Why does li have a larger

The van der Waals radius also known as the nonbonding atomic radius is the radius of an atom which is not bonded to other atoms; this is determined by measuring the distance between atomic nuclei which are in direct but nonbonding contact with each other in a crystal lattice.

The covalent atomic radius also known as the bonding atomic radius is determined for metals by taking one-half of the distance between two adjacent atoms in a metallic crystal, or one-half the distance between like bonded atoms for nonmetals. Unfortunately, it is not possible to determine the radius for every element on the periodic table in the same way, and consequently, it is sometimes difficult to make comparisons between different sets of data.

In the table above, most of the atomic radii listed are average atomic radii, while for the halogens Group 7A and the noble gases Group 8A the covalent radius is used. Atomic radii vary in a predictable way across the periodic table. As can be seen in the figures below, the atomic radius increases from top to bottom in a group , and decreases from left to right across a period. Thus, helium is the smallest element, and francium is the largest.

The sizes of cations and anions follow similar trends to those of neutral atoms. In general, anions are larger than the corresponding neutral atom, since adding electrons increases the number of electron-electron repulsion interactions that take place. Cations are smaller than the corresponding neutral atoms, since the valence electrons, which are furthest away from the nucleus, are lost.

Taking more electrons away from the cation further reduces the radius of the ion. The table below illustrates these trends for the main group elements. For elements which form more than one cation, the cation charges and sizes are listed in two separate columns. The transition metals and inner transition metals have been omitted, since almost all of those elements can form two or more possible cations.

Number Name Neutral. Charge Cation1. Charge Cation2. Charge Anion. Uub n. He Li Be Ne Na Mg Al Si P S Cl Ar K Ca Since outermost electrons are located on the same energy level in both the lithium and the fluorine atoms, it follows that the nucleus that has more protons will attract these electrons more - this is known as effective nuclear charge.

So even if the outermost electron of lithium and the outermost electrons of fluorine reside on the same energy level , the latter will be more attracted to the nucleus, since fluorine has 9 protons there. As a result, the atomic size of fluorine will be smaller than that of lithium, or, in other words, lithium will have larger atomic radius than fluorine. Does Li or F have the larger atomic radius?

Stefan V. Oct 31, Explanation: Start by taking a look at a periodic table and making a note of where lithium, "Li" , and fluorine, "F" , are located. Here is where the number of protons that each atom has in its nucleus becomes important. May 8, Here's what's going on here. Explanation: Lithium , "Li" , and beryllium , "Be" , are both located in period 2 of the periodic table , in group 1 and group 2, respectively.

Related questions Why do periodic trends exist for electronegativity? Why does atomic size increase down a group? What do periodic trends of reactivity occur with the halogens? How can I determine atomic size of ions? The ionization energy of an atom is the amount of energy that is required to remove an electron from a mole of atoms in the gas phase:.

Ionization energies are always positive numbers, because energy must be supplied an endothermic energy change to separate electrons from atoms. The second ionization energy is always larger than the first ionization energy, because it requires even more energy to remove an electron from a cation than it is from a neutral atom.

The first ionization energy varies in a predictable way across the periodic table. The ionization energy decreases from top to bottom in groups , and increases from left to right across a period.

Thus, helium has the largest first ionization energy, while francium has one of the lowest. There are some "fluctuations" in these general trends. For instance, the first ionization decreases from beryllium to boron, and from magnesium to aluminum, as electrons from the p-block start to come into play.

In the case of boron, which has an electron configuration of 1s 2 2s 2 2p 1 , the 2s electrons shield the higher-energy 2p electron from the nucleus, making it slightly easier to remove. A similar effect occurs in aluminum, which has an electron configuration of 1s 2 2s 2 2p 6 3s 2 3p 1. Even though oxygen is to the right of nitrogen in period 2, its first ionization energy is slightly lower than that of nitrogen.

Nitrogen has an electron configuration of 1s 2 2s 2 2p 3 , which puts one electron in each p orbital, making it a half-filled set of orbitals:. Half-filled sets of p orbitals are slightly more stable than those with 2 or 4 electrons, which makes it slightly harder to ionize a nitrogen atom.

Oxygen has an electron configuration of 1s 2 2s 2 2p 4 , which puts another electron in one p orbital; since this is one electron away from being half-filled, it is slightly easier to remove this additional electron:. Uub n. H He Li Be B C N O F Ne Na Mg Al Si P S