C4 & C5 chemistry

carbon dioxide test

• add lime water (calcium hydroxide)

• if present, cloudy and bubbling

• CO2 reacts with lime water & forms insoluble carbonate

Chlorine gas test

• tap water on blue litmus paper so damp

• bleaches damp blue litmus paper turning it red (due to acidity) then white

Hydrogen test

• lit splint over the open end of a test tube containing hydrogen gas

• will create a squeaky pop noise

• made because hydrogen burning rapidly with oxygen to make water

Oxygen test

• glowing splint over a test tube containing oxygen gas

• glowing splint inside a test tube will relight if oxygen is present

Flame test

• when metal ions are heated, energy is transferred to their electrons

• This makes the electrons become excited and move up to higher electron shells

• at these higher energy levels, they are unstable and move back down to their normal electron shells

• as they move back down, energy is transferred to the surroundings as radiation, which is seen as light

• Different metal ions produce different colours

Flame colours for common metal ions

• Lithium →red

• Sodium → yellow

• Potassium → lilac

• Calcium → orange-red

• Copper → green-blue

Group 1 metals

alkali metals

• form alkaline solutions when reacted with water

Physical properties of Group 1

• soft metals

• increase in softness going down group

• low densities

• very reactive

• good conductors of electricity

why do Group 1 metals have relatively low MPs decreasing going down group?

• gets further away from nucleus- atoms get larger due to increasing number of shells, creating a greater space away from nucleus. this distance decreases attractive forces between outermost electron and nucleus’ positive charge

reactivity trend of Group 1

• increases going down group (weaker FOA to overcome)

• only one electron lost, so easier to lose electron

• then obtains noble gas configuration

Alkali metal + oxygen

_____ oxide (superoxide for K)

Alkali metal + Chlorine

____ chloride + a white precipitate

Alkali metal + Water

____ Hydroxide + Hydrogen

Li + H2O observation

• slow reaction

• Li doesn’t melt

• fizzing can be seen and heard

Na + H2O Observation

• large amounts of heat causes Na to melt

• Hydrogen catches fire

• Na dashes on the surface

K + H2O Observation

• reacts violently

• enough heat released so hydrogen burns & produces a lilac flame

• K melts into a shiny ball

• K dashes on the surface

Why does reactivity increase going down the group?

• Alkali metals only lose 1 electron to gain noble gas configuration

• Going down, number of shells increase by 1

• further away from nucleus= weaker forces of attraction

• less energy required to overcome FOA so electron is lost more easily

Group 7

Halogens

• diatomic elements which form -1 halide ions

• formed by a single covalent bond

Group 7 State & Appearance at room temp

• F - yellow gas

• Cl - pale yellow/green gas

• Br - red/brown liquid

• I - purple/black solid

Characteristics of Group 7 elements

• F- very reactive, poisonous gas

• Cl- reactive, dense, poisonous gas

• Br - dense red-brown volatile liquid

• I - shimmery, crystalline solid, sublimes to form purple vapour

Boiling and Melting point trends

increase going down the group

• intermolecular forces strengthen as atoms get larger, more energy required to overcome forces

reactivity trend in Group 7

decreases going down

• fluorine- smallest halogen and closest to nucleus. ability to attract an electron is strongest in F2, making it most reactive

Halogen displacement reaction

• when a more reactive halogen displaces a less reactive halogen from an aqueous solution

• NEED TO KNOW: Cl, Br, I (most → least reactive)

Displacements of Cl, Br, I

Chlorine displaces: Br (yellow/orange colour seen), I (brown colour is seen)

Bromine displaces: I (brown colour is seen)

Iodine displaces: none

Chlorine + Bromine half equation

Cl2 + 2br- → 2Cl- + Br2

Chlorine + Iodine half equation

Cl₂ + 2I^- → 2Cl^- + I₂

Bromine + Iodine half equation

Br2 + 2I- → 2Br- + I2

Metal Halide reactions

• Cl, Br, and I react with metals to form compounds

• create metal halide salts

• halides hold a -1 charge

• rate of reaction is slower for halogens which are further down the group

Non metal halides

• halides react with non metals to form simple molecular covalent structures

Group Zero

Noble gases

• monatomic, colourless, non flammable gases at room temp

• unreactive

Group 0 characteristics

• low density

• increasing density down group/ get heavier

• non metals

uses of Group 0 elements

• chemically inert

• helium fills balloons as it doesnt burn and is less dense

• Neon, argon & Xenon are used in signs

• Ar is used to fill light bulbs

• Ar creates inert atmosphere for welding

gases at room temp

• individual atoms are widely spaced apart and so they have low densities

Why do group 0 have low MPs and BPs?

• atoms get larger moving down group

• BP increases going down (although still below 0*c)

• increase in intermolecular forces increasing amount of energy needed to overcome these forces

Transition metals properties

• hishly dense

• good conductors of heat & electricity

• lustrous

• malleable

• shiny when freshly cut

• form coloured compounds

• hard, strong metals

• multiple oxidation states

use of transition metals

• catalysts

• they don’t take part in the reaction

• catalytic ability stems from their ability to interchange oxidation states

• can form complexes with reagents which can easily donate and accept electrons

COMMON METAL CATALYSTS*****

• iron- HABER process

• vanadium pentoxide- contact process

• nickel- hydrogenation of alkenes

oxidation state facts

• compounds containing transition metals in different oxidation states will have different properties & colours in aqueous solutions

PREDICTING REACTIVITY OF DIFFERENT GROUPS

• Group 1-2: highly reactive- increasing reactivity, forms ionic compounds with NM

• Group 7: get less reactive going down group

• Group 0: elements are unreactive

Transition metals VS Group 1

• G1- +1 ions, TM- ions with variable charges

• G1- soft metals, TM- hard, strong metals

• G1- low MPs, TM- much higher MPs

• G1- very reactive, TM- less reactive

• G1- reacts with O2,H2, and G7, TM- react slow or not at all

• G1- tarnish in presence of oxygen, TM- takes several weeks to form metal oxides & require water

Reactivity series facts

• most metals react with dilute acids like HCl

• metal atoms form positive ions by loss of e- when reacted

• tendency of a metal to lose e- is measure of its reactivity → more reactive= easier it is to lose electrons

• metals that react with cold water form metal hydroxide + hydrogen

Metal + Acid…

→ metal salt + hydrogen

• iron + hydrochloric acid   →  iron(II)chloride + hydrogen

• METAL BECOMES A POSITIVE ION

reactivity series in order

K (reacts violently with H2O and acid)

Na (reacts quickly with H2O, violent w/ acid)

Ca

Mg

Zn

Fe

H

Cu

Ag

Au

rate of H2 production

more reactive a metal is = greater rate of H2 production so reaction is more vigorous

When do metal + acid/ water reactions take place?

• if the metal is able to displace the hydrogen in them

Metal Cations

can be identified by the colour of the precipitate they form on addition of NaOH

PRACTICAL

• few drops of NaOH slowly

Metal Cation colours

• Iron (II): green

• Iron (III): orange-brown

• Copper (II): blue

• Calcium: white

• Zinc: white

• Aluminium: white

• Magnesium: white

How to differentiate the metal cations which create white precipitates?

• add NaOH in excess

• Zinc and Magnesium will dissolve into colourless solutions. Calcium and Aluminium won’t.

• They can also be Flame tested to establish identity.

Test for Carbonate ions

1. add dilute acid

   • if carbon is present, CO2 will be formed, bubbles will be seen.

2. Add gas with limewater, Ca(OH)2, and IF CARBONATE PRESENT: white precipitate and CaCO3 will be formed.

Sulfate ion test

1. acidify sample with dilute HCl

2. add barium chloride (or nitrate)

IF SULFATE PRESENT:

• white precipitate is formed

Halide testing

1. acidify sample with dilute nitric acid (HNO3)

   • this removes carbonate ions which may give a false positive

2. add silver nitrate (AgNO3)

IF PRESENT: Silver halide precipitate is formed (AgX)

Flame Testing

1. dip loop of unreactive metal in dilute acid, and hold it in blue flame until there is no colour change (sterilisation step)

2. Dip loop into solid sample

3. Place loop at the edge of bunsen blue flame

4. colour can be observed

Flame Test Colours

• Li: Red

• Na: Yellow

• K: Lilac

• Ca: Orange/red

• Cu: Blue/green

Flame emission spectrometer

• detailed analysis

• used to identify multiple ions present

How does a Flame Spectrometer work

• exposes sample to a very hot flame and then measuring the intensity and wavelength of light emitted

• image created is viewed as a line emission spectra and each element has a characteristic pattern

What can affect rate of reaction?

• concentration of reactants in solution

• pressure

• temperature

• surface area

• catalyst

economic interest

• higher rate of reaction

• high atom economy

• high percentage yield

Effect of increase concentration or pressure

increases rate of reaction

• on a graph, line will be steeper.

Effects of increasing surface area and temperature

same effect as concentration and pressure.

Collision theory

• chemical reactions only occur when reactant particles collide with sufficient energy to react

• rate of reaction is dependent on the energy and numbers of the collisions

Why does increasing concentration or pressure of a solution increase rate of reaction?

• more reactant particles in a given volume- more frequent and successful collisions per second

Why does an increase in temperature increase rate of reaction

• particles have more kinetic energy than required activation energy

• more successful and frequent collisions per second

Why does a larger surface area increase rate of reaction

more room for reaction to take place so higher rate of reaction

Haber process

catalyst of iron

production of ammonia

450 degrees

200 atm

Sulphuric acid Contact process

• detergents, paints, fibres

• 450 degrees

• 2atm

• vanadium pentoxide

Calibration curve

• light intensity produced is directly proportional to number of ions vapourised

• used to determine concentration of metal ions in a solution by reference to a standard solution of known concentration on a calibration curve

mass spectrometer

• powerful analytic technique. The most useful instrument for accurate determination of the relative atomic mass of an element based on the abundance & mass of each of its isotopes

• also used to find the relative molecular mass of molecules

• a spectrum is produced of the mass/charge ratio against abundance

• There are several types, but all are based on the ratio of their charge to their mass

Mass spectrometry in identifying isotopes

• The height of the peaks shows the proportion of each isotope present

• molecules- not just atoms- can be analysed

• The molecular ion peak can identify molecular mass of a compound, however, different compounds may have the same molecular mass

• molecules can fragment as they are ionised and the fragments can pass through to give a range of different peaks- can fragment due to formation of characteristic fragments, or loss of small molecules

Advantages of instrumental analysis

-can analyse chemical substances due to advancements in technology

• X-rays, Infra-red, Mass Spectroscopy, Gas Chromatography, Flame photometry

-provide greater accuracy

-faster and easier to use

-automated and can perform multiple simultaneous sampling and testing

-very sensitive and can work with very small sample sizes

concentration in g/dm³

mass of solute (g) / volume of solution (dm³)

concentration

• measure of how much of a substance is present in a given volume

converting between g/dm³ and mol/dm³

from g/dm³ → to mol/dm³ = DIVIDE BY MOLAR MASS IN GRAMS

from mol/dm³ → to g/dm³ = MULTIPLY BY MOLAR MASS IN GRAMS

cm³ to dm³

• divide by 1000

dm³ to cm³

multiply by 1000

concentration in mol/dm³

• moles per unit volume

• concentration = moles / vol(dm³)

titration

• analysing the concentration of a solution

• acid base titrations are commonly used to determine exactly how much alkali is needed to neutralise a quantity of acid

• used to prepare salts or other precipitates in redox reactions

• indicators show the end point in a titration

• PHENOLPHTHALEIN being a popular choice

indicator choice in titrations

• PHENOPHTHALEIN is a popular choice- distinct colour (pink) shows

• wide range indicators, like litmus, aren’t suitable as they don’t give a sharp enough colour change at the end point

• universal indicator isn’t suitable- mix of indicators and has too many subtle colour changes

health and safety in titrations

• dilute HCL- may cause harm to eyes or skin

• Acids & alkalis are corrosive and should be handled with care

• avoid contact with the skin & use safety goggles with both substances

• pipette should always be used with a safety filler to avoid contact

EQUIPMENT in titrations

• 25cm³ volumetic pipette

• pipette filler

• 50cm³ burette

• 250cm³ conical flask

• small funnel

• 0.1 mol/dm³ NaOH solution

• sulphuric acid- concentration unknown

• phenolphthalein indicator

• clamp stand, clamp & white tile

titration practical

1. use pipette to place exactly 25cm³ NaOH solution into conical flask

2. place conical flask on a white tile soo tip of burette is inside of flask

3. add a few drops of a suitable indicator to the solution in conical flask

4. perform rough titre by taking burette reading and running solution in 1-3 portions while swirling flask continually

5. close tap when colour change is reached, and record volume, placing eye level with meniscus

6. now repeat with fresh NaOH

7. as rough end point volume is approached, add solution from burette one drop at a time until indicator changes colour

8. swirl after each addition and rinse the sides of the flask down with distilled water to make sure that all that was added was reacted

9. finish at first sign of colour change & record volume to nearest 0.05

10. repeat until 2 concordant results

REMOVE FUNNEL AFTER FILLING BURETTE- CAN DROP SOLUTION INTO BURETTE- LEADING TO ERROR

GAS VOLUME CALCULATIONS

volume = amount of gas (g) x 24 dm³mol^-1

avogardro’s law

at the same conditions of temp and pressure, equal amounts of gases occupy the same volume of space.

• at RTP and pressure, volume CCEPTED BY 1 MOLE OF ANY GAS is found to be 24dm³.

• 20 degrees celcius, 1 ATM

reaction yield

(actual / theoretical) x 100

reaction yield explained

amount of products retrieved from a reaction- you never get 100% yield in a chemical process for several reasons.

reasons in theoretical yield

• reactants left behind

• reactions may be reversible

• a high yield is never possible

• products may be lost during separation and purification

• may be side reactions - gas, precipitates

• can be lost in transfer of containers

actual yield & theoretical yield

• actual: recorded amount

• theoretical: amount obtained under perfect practical & chemical conditions

• percentage yield compares the two

economics of yield

• businesses look at yield to check out how successful chemical processes are, and will try it out with different reaction pathways, which are compared and evaluated, so a manufacturing process can be chosen

• - companies look for a high percentage yield as possible to increase profits and reduce costs and waste- COST EFFICIENT

atom economy

• analyses the efficiency of reactions

• studies the amount of reactants that get turned into useful products- illustrates what percentage mass is turned to useful product, and is used to obtain sustainable development

atom economy formula

100 x (total Mr of desired product / total Mr of all products)

choosing a reaction pathway

• reactions which have low atom economy use up lots of resources and produce a lot of waste material, which then needs to be dispose of, which is very expensive. UNSUSTAINABLE AND NOT ECONOMICALLY ATTRACTIVE.

companies AE

companies analyse different reaction pathsways to improve efficiency.

• atom economy, percentage yield, and efficiency are important and need to be considered when choosing a reaction pathway. High Yield & fast rate of reaction are desirable

• in reversible reactions, position of equillibrium may need to be changed in favour of the products by altering reaction conditions

measuring rates

reactant used(or product formed) = rate of reaction x time taken

reaction times

• different reactions take place at different rates: rusting= slow, explosions= fast

rate of reaction

• rate of reactions can be measured either by how fast a reactant is used up, or by how fast the product is made

product made equation

product made = rate of reaction x time taken

measuring rates units

cm³, dm³, or volume.

usually measuring time in seconds

measuring mass

if product is a gas→ reaction can be performed in an open flask on a scale to measure loss in mass of reactant.

-cotton wool is placed in the mouth of the flask, which allows gas out, but prevents any molecules from being ejected- not good for H2 or any gas with a small Mr

measuring volume of gas

• gas trapped and volume is measured in pushed out gas syringe.

• (do the one tub one if gas is not water soluble)

• exampls: Mg + HCl → H2 + MgCl

measure of precipitates

*for example, sodium thiosulfate + hydrochloric acid (Na2S2O3 + 2HCl)*

• time as you add acid

• watch until you cannot see X

rate graphs

• useful for calculating mean rate of reaction

• rate at specific point

• time a reaction until it reaches completion

  → reactants will decrease, as concentration of products increase