Chemistry UNIT 1 periodic trends

Atomic radius

the size of an atom, typically defined as the distance from the nucleus to the outermost electron

effective nuclear charge

The net positive charge (of the protons) pulling the electrons towards the nucleus

what happens to the atomic radius in a period (left to right)

as you go left to right the number of protons in an atom increases as the atomic number increases. As a result the effective nuclear charge increases as well as the increasing number of protons begin to pull the electrons closer to the nucleus. Thus, the distance between the nucleus and the outermost electron decreases and atomic radius becomes smaller

Atomic radius with group/family

as you go down a group the # of electron shells increases, because there is now more shells it acts as shielding from the outermost electrons which leads to the electrons facing less attraction to the protons causing them to not shift and if not stay farther away from the nucleus. because of this going down a group atomic radius becomes larger

Trend in atomic radius

left to right atomic radius decreases, top to bottom atomic radius increases

The trend in Electronegativity

Increases across a period and increases up a group

What is electronegativity?

The tendency for an atom to attract a bonding pair of electrons in a covalent bond

What is electronegativity dependent on?

atomic radius

The trend of electronegativity across period

as you move across a period the effective nuclear charge increases making atoms more attractive towards electrons

Electronegativity down a group

The distance between the outer electrons and the nucleus increases, leading to lower electronegativity (smaller effective nuclear charge thus atoms are not as attractive to electrons)

Ionization energy

the energy required to remove an electron from an atom

What is the trend for ionization energy?

left to right: increase, bottom to top: decrease

Ionization energy in period

Moving across a period the effective nuclear charge increases making it more difficult to remove electrons thus ionization energy increases as the effective nuclear charge is strongly holding the electrons in place

Ionization energy moving down a group

moving down a group electrons are farther away from the nucleus as shells are increasing due to increasing periods therefore the ionization energy decreases as there is a weak effective nuclear charge and the outer electrons can be easily removed as it requires less energy

Electron affinity

the energy change when an atom gains an electron to form a negative ion

Electron Affinity in Period

Across a period (L to R), the effective nuclear charge increases, making it more favorable for an atom to accept an electron to gain a stable/full octet

Electron affinity in a group

down a group the added electron is farther from the nucleus, resulting in a lower electron affinity this is because sheilding causes the valence electrons to not feel a strong nuclear charge resulting in a weak pull towards the nucleus. Because there is a weaker pull it causes there to be less energy released when the electron is added

Why do noble gasses have a ionization energy?

noble gasses have a high ionization energy because their octet is already stable. These are atoms that already have a full number of electrons and do not want or need more. This causes them to have high ionization energy because they would need a very VERY strong pull of energy to pull/remove an electron.