atomic structure Flashcards
(31 cards)
Periodic table basic properties [brief]
Period - LEFT TO WRITE
→ Atomic number increases
→ Electrons are added to the same energy level or shell
→ Nuclear charge increases ( refers to the total positive charge present in an atom’s nucleus, which is directly determined by the number of protons it contains; essentially, the higher the atomic number = the greater the nuclear charge = stronger attraction between the nucleus and the surrounding electrons)[The effective nuclear charge (Z effective or Zeff) is defined as the net positive charge pulling these electrons towards the nucleus.]
Periodic table basic properties Periodic table basic properties
→ Atomic Number Increases
→ The number of electron shells (or energy levels) increases
→ Shielding effect increases(occurs when electrons in the inner shells of an atom repel electrons in the outer shells) [ Happens due to the increase no. of shells + atomic number] [more inner electron shell which repels the outer electron]
charge of a single electron
-1.602 x 10-19
Atomic radius:
It is half the distance between the two nuclei of two covalently bonded atoms of the same type.
Trend across period
→ They generally decrease across each Period
Atomic radii decrease as you move across a Period as the atomic number increases which increases the positive nuclear charge but at the same time extra electrons are added to the same principal quantum shell.The larger the nuclear charge, the greater the pull of the nuclei on the electrons which results in smaller atoms
Trend down group
→ They generally increase down each Group
Atomic radii increase moving down a Group as there is an increased number of shells going down the Group
The electrons in the inner shells repel the electrons in the outermost shells, shielding them from the positive nuclear charge
This weakens the pull of the nuclei on the electrons resulting in larger atoms
Ionic radius:
The ionic radius of an element is a measure of the size of an ion
Trend of ionic radius with increasing negative charge
→ Ionic radii increase with increasing negative charge
Ions with negative charges (anions) are formed when atoms gain extra electrons. The nuclear charge remains the same because the number of protons doesn’t change. However, as the atom gains more electrons, the outermost electrons are pushed farther from the nucleus, and the attraction between them and the nucleus weakens. This causes the ionic radius to increase. The more negative charge an ion has, the larger its ionic radius, due to the increased electron-electron repulsion(shielding effect) and weaker pull from the nucleus.
Trend of ionic radius with increasing positive charge
→ Ionic radii decrease with increasing positive charge
Positively charged ions are formed by atoms losing electrons
The nuclear charge remains the same but there are now fewer electrons which undergo a greater electrostatic force of attraction to the nucleus which decreases the ionic radius
The greater the positive charger, the smaller the ionic radius
ISOTOPE
Chemical properties
Isotopes of the same element display the same chemical characteristics
This is because they have the same number of electrons in their outer shells
Electrons take part in chemical reactions and therefore determine the chemistry of an atom
ISOTOPE
Physical properties
The only difference between isotopes is the number of neutrons
Since these are neutral subatomic particles, they only add mass to the atom
As a result of this, isotopes have different physical properties such as small differences in their mass and density
Free radicals
A free radical is a species with one or more unpaired electron
- Full-Fill Orbitals:
Full-fill refers to an orbital that is completely filled with electrons.
Each orbital can hold a maximum of 2 electrons (due to the Pauli Exclusion Principle, which states that two electrons in the same orbital must have opposite spins).
A full orbital means that both available electron spots (with opposite spins) in the orbital are occupied.
- Half-Fill Orbitals:
Half-fill refers to an orbital that is partially filled with electrons, specifically 1 electron in an orbital.
This is commonly observed in the p, d, or f orbitals where each orbital can hold more than 2 electrons. According to Hund’s Rule, electrons will fill degenerate orbitals (orbitals of the same energy level, such as the 3 p-orbitals or the 5 d-orbitals) one by one before pairing up, to maximize the number of unpaired electrons.
When an orbital has exactly one electron, it is said to be half-filled.
IONISATION ENERGY
The ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions
Ionisation energies are measured under standard conditions which are 298 K and 101 kPa
The units of IE are kilojoules per mole (kJ mol-1)
The values for ionisation energies are always positive as this is an endothermic process
→This is because energy is required to break the force of attraction between the electron and the central positive nucleus
First ionisation energy
The first ionisation energy (IE1) is the energy required to remove one mole of electrons from one mole of atoms of an element to form one mole of 1+ ions
Trends in Ionisation Energy
As could be expected from their electronic configuration, the group I metals have a relatively low ionisation energy, whereas the noble gases have very high ionisation energies
[ Group 1 metals have 1 electron in their valence which is easier to lose electron and thus require less energy to remove electron while noble gases have full orbitals making it harder to lose electron and thus requiring more energy to remove electron. ]
The size of the first ionisation energy is affected by which three factors:
Size of the Nuclear Charge:
Spin-Pair Repulsion:
Distance of Outer Electrons from the Nucleus & Shielding Effect of Inner Electrons:
Trends in Ionisation Energy
ACROSS THE PERIOD
→ ACROSS THE PERIOD - INCREASES
As atomic number increases, the nuclear charge becomes stronger, resulting in a greater attraction between the nucleus and electrons, thus requiring more energy to remove an electron.
Spin-Pair Repulsion:
Increases across the period when orbitals are filled.
Electrons in the same orbital repel each other more than those in different orbitals, making it easier to remove an electron, resulting in lower first ionisation energy.
Trends in Ionisation Energy
DOWN THE GROUP -
→ DOWN THE GROUP - DECREASES
Distance of Outer Electrons from the Nucleus & Shielding Effect of Inner Electrons:
Increases down the group.
As the number of electron shells increases, the outer electrons are farther from the nucleus and experience a weaker attraction. Additionally, the inner electrons shield the outer electrons from the full effect of the nuclear charge.
Z_eff stays relatively constant or increases slightly, but with more electron shells added, the outer electrons are farther from the nucleus and experience more shielding. As a result, ionisation energy decreases down the group.
Ionisation energy decreases down the group.
Successive Ionisation Energies of an Element
The successive ionisation energies of an element increase
This is because once you have removed the outer electron from an atom, you have formed a positive ion
Removing an electron from a positive ion is more difficult than from a neutral atom
As more electrons are removed, the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio
The increase in ionisation energy, however, is not constant and is dependent on the atom’s electronic configuration
Why is the second ionisation energy of sodium much higher than the first?
The second electron is removed from a lower energy level (2p) closer to the nucleus, with less shielding and higher nuclear attraction.
: Why does aluminium have a lower first ionisation energy than magnesium, despite a greater nuclear charge?
Al’s outer electron is in a 3p orbital (higher energy, more shielding) compared to Mg’s 3s.
Why do noble gases have high first ionisation energies?
Full outer shells → very stable; high nuclear attraction and minimal shielding.
