topic 7 - periodic trends

atomic radius is defined to be __

half the distance between the nuclei in a molecule consisting of identical atoms.

how can you describe the trend of atomic radius on a periodic table?

Atomic radius increases down a group and from right to left across a period.

why does the atomic radius increase down a group?

atomic radius increases down a group because the number of shells an atom has increases —> larger radius

why do atomic radius decrease from left to right across a period?

Because the interactions between the protons and electrons —> a smaller electron cloud.

the number of protons an atom has increases —> stronger pull placed upon valence electrons toward nucleus —> smaller electron cloud —> decrease radius

which main group element has the largest atomic radius?

Cs

What law can be used to explain the periodic trend when it comes to atomic radius?

Coulomb’s Law: opposite charges attract whilst like charges repel

electrostatic force is proportional between the product of 2 point charges (q1 and q2) over the radius squared

Cations __ in size compared to their atom counterpart because _

decrease

they hold less electrons compared to the atom counterpart + outer shell electron(s) are lost

anions __ in size compared to their atom counterpart because _

increase

they have more electrons than their atom counterpart —> electron cloud gets bigger

define isoelectronic species

ions that have the same number of electrons, but not of the same element/don’t have the same number of protons

atomic radius trends for isoelectronic species

radius decreases as atomic number/number of protons increases

define ionization energy

required amount of energy to remove an electron from an atom in the gas phase

write chemical reaction/equation for the third ionization energy of an atom (X)

I3

X²+ (g) —> X³+ (g) + e^-

what does the first ionization energy do?

it shows the amount of energy required to remove the outermost (first) electron in an atom in its gas state

successive ionization energies of the same atom are …

larger.

why are electrons in a cation more difficult to remove than an electron in an atom

because electrons in a cation experience fewer electron-electron repulsions → they are more difficult to remove.

in graphs showing the ionization energy of different elements, sudden jumps between ionization energies indicate:

that the electron being removed is a core electron. This can help to show what group the atom/ion belongs to

how to describe the periodic trend of ionization energy?

as you move across a period from left to right, atoms get smaller → ionization energy increases

as you move up a group, ionization energy increases.

both of these occur because the atom gets smaller when moving through the periodic table in these directions. The closer the electron is to the nucleus, the more difficult they are to remove.

define shielding effect

the phenomenon where inner-core electrons repel outer electrons → reduces the net attraction exerted by the nucleus

ex: Mg and Al

Mg: [Ne] 3s²

Al: [Ne] 3s² 3p^1

the 3p^1 electron in Al is easier to remove/requires less energy than one in the s orbital

list the atoms that are exceptions to the general ionization energy periodic trend due to electron shielding

Be, B

Mg, Al

Ca, Ga

define electron-electron repulsion

paired electrons are required less energy to remove

exceptions to the general periodic trend of ionization energy because of electron-electron repulsions

N,O

P,S

As,Se

define electron affinity

the heat released when a gaseous atom gains an electron to create an anion

E = -delta H

larger electron affinity indicates that the ion will be

more stable

describe the periodic trends of electron affinity

• Cl has the greatest EA

• Halogens (group 17) have the largest EAs

• EA decreases down a group as atomic size gets bigger.

  • bigger atom → electrons added are further away from the nucleus → binding energy is weaker

• elements in the P block of period 2 have bigger EAs than expected (B and O)

• Group 1 + 2 metals have lower EA → don’t readily form anions’

• noble/inert gases + N, Be, and Mg have no EA values listed.

define Zeff (effective nuclear charge)

the number of elementary charges an electron experiences by the nucleus

describe periodic trend of effective nuclear charge

Zeff increases down a group and from left to right across a period