Photography

Types of Chemical Reactions

Precipitation Reactions —Idea is that two soluble reactants are mixed together

Solutions

Solute - substance there is less of

Solvent - substance there is more of

Solution - combination of solute and solvent

Aqueous solutions - water is the solvent

Water, the common solvent

Polar molecule

Hydration - ionic solids dissolves in water, process aided by polar ends of the water molecules attracting the positive and negative ions in the solute

MX (s) dissolves in water → M⁺ (aq) + X^-(aq)

M⁺ (aq) means aquated M⁺ ion - negative, oxygen end of water attracted

X^-(aq) means the aquated X^- ion - positive, hydrogen ends of water molecule attracted

Strength (strong/weak) and solubility (concentrated/dilute) are two different concepts

Strength

If a given reactant is a strong electrolyte, it ionizes 100%. Reactant should be written as ions.

Salts: NaCl, KCl, KNO₃, AgNO₃, FeCl₃

Acids: HCl, HNO₃, H₂SO₄(for loss of one proton)

Bases: NaOH, KOH, Ca(OH)₂

Anything that is somewhat ionized but not 100% ionized is a weak electrolyte; often a whole lot less than 100%!

Salts:

Acids: Acetic acid, HSO_4^-, H₂SO₃

Bases: Ammonia

A reactant may be soluble but a non-electrolyte, i.e., no ions at all are formed: sugar, ethanol

Solubility, usually expressed in moles of solute per liter of solution

A concentrated solution has a lot of solute for a given volume of solution

A dilute solution has a small amount of solute for a given amount of solution

Is a solid (precipitate) formed?

Learn by seeing and doing

Microchemical tests of pigments

Formation of Prussian Blue

Formation of photosensitive papers - AgNO₃ + KCl

Simple solubility rules of salts in water

Acid-Base Reactions

Arrhenius model— Arrhenius defined acids as substances that produce hydrogen ions, H⁺, in water and bases as substances that produce hydroxide ions, OH^-, in water.

Bronsted-Lowry model

Acid is a proton (H nucleus) donor

Base is a proton acceptor

Must have at least one of each!

Example of a strong acid and a strong base:

HCl (aq) + NaOH (aq) → H₂O + Na⁺ (aq) + Cl^- (aq)

Example of water as the acid and the base!

Autoionization

Amphoteric substance can act either as an acid or a base. Water common ex.

H⁺ is so tiny that it combines with one or more solvent water molecules and often written as a hydronium ion, H₃O⁺

H₂O + H₂O H₃O⁺+ OH^-

Or better, H₂O + H₂O H₃O⁺(aq) + OH^-(aq)

aq means aquated, surrounded by water

Example of weak acid

HA + H₂O H₃O⁺+ A^-

H₂CO₃ + H₂O H₃O⁺+ HCO₃^-

HCO₃^- + H₂O H₃O⁺+ CO₃^2-?

Example of weak base

B^- + H₂O HB + OH^-

CO₃^2-+ H₂O HCO₃^- + OH^-

HCO₃^- + H₂O H₂CO₃+ OH^-?

Carbonate system:

H₂CO₃ + H₂O H₃O⁺+ HCO₃^-

Acid Base

CO₃^2- + H₂O HCO₃^- + OH^-

Base Acid

HCO₃^- + H₂O H₃O⁺?orOH^-?

Base/Acid Acid/Base

pH scale is based on the Bronsted-Lowry model

NB! There is a lower case p and an upper case H (for hydrogen ion)

"…pH is the negative log(arithm) of the hydronium ion concentration"

Demonstrations/analogies

Idea is of a continuum - a gradual change over a large range

Sweet and sour sauce

Potassium permanganate (KMnO₄) solutions

1/2

128

256

>256

Yellow to red food coloring

Yellow	10	10	10	15	15	5	0
Red	0	2	6	15	30	30	30
Yellow/Red	large!	5	1.67	1	0.5	.016	0

How can you display a large range of values on a linear scale?

Linear	10x	Log	-Log
0.00001	-5	-5	5
0.0001	-4	-4	4
0.001	-3	-3	3
0.01	-2	-2	2
0.1	-1	-1	1
1	0	0	0
10	1	1	-1
100	2	2	-2
1000	3	3	-3
10000	4	4	-4

Natural #	1	2	3.2	5	7	10
Log	0.00	0.30	0.51	0.70	0.85	1.00

Concentration usually in units of moles/liter, designated by [ ]

Thus concentration of the hydronium ion, H₃O⁺, is written as [H₃O⁺]

pH is DEFINED as -log [H₃O⁺],

i.e., the negative, base-ten logarithm of the hydronium ion concentration.

[H₃O⁺][OH^-] = K_w = 1.0x10^-14

Thus, if [H₃O⁺] increases (decreases), [OH^-] MUST decrease (increase), but as with the demonstrations, neither can ever = 0.

[H₃O⁺] = [OH^-]

means a neutral solution;

[H₃O⁺] = 10^-7; [OH] = 10^-7

pH = 7.0

[H₃O⁺] > [OH^-]

means acidic solution;

[H₃O⁺] > 10^-7; [OH] < 10^-7

pH < 7.0

[H₃O⁺] < [OH^-]

means basic solution;

[H₃O⁺] < 10^-7; [OH^-] > 10^-7

pH > 7.

Lewis model

Acid is an electron pair acceptor, i.e., there is an empty atomic orbital that can share/accept a pair of electrons

Base is an electron pair donor

Again, need an acid and a base

Lewis model more general than Bronsted-Lowry and can explain all Bronsted-Lowry reactions

See examples in text - particularly metal cations (acids) being surrounded by water (bases) or other species (ligands)

Oxidation-Reduction Reactions (RedOx)

Like Lewis acid-base reactions, emphasis is on the electrons, but now electrons are transferred (lost/gained) rather than just shared

Terminology somewhat historic:

4 Fe + 3 O₂ → 2 Fe₂O₃

4 Na + O₂ → 2 Na₂O

2 Na + Cl₂ → 2 NaCl

In all cases, products found to contain ions (O^-2, Cl^-, Fe⁺³, Na⁺)

Oxygen is an oxidizing agent!

By analogy, chlorine is also an oxidizing agent

Oxygen (chlorine) gains electrons and is said to be reduced

O₂ → O^-2via O₂ + 2 e^-→ O^-2

Cl₂ → 2 Cl^- via Cl₂ + 2 e^- → 2 Cl^-

The Fe and the Na are oxidized (by the oxidizing agent) and thus are themselves reducing agents and lose electrons:

Fe → Fe⁺³ via Fe - 3e^- → Fe⁺³ , or, more commonly,
Fe → Fe^{+3 +}3e^-

Na → Na⁺ via Na - e^-→ Na⁺ , or, Na → Na⁺ + e^-

Oxidation numbers - valences (?)

Essentially a mutually-agreed-upon bookkeeping system.

Oxidation states/oxidation numbers are defined by a set of rules

But, periodic table predicts a lot about oxidation numbers/valence states. Examples from the Representative Elements, groups 1A - 8A

"Oxidation states/oxidation numbers of the atoms in a covalent compound (are defined) as the imaginary charges the atoms would have if the shared electrons were divided equally between identical atoms bonded to each other or, for different atoms, were all assigned to the atom in each bond that has the greater attraction for electrons. … for ionic compounds containing monatomic ions, the oxidation states of the ions are equal to the ion charges."

Can now make some sense of the term 'reduction'

From above we know that oxidation is loss of electrons and reduction a gain of electrons

Now can say that oxidation is a gain in the oxidation number and reduction, not surprisingly, is a REDUCTION or lowering of the oxidation number.

Substance that is:	Agent	Electrons	Oxidation Number
Oxidized	Reducing	Lost/donated	Raised
Reduced	Oxidizing	Gained/accepted	Lowered

where valence and oxidation number are taken as synonyms

What is going on here:

2 Fe₂O₃ + 3 C → 4 Fe + 3 CO₂

CuO + H₂ → Cu + H₂O

Redox equations:

Main idea is that the number of electrons lost by the species being oxidized is equal to the number of electrons gained by the species being reduced.

More difficult in practice sometimes - particularly in water solutions.

In this course, you are not expected to be able to balance complicated redox reactions - simple ones - yes.

You should be able to assign oxidation numbers and identify which species are being oxidized/reduced

Transition Metal Chemistry

Similar properties in given horizontal period as well as vertical groups

Why? - from aufbau filling, last electrons added are inner electrons and the outer, s electrons are the ones that are lost first

Typical characteristics when combining with nonmetals to form ionic compounds:

More than one common oxidation state.

Complex ions: Lewis bases (ligands) surrounding central, positive cation

Bases can be neutral, e.g., NH₃, OH₂ (written this way to emphasize that the O is the part where bonding will take place, i.e., where the pair of electrons is that will be donated)

Bases can be negatively charged, e.g., Cl^-, CN^- (a pseudo halide), SO₄^2-, S₂O₃^2-, SCN^-

Colored

Electron configurations

Outer s first in - first back out

At least for first row transition elements, 4s and 3d orbitals of the neutral atoms are at about the same energy -> some configurations are a little different than you might expect.

However in the cations, the energies of the 3d orbitals are less than the energy of the 4s orbital

(Look at Table 20.2 - replace)and correlate these common states to electron configurations

Ionization energies increase slightly as you go across the row

Size decreases across the row - higher effective nuclear charge

Coordination compounds and complex ions - more

A compound, obviously, is neutral

Coordination compound usually contains a complex ion and counter ions (anions or cations as needed to produce an overall 0 charge)

Example

Werner's original ideas of primary and secondary valence now called oxidation state/number and coordination number - the number of Lewis bases/ligands surrounding the central cation

6 and then 4 are the most common coordination numbers

Ligands

Lewis base/ligand has one (or more) pairs of electrons to donate to the central cation

One bond - monodentate or unidentate ligand, e.g., ________

More than one atom in ligand that can form a coordinate covalent bond

Called chelating ligands/agent or chelates - Greek claw

Bidentate - two sites

More - polydentate

Nomenclature

Original names a zoo, e.g., color names or names of discoverers, or, slightly better: ferric ferricyanide, ferric ferrocyanide, ferrous ferricyanide, etc.

Black and White Photography

See Gelatin Silver processing chemistry handout from Pradip in lab

For each reaction described below, be able to classify the type (see all of the above), be able to deal with ionic and net ionic equations, etc.

Light sensitive silver-halides are starting point

We'll assume AgBr, though AgI, AgCl possible

AgBr insoluble - so how get it on the paper?

AgNO₃ (aq) + KBr (aq) → AgBr (s) + KNO₃ (aq)

Goddard, 1840, exposed Daguerre's silver-copper coated plates to bromine and/or iodine vapors:

2 Ag _(s) + I_{2 →} 2 AgI _(s)

Similar reactions for Br₂ and also using Cu

Exposure of light sensitive material, e.g. AgBr, to form latent image

Idea simple: 2 AgBr(s) 2 Ag (s) + Br₂ s

Remember ∆E = hν

h is Planck's constant

ν is the frequency of the light

hν is used to indicate a bundle or packet of energy - a photon

The photon can knock off an electron from a species - more below

Chemists use hν over the arrow in a chemical equation to indicate a photochemical reaction, i.e., that light is involved in the process.

Another possibility would be 2 AgBr(s) + hv → 2 Ag (s) + Br₂but this is usually not done

AgBr on the film is in units/crystal units called grains

Size of grains can be carefully controlled in film manufacture

Crucial idea: several Ag atoms formed from the photochemical process sensitize the entire grain

During development the whole grain will 'react'

What is 'a few'? Perhaps 3-4 silver atoms out 10¹⁰ Ag nuclides (atoms or ions)!

At this point, after the exposure to light, the latent image is on the film

Developing (the latent image)

Latin: to lie hidden

Greek: to escape notice

Spanish word to have pictures developed:

Webster's New Collegiate Dictionary: Present and capable of becoming though not now visible or acitve; dormant; ptotential

Make the latent image visible

Exposed AgBr grains changed to visible metallic deposit: Ag⁺ +e^- → Ag_(s)

Since the process is one of reduction we need a reducing agent.

Fox Talbot, in the 1840s, used an idea of J. R. Reade and used gallic acid as a mild reducing agent

Reducing agents used today are not a whole lot different!

On a separate page (Orna and Goodstein, pp 337-339) are some common reducing agents used today and a sample reaction.

Other components of the developer

Accelerator - speed up the otherwise rather slow process

Several alkaline (basic) materials used that speed up the reductions without consuming the gelatin substrate

Examples:

Sodium carbonate, Na₂CO₃ , and potassium carbonate - CO₂formed in reaction - gives rise to bubbles and thus blisters on the film

Sodium tetraborate (borax) Na₂B₄O₇ and sodium metaborate, NaBO₂ - these are milder

Preservative - preserve/protect the reducing agent

Reducing agents tend to be oxidized!

Atmospheric oxygen is always around

Want a substance that preferentially will react with atmospheric oxygen - sacrificial lamb

Sodium sulfite used extensively for this purpose

2 Na₂SO₃ + O₂ → 2 Na₂SO₄

Sodium sulfate inactive and very soluble

Sodium sulfite is basic: H₂SO₃ + H₂O HSO₃^-+ H₃O⁺

HSO₃^-+ H₂O SO₃^-² + H₃O⁺

SO₃^{- 2}+H₂O HSO₃^-+ OH

Restrainer - keep the alkaline materials present (accelerator and preservative) from fogging the negative - need some acid

Somehow (?) halides are good for this, and Br^-often added

The solvent - usually water

Temperature also important

Stop further reaction

Lots and lots of water would stop further reaction

Often a 3-4 weight percent solution of acetic acid is used to neutralize the alkaline conditions necessary for the developing process to take place

If indicator added, can check to see what the pH is at any time

Fixing - remove the unexposed (thus, unreduced) AgBr

If left, the AgBr would be eventually exposed to light and become darkened

AgBr insoluble, so just washing isn't sufficient

Need to dissolve the AgBr via forming a complex ion.

Most common substance is sodium thiosulfate, older name was sodium hyposulfate, and thus 'hypo': Na₂S₂O₃

Another is sodium ammonium thiosulfate

AgBr (s) + S₂O₃^-² V AgS₂O₃+ Br^-

^{AgBr (s) + 2 S₂O₃^-² V Ag(S₂O₃)₂₃^-+ Br-}

The resulting complex ion is soluble!

Wash and dry

Get rid of all fixer and original AgBr present

Drying removes the water solution, further aiding the preservation of the image

Reversal processing

Instead of making a negative, make a positive (slide)

After developing and removing from the stop bath:

Unexposed AgBr is still on film, and it would constitute the positive image if it could be developed

Instead of fixing, put in a bath that dissolves away the newly formed silver image - the negative portion

Sodium dichromate, potassium dichromate, or potassium ferricyanide used

Once the silver image is removed, the remaining silver halide is exposed to light, or chemical means can be used

Then developed and fixed as above

Result is a positive

Dye sensitizers

Untreated AgBr only responds to short wavelengths - from UV to cyan (curve A) in figure from Williamson and Cummins, pg 298

In 1873, H. W. Vogel found that an appropriate organic dye in the emulsion could extend the wavelength range to green

The dye molecules somehow become attached to the crystals, absorb energy from the light, and transfer electrons to the crystals

Curve B - called orthochromatic dye/emulsion

Panchromatic dyes, C and D, extend the range through the red

Cyanines - alternating single and double bonds (conjugated) along the chain, with resulting positive N sites that can attach to the Br^- in the crystal

Still other dyes, curve E, can extend the range into the IR

Suppose want only the IR - add a filter to filter out the visible

Redox in cyanotypes and VanDyke prints - see lab writeup

Photography

Types of Chemical Reactions

Precipitation Reactions —Idea is that two soluble reactants are mixed together

Solutions

Solute - substance there is less of

Solvent - substance there is more of

Solution - combination of solute and solvent

Aqueous solutions - water is the solvent

Water, the common solvent

Polar molecule

Hydration - ionic solids dissolves in water, process aided by polar ends of the water molecules attracting the positive and negative ions in the solute

MX (s) dissolves in water → M+ (aq) + X- (aq)

M+ (aq) means aquated M+ ion - negative, oxygen end of water attracted

X- (aq) means the aquated X- ion - positive, hydrogen ends of water molecule attracted

Strength (strong/weak) and solubility (concentrated/dilute) are two different concepts

Strength

If a given reactant is a strong electrolyte, it ionizes 100%. Reactant should be written as ions.

Salts: NaCl, KCl, KNO3, AgNO3, FeCl3

Acids: HCl, HNO3, H2SO4 (for loss of one proton)

Bases: NaOH, KOH, Ca(OH)2

Anything that is somewhat ionized but not 100% ionized is a weak electrolyte; often a whole lot less than 100%!

Salts:

Acids: Acetic acid, HSO4-, H2SO3

Bases: Ammonia

A reactant may be soluble but a non-electrolyte, i.e., no ions at all are formed: sugar, ethanol

Solubility, usually expressed in moles of solute per liter of solution

A concentrated solution has a lot of solute for a given volume of solution

A dilute solution has a small amount of solute for a given amount of solution

Is a solid (precipitate) formed?

Learn by seeing and doing

Microchemical tests of pigments

Formation of Prussian Blue

Formation of photosensitive papers - AgNO3 + KCl

Simple solubility rules of salts in water

Acid-Base Reactions

Arrhenius model— Arrhenius defined acids as substances that produce hydrogen ions, H+, in water and bases as substances that produce hydroxide ions, OH-, in water.

Bronsted-Lowry model

Acid is a proton (H nucleus) donor

Base is a proton acceptor

Must have at least one of each!

Example of a strong acid and a strong base:

HCl (aq) + NaOH (aq) → H2O + Na+ (aq) + Cl- (aq)

Example of water as the acid and the base!

Autoionization

Amphoteric substance can act either as an acid or a base. Water common ex.

H+ is so tiny that it combines with one or more solvent water molecules and often written as a hydronium ion, H3O+

aq means aquated, surrounded by water

Example of weak acid

HA + H2O H3O+ + A-

H2CO3 + H2O H3O+ + HCO3-

HCO3- + H2O H3O+ + CO32-?

Example of weak base

B- + H2O HB + OH-

CO32- + H2O HCO3- + OH-

HCO3- + H2O H2CO3 + OH-?

Carbonate system:

H2CO3 + H2O H3O+ + HCO3-

Acid Base

CO32- + H2O HCO3- + OH-

Base Acid

HCO3- + H2O H3O+ ?orOH- ?

Base/Acid Acid/Base

pH scale is based on the Bronsted-Lowry model

NB! There is a lower case p and an upper case H (for hydrogen ion)

"…pH is the negative log(arithm) of the hydronium ion concentration"

Demonstrations/analogies

Idea is of a continuum - a gradual change over a large range

Sweet and sour sauce

Potassium permanganate (KMnO4) solutions

Yellow to red food coloring

How can you display a large range of values on a linear scale?

Concentration usually in units of moles/liter, designated by [ ]

Thus concentration of the hydronium ion, H3O+, is written as [H3O+]

pH is DEFINED as -log [H3O+],

i.e., the negative, base-ten logarithm of the hydronium ion concentration.

[H3O+][OH-] = Kw = 1.0x10-14

Thus, if [H3O+] increases (decreases), [OH-] MUST decrease (increase), but as with the demonstrations, neither can ever = 0.

[H3O+] = [OH-]

means a neutral solution;

[H3O+] = 10-7; [OH­] = 10-7

MX (s) dissolves in water → M⁺ (aq) + X^-(aq)

M⁺ (aq) means aquated M⁺ ion - negative, oxygen end of water attracted

X^-(aq) means the aquated X^- ion - positive, hydrogen ends of water molecule attracted

Salts: NaCl, KCl, KNO₃, AgNO₃, FeCl₃

Acids: HCl, HNO₃, H₂SO₄(for loss of one proton)

Bases: NaOH, KOH, Ca(OH)₂

Acids: Acetic acid, HSO_4^-, H₂SO₃

Formation of photosensitive papers - AgNO₃ + KCl

Arrhenius model— Arrhenius defined acids as substances that produce hydrogen ions, H⁺, in water and bases as substances that produce hydroxide ions, OH^-, in water.

HCl (aq) + NaOH (aq) → H₂O + Na⁺ (aq) + Cl^- (aq)

H⁺ is so tiny that it combines with one or more solvent water molecules and often written as a hydronium ion, H₃O⁺

HA + H₂O H₃O⁺+ A^-

H₂CO₃ + H₂O H₃O⁺+ HCO₃^-

HCO₃^- + H₂O H₃O⁺+ CO₃^2-?

B^- + H₂O HB + OH^-

CO₃^2-+ H₂O HCO₃^- + OH^-

HCO₃^- + H₂O H₂CO₃+ OH^-?

H₂CO₃ + H₂O H₃O⁺+ HCO₃^-

CO₃^2- + H₂O HCO₃^- + OH^-

HCO₃^- + H₂O H₃O⁺?orOH^-?

Potassium permanganate (KMnO₄) solutions

Thus concentration of the hydronium ion, H₃O⁺, is written as [H₃O⁺]

pH is DEFINED as -log [H₃O⁺],

[H₃O⁺][OH^-] = K_w = 1.0x10^-14

Thus, if [H₃O⁺] increases (decreases), [OH^-] MUST decrease (increase), but as with the demonstrations, neither can ever = 0.

[H₃O⁺] = [OH^-]

[H₃O⁺] = 10^-7; [OH] = 10^-7

[H₃O⁺] > [OH^-]

[H₃O⁺] > 10^-7; [OH] < 10^-7

[H₃O⁺] < [OH^-]

[H₃O⁺] < 10^-7; [OH^-] > 10^-7

4 Fe + 3 O₂ → 2 Fe₂O₃

4 Na + O₂ → 2 Na₂O

2 Na + Cl₂ → 2 NaCl

In all cases, products found to contain ions (O^-2, Cl^-, Fe⁺³, Na⁺)

O₂ → O^-2via O₂ + 2 e^-→ O^-2

Cl₂ → 2 Cl^- via Cl₂ + 2 e^- → 2 Cl^-

Fe → Fe⁺³ via Fe - 3e^- → Fe⁺³ , or, more commonly,
Fe → Fe^{+3 +}3e^-

Na → Na⁺ via Na - e^-→ Na⁺ , or, Na → Na⁺ + e^-

2 Fe₂O₃ + 3 C → 4 Fe + 3 CO₂

CuO + H₂ → Cu + H₂O

Bases can be neutral, e.g., NH₃, OH₂ (written this way to emphasize that the O is the part where bonding will take place, i.e., where the pair of electrons is that will be donated)

Bases can be negatively charged, e.g., Cl^-, CN^- (a pseudo halide), SO₄^2-, S₂O₃^2-, SCN^-