Presentation on theme: "Physical Properties. Syllabus statements 3.2.1 Define the terms first ionization energy, and electronegativity. 3.2.2 Describe and explain the trends."— Presentation transcript:
Syllabus statements 3.2.1 Define the terms first ionization energy, and electronegativity. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities and melting points for the alkali metals, and the halogens. 3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities for the elements across period 3. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions on the periodic table.
The first ionisation energy of an atom tells us how difficult it is to remove an electron from that atom. We only talk about first ionisation energy when we make a positive ion – this will be important to remember when we study energetics! A definition (learn it!) The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. Every part of the definition is important!
First ionization energies are measured in kJ mol -1 You don’t need to remember the numbers – they are given in the data booklet. You DO need to remember the trends.
Trends in Properties We will look at the chemical properties later; first we need to think about some of the physical properties:
Atomic radius The atomic radius is the distance from the nucleus of an atom to its outermost electron.
BUT Electrons don’t stand still (and that’s without even mentioning the Heisenberg Uncertainty Principle!) So atomic radius is sometimes defined as half the distance between neighbouring nuclei
For group 1 elements: ElementElectron arrangement Atomic radius (10 -12 m) Lithium Li2,1152 Sodium Na2,8,1186 Potassium K2,8,8,1231 Rubidium Rb2,8,8,...,1244 Cesium Cs2,8,8...,..., 1262
The numbers are in the data booklet – you don’t need to learn them. The atomic radius increases down a group as the number of occupied electron shells increases. The data booklet doesn’t give atomic radii for the noble gases. Why?
All the group 1 elements lose a single electron when they are ionized. As we would expect, positive ions are smaller than the atoms from which they were formed. Positive ions have fewer occupied electron shells, and greater electrostatic attraction than the atoms from which they were formed.
There is a similar trend in atomic radii going down group VII. Fluorine has more electrons than Li. So why does it have a smaller atomic radius? We will answer that question in a minute! ElementAtomic Radius 10 -12 m F64 Cl99 Br114 I133
All the group VII elements gain one electron when they form ions. The ions have more electrons than the atoms from which they were formed. Hence negative ions are larger than the atoms from which they were formed.
Cations contain _______ electrons than protons, so they are ______ than their parent atoms; Anions contain _______ electrons than protons, so they are ______ than their parent atoms.
Cations contain _fewer_ electrons than protons, so they are smaller than their parent atoms; Anions contain __more_ electrons than protons, so they are bigger than their parent atoms.
Now consider period 3 This is the row which contains: Na Mg Al Si P S Cl Ar How does atomic radius change between Na and Cl? It gets smaller, even though Cl has more electrons. Why?
All the elements in period 3 contain electrons in the 3 rd shell. I.e. they all have the same outer shell. But, as we go across period 3 each element has one more proton than the previous element. This extra positive charge pulls the outer shell of electrons slightly. As we go across a period, atomic radius decreases because of increased electrostatic attraction between the nucleus and the outermost electrons.
The situation is a little harder for ionic radii. We need to consider positive ions (cations) and negative ions (anions) separately. All the ions across a period contain the same number of electrons They are “isoelectronic” with a noble gas
For cations: As we go across a period, we add protons. The increased nuclear charge pulls the outer electrons in closer. Ionic size decreases across a period.
When we start forming negative ions (anions) There is suddenly a large increase in size. This is because electrons have been added to the outer shell, resulting in increased repulsion with the nucleus. BUT as we continue across the group we add further protons (ie the number of protons gets closer to the number of electrons) Hence the anions get smaller as we add more protons.