Week 6/7 Core and valence electrons

Question 1

What is the outermost shell and electrons called

Answer 1

The outermost shell of electrons in any atom is called the valence shell and these electrons are called valence electrons. All electrons in an atom, which are not valence electrons are called core electrons.

Question 2

Difference between core and valence electrons

Answer 2

Core electrons

Core electrons remain bound to the same nucleus through all chemical reactions and
core electrons provide shielding between protons in the nucleus and electrons in outer shells.
Valence electrons

Valence electrons in the outermost shell of at atom, are the only electrons which can be lost when an atom becomes a cation and
the valence shell of an atom is the only shell which can accept electrons allowing an atom to form an anion.

Question 3

Actual versus effective nuclear charge

Answer 3

The actual nuclear charge of an atom is defined as the number of protons in the nucleus of that atom, Z.

A valence electron in a multi-electron element will however be shielded by any electrons in shells closer to the nucleus, meaning the effective nuclear charge (Zeff) felt by a valence electron is less than the actual nuclear charge of nucleus:

Zeff = Z - S

Where Z is the actual nuclear charge and S is the shielding (or screening) constant.

Question 4

Effective nuclear charge

Answer 4

Effective Nuclear Charge

In a sodium atom, the effective nuclear charge felt by the valence electron is diminished by the 10 core electrons in the atom.

Using the equation above:

Zeff = 11 electrons in Na - 10 core electrons = +1

Question 5

Atomic radius definition

Answer 5

Atomic radius is generally stated as being the total distance from an atom’s nucleus to the outermost orbital of electrons. As you begin to move across or down the periodic table, trends emerge that help explain how atomic radii change.

Unlike a ball, or other fixed sphere, an atom doesn’t have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.

Question 6

Ionic Radii

Answer 6

The size of an ion, from the nucleus to the outer electrons, is called the ionic radius. There is a size change when a neutral atom is converted to an ion and the change in size is different for anions and cations as we gain or lose electron(s).

Cations

The ionic radius of a cation is smaller than the atomic radius of the corresponding ground state atom. On forming a cation, the nuclear charge stays the same but the number of electrons decreases. The remaining electrons are therefore attracted more strongly to the nucleus.

Anions

The ionic radius of an anion is larger than the atomic radius of the corresponding ground state atom. On forming an anion, the nuclear charge stays the same but the number of electrons increases. This means the electrons are less strongly held by the nucleus and the orbitals enlarge.

Question 7

ionisation energy definition

Answer 7

‘the energy required to remove an electron from a gaseous atom in the ground state’.

Question 8

ionisation energy equation example

Answer 8

Below is the equation for removing an electron from sodium. The amount of energy needed to achieve this is 495 kJ mol-1. This energy is called the 1st ionisation energy i.e. the amount of energy required to remove the 1st electron.

Na (g) + energy → Na+ (g) + e-

To remove another electron from sodium requires much more energy. This is because the electron configuration of Na+ is the same as Neon i.e. a full valence electron shell. Removing an electron from a full shell requires a much higher amount of energy; for sodium the 2nd ionisation energy is 4560 kJ mol-1.

Na+ (g) + energy → Na2+ (g) + e-

Question 9

Electron affinity

Answer 9

Electron affinity is the energy change that occurs when a gaseous atom accepts an electron to form a negatively charged ion. In the example below, adding an electron to a chlorine atom, to form a chloride ion, results in the release of energy; it is an exothermic process.

Cl (g) + e- → Cl- (g) -349 kJ/mol (exothermic)

In general, electron affinities become more exothermic (give out more heat) on moving towards the right-hand side of the periodic table i.e. electrons bind more strongly. There is little change in electron affinities down a group.

Question 10

Characteristics of Metals, Metalloids and Non-Metals

Answer 10

Metals

Shiny lustre, conduct electricity, malleable
All are solids (except mercury)
Low ionisation energies so they form cations
Metalloids

Have properties between those of metals and non-metals
Silicon looks like a metal but is brittle and a poorer conductor of heat and electricity
Compounds of metalloids can have characteristics of the compounds of metals or non-metals depending on the compound.
Non-metals

Vary greatly in appearance
Poor conductors of heat and electricity
Melting points are generally lower than the metals

Question 11

3 types of chemical bond that hold matter together

Answer 11

Whilst matter exists in a number of forms, there are only three types of chemical bonds that hold matter together; ionic, covalent and metallic bonds.

Question 12

Ionic bonding

Answer 12

Copper

The expected ground state electron configuration of copper is [Ar] 4s2 3d9 however, as there is stability in half-filled or full sub-shells it is usually written as:

[Ar] 4s1 3d10

This movement of electrons from one value of n to another is possible as the energy difference between n=3 and n=4 is small. In this configuration the 3d orbital is full and the 4s orbital is half-full.

To form a cation, the one 4s electron is lost resulting in a 1+ cation, Cu+; with the electron configuration [Ar] 3d10. Experimentally however, we know that copper can also form a +2 cation (Cu2+). In this case the electron configuration of the Cu2+ cation would be:

[Ar] 3d9

This results from the loss of the two 4s electrons we had in the initial electron configuration we considered above.

Question 13

Electronegativity

Answer 13

Electronegativity is the ability of an atom to attract electrons. This determines that metals will release an electrons and non-metals will accept one; ionic bonding. Linus Pauling devised a scale of electronegativity for the elements; Fluorine is the most electronegative and Cesium is the least.

The differences in electronegativity between every pair of elements determines the nature of the bond which can form between the atoms of those elements; ionic, polar covalent or covalent.

Question 14

Dipole moments

Answer 14

The electronegativity difference between hydrogen and fluorine in HF means more of the negative charge in the bond is concentrated towards the fluorine atom rather than the hydrogen atom in the polar covalent bond.

Consequently we can describe hydrogen fluoride as a polar molecule. We use the delta (δ) symbol to denote a partial charge; δ+ and δ-. In the image above the hydrogen has a delta positive partial charge and the fluorine and delta negative partial charge.

Polar molecules are attracted to one another and to ions in solution; the positive end is attracted to a negative ion and the negative end to a positive ion which accounts for many of the properties of materials in solution.

When two electrical charges, of the same magnitude but opposite signs, are separated by a distance, a dipole is formed.