(Green) Bonding I ionic Flashcards
(Green) Bonding I ionic
what is the definition of ionic bond
DEFINITION: An ionic bond is an electrostatic attraction between oppositely charged positive cations and negative anions. It is usually formed by electron transfer from a metal atom to a non metal.
Which group of elements will have the greatest tendency to lose their outer shell electrons?
Group I
Which is the second most likely group to lose outer shell electrons?
Explain why
Group II
Group I have a large atomic radius, outer electrons are readily lost, DHi(1) is low.
Group II have a fairly large atomic radius, outer electrons are readily lost, DHi(1+2) is fairly low.
One element in this second group does not in fact readily give up electrons. Identify this element and explain why it behaves differently.
Be – very small ionic radius, a Be2+ ion has a very high charge density and polarises the larger anion (pulls electron cloud towards itself and away from the anion) —-> a covalent bond.
Which group of elements will have the greatest tendency to gain outer shell electrons?
Group VII
Which is the second most likely group to gain outer shell electrons?
Explain why
Group VI
Fairly small atomic radius, if they gained one or two electrons they would be ISOELECTRONIC with a noble gas and stable.
info card (read and recite )
These oppositely charged ions are attracted to one another by strong electrostatic forces called ionic bonds. They form a giant ionic lattice where each ion becomes surrounded by a number of ions of the opposite charge.
What 2 factors does the co-ordination numbers in a lattice depend on?
- charge on the ions involved
- radius of the ions involved (actually the sum of the ionic radii)
what is the co-ordination number
it is the number of atoms that surround atom 1 : atom 2
Compare the sizes of atoms and the cations that they form. State and explain the difference.
Cations are smaller. Loss of electron(s) may result in loss of outer shell
Cation now has more positive protons than electrons so electrons will be attracted more.
Likewise compare the sizes of atoms and respective anions. State and explain this difference.
Anions are larger. Gain of electron(s) so e- > p+. electrons repel each other, held less strongly.
Why do the cations and anions increase in radius as a group is descended?
· More protons
· But more shielding by complete inner shells of electrons so less attraction of e-to p+.
· Another complete inner shell with successive periods
what does Isoelectronic ions mean
Isoelectronic ions are those having same electron configuration
what colour is Cu 2+
blue
what colour CrO4 2-
is yellow
what are the 4 physical propertys of ionic compounds
High melting temperatures
Brittleness
Poor electrical conductivity when solid but good when molten or in solution
Often soluble in water
Explain why ionic solids:-
Have high melting points
When heated ions vibrate more in fixed positions until enough E allows them to separate.
A lot of heat energy is needed to overcome the strong electrostatic forces of attraction between oppositely charged ions to melt the solid.
Explain why ionic solids:-
Have a hard and brittle structure
Hard – strong electrostatic forces of attraction in 3D.
Brittle - When hit with enough force ions are knocked out of alignment so ending next to similarly charged ions which then repel each other.
Explain why ionic solids:-
They usually dissolve in water:
Water molecules are polar, the d- side clusters around cations and the d+ side around anions so hydrating the ions releasing energy; this is enough energy to overcome the electrostatic attractions so separating the ions from the lattice.
Explain why ionic solids:-
Conduct electricity when molten, but not in the solid state
When molten or dissolved in solution ions are free to move, +ve ions to the cathode, -ve ions to the anode, so conducting electricity.
When solid the ions are held in fixed positions in the lattice …
what is the definition of an ionic radious
When you write an ionic equation you include only the ions which actually take part in the reaction to form precipitates, elements, (e.g. metals) gas molecules (e.g H2, CO2) and liquid molecules (e.g. H2O)
How do you know which ions participate in the bond forming process and which don’t?
what are the four things that you have to look for
· Nitrate(V) ions are never involved – all nitrates are soluble in water.
· Group I ions are never involved – all group I salts are soluble in water.
· Look for a precipitate forming– this will be formed from oppositely charged ions reacting together to form an uncharged insoluble solid.
· Coloured precipitates usually either have a transition metal cation or the anion contains a transition metal e.g. the dichromate(VI) ion Cr2O72-
what is the full equation and ionic equation
and the observations of adding :
Add barium chloride solution drop by drop
to 0.5 cm3 of sodium sulfate solution
Precipitation reactions
Immediate thick White Ppt formed
Full Equation BaCl2(aq) + Na2SO4(aq) —> BaSO4(s) + 2NaCl(aq)
Ionic equation: Ba2+(aq) + SO42-(aq) —-> BaSO4(s)
what is the full equation and ionic equation
and the observations of adding :
Add silver nitrate(V) solution
drop by drop to 0.5 cm3 of sodium
chloride solution
Precipitation reactions
Fine White ppt formed turned grey in light
Full Equation NaCl(aq) + AgNO3(aq) —-> NaNO3(aq) + AgCl(s)
Ionic equation: Cl-(aq) + Ag+(aq) —-> AgCl(s)
what is the full equation and ionic equation
and the observations of adding :
Add silver nitrate solution
drop by drop to 0.5 cm3 of potassium
bromide solution.
Precipitation reactions
Fine Cream ppt formed
Full Equation KBr(aq) + AgNO3(aq) —->NaNO3(aq) + AgBr(s)
Ionic equation: Br-(aq) + Ag+(aq) —-> AgBr(s)
what is the full equation and ionic equation
and the observations of adding :
Add silver nitrate solution
drop by drop to 0.5 cm3 of potassium
iodide solution
Precipitation reactions
Yellow ppt formed
Full Equation KI(aq) + AgNO3(aq) —>NaNO3(aq) + AgI(s)
Ionic equation: I-(aq) + Ag+(aq) —-> AgI(s)
what is the full equation and ionic equation
and the observations of adding :
Add potassium iodide solution drop by
drop to 0.5 cm3 of lead(II)
nitrate(V) solution.
Precipitation reactions
Bright Yellow ppt formed
Full Equation 2KI(aq) + PbNO3(aq) —–> 2KNO3(aq) + PbI2(s)
Ionic equation: 2I-(aq) + Pb2+(aq) —-> PbI2(s)
what is the full equation and ionic equation
and the observations of adding :
Add sodium hydroxide
solution drop by drop
to 0.5 cm3 of copper(II)
sulfate(V1) solution
Precipitation reactions
Blue jelly-like ppt formed
Full Equation CuSO4(aq) + 2NaOH(aq) —-> Cu(OH)2(s)+Na2SO4(aq)
Ionic equation: Cu2+(aq) + 2OH-(aq) —-> Cu(OH)2(s)
what is the full equation and ionic equation
and the observations of adding :
Add a small spatula of Zinc to 2 cm3 of
copper II sulfate solution
DISPLACEMENT REACTIONS (REDOX)
Blue copper sulfate solution colour fades, solution eventually turns colourless, a pink/ black solid formed
Full equation Zn(s) + CuSO4(aq) —-> ZnSO4(aq) + Cu(s)
Ionic equation: Zn(s) + Cu2+ (aq) —-> Zn2+(aq) + Cu(s)
what is the full equation and ionic equation
and the observations of adding :
Add a small piece of Mg ribbon to 2cm3 of
a zinc II sulfate solution
DISPLACEMENT REACTIONS (REDOX)
Solution gets warm with effervescence magnesium ribbon gets smaller
Full equation Mg(s) + ZnSO4(aq) —> MgSO4(aq) + Zn(s)
Ionic equation: Mg(s) + Zn2+(aq) —-> Mg2+ (aq) + Zn(s)
what is the full equation and ionic equation
and the observations of adding :
Add a small piece of magnesium ribbon to
2cm3 of copper II sulfate solution
DISPLACEMENT REACTIONS (REDOX)
Solution gets warm with effervescence, blue copper sulfate solution fades, eventually turns colourless and a pink/ black solid is formed
Full equation Mg(s) + CuSO4(aq) —-> MgSO4(aq) + Cu(s)
Ionic equation: Mg(s) + Cu2+(aq) —> Mg2+ (aq) + Cu(s)
what is the full equation and ionic equation
and the observations of adding :
Put a small spatula of sodium carbonate
into a boiling tube, carefully add sulphuric
acid until no further changes occur
Effervescence, white solid ‘disappears’ to leave a colourless solution
Full Equation Na2CO3(s) + H2SO4(aq) —-> CO2(g) + H2O(l) + Na2SO4(aq)
Ionic equation (2 options) Na2CO3(s) + 2H+ (aq) —-> CO2(g) + H2O(l) + 2Na+(aq)
CO32- (s) + 2H+(aq) —-> CO2(g) + H2O(l
what is the full equation and ionic equation
and the observations of adding :
To 2 cm3 of Sodium hydrogencarbonate
add drop by drop sulphuric acid until no
further changes occur
REACTIONS OF ACIDS
Effervescence, white solid ‘disappears’ to leave a colourless solution
Full Equation NaHCO3(aq) + H2SO4(aq) —-> Na2SO4(aq) + H2O(l) + CO2(g)
Ionic equation HCO3-(aq) + 2H+ —> H2O(l) + CO2(g)
what is the full equation and ionic equation
and the observations of adding :
To 2cm3 of hydrochloric acid add 2cm3 of
sodium hydroxide solution, note any
temperature change
REACTIONS OF ACIDS
Becomes warmer no other observations
Full Equation HCl(aq) + NaOH(aq) —> H2O(l) + NaCl(aq)
Ionic equation H+(aq) + OH- (aq) —> H2O(l)
what is the full equation and ionic equation
and the observations of adding :
To 2 cm3 of sulphuric acid add a small
strip of magnesium ribbon
REACTIONS OF ACIDS
Effervescence and became warmer, magnesium metal gets smaller ‘disappears’
Full Equation Mg(s) + H2SO4(aq) —-> MgSO4(aq) + H2(g)
Ionic equation Mg(s) + 2H+(aq) —> Mg2+(aq) + H2(g)
