Topic 2-Bonding and Structure Flashcards
Name 3 Physical properties of ionic compounds.
High melting point
Conduct electricity when molten/ dissolved
Soluble
Name 3 properties of covalent compounds.
Low melting point
Cannot conduct electricity -they’re insulators
Solubility varies based on the molecule e.g. two water are soluble but hydrocarbons aren’t.
What is metallic bonding?
Positive metal ions electrostatically attracted to the delocalised electrons. This forms a lattice of closely packed positive ions and sea of delocalised electrons.
Why are melting points of metals generally high?
Due to the strength of metallic bonding (more energy needed to overcome)
What two things affect the melting point (metallic bonding)
The number of delocalised electrons per atom
The size of the metal ion and the lattice structure.
Why are metals malleable and ductile?
There are layers of ions separated by layers of electrons and these layers can slide over each other without disrupting the attractive between the ions and electrons.
What does ductile mean?
the ability to be drawn into wire
Why are metals good thermal conductors?
The delocalised electrons can pass kinetic energy to each other.
Why are metals good electrical conductors?
Because thee delocalised electrons are free to move and can carry a current.
What can reduce the electrical conductivity of metals?
Impurities because they can reduce the number of electrons that are free to move and carry charge- electrons transfer to impurities and form anions.
Metals are insoluble except in ……… ……. because ……..
Metals are insoluble except in liquid metals because of the strength of the metallic bonds.
Define a Covalent bond.
Strong electrostatic attraction between 2 positive nuclei & the shared electrons in the bond.
In covalent bonding two atoms ………. electrons so they’ve both got full outer shells
In covalent bonding two atoms share electrons so they’ve both got full outer shells
What is an Ionic bond?
An ionic bond is the strong electrostatic attraction between two oppositely charged ions.
What does ionic bonding involve?
The TRANSFER of electrons from one atom to another!
In general the greater the charge on an ion the ……… the ionic bond and therefore the ………. the melting/boiling point.
In general the greater the charge on ion the stronger the ionic bond and therefore the higher the melting/boiling point.
Name the 2 things that affect the strength of an ionic bond.
Ionic charges
Ionic radii
Why do smaller ions have stronger ionic bonding than larger ions?
Smaller ions can pack closer together than larger ions and electrostatic forces get weaker with distance.
What does the size on an ion depend on (2 things)
Its electronic shells and atomic number.
Ionic compounds can form …… ……. ……. structures
Ionic compounds can form Giant ionic lattice structures.
What is evidence for the presence of charged particles ?
The migration of ions shown using wet filter paper, copper II chromate solution, a microscope slide AND then pass a current through the solution (anode & cathode)
Describe the nature of a sodium-chloride BOND.
The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges.
For a molecule with 4 electron pairs and 0 lone pairs what is this shape described as and what is it’s bond angle?
Tetrahedral = 109.5
For a molecule with 4 electron pairs and 1 lone pair what is this shape called and what is it’s bond angle?
Pyramidal = 107 degrees
lone pair squashes it together more
For a molecule with 4 electron pairs and 2 lone pairs what is this shape called and what is it’s bond angle?
V-shaped= 104.5
How will a double bonded molecule with 2 total electrons (0 lone) differ from a singularly bonded molecule with 2 electrons?
There is NO DIFFERENCE, double bonds are treated the same as single bonds!
Therefore it’s linear and has a 180 bond angle.
Why does the atomic absorption spectrum of hydrogen contain discrete lines?
As only certain energy levels in atoms are allowed.
Positive ions are ……. than parent atom
Negative ions are ……… than parent atom
Positive ions are smaller than parent atoms.
Negative ions are bigger than parent atoms
For a substance to dissolve, what things must all happen?
Bonds in the substance have to break
Bonds in the solvent have to break
New bonds have to form between the substance and the solvent.
Usually a substance will only dissolve if the strength of the …… …… ….. is about the same or greater than the strength of the ……. ……….
Usually a substance will only dissolve if the strength of the new bonds formed is about the same or greater than the strength of the bonds broken.
The ionic radius of a set of isoelectronic ions decreases as the …….. ….. ……..
The ionic radius of a set of isoelectronic ions decreases ad the atomic number increases.
Which dot and cross diagrams show the charges (where electrons come from).
IONIC!
Give an example of a triple bonded molecule
N2
Give an example of a double bonded molecule
O2
Bonds are polar if the difference in electronegativity values is more than about ………
Bonds are polar if the difference in electronegativity values is more than about 0.4
What can you use the Pauling scale to do?
Work out the percentage ionic character.
Intermolecular bonds are not as ……. as our regular bonds e.g. ionic
Intermolecular bonds are not as strong as our regular bonds.
Give the order of strength of the intermolecular bonds
Induced dipole-dipole (London forces) –> Permanent dipole-dipole –> H bonds
What are London forces?
Electrons in charge clouds are likely to be more one side than the other, creating a temporary dipole.
Which molecules will have stronger London forces?
Molecules with larger electron clouds or greater surface areas (bigger exposed e- cloud)
Where are London forces found?
In all atoms + molecules
Stronger London forces mean higher …….. ……. …… …..
Melting and boiling points.
What if molecules have similar London forces but one has Permanent dipole-dipole forces as well? What happens to melting and boiling points?
The molecule with London+ Permanent will have higher melting and boiling point
(more energy needed to overcome the forces)
Name the 3 elements that Hydrogen bonding can occur with?
Oxygen, Fluorine and Nitrogen.
Examples of substances containing Hydrogen bonding are:
Ammonia
Water
Hydrogen Fluoride
Why does the boiling point of HF massively drop down the HCl, but then steadily go up at HBr and HI?
As HF is the only one containing H bonds(strongest intermolecular force)
Cl has the least amount of electrons so will have the weakest London forces AND no H bonds.
Why does the boiling point of H2O massively drop down to H2S, then increase at H2Se and H2Te?
Again, H2O is the only molecule that contains hydrogen bonding so is the highest. However H2S is the lowest as the increase in London forces is overridden by the permanent dipole interactions
Why does Butan-1-ol have a much higher boiling point than butane, a similar alkane?
Because Butan-1-ol contains hydrogen bonds in it’s OH group. H bonds are the strongest intermolecular bond and the alcohol also has London forces. Butane ONLY has London forces. Therefore the forces are much weaker and easier to break.
What is Hydration?
Ions pulled away from ionic lattice by water molecules surrounding the ions.
What is an exception to polar bonded molecules dissolving in water?
Halogenoalkanes as their dipoles aren’t strong enough to form hydrogen bonds with water.
Substances usually dissolve best in solvents with similar what?
Substances usually dissolve best in solvents with similar intermolecular forces.
How soluble a substance is in water depends on what?
The type of particles it contains e.g. water able to form h bonds so substances that can also form h bonds or are CHARGED dissolve in it well.
(non-polar or uncharged substances won’t)
What is bond enthalpy?
A measure of how strong a chemical bond is
Bond enthalpy is related to the …….. of a bond
Bond enthalpy is related to the length of a bond.
What is the distance between the two nuclei always?
covalent bonding
The distance between the two nuclei is where the attractive and repulsive forces balance each other. This is the bond length.
C=C has a greater ……. ……… and is …….. than a C-C bond. Why?
C=C has a greater bond enthalpy and is shorter than a C-C bond because the electron density is greater, so the bond is shorter.
In covalent bonding, are single, double or triple bonds usually the shortest?
Triple bonds as they have the higher bond enthalpy (greater electron density) so it is shorter due to a stronger attraction between the atoms.
Most substances with …… bonds dissolve in water
Most substances with polar bonds dissolve in water.
When ionic substances mix with water the ….. are ……. to ….. ………. ……. of the water molecule
When Ionic substances mix with water the ions are attracted to oppositely charged ends of the water molecules.
(make hydrated ions)
Why may some ionic substances not dissolve?
If bonding between ions is too strong
More electronegative elements have…
A smaller atomic radius
Higher nuclear charge. (more protons)
If polar bonds all point in the SAME rough direction then the molecule will be ………
POLAR
What does 𝛿+ or 𝛿- mean?
partial positive or partial negative charge
What’s an important general rule to remember surrounding solubility?
Like dissolves like- substances usually dissolve best in solvents with similar intermolecular forces.
What is an isoelectric ion?
Ions of different atoms with the same number of electrons e.g. Na+ and F-
What is an orbital?
A region within an atom that can hold up to 2 electrons with opposite spins
What experiment could you do to tell if something is polar?
Stream of substance (e.g. ethanol) from burette
Bring charged rod near….
+ stream will deflect.
Give 2 reasons why oxygen is more electronegative than carbon.
Smaller atomic radius
More protons
Why is N-N weaker than P-P, whereas N≡N is stronger than P≡P? [4]
In N-N lone pairs closer to each other than in the P-P (as P has extra shell)
so repulsion = significant which weakens the bond
N≡N and P≡P both have lone pairs at end of molecule so repulsion is insignificant
N≡N pi bonds stronger as sideways overlap between 2p orbitals compared to 3p orbitals of phosphorus
Describe how London forces form between halogen molecules. [3]
setting up of dipole from
- Uneven distribution of electrons
- Resulting in an instantaneous dipole on one molecule
- which induces a dipole of the other molecule
