Lecture 3 - Atomic properties: electronegativity, size, metals to non- metals Flashcards
(17 cards)
Learning objectives
3.1 Give a formal definition of electronegativity
3.2 Describe trends (and discontinuities) in electronegativity through periodic table
3.3 Discuss size changes through the periodic table
3.4 Describe and rationalise the trend from metal to non-metal across the periodic table
3.5 Explain how size can influence chemical behaviour
What is the definition of electronegativity
The power of an atom to attract electrons to itself
Eg A—:-B where A is now delta + and B is delta -
What happens to electronegativity across a row?
Zeff increases
Radius decreases
Electronegativity increases
What happens to electronegativity down a group?
Zeff is relatively constant after initial rise
Valence electrons are further away from the nucleus
EN decreases
Exception in periods 3 and 4 after d block (eg Ga>Al in energy) this is because of the presence of the d orbital and the poor shielding effect of the d-electrons in Ga, leading to lower atomic radius and higher Zeff
Same occurs for Tl and Pb due to presence of the 4f subshell
Bond energy in KJ/mol (amount of energy required to break a bond)
Cl-Cl: 240
F-F: 152
Cl-F: 253
However would expect Cl-F to be (240+152)/2 = 196 why is it not?
The “extra” energy of 57 KJ/mol is due to the difference in EN causing Cl to be delta positive and F to be delta negative. This leads to a polar bond and this charge separation causes the bond to be stronger
What is the definition of covalent radius
1/2 the length of an X-X homonuclear bond
Size of atom across a row
Zeff increases
Electrons pulled closer to the nucleus (maximum in RDF closer to nucleus)
Radius decreases
Size of atom down a group
Zeff is approx constant except for first jump
Valence electrons in higher “n” orbital
RDF shows us radius increases as maximum gets further away
General pattern is that Fr is biggest (bottom left corner) and smallest is top right
Rough EN of:
Metals
Non-metals
Metalloids
Metals: X<1.9
Non metals: X>2.2
Metalloids: 1.9<X<2.2
A more formal way of describing metallic bonding
Good diagram on slide 27
Think of X₂ to X₁₀₀ where X has 2e- and 100e-
Energy levels get closer and closer together until it bottom half of orbitals are filled and top half are empty. This bottom half is a “band” of crystal orbitals (like MO). It is a conductance band where electrons move freely from filled to unfilled orbitals
This is metallic bonding
What is a metal?
What favours metallic bonding
Metal: infinite lattice of ions with electrons in a partially filled conductance band. Metals have delocalised covalent bonding
Across a row:
EN increases, orbitals get smaller
No. Of valence electrons increase
Disfavours metallic bonding
Down a group:
Orbitals get larger
Number of valence electrons are constant
Favours metallic bonding Down
Stability of ionic compounds based on ionic size:
1. Small A+, small X-
2. Small A+, big X-
3. Big A+, small X-
4. Big A+, big X-
- and 4. are stable, 2. and 3. are not stable
Solubility of group 2 sulfates due to ionic sizes
ie MgSO4 -> BaSO4
SO4 is a anion
Size of group 2 metal increases down the group
Therefore BaSO4 is more stable and less likely to break up as big big. Solubility decreases
Decomposition of group 2 carbonates due to ionic sizes
400 C
MgCO3 ———> MgO + CO2
S B S S
Goes from SB —> SS and as SS is more favourable reaction is energetically favourable
1360 C BaCO3 ———> BaO + CO2 B B B S Goes from BB to BS which is unfavourable
Therefore reactions become more endothermic down group 2
Stability of group 1 oxides: comparing Li to Na
Li forms Li2O when burnt in air
Na forms Na2O2 when burnt in air ( 2Na+ with O₂²⁻)
Ie Na+ -O—O- +Na
Li202 -> Li2O + 1/2 O2
S B S S
Li202 -> Li2O + 1/2 O2
B B B S
Therefore delta H becomes more positive down the group
What 6 things depend on Zeff
- IE
- Electron affinity
- EN
- Size of atoms/ions
- Metal to non metal trend
- Size and chemistry
