Note: You might argue that the fall doesn't apply throughout the Group because both potassium and rubidium have an electronegativity of 0.8. The atoms become less and less good at attracting bonding pairs of electrons. (Remember that the most electronegative element, fluorine, has an electronegativity of 4.0.) Notice that electronegativity falls as you go down the Group. If you choose to follow this link, use the BACK button on your browser to return quickly to this page.Īll of these elements have a very low electronegativity. Note: You will find electronegativity covered in detail in another part of this site. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0. However, as you go down the Group, the distance between the nucleus and the outer electrons increases and so they become easier to remove - the ionisation energy falls.Įlectronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Just as when we were talking about atomic radius further up this page, in each of the elements in this Group, the outer electrons feel a net attraction of 1+ from the centre. The distance between the outer electrons and the nucleus.Īs you go down the Group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons. The amount of screening by the inner electrons, Notice that first ionisation energy falls as you go down the group.Įxplaining the decrease in first ionisation energy
![does atomic radius increase down a group does atomic radius increase down a group](http://es1.ph.man.ac.uk/AJM2/e2e_hardware_files/NIM1.jpg)
If you choose to follow this link, use the BACK button on your browser to return quickly to this page.
![does atomic radius increase down a group does atomic radius increase down a group](http://education-portal.com/cimages/multimages/16/period-electronegativities.jpg)
Note: You will find ionisation energy covered in detail in another part of this site. If you don't get into the habit of thinking about all the possible factors, you are going to make mistakes.įirst ionisation energy is the energy needed to remove the most loosely held electron from each of one mole of gaseous atoms to make one mole of singly charged gaseous ions - in other words, for 1 mole of this process: That isn't true if you try to compare atoms from different parts of the Periodic Table. If you are talking about atoms in the same Group, the net pull from the centre will always be the same - and you could ignore it without creating problems. It is a matter of setting up good habits. Why, then, bother about exploring the net pull on the electrons from the centre of the atom? Note: You may think that this is all a bit long-winded! It is, after all, fairly obvious that atoms will get bigger if you add more layers of electrons. That means that the atoms are bound to get bigger as you go down the Group. Obviously, the more layers of electrons you have, the more space they will take up - electrons repel each other. The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom.
![does atomic radius increase down a group does atomic radius increase down a group](https://s3-us-west-2.amazonaws.com/courses-images/wp-content/uploads/sites/752/2016/09/26195606/halogens.jpeg)
Work it out for potassium if you aren't convinced. This is equally true for all the other atoms in Group 1. The positive charge on the nucleus is cut down by the negativeness of the inner electrons.
![does atomic radius increase down a group does atomic radius increase down a group](https://venturebeat.com/wp-content/uploads/2020/03/microkingdoms2.jpg)
In each case, the outer electron feels a net pull of 1+ from the nucleus. Use the BACK button on your browser to return quickly to this page. Note: If you aren't sure about writing electronic structures using s and p notation it might be a good idea to follow this link before you go on. The pull the outer electrons feel from the nucleus. The number of layers of electrons around the nucleus You can see that the atomic radius increases as you go down the Group. Note: You will find atomic radius covered in detail in another part of this site. The same ideas tend to recur throughout the atomic properties, and you may find that earlier explanations help to you understand later ones. You will find separate sections below covering the trends in atomic radius, first ionisation energy, electronegativity, melting and boiling points, and density.Įven if you aren't currently interested in all these things, it would probably pay you to read the whole page. This page explores the trends in some atomic and physical properties of the Group 1 elements - lithium, sodium, potassium, rubidium and caesium. Atomic and physical properties of Periodic Table Group 1ĪTOMIC AND PHYSICAL PROPERTIES OF THE GROUP 1 ELEMENTS