3.1 the periodic table classification of elements

A column in the periodic table. elements in same group have similar properties

horizontal row in the periodic table; the period number tells you how many electron shells an element has.

An element that is located between metals and non-metals in the periodic table and shows intermediate properties between a metal and a non−metal. In the “stairs”.

by seeing which sublevel does its valence electrons occupy group 1 and 2= s-level transition metals= d block non metals= p-level

The distance from an atom’s nucleus to the outermost orbital of one of its electrons. Usually measured as the half-distance between the nuclei of two of the same atoms bonded together.

The net positive charge from the nucleus experienced by outer valence shell electrons once the screening effect of the core electrons has been considered.

elements have a greater number of protons in the nucleus and a greater number of electrons occupying the outermost energy level. This greater number of protons and valence electrons increases the effective nuclear charge

increases because going down group elements have more electrns that occupy higher energy levels which are found further away from the nucleus, meaning electrostatic force of attraction of nucleus to outermost electron is weaker/ less resulting in a bigger atomic radii

decreases as; the electrons occupy the same energy level however there is an increase in number of protons, meanig increase in the effective nuclear charge, bigger force o fattraction from nucleus to outermost electrons decreasing atomic radii

measure of size of ion, measured by distance between nucleus and outermost ion shell

positive ion=loss of electrons from outermost shell positive ion is smaller than its atom increase attraction between nucleus and outermost electron (less hsileding) bigger effective nuclear charge thus smaller the radii more electrons removed the smaller the radii less electron electron repulsion

negative ion=gained electrons radii is bigger than atom gain of electrons increases electron electron repulsion decreased attraction between the nucleus and outermost electron (decrease in effective nuclear charge) greater number of electrons gained bigger the radii

Molecules with the same electron configuration.

use effective nuclear charge

isolectronic species have same number of electrons but different number of protons bigger number of protons= bigger the effective nuclear charge, means; stronger force of attraction from nucleus to outermost elecctrons therefore smaller radii lower effective nuclear charge= weaker force of attraction from nucleus to outermost electron, bigger radii

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

The energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms in the ground state.

X(g) → X+(g) + e–

-Going down a group the first ionisation energy decreases because atomic radii increases and electrons occupy higher energy levels, Electrons in higher energy levels have a weaker attraction to the nucleus so Less energy is required to remove an electron.

Going across a period the first ionisation energy increases, nuclear charge increases (there are more protons in the nucleus). so the atomic radius decreases (atoms get smaller except for Ar which increases slightly compared to Cl). More energy is required to remove an electron.

atomic radii gets smaller accros a period except for argon which is slightly bigger compared to chlorine

The amount of energy released when a neutral atom gains an electron to form a negatively charged ion. in the gaseous state

X(g) + e– → X-(g)

The release of energy that occurs when a neutral, gaseous atom gains one electron. exothermic