Topic 3: Periodic Trends and Atomic Spectroscopy

Transcription

1 Topic 3: Periodic Trends and Atomic Spectroscopy Introduction Valence Electrons are those in the outer most shell of an element and are responsible for the bonding characteristics of that element. Core Electrons are the other electrons of an element and generally play no part in the reactivity and bonding of the element. 1. Write the electron configuration of the following elements and ions and identify which electrons are core electrons and which are valence electrons. Electron configuration Core electrons Valence electrons C 1s 2 2s 2 2p 2 1s 2 2s 2 2p 2 O O 2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 Na Ca Ca 2+ Fe 3+ Se 2 Periodic Trends First Ionisation Energy increases across a period and decreases down a group. The first ionisation energy is: M (g) M + (g) + e - The trend reflects the effective nuclear charge experienced by the electron being removed. Blackman Figure

2 2. For each of the following pairs, predict which will have the largest ionisation energy. a. H or He b. He or Ne c. K or Ca d. Br or Kr e. N or P 3. Order these elements in terms of increasing first ionisation energy: As, Cs, N, Ne, Pb 4. The first ionisation energies of Mg, Al and Si are shown below. Element Atomic number E i (kj mol - 1 ) Atom configuration Mg Al Si Cation configuration a. Fill in the electronic configurations for the atoms and the 1+ cations. b. Which ionisation energy does not fit in with the trends we have learnt? c. Provide an explanation for this anomaly. The detail can often provide valuable information about orbital energies I.E. (kj/mol) H He Li Be B C N O F Ne Na Mg Al Si P Si Cl Ar 23

3 Electron Affinity is the energy associated with X (g) + e - X - (g). It is usually exothermic but the trends are complex. What can be said is that the electron affinity generally becomes more negative across a period. Blackman Figure 4.51 Atomic Radius decreases across a period and increases down a group. It is difficult to determine the exact size of an isolated atom, so the atomic radius is defined as half the distance between two atoms of the same element. Cations are always smaller, and anions larger, than the neutral atoms from which they are formed. Radius in nanometres: Li C Li C F C F For each of the following pairs, predict which atom is larger. Blackman Figure 4.47 a. Si or Cl b. S or Se 6. Predict whether the following atoms will be smaller or larger than As. a. P b. Ge c. Se d. Sb 7. For each of the following pairs, predict which is larger. a. Li or Li + b. F or F c. Na + or F - d. Ca 2+ or S 2 8. Order these elements in terms of increasing atomic radius: Bi, Ca, F, S, Se 24

4 9. Complete the following table: Element N 3 O 2 F Ne Na + Mg 2+ Al 3+ Atomic No No of electrons Relative size Electronegativity increases across a period and decreases down a group. Electronegativity is an empirical scale that represents the ability of an atom, when in a compound, to attract the electrons of a chemical bond towards itself. It was introduced by Linus Pauling who assigned values to the elements on an arbitrary scale from 0-4. Blackman Figure Order these elements in terms of increasing electronegativity: Co, F, Ge, Rb, S 11. Explain why Cl has a higher electron affinity than Al. 12. Why are there no values of electronegativity assigned to the Noble gases? 13. Why are anions always larger than the neutral atom from which they are derived? 25

5 14. Identify the elements based on the following clues. a. Has a smaller atomic radius than hydrogen. b. Not a noble gas. Has a first ionisation energy larger than both nitrogen and oxygen. c. Has the smallest atomic radius of the metalloids. d. Has a partially filled 3p orbital and prefers to exist as the 2- anion. e. Has one electron in its valence s orbital. Has a lower first ionisation energy than rubidium. Spectroscopy and Electronic Transitions When an electron changes from one energy level to another light is either absorbed (electron moves from a low energy to higher energy orbital) or emitted. The energies of the orbitals are unique to each element (it depends on the number of electrons present and the number of protons in the nucleus). Studying these transitions is called spectroscopy. UV-Visible versus X-ray Spectrometry The visible and UV wavelength range corresponds to changes in outer electron configurations for most atoms. The energies involved are similar to or less than the ionization energy of the element. When an atom bonds to form a compound, the valence electrons are involved and this may change the energy of the electrons involved. Consequently molecules also display a unique spectrum. Emission spectroscopy records the energy change when an electron of an excited state atom falls to a lower energy orbital. This gives rise to the colours seen in fireworks or the flame tests for certain elements. Absorption spectroscopy records the absorption of energy by an atom when an electron is promoted to a higher energy orbital. Atomic absorption spectroscopy (AAS) is more sensitive than emission spectroscopy as there are a vastly greater number of atoms with electrons in the ground state (the starting point for AAS) than in the excited state (the starting point for emission spectroscopy). The colour Blackman Figures 41.5 and

6 observed is the complimentary colour to that absorbed and appear in the opposite side of the colour wheel. X- rays probe much higher energy changes in core electron configurations. These are insensitive to bonding (which mainly effects outer shell electrons), so elaborate preparations are not required. Other restrictions arise when working with x- ray and higher energies. 15. The absorbance spectrum of methyl orange in acidic and basic forms is shown below. Basic form Acidic form Absorbance Wavelength / nm 400 nm nm Violet Blue Green Yellow Orange Red a) Predict the colour of methyl orange in acidic and basic forms. b) What do you observe when a green laser pointer is pointed at the two solutions? 27