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1

Transition metal definition

D-block elements which form one or more stable ions which have incompletely filled d-orbitals

2

In the first row of d-block elements, which elements are not transition metals?

Sc and Zn

Sc only forms 1 ion, which has an empty d subshell

Zn only forms 1 ion, which has a full d-subshell

3

Electronic configuration of d-block elements:

- 4s fills before 3d as 4s has a lower energy

- Cr and Cu are exceptions- an electron in the 4s orbital moves into the 3d as this gives a lower energy

- Electrons then fill 3d subshells singly before pairing up

- When ions are formed, the 4s orbital always loses its electrons first

4

Determining the iron content using a redox titration:

- Fe (II) ions can be oxidised to Fe (III) ions by potassium manganate (VII) in acidic solution

- Known volume of Fe (II) solution is titrated with potassium manganate (VII) solution of known concentration

- Potassium manganate in burette

- End point is when the first permanent pink colour is observed

5

Common oxidation states of iron

Fe (II)

Fe (III)

6

Common oxidation sates of copper

Cu (I) - unstable

Cu (II)

7

Colour of iron (II) hydroxide

Colour of Iron (III) hydroxide

Dark green

Red/orange

8

Colour of aqueous iron (II) ions

Colour of aqueous iron (III) ions

Pale green

Yellow

9

Colour of aqueous copper (I) ions

Colour of aqueous copper (II) ions

Unstable - no colour

Pale blue

10

Colour of copper (II) hydroxide

Colour of copper (II) ammonia complex

Blue

Dark blue

11

Describe what happens when ammonia solution is added to copper (II) hydroxide

- Ammonia = source of OH- ions

- Pale blue precipitate of copper (II) hydroxide dissolves to give a deep blue solution of a copper ammonia complex ion

PALE BLUE------------------DARK BLUE

12

Ligand definition

A negatively charged ion or neutral molecule with a LONE PAIR of electrons which it donates to a central TRANSITION METAL ION to form a COORDINATE bond

13

Complex ion definition

A transition metal ion or atom surrounded by a number of ligands

14

Ligand substitution definition

+ when it occurs

When a ligand is swapped for another ligand

Usually results in a colour change

Occurs if the new complex formed is more stable than the previous complex

Sometimes substitution is only partial

15

Writing complex ions

The overall charge of the complex ion is the sum of the charge on the central metal ion + the charges on the ligands

(Fe(H2O)6)^2+

(Fe(H2O)6)^3+

(Cu(H2O)6)^2+

(Cu(NH3)4)^2+

(CuCl4)2-

16

Another word to describe a coordinate bond:

Dative

17

Monodentate ligand definition

a ligand that attaches to a transition metal through one atom only (eg OH-, H2O)

18

Bidentate ligand definition

a ligand with 2 atoms with lone pairs or negative charges, which forms 2 bonds to a metal ion

19

Structure of ethanedioate

How does it act as a ligand?

-^OO=C-C=OO^- (see notes)

Bidentate ligand- has 2 atoms with lone pairs which forms 2 coordinate bonds to a metal ion

20

Iron (II) hydroxide and iron (III) hydroxide reaction/colour change with ammonia solution

Both do not form complexes with ammonia

21

Coordination number definition:

the number of coordinate bonds that are formed with the central metal atom

Usual coordination numbers are 4 and 6

22

Shape of complex ion and bond angle when coordination number is 6:

Octahedral

90 degrees

23

Shape of complex ion and bond angle when coordination number is 4

Tetrahedral - 109.5

OR

Square planar - 90

24

Why shape of complex ion differs when coordination number is 4?

Shape either tetrahedral or square planar

Depends on ligands

If ligands are small (eg water or ammonia) 6 can fit around central metal ion

If ligands are larger (eg Cl-) only 4 can fit around central metal ion

25

How ligands lead to colour in transition metal complexes:

- Normally the 3d orbitals of transition metal ions all have the same energy

- When ligands bond to the ions- orbitals close to the ligands are pushed to slightly higher energy levels

- As a result- some of the orbitals gain energy which splits the 3d orbitals into 2 different energy levels

- When visible light absorbed- electrons can be excited to a higher energy level

- Light is only absorbed if the energy of the light matches the energy gap between the 2 energy states in the atom

- The frequency of the light absorbed is proportional to the gap between the energy levels (E=hv)

- The complementary colour is transmitted (the colour we see)

26

Why do transition metal ions of the same element often have different colours in aqueous solutions?

Transition metal ions have different oxidation states

The energy needed to excite a d electron to the higher energy level depends on the oxidation state of the metal

27

What does the colour of a transition metal complex depend on?

- Number of d-electrons present in the transition metal ion

- The arrangement of ligands around the ion - this affects the splitting of the d sub-shell

- The nature of the ligand- different ligands have a different effect on the relative energies of the d orbitals in a particular ion

28

Complementary colours:

- Complementary colours- opposite each other in a colour wheel

- Mix 2 complementary colours of light = white light

- The colour of a transition metal solution is the complement of the colour that's absorbed

- eg- blue solution- absorbs the red part of the spectrum

29

Colorimetry definition

A technique that can be used to find the concentration of a coloured solution

30

Explain how colorimetry can be used to find the concentration of a transition metal ion solution

- Intensity of colour depends on the concentration of the solution

- The more concentrated the solution, the darker its colour and the more light it absorbs

- Absorbance of a solution is proportional to its concentration

- High absorbance reading means a lot of light has been absorbed + sample is very concentrated

31

Describe how to use a colorimeter to measure concentrations of solutions

1) Make up a series of standard solutions of the test solution of known concentration

2) Select a filter with the complementary colour to the test solution

3) Zero the colorimeter with a tube of pure solvent

4) Measure the absorbance of the standard solution + plot a calibration curve, showing the concentration (x) against absorbance (y)

5) Measure the absorbance of the test solution and read off the concentration from the calibration curve

32

Describe how to use a visible spectrophotometer to measure concentrations of solutions

Uses a beam of light of a single colour (monochromatic light)

1) The beam of monochromatic light is passed through a solution of the complex

2) A detector measures the light intensity before and after it's passed through the solution- from this can work out absorbance

3) Different complexes absorb different colours- so produce characteristic visible absorption spectra

4) Use calibration curve to convert absorbance data to find concentration data

33

Balancing redox equations rules:

Check:

1) Atoms balance

2) Oxidation states balance

3) Charges balance

If oxidising agent contains oxygen- might have to add H+ ions and water to make equation balance

34

Components of an electrochemical cell:

- 2 half-cells joined together

- Half cell- a metal dipped in a salt solution of the metal's ion

- Voltmeter- measures the maximum potential difference between 2 half cells (Ecell)

- Salt bridge- provides an ionic connection between 2 half cells

35

Ecell definition

The potential difference between 2 half cells

36

Electrochemical cells involving metal ion/metal half cells

Eg- zinc/copper electrochemical cell

- Zinc has a more negative electrode potential (more reactive metal). Loses electrons more easily- oxidised

- Copper has a more positive electrode potential- (less reactive metal)- accepts electrons- reduced

37

What direction do electrons flow in an electrochemical cell?

From negative half cell to positive half cell

38

Half cell with more negative electrode potential- metal is....................

Half cell with more positive electrode potential- metal is....................

Oxidised

Reduced

39

Electrochemical cells involving half-cells based on different oxidation states of the same element in aqueous solution:

Eg- Fe 2+, Fe 3+

- Conversion between these ions occurs on surface of the electrode

- As neither reactants nor products are solids, something else needs to be the electrode

- Platinum or graphite used- conducts electricity, inert

40

Standard electrode potential definition:

Standard electrode potential of a half cell is the voltage under standard conditions when the half cell is connected to a standard hydrogen electrode

41

What are the standard conditions for measuring the standard electrode potential of a half cell?

- Concentration of solutions = 1 mol dm-3

- Temperature = 298 K

- Pressure = 100 kPa

42

Using a hydrogen electrode half cell to find the standard electrode potential of another metal ion:

- Measure electrode potential of half cell against standard hydrogen electrode

- Standard hydrogen electrode half-cell has value - 0.00V

- Means voltage reading will be equal to the standard electrode potential of other half cell

43

Calculating Ecell:

Ecell = (more positive electrode potential - more negative electrode potential)

44

Using electrode potentials to predict whether a reaction will happen:

- Write the 2 half equations as reduction

- Look at electrode potential of each half equation- one with more negative electrode potential- flipped to oxidation

- Combine half equations + make sure electrons are balanced

45

Techniques and procedures to set up and use electrochemical cells

1) Get a strip of each metal being investigated

2) Clean surface of metals using a piece of emery paper + clean grease from electrodes using propanone

3) Place each electrode into a beaker filled with a solution containing ions of that metal

4) Create a salt bridge to link the solutions together (eg soak a piece of filter paper in salt solution + drape it between 2 beakers)

5) Connect electrodes to voltmeter using crocodile clips + wires

46

Despite predicting a reaction will occur through an experiment using electrochemical cells, explain why it may not occur

- Rate of reaction may be so slow that reaction may not appear to happen

- Too high activation energy- stops a reaction from occuring

- Prediction could be wrong if conditions not standard

47

Electronic configuration of Chromium (24 electrons)

Chromium prefers to have 1 electron in each orbital of the 3d sub-shell and just 1 electron in the 4s sub-shell

Gives it more stability

1s2 2s2 2p6 3s2 3p6 3d5 4s1

48

Electronic configuration of Copper (29 electrons)

Copper prefers to have a full 3d subshell and just 1 electron in the 4s subshell

Gives it more stability

1s2 2s2 2p6 3s2 3p6 3d10 4s1

49

Characteristics of transition metals:

- Exist in variable oxidation states

- Form coloured ions in solution

50

Equation involved in the formation of rust:

Fe2+ + 2e- ----------------------- Fe

2H2O + O2 + 4e- -------------------------- 4OH-

As the iron half equation has a more negative electrode potential, the iron is oxidised and oxygen is reduced

Overall: 2H2O + O2 + 2Fe ---------- 2Fe 2+ + 4OH-

Fe2+ and OH- ions formed combine to produce precipitate of iron (II) hydroxide

Further oxidised to iron (III) hydroxide

This gradually turns into hydrated iron (III) oxide (rust)

51

Where oxidation and reduction occur when a water droplet causes iron to rust

Edges of droplet- oxygen concentration is higher so oxygen is reduced to hydroxide ions (cathode)

Centre of droplet by iron- oxygen concentration is low so iron is oxidised (anode) This is where iron corrodes. Electrons flow away from this site to areas with higher oxygen concentrations

(This explains why corrosion is always greatest at the centre of a water droplet - where oxygen supply is limited)

52

Rusting prevention:

- paint/apply oil- provides barrier between metal + atmospheric oxygen

- Sacrificial method- involves adding a metal with a more negative electrode potential to prevent rusting. Other metal is oxidised + corrodes first. Release of electrons promotes reduction in iron half cell to ensure iron in solid state always exists

eg- coating of zinc can be sprayed onto object (galvanising) or blocks of zinc can be bolted to the iron

53

Why is iron less likely to rust in alkaline conditions?

- More OH- ions

- Shifts position of equilibrium of water half equation left to try to reduce number of OH- ions

- Electrons produced

- Shifts position of Fe half equation towards Fe(s)

- Less Fe 2+ so less rust produced

54

Why do transition metals and their compounds make good catalysts?

- Can change oxidation states by gaining or losing electrons within their d-orbitals

- Means they can transfer electrons to speed up reactions

55

Transition metals and their compounds as hetrogeneous catalysts:

- Catalyst and reactants- different physical states

- Transition metals can use 3d and 4s electrons of the atoms on the catalyst surface to form weak bonds to reactants

56

Transition metals and their compounds as homogeneous catalysts:

- Catalyst and reactants- same physical state

- Change oxidation states by gaining or losing electrons within their d-orbitals