Periodic Trends and Bonding Higher Supported Study - Week 2 Periodic Trends and Bonding

Key Areas – Periodic Trends The trends in covalent radius across periods and down groups. The trends in ionisation energies across periods and down groups. The trends in electronegativity across periods and down groups.

Why is there no covalent radii information for noble gases? Covalent Radius The covalent radius is half the distance between the two nuclei in a covalent bond. There is no information for the noble gases because they do not form covalent bonds. Why is there no covalent radii information for noble gases?

Covalent Radius Trends Across a period Covalent radius decreases due to INCREASING NUCLEAR CHARGE pulls electrons in closer to the nucleus Down a group Covalent Radius increases due to increasing number of electron shells Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening to the size of the atom and why. Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

Ionisation Energies Definition Energy required to remove one mole of electrons from one mole of gaseous atoms This is a definition which you should be able to recall. The equation for first ionisation energies and second ionisation energies are given on page 11 of data book. The elements on this page are arranged in periods.

Ionisation Energy Trends Across a period Ionisation Energy increases due to INCREASING NUCLEAR CHARGE pulls electrons in closer to the nucleus, meaning more energy is required to remove these electrons Down a group Ionisation Energy decreases due to the electron being further away from the nucleus due to additional energy levels and the inner electrons screening the outer electrons from full effect of nuclear charge. Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why. Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

Ionisation Energy Trends Always write out the electron arrangements of atoms/Ions. Compare nuclear charge and no. of electron shells. Across a period - INCREASES Increasing nuclear charge Down a group – DECREASES additional energy levels and screening Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why. Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

Ai) K(g)  K+(g) + e- (this should be easiest mark to get in exam as general equation given in data book, STATE SYMBOLS must be included. Aii) Write out Nuclear charge and electron arrangements for the two things which are being compared. K (19) 2, 8, 8, 1 Cl (17) 2, 8, 7 K outer electron is further away from nucleus, due to greater number of electron shells, also inner electrons shield the outer electrons from full effect of positive nuclear charge.

B C

Other Styles of Questions Explain why potassium has a lower first ionisation energy than lithium Explain why the second ionisation energy of sodium is much higher than the first ionisation energy of sodium. Explain why the P3- ion is larger than the Al3+ ion Explain why the Ca2+ ion is smaller than the P3- ion. K – 2, 8, 8, 1 Li- 2, 1 Potassium has more electron shells, the electron to be removed is further away from the nucleus and the inner electrons shield the outer electrons from the full effect of the nuclear charge. (Notice I link response to definition of first Ionisation energy) Na – 2,8,1 The second electron is being removed from a shell closer to the nucleus. This electron has a much greater nuclear charge to overcome, as is closer to nucleus. P3- – 2, 8, 8. Al3+ - 2, 8. Al only has 2 electron shells, while P has 3 making the P ion larger. Ca2+ - 2, 8, 8 P3- – 2, 8, 8 Have same number of electron shells, so change in size must be due to difference in nuclear charge. Ca has a greater nuclear charge, drawing electrons in closer making the ion smaller.

for five elements A, B. C, D and E, in kJ mol-1. The following table shows the approximate first ionisation energies (IE) for five elements A, B. C, D and E, in kJ mol-1. Which of these elements; (a) could be in group 2 of the periodic table? (b) could be in the same group of the periodic table? (c) would require the least amount of energy to convert one mole of gaseous atoms into ions carrying a three positive charge. Explain your answers. 3. Write equations corresponding to; (a) The second ionisation energy of magnesium. (b) The third ionisation energy of aluminium. element first IE second IE third IE fourth IE A 520 7300 11500   B 2100 3900 6100 9400 C 580 1800 2800 D 740 1450 7700 10500 E 420 3050 4500 5900 A – Group 1 B – Group 3 C- Group 3 D – Group 2 E – Group 1 (c) X(g)  X3+(g) Need to add together (1st IE + 2nd IE + 3rd IE) energy required to remove 3 electrons. 3a) Mg+(g)  Mg2+ + e- b) Al2+ (g)  Al3+ + e-

I.E. Questions State the energy required for the following changes. a) Al(g)  Al+(g) + e- b) Al(g)  Al2+(g) + 2e- c) Al+(g)  Al3+(g) + 2e- First IE from data book First IE and Second IE added together, as 2 electrons being removed. One electron already lost before started, so Second IE and third IE added together.

Electronegativity Atoms of different elements have different attractions for bonding electrons. Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond. Another key definition to learn – electronegativity

Electronegativity Trends Across a period Electronegativity increases due to INCREASING NUCLEAR CHARGE causes atoms to attract bonding electrons more strongly Down a group Electronegativity decreases due to increasing number of electron shells, the inner electrons shield the outer electrons from the nuclear charge, meaning the electrons are less strongly attracted to the atom Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why. Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

Ionisation Energy and ELECTRONEGATIVITY Trends Always write out the electron arrangements of atoms/Ions. Compare nuclear charge and no. of electron shells. Across a period - INCREASES Increasing nuclear charge Down a group – DECREASES additional energy levels and screening Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why. Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

Key Areas – Bonding Covalent bonding. Polar covalent bonding. The bonding continuum. VDWFs Properties linked to Bonding

Recap From Nat 5 Bonding in Elements Bonding in Compounds Metallic Covalent molecular Covalent network Monatomic (group 8) Bonding in Compounds Ionic Covalent

Covalent Bonding This is listed in Higher Arrangement documents, so could be asked at Nat 5.

3 types of Bonding Covered at Nat 5.

Intramolecular Covalent molecules Non polar – no difference in electronegativity - electrons shared equally Polar – DIFFERNCE in electronegativity - electrons NOT shared equally LARGER THE DIFFERENCE IN ELECTRONEGATIVTY THE MORE POLAR THE BOND WILL BE The key here is electronegativities.

Bonding Continuum Depending on difference in electronegativities, bonds are classed as Pure covalent, polar covalent or ionic. This scale is called the Bonding continuum.

Bonding Continuum If the electronegativity difference is large then the movement of bonding electrons from one atom to another results in formation of ions. Compounds formed between metals and non-metals are often, but not always ionic. If the electronegativity difference is less than 0.4 we class the bonds as pure covalent. If the electronegativity dofference is between 0.4 and 1.7 it is classed as polar covalent If the electronegativity difference is greater than 1.7 it is ionic. The greater the ionic character the greater the difference in electronegativity (common exam question) Most covalent character the smaller the difference in electronegativity.

Intermolecular Bonding All molecular elements and compounds and monatomic elements condense and freeze at sufficiently low temperatures. For this to occur, some attractive forces must exist between the molecules or discrete atoms. Van Der Waals Forces London Dispersion Forces Permanent dipole - permanent dipole Hydrogen Bonding HINT – Always useful to include diagrams when trying to explain these types of interactions with ∂+ and ∂- signs when appropriate First two sentences lifted straight from Higher arrangement documents. Remember these forces only found between molecules and monatomic. Ionic/metallic/network are a continuous structure of atoms/ions. There is no individual atoms/molecules so forces of attraction can exist “between”.

Properties of Molecules Viscosity Melting/Boiling Point Can all be explained by comparing the types of VDWF’s present. If comparing two molecules state the Type of VDWF’s in each molecules The strength of the VDWF A detailed explanation of how the VDWF arises.

ICE H- Bonding can be used to explain why ice is less dense than water. Ice contains an OPEN-LATTICE ARRNAGEMENET OF WATER MOELCULES

Solubility Ionic compounds and polar molecules tend to be soluble in polar solvents such as water and are insoluble in nonpolar solvents. Non-polar molecules tend to be soluble in non-polar solvents and insoluble in polar solvents. “Like dissolves like”

Diagrams can help explain solubility

2012 10. Atoms of nitrogen and element X form a bond in which the electrons are shared equally. Element X could be A carbon B oxygen C chlorine D phosphorus. Answer C

11. Which line in the table represents the solid in 2012 11. Which line in the table represents the solid in which only van der Waals’ forces are overcome when the substance melts? Answer - D

14. In which of the following solvents is lithium 2012 14. In which of the following solvents is lithium chloride most likely to dissolve? A Hexane B Benzene C Methanol D Tetrachloromethane Answer – C Note the use of the language “most likely”

Answer - C