Trends in the Periodic Table Flashcards Preview

LC Chemistry > Trends in the Periodic Table > Flashcards

Flashcards in Trends in the Periodic Table Deck (23)
Loading flashcards...

What happens when you go along a period?

Atomic radius decreases
Increase in effective nuclear charge
No increase in screening effect


What happens when you go down a group?

Atomic radius increases
New energy level
Screening effect


Reasons the atomic radius increases down the groups

1. Additional electrons are going into a new energy level (shell) which is further from the nucleus - atomic radius increases
2. Screening effect of inner electrons. Increase in nuclear charge is lessened bc of what is called a screening effect. Even though nuclear charge increases, increase is counteracted by shielding effect of inner energy levels of electrons.


Screening effect

In atoms w/ many electrons, electrons in inner energy level or levels help to screen/shield outer electrons from positive charge in nucleus. Attractive force of positive charges in nucleus is lessened by inner energy levels of electrons.


Reasons atomic radius decreases across a period

-Increase in effective nuclear charge. Number of protons increases from left to right, greater attractive force on the outer electrons. Draws energy levels closer to nucleus. Atoms become smaller, ie. atom radius decreases

-No increase in screening effect. Extra electron is going into same outer energy level. No increase in screening effect. No additional energy level of electrons to counteract increasing attractive force of the nucleus. ie. atomic radius decreases


First ionisation energy of an atom

Minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state


Reasons for why first ionisation energy decreases down a group

-Increasing atomic radius. Outermost electrons are becoming further away from attractive force of nucleus. Easier to remove electron from outer energy level.

-Screening effect of inner electrons. Inc. in nuclear charge cancelled out by intervening energy levels. Outermost electrons shielded from attractive force of nucleus + so are easier to remove


Reasons for why first ionisation energy increases across a period

-Increasing effective nuclear charge. Left to right across Periodic Table, number of protons increasing. Attraction between nucleus + outer electrons increasing without any increase in screening. Electrons held more firmly, required more energy to remove electron.

-Decreasing atomic radius. Atomic radius decreases from left to right. Electron in outermost level becoming closer to nucleus. More difficult to remove electron due to increased attraction.


Exception to the general trend across a period

-First ionisation energies of beryllium + nitrogen are higher than expected

-Irregularities in first ionisation energy trends are explained by fact that any sublevel that is completely filled or exactly half filled has extra stability

-Their first ionisation energy values are higher than expected since it required more energy than expected to break up these stable arrangements by removing an electron


Second ionisation energy

The energy required to remove an electron from an ion with one positive charge in the gaseous state


why there is a general increase in electronegativity values across periods in periodic table of elements

-bigger nuclear charge
-decreased atomic radius


why is there an increase in electroneg value moving from gallium to germanium in periodic table?

-nuclear charge increasing
-atomic radius decreasing


why first ionisation energy of smth can be greater than another's

-greater nuclear charge
-smaller atomic radius


why first ion energy of smth can be less than another's

-greater atomic radius
-most loosely-bound electron more shielded from nucleus


how to explain how a graph of ionisation energies provides evidence for energy levels

-(number) groups of ionisations with gradual energy differences between them bc they involve electrons in same energy level/shell
-(number) bigger energy differences between these groups of ionisations bc the (number) energy levels have significantly different energies


why there might be a (/large) decrease in first ion energies between two elements

-Has stable outer octet of electrons
-half-full x sublevel has associated stability
-more stable electron configuration
(eg. Oxygen has 2pₓ² 2pᵧ¹ 2p𝓏¹ while Nitrogen has 2pₓ¹ 2pᵧ¹ 2p𝓏¹

-new energy level (inc atomic radius)


unit used to measure first ionisation energy

kilojoules per mole
kJ mol⁻¹


using X to represent an element, express the definition of first ionisation energy in the form of a balanced equation

X₍𝓰₎ --> X⁺₍𝓰₎ + e⁻


would it take more or less energy to remove the most loosely bound electron from an atom if that electron were not in its ground state?

-already gained energy
-further from nucleus


an element has low first ion energy and electroneg. How reactive is it likely to be, and what is likely to happen to it when they react?

-very reactive
-lose electrons/oxidised/ become positively charged


the bond energy vaues of the C-Cl bond, the C-Br bond, and the C-I bond are 338, 276, and 238 respectively. Suggest a reason for this trend.

-increased atomic radius
-atoms getting bigger
-decreased electroneg difference


which element shows behaviour closes to that of an ideal gas?



why are the atomic radii of chromium, copper and zinc similar?

no change in effective nuclear charge / offset by changed shielding / all d-block