increase in shielding as more energy levels added
- outer electrons is less attracted to the nucleus and thus easier to lose
Describe the ionisation energy trends between group 2 and 3
- dip in first ionisation energy down
- move to new energy level with higher energy. S —> P subshell
- further away from the nucleus and more shielding
Describe the ionisation energy trends between groups 5 and 6
- dip in first ionisation energy
- group 5: p subshell has one electron in each orbital, no repulsion
- group 6: p subshell has one orbital with 2 electrons in
- the electrons in same orbital repel each other so less energy needed to remove outer electron
What does successive ionisation energies refer to
- more and more electrons are removes, each from an ion that is becoming increasingly positive
How does successive energies increase between shells?
- each time new shell broken into, there’s sudden rise in ionisation energy
- sodium - one outer electron: big jump in ionisation from first to second because second electron is being removed from a shell that’s closer to nucleus —> stronger attraction from nucleus
How does successive ionisation energies increase within each shell
- each time an electron removed, the ionisation energy increases as you are removing an electron from an increasingly positive ion
- like magnesium —> two outer electrons, second ionisation energy greater than the first, because it is harder to remove an electron from Mg2+ ion than Mg+ atom —> stronger attraction back
