Chapter 8 Flashcards by kiertan Varaitch

What is a reducing agent

An element that reduces other substances by losing electrons to another species

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group 2

Redox reactions with oxygen

2Mg + O2 —> 2MgO

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group 2

Redox reactions with water

Forms alkaline hydroxide and hydrogen

Mg + 2H2O —> Mg(OH)2 + H2

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group 2

Redox reactions with dilute acids

Metal + acid —> salt + hydrogen

Mg + HCl —> MgCl + H2

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Why does reactivity increase down group 2

The ionisation energies decrease down the group

Because attraction between nucleus and outer electrons decrease

atomic radius increase
more shielding

Less energy needed to remove electron
More easily lost
More reactive

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group 2

Redox reactions with

Metal oxides and water

Releases hydroxide ions, OH-
Forming alkaline solutions

MgO (s) + H2O —> Mg2+(aq) + 2OH-(aq)

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Since group 2 hydroxides are only slightly soluble in water

What happens when the solution becomes saturated

Further metal and hydroxide ions will form a solid precipitate

Mg2+ (aq) + 2OH- (aq) —> Ca(OH)2

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What is the solubility of hydroxides in water

down a group

Alkalinity increases
Solubility increases
pH increases
More OH- ions

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Testing the solubility and alkalinity if group 2 down a group

1) add partial of each group 2 oxide to water
2) shake mixture
3) insufficient water to dissolve all metal hydroxide, undissolved white solid at bottom - saturated solution
3) measure pH

Should increase down group

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How can metal hydroxides be used in everyday

Used to neutralise acidic fields
Ça(OH)2 (s) + H+ (aq)—> Mg2+(aq) + 2H2O

Used in antacids for treating acid indegstion
CaCO3 (s) + HCl (aq) —> CaCl2 (aq) + H2O + CO2
Mg(OH)2

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Boiling point trend for

Halogens

Boiling points increase down group
Because

More electrons
Stronger London forces
More energy required to break London force

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Halogens are oxidising agents - reduced

Definition

Substance that has the ability to oxidise other substances by gaining electrons

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How do displacement reactions show the reactivity of halogens

Halogen reactivity decrease down a group

1) a solution of each halogen is added to aqueous solutions of the other halides

If the halogen added is more reactive than halide then

2) a reaction takes place - displacing
3) the solution changes colour

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Colours of halogens in water

I2 - brown
Br2- orange
Cl2 - pale green

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Bromine and iodine colours are similar

How to overcome confusion

Add a non polar solvent = cyclohexane
Shake

Halogens dissolve more readily
Colours are easier to tell apart

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Halogen colours in cyclohexane

I2- pale green
Br2- orange
Cl2- violet

Equation of displacement between

Cl and Br

Cl2 + Br- —> Cl- + Br2 (aq)

What about fluorine and astanine

Fluorine - Pale yellow gas

Reacts with all substances it comes into contact with

Astanine - has not been seen

It’s radioactive and decays rapidly

Trend in reactivity in halogens

Halogens react by gaining electrons
As you go down the group tendency to gain a electron decreases
Because

Atomic radius and shielding increases
Less nuclear attraction to capture an electron from another species
Reactivity decreases

What is disproportionation

Example

Redox reaction in which the same element is both oxidised and reduced

Chlorine with water and cold dilute sodium hydroxide

Chlorine with dilute sodium hydroxide
- dissolve more readily than in water

Cl2 + 2NaOH -> NaClO + NaCl + H2O

Identifying a carbonate

Add dilute nitric acid (HNO3)

Bubble gas through limewater (Ca(OH)2)

CO2 (g) + Ca(OH)2 —> CaCO3 (s) + H2O

White precipitate - milky

Identifying sulfate

Add barium chloride to a solution

If white precipitate formed then sulfate is present

Ba2+ + SO42- —> BaSO4

Identifying halides

Aqueous silver ions react with halide ions to from precipitate

0) add dilute nitric acid - remove ions that might interfere
1) add silver nitrate AgNO3 To solution
2) White cream yellow - Cl2, Br2, I2 precipitate formed
3) add aqueous ammonia to gets solubility
4) dissolves in :
Cl2 - dilute ammonia
Br2 - concentrated ammonia
I2 - insoluble in concentrated

Correct order for analysing unknot inorganic compound

Carbonate CO3 2-
Sulfate SO4 2-
Halide Cl-

Why is there an order for tests

You have to carry out carbonate before sulfate - because it is the only one with effervescence - barium carbonate also forms a white precipitate Do halide last - Ag2CO3 + Ag2SO4 will both form precipitates

Testing for a mixture of ions Carbonate test

Complete until effervescence stops Carbonate ions will be removed Use dilute HNO3

Testing for a mixture of ions Sulfate test

Add excess Ba(NO3)2 Filter out white precipitate Can’t use BaCl2 - will show in halide test

Testing for a mixture of ions Halide test

Add AgNO3 Add NH3 to confirm

Testing for cations NH4 + Aqueous ammonium ions will react with hydroxide ions

1) add NaOH to solution of ammonium ions 2) ammonium gas produced 3) mixture is warmed to release gas 4) test gas with moist pH indicator paper Ammonia is alkaline Will turn paper blue

Why does chlorine have a lower boiling point than iodine

The strength of the London forces between molecules of chlorine is weaker than the forces between iodine Because The size and relative mass of chlorine is less than iodine And contains fewer electrons

Melting point of metals across period

Metal melting points increases across a period Giant metallic structure with metallic bonding Increases because metallic bonding increases across period, Because of Increasing positive charge More delocalised electrons Decreasing atomic radius

Melting point of giant covalent structures

Have the highest melting point Strong covalent bonds holding together atoms Large energy needed to overcome bonds

Simple molecules boiling point across period

Bp depends on strength of London forces between the molecules London forces are weak, easily overcome Bp are very low Bigger molecules with more atoms means stronger London forces Sulfur is biggest so has strongest London forces and highest bp Argon is lowest because exists as monoatomic with very weak London forces

State the effect of using cyclo hexane

Halogen will dissolve Form a separate layer Yellow will be organic - bromine Brown will be violet - iodine

Benefits of chlorine

Kills disease causing microorganism Stops further reinfection Makes water safe Prevents algae, bad taste and smell

Risks of chlorine

Respiratory irritant If reations with organic hydrocarbons - chlorinated hydrocarbons cause cancer Chemical burns

What are 2 alternatives of using chlorine to treat water

Ultraviolet and ozone

Benefits and disadvantages of ozone

Strong oxidising agent Kills diseases causing microorganisms Expensive to produce Short half life Not permanent

Benefits and disadvantages of uv

Kills microorganisms - damaging DNA Ineffective in cloudy water Won’t stop from being contaminated further down