Periodic Trends and Bonding

Higher Supported Study - Week 2

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.

Covalent RadiusThe covalent radius is half the distance

between the two nuclei in a covalent bond.

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 groupCovalent Radius increases due to increasing number of electron shells

Ionisation Energies

Definition Energy required to remove one mole of electrons from one mole of gaseous atoms

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 groupIonisation 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.

Ionisation Energy Trends Always write out the electron arrangements of

atoms/Ions. Compare nuclear charge and no. of electron

shells. Across a period - INCREASESIncreasing nuclear charge Down a group – DECREASES additional energy levels and screening

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.

element first IE second IE third IE fourth IE

A 520 7300 11500

B 2100 3900 6100 9400

C 580 1800 2800 11500

D 740 1450 7700 10500

E 420 3050 4500 5900

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.

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-

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.

Electronegativity Trends • Across a period Electronegativity increases due to INCREASING NUCLEAR CHARGE causes atoms to attract bonding electrons more strongly

• Down a groupElectronegativity 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

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 - INCREASESIncreasing nuclear charge Down a group – DECREASES additional energy levels and screening

Key Areas – Bonding

• Covalent bonding. • Polar covalent bonding. • The bonding continuum.• VDWFs• Properties linked to Bonding

Recap From Nat 5

• Bonding in Elements – Metallic – Covalent molecular – Covalent network – Monatomic (group 8)

• Bonding in Compounds– Ionic – Covalent

Covalent Bonding

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

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.

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

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

10. Atoms of nitrogen and element X form a bond in which the electrons are shared equally.

Element X could be

A carbonB oxygenC chlorineD phosphorus.

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11. Which line in the table represents the solid inwhich only van der Waals’ forces are overcomewhen the substance melts?

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14. In which of the following solvents is lithiumchloride most likely to dissolve?

A HexaneB BenzeneC MethanolD Tetrachloromethane

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