Topic 2 - Bonding & Structure Flashcards
What is ionic bonding?
The strong electrostatic attraction between oppositely charged ions.
What effects does the ionic radius have of the strength of ionic bonding?
As ionic radius increases, charge density decreases ∴ weaker bonds are made. (enthalpy = less exo)
What effects does the ionic charge have on the strength of ionic bonding?
The bigger the charge difference the stronger the bond will be.
What are the trends in ionic radii down a group for a set of isoelectronic ions?
(+ve) ions are smaller than their atoms as they have less electrons and a greater ration of protons to remaining electrons. (-ve) ions are larger as there are more elctrons fro the same amount of protons, so the pull is less.
How does the migration of ions explain the existence of ions?
The (-ve) ions move towards the (+ve) electrode and the (+ve) ions move towards the (-ve) electrode.
How do the physical properties of ionic compounds explain the existence of ions?
- high mp (strong attractive forces between ions)
- non conductor as solid (ions held together tightly)
- conductor when molten/in aq (ions free to move)
- brittle
What is a covalent bond?
The strong electrostatic attraction between 2 nuclei & the shared pair of electrons between them.
What are diative covalent bonds?
When the shared pair of electrons in the covalent bond come from only 1 bonding ato. (co-ordinate bonding, shown by an arrow from sharing atom)
What is the relationship between bond lengths & strengths in covalent bonding?
If their is a greater force of attraction (x2/x3 bonds) the bond will be shorter but stronger.
How do lone pairs and bonding pairs arrange themselves around a central atom & determine the molecules shape?
Lone pairs repel more than bonding pairs, which reduces bond angles.
What is a linear molecule?
- 2 bonding pairs, no lone pairs
- 180° bond angle
- CO₂, BeF₂
What is a trigonal planar molecule?
- 3 bonding pairs, no lone pairs
- 120° bond angle
- SO₃, NO₃⁻, CO₃²⁻
What is a tetrahedral molecule?
- 4 bond pairs, no lone pairs
- 109.5° bond angle
- SiCl₄, SO₄²⁻, NH₄⁺
What is a trigonal pyramidal molecule?
- 3 bond pairs, 1 lone pair
- 107° bond angle
- PF₃,H₃O⁺
What is a bent/v-shaped molecule?
- 2 bond pairs, 2 lone pairs
- 104.5° bond angle
- H₂S, SCl₂
What is a trigonal bipyramidal molecule?
- 5 bond pairs, no lone pairs
- 120° & 90° bond angles
- PCl₅
What is an octahedral molecule?
- 6 bond pairs, no lone pairs
- 90° bond angle
- SF⁶
What is a square planar molecule?
- 4 bond pairs, 2 lone pairs (variation or octahedral)
- 90° bond angle w/ lone above & below
- XeF₄
What is the electron-pair repulsion theory?
VSEPR, used to predict the shapes of molecules as they are arranged to produce the minimum amount of repulsion/maximum separation.
What is electronegativity?
The ability of an atom to attract the bonding electrons in a covalent bond to itself.
What are the factors affecting electronegativity?
- ↑ across a period as protons ↑ ∴ radius ↓
- ↓ down a period as distance & shielding ↑
How does electronegativity affect bonding?
- ↓ elec.neg. difference = covalent
- ↑ elec.eg. difference = ionic (>1.7)
What is a polar covalent bond?
Not purely ionic or covalent. The least elec.neg. element = δ⁺ & the most elec.neg. = δ⁻
How do you know if a bond is polar or non-polar?
- if symmetrical (all bonds same & no lone pairs) = non-polar even if individual bonds are (cancel out)
- CO₂ = non, CH₃Cl = polar
What are London forces?
Instantaneous, induced dipole-dipole interactions, occur when fluctuations in e⁻ density causes temporary dipoles to form, which then = induced dipoles.
What affects the strength of London forces?
↑ electrons = ↑ chance of temporary dipoles = stronger L forces which need ↑ energy to overcome. Seen in ↑ bp down group 7. (I₂ = solid & Cl₂ = g)
How does the chain length of alkanes affect the size of L-forces?
↑ chain length = ↑ surface area of contact ∴ stronger L-forces. Branches in c-chain = ↓ points of contacts ∴ ↓ bps.
What are permanent dipole-dipole forces?
- occur between polar molecules, stronger than L-forces
- bond has a significant difference in elec.neg.
- can occur in addition to L forces. (δ⁻ Cl in H-Cl attracted to δ⁺ H in another H-Cl)
What are hydrogen bonds?
Bonds that occur in compounds that have a H atom attached to F, N or O, and an available lone pair.
-occurs in addition to L-forces
What is the reason for the shape of H-bonds?
They form a 180° angle around the H atom as the H gives 2 electrons which repel to min repulsion/max separation.
How are H-bonds formed in H₂O?
Water forms 2 H-bonds per molecule as δ⁻ oxygen has 2 lone pairs ∴ = stronger H-bonds & ↑ pb.
Why does ice have a lower density than water?
The H-bonds create interlocking rings of 6 water molecule when water freezes, creating distance between the molecules which decreases its density.
Why do H₂O, NH₃ & HF have abnormally high bps?
They form H-bonds as well as L-forces which needs ↑ energy to break.
Why do alcohols have lower volatility compared to alkanes?
They can form H-bonds ∴ have ↑ bps so need ↑ energy to overcome.
What are the trends in bp of the hydrogen halides from HF to HI?
- HF = ↑due to H-bonds
- ↑ from HCl to HI as L-forces ↑ due to ↑ electrons
What does ‘like dissolves like’ mean?
Compounds with similar intermolecular forces to those in the solvent will generally dissolve. (Iodine = only L-forces dissolves hexane = only L-forces.)
Why is propane a useful solvent?
It has polar & non-polar characteristics, so can dissolve polar and non-polar substances.
How is water used to dissolve some ionic compounds?
Water can hydrate the ions, the δ⁺ H bonds to the (-ve) ion in the bond & the δ⁻ O bonds to the (+ve) ion.
How is hydration affected by charge density?
↑ charge density = ↑ hydration enthalpy, as ions attract water molecules more strongly.
How does water dissolve simple alcohols?
↓ alcohols = soluble in water as can form H-bonds, as hydrocarbon chain ↑ the less soluble it becomes.
When is water not useful as a solvent?
If substances cant form H-bonds (halogenoalkanes) or are non-polar (hexane) they are insoluble in water.
What is metallic bonding?
The strong electrostatic attraction between metal ions & delocalised electrons.
What are the main factors that affect the strength metallic bonding?
- no. of protons (↑ protons = stronger bond)
- no. of delocalised electrons per atom (↑ = stronger)
- size of ion (↓ ion = stronger bond)
What are giant lattices present in?
- ionic solids (giant ionic lattices)
- covalently bonded solids (giant covalent lattices)
- solid metals (giant metallic lattices)
What is the structure of covalently bonded substances like I₂ & H₂0?
simple molecular
What are the allotropes of carbon?
- diamond, tetrahedral, 4 bonds per atom
- graphite, planar, 3 bonds per atom, layer of delocalised electrons
- graphene, 1 layer of graphite, ↑ tensile strength & good conductor of electricity
What are the physical properties of ionic structures?
- ↑ m/bp due to giant lattice w/ strong electrostatic forces between oppositely charged ions
- soluble in water
- poor conductor as solid, good when molten (ions free)
- crystalline solids
What are the physical properties of simple molecular structures?
- ↓ m/bp due to weak intermolecular forces
- generally poor solubility in water
- poor conductor as solid & molten (no ions)
- mostly gases & liquids
What are the physical properties of macromolecular structures?
- ↑ m/bp due to many strong covalent bonds
- insoluble in water
- diamond poor conductor, graphite good (delocalised)
- solids
What are the physical properties of metallic structures?
- ↑ m/bp due to strong electrostatic forces
- insoluble in water
- good conductor as solid & molten
- shiny & malleable
