Chemistry Flashcards
RPA -Making salts
Made by the neutralization of an acid from a base.
Steps:
Take a measured amount of the acid using a Measuring cylinder Measure 40cm.
Pour into Beaker, Add a temperature source Bedson Burner and Gauze To speed up the reaction
Leave on the gentle heat for 30 seconds or so.
Base is any material that will neutralize an acid and make a salt.
Add Copper oxide (Base) Adding an excess to ensure all acid has reacted, To tell if its an excess add to the beaker and stir until their is reactants at the bottom of the beaker.
filter the excess out using filter paper.
Solution should be a blue at this point.
- Evaporate the solution out, add to evaporating basin, Add the heat over a beaker of water to avoid too much temperature being applied. Heat until all water has evaporated.
Scrap from evaporating dish onto paper
RPA -Rates of reaction
(Sodium Thiosulphate)
Between Sodium thiosulphate and hydrochloric acid, Will create an illsoluble precipitate of Sulphur.
Time how long it takes for the colour change until cross is no longer visable
50 cm of Sodium Thiosulphate and 10 Hydrochloric acid]
vary the concentration of the Sodium Thiosulphate to see if it has an effect on the time. Using distilled water.
Measure 50 cm into measuring cylinder Pour into the beaker
Measure 10 of the acid
Reaction begins when they are first mixed, A swirl to mix.
After a few seconds the percipate changes the colour, Stop the time when the cross is no longer visible
Keep the total volume at 50cm with different concentrations of Sodium Thiosulphate diluted with water.
Less concentration means slower rate of reaction. Measuring the time taken each time.
Take a mean average plotting a graph of the concentration against the time . Creates a curve
Variables
- Independent - Concentration of Sodium Thiosulphate
- Dependent - Time taken to no longer see the cross.
- Control - Temperature, amount of hydrochloric acid.
RPA - Neutralization
Equipment List
* 25 cm3 volumetric pipette and pipette filler
* burette
* small funnel
* clamp stand
* 250 cm3 conical flask
* white tile
* dilute sulfuric acid of unknown concentration
* 0.1 mol/dm3 sodium hydroxide solution
* methyl orange indicator.
Method
1. Use the pipette and pipette filler to put exactly 25cm3 sodium
hydroxide solution into the conical flask. Your teacher will show you
how to do this. Stand the flask on a white tile.
2. Clamp the burette vertically in the clamp stand about halfway up its
length. There should be just enough room underneath for the
conical flask and tile.
3. Close the burette tap. Use the small funnel to carefully fill the burette with dilute sulfuric acid to the
0cm3
line. You should do this at a low level so that you are not pouring acid from above head height.
For example put the clamp stand temporarily on a lab stool or the floor.
4. Put 5–10 drops of methyl orange indicator into the conical flask. Swirl to mix and place under the
burette with the tile.
5. Carefully open the tap so that sulfuric acid flows into the flask at a drop by drop rate. Constantly swirl
the flask when adding the acid. Look for a colour change from yellow to red in the indicator.
6. There will be signs that the colour change is close to being permanent. When this happens use the tap
to slow the drops down. You need be able to shut the tap immediately after a single drop of acid
causes the colour to become permanently red.
7. Read the burette scale carefully and record the volume of acid you added.
8. Repeat steps 1‒7 twice more and record the results in a table.
9. Calculate the mean value for the volume of acid needed to neutralise 25 cm3 of the sodium hydroxide
solution. Record this value in the final space in the table. Use your mean result to calculate the
concentration of the acid in mol/dm3 and g/dm3 using the following calculation steps.
Variables
**Independent ** - Concentration of Sodium Hydroxide
Dependent - Amount needed to dilute the solution
Control - Temperature
RPA - Electrolysis
Equipment List
* copper(II) chloride solution
* copper(II) sulfate solution
* sodium chloride solution
* sodium sulfate solution
* 100 cm3 beaker
* petri dish lid
* two carbon rod electrodes
* two crocodile / 4 mm plug leads
* low voltage power supply
* blue litmus paper
* tweezers
Method
1. Pour copper (II) chloride solution into the beaker to about 50cm3
.
2. Add the lid and insert carbon rods through the holes. The rods
must not touch each other. Attach crocodile leads to the rods.
Connect the rods to the D.C. terminals of a low voltage power
supply.
3. Select 4 V on the power supply and switch on.
4. Look at both electrodes. Is there bubbling at neither, one or both
electrodes?
5. Use tweezers to hold a piece of blue litmus paper in the solution
next to the positive electrode (the one connected to the red
terminal). You will need to lift the lid temporarily to do this. Write
your observations in the first blank row of the table below. What
is this element?
6. After no more than five minutes, switch off the power supply. Examine the negative electrode (the one
connected to the black terminal). Is there evidence of a metal coating on it? What could it be? Record
your results in the table.
7. Clean the equipment carefully. Repeat steps 1‒6 using solutions of copper (II) sulfate, sodium chloride
and sodium sulfate.
Variables
**Independent ** - Solutions of copper sulfate, sodium chloride and sodium sulfate
Dependent - Metal or no metal coating
Control - Temperature
RPA - Temperature Change
Equipment List
* 2 M dilute hydrochloric acid
* 2 M sodium hydroxide solution
* expanded polystyrene cup and lid
* 250 cm3 beaker
* 10 cm3 measuring cylinder
* 50 cm3 measuring cylinder
* thermometer.
Method
1. Use the 50 cm3 measuring cylinder to put 30 cm3 dilute hydrochloric acid into the polystyrene cup.
2. Stand the cup inside the beaker. This will make it more stable.
3. Use the thermometer to measure the temperature of the acid. Record it in a table.
4. Put 5 cm3 sodium hydroxide solution into the 10 cm3 measuring cylinder.
5. Pour the sodium hydroxide into the cup. Fit the lid and gently stir the solution with the thermometer
through the hole. When the reading on the thermometer stops changing, write the temperature in the
table.
6. Repeat steps 4 and 5 to add further 5 cm3 amounts of sodium hydroxide to the cup. A total of 40 cm3
needs to be added. The last few additions should produce a temperature fall rather than a rise.
7. Repeat steps 1–6 and record the results in the table.
8. Calculate the mean maximum temperature reached for each of the sodium hydroxide volumes. Record
these means in the table.
9. Plot a graph with ‘Mean maximum temperature in oC’ on the y-axis and ‘Total volume of sodium
hydroxide added in cm3
’ on the x-axis.
10. Draw two straight lines of best fit one through the points which are increasing and one through the
points which are decreasing. Ensure the two lines are extended so they cross each other.
11. Use the graph to estimate how much sodium hydroxide solution was needed to neutralise 25cm3 dilute
hydrochloric acid.
Variables
**Independent ** - Amount of sodium Hydroxide
Dependent - Tempreture reached
Control - External Temperature sources, Thickness of polytene
RPA - Rates of Reaction (Gas)
Equipment List
* Safety goggles
* conical flask (100 cm3
)
* single-holed rubber bung and delivery tube to fit conical flask
* trough or plastic washing-up bowl
* two measuring cylinders (100 cm3
)
* clamp stand, boss and clamp
* stop clock
* graph paper
* magnesium ribbon cut into 3 cm lengths
* dilute hydrochloric acid, (2.0 M, and 1.0
M).
Method
1. Measure 50 cm3 of 2.0 M hydrochloric acid using one of the measuring cylinders. Pour the acid into the
100 cm3 conical flask.
2. Set up the apparatus as shown in the diagram. Half fill the trough or bowl with water.
3. Fill the other measuring cylinder with water. Make sure it stays filled with water when you turn it upside
down.
4. When you are ready, add a 3 cm strip of magnesium ribbon to the flask, put the bung back into the
flask as quickly as you can, and start the stopclock.
5. Record the volume of hydrogen gas given off at suitable intervals (eg 10 seconds) in a table. Continue
timing until no more gas appears to be given off.
6. Repeat steps 1-5 using 1.0 M hydrochloric acid.
7. Plot a graph with ‘Volume of gas produced in cm3
(for 2.0 M hydrochloric acid)’ on the y-axis and ‘Time
in seconds’ on the x-axis.
8. Draw a smooth curved line of best fit
9. Plot a curve for 1.0 M hydrochloric acid on the same graph.
10. Use this graph to compare the rates of reaction of 1.0 M and 2.0 M hydrochloric acid with magnesium
11. Compare your results with the data collected in Activity 1.
Variables
**Independent ** - 1.0 M hydrochloric acid
Dependent - Volume of hydrogen per 10 seconds
Control - Temperature
RPA - Chromatography
Equipment List
* 250 cm3 beaker
* Glass rod
* A rectangle of chromatography paper
* Four known food colourings labelled A-D
* An unknown mixture of food colourings labelled U
* Glass capillary tubes.
Method
1. Use a ruler to draw a horizontal pencil line 2 cm from a short edge of the chromatography paper. Mark
five pencil spots at equal intervals across the line. Keep at least 1 cm away from each end.
2. Use a glass capillary tube to put a small spot of each of the known colourings on four of the pencil
spots. Then use the glass capillary tube to put a small spot of the unknown mixture on the 5th pencil
spot. Try to make sure each spot is no more than 5 mm in diameter. Label each spot in pencil.
3. Pour water into the beaker to a depth of no more than 1 cm.
4. Tape the edge of the chromatography paper to the glass rod. The paper needs to be taped at the end
furthest from the spots.
Rest the rod on the top
edge of the beaker. The
bottom edge of the
paper should dip into the
water. Ensure that the
pencil line is above the
water surface and that
the sides of the paper
do not touch the
beaker wall.
5. Wait for the water
solvent to travel at least
three quarters of the
way up the paper. Do
not disturb the beaker
during this time.
Carefully remove the
paper. Draw another
pencil line on the dry
part of the paper as
close to the wet edge as
possible.
6. Hang the paper up to dry thoroughly.
7. Measure the distance in mm between the two pencil lines. This is the distance travelled by the water
solvent.
8. Measure and record the same distance for each food colouring in a table.
9. For each of the four known colours, measure the distance in mm from the bottom line to the centre of
each spot. Write each measurement in the table.
10. Use the following equation to calculate the Rf value for each of the known colours. Write the calculated
values in the table.
Rf =
distance moved by substance
distance moved by solvent
10. Match the spots in mixture U with those from A–D. Use the colour and distance travelled to help you.
Which of colourings A–D are in mixture U? Are there any other colourings in mixture U which do not
match A–D?
Variables
**Independent ** - Types of Ink
Dependent - Distance travelled
Control - Water amount/ concentration of ink
RPA - Identify ions
Equipment List
* Bunsen burner
* test tubes and test tube rack
* teat pipette
* nichrome wire mounted in handle
* limewater
* 0.4 M dilute hydrochloric acid
* 0.1 M barium chloride solution
* 0.4 M dilute nitric acid
* 0.05 M silver nitrate solution
* Known labelled solutions: chlorides of lithium, sodium, potassium, calcium and copper
* known labelled solutions: sodium salts containing carbonate, sulfate, chloride, bromide and iodide
* salt solution labelled ‘unknown’.
Methods
Flame Tests
1. Pour around 1 cm depth of each of the labelled chloride solutions into five test tubes in the rack.
2. Dip the nichrome wire into the first solution. Then hold the tip of the wire in a blue Bunsen burner
flame.
3. Record your observation in a table.
4. Clean the wire carefully.
5. Repeat steps 2‒4 for each of the other four solutions.
6. Empty and clean the test tubes.
Carbonate test
1. Pour around 1 cm depth of each of the labelled sodium solutions into five test tubes in the rack.
2. Place 2 cm depth of limewater in a sixth test tube.
3. Add 1 cm depth of dilute hydrochloric acid to each sodium salt in turn. Only if you see bubbles,
quickly use the teat pipette to transfer the gas produced to the limewater. Your teacher may show you
how to do this. You will need to take several pipettes of the gas to get a change in the limewater.
4. Record your results in a separate table.
5. Empty and clean the test tubes.
Sulfate test
1. Pour around 1 cm depth of each of the labelled sodium solutions into five test tubes in the rack.
2. Add a few drops of dilute hydrochloric acid to each solution. Then add 1 cm depth of barium
chloride solution.
3. Record your observations in the table.
4. Empty and clean the test tubes.
Halide test
1. Pour around 1 cm depth of each of the labelled sodium solutions into five test tubes in the rack.
2. Add a few drops of dilute nitric acid to each solution. Then add 1 cm depth of silver nitrate solution.
3. Record your observations in the table.
Unknown compound
1. Repeat the flame, carbonate, sulfate and halide tests on the unknown salt solution.
2. Use your results to identify the positive metal ion in the unknown compound, as well as the negative
non-metal ion.
Organic Chemistry : Contents
Chemistry which includes the usage of Carbon in the making of compounds
- Mainly bonded to other carbon atoms or to hydrogen atoms
Organic Chemistry : Hydrocarbons
Any compound that is formed from carbon and Hydrogen Only.
Must be strictly Hydrogen and Carbon only.
Alkanes
Follow format of Cn H2n+2
- Methane
CH4 - Ethane
C2H6 - Propane
C3H8 - Butane
C4H10
Alkanes are Saturated
Organic Chemistry : Properties’ of Alkanes
All alkanes have similar properties.
- Chain length is the amount of atoms attached
Boiling Point increases with longer chain lengths
First Four are gas at room temperature, Then Liquid then Solid
Shorter Alkanes Are more volatile - Evaporates easier
** Longer Chains** Means more viscous which means thick and sticky like
Shorter chain are more flammable than longer making them easier to burn
Organic Chemistry : Combustion Reaction
Main usage of Hydrocarbons is Fuel
- Complete combustion will happen when there is enough Fuel and Oxygen
- Combustion is an Exothermic Reaction
- Carbon and Hydrogen are being oxidized
- Leads to CO2 and H2O being Produced
hydrocarbon +
oxygen
➔ carbon dioxide +
water
Organic Chemistry : Crude Oil
Crude Oil is a fossil fuel and is a mixture of different compounds.
- nearly all are Hydrocarbons
Crude oil is formed naturally from dead plants and animals mainly plankton and have been burried under earth.
- Left for millions of years
- Temperature and the pressure has lead to a release into crude oil
- As it formed soaks into the rocks and is stored into millions of years
- Can be taken out by “Fracking” the rocks
- Finite Resource
- Needs to be separated into different hydrocarbons using “Fractional Distilliation”
Organic Chemistry : Fractional Distiliation
Heats it up and separates using the fact that the compounds have different boiling points.
- Heat till most has had a gas
- Pass into a fractionating column
- Gasses will rise through the Colum reaching their condensing point.
- Longest chains will condense earliest
- Short chains will rise up the Column
- Shortest will remain as a gas for the longest and will just continue to stay as a gas.
Shortest Chain Are MOST economical
Organic Chemistry : Cracking and Alkenes
Cracking is a Thermal Decomposition Reaction
- Breaking down Molecules by heating
Catalytic Cracking
- Heat the long hydrocarbons, till vaporised
- Get Hot Powdered Aluminum Oxide
- When the Hydrocarbon vapor and Powder touch the chain will split into two
Steam Cracking
- Heat to a very very high temperature.
- Will lead to them also cracking
Formula
Long chain Alkane -> Shorter Chain Alkane + Alkene
- Must remain the same with a balanced equation
Alkene Formula
Cn H2n
Alkenes are Unsaturated
Organic Chemistry : Bromine Water Test
Adding Alkenes to Bromine water will make the water go from Orange to Colorless
Homologous
Homologous series is compounds with similar properties
Alkenes and Alkanes are examples
Energy Changes : Endothermic vs Exothermic
Different amounts of energy in different bonds in molecules
Total Reactant energy compared to the total product energy
Conservation of Energy so if energy is lost it is to the surroundings.
If energy is lost to surroundings its usually via thermal stores.
Exothermic - Transferring energy to the surroundings (Larger Reactants Energy)
Endothermic
- Transferring thermal energy into the surroundings (Larger Product energy)
Energy Changes : Activation energy
Activation Energy is
- The minimum amount of energy the reactant particles need in order to collide with each other and react
Energy Changes : Energy Change
Bond energy
is the amount of energy needed to break one mole of a particular Covalent Bond
Endothermic Process to break bonds
Exothermic Process is to form bonds.
To know if its endo or exo then you need to compare
Equation
- Energy required to break bonds - Energy released by forming bonds
Negative value means exothermic
During chemical reactions bonds are formed
Molecules, Compounds and Mixtures
Molecules
- 2 or more atoms held together by chemical bonds
e.g oxygen (O2)
Compound
- Two or more different elements held together chemically
E.g Carbon Dioxide
CO2
Mixtures
Two or more substances not chemically bonded
Evolution of the Atmosphere
Currently
80% nitrogen 20% oxygen and a spread of gasses less than one
Only a theory, not as much evidence as possible so considered a theory.
first Billion years
- Really dry
- a lot of Carbon dioxide
- water vapor
- nitrogen
A very dry atmosphere with intense volcanic activity
water vapor’s condensation formed seas and oceans
2.7 Billion years ago
algae and green plants appeared due to photosynthesis
- Carbon dioxide declined
- oxygen has started to build up
Sedimentary rocks trapped the carbon.
Global warming and Climate change
Atmosphere works as a greenhouse to maintain a level temperature
As the suns energy in short wavelength radiation hits the atmosphere and some is absorbed by earth
Most is reflected in longer wavelengths
Some make it back out to space others reflect of gas in the atmosphere
Heat energy stays close to earth for a longer time than if it was just to be reflected into space.
Atmosphere made up of Carbon dioxide and water vapor known as the Greenhouse gasses
due to more gasses the greenhouse effect has got stronger
Deforestation is a large factor for CO2 over production
Methane from farm animals and waste
Consequences are unknown
Main consequence will be climate change which will be how long term weather patterns are effected
rare weather events more common and more severe
Biodiversity will struggle
Carbon footprint
Carbon footprint
- The total amount of Co2 and other greenhouse gasses emitted over somethings entire lifecycle.
refers to products such as phones, events or services.
Used to measure the cost effectiveness and ability to measure if they were sustainable
Very very hard to track
A rough calculation is what is needed to better us off at least
sensible plans to reduce the emissions
Usage of renewable methods will minimize the energy waste
entire economy based on fossil fuels
international agreements are hard to not weaken the countries economy and ideas of the betrayal.
Carbon capture technology allows us to trap carbon and store it deep underground. A downside is that it is very expensive.
Air Pollution
Pollutants that are as a result of not completely burning harmful gasses
Complete combustion
Hydrocarbon + oxygen ➔ water + carbon dioxide
lead to particulates(soot) and carbon monoxide
Particulate
- respititory problems
- Smog
- Global Dimming
Carbon monoxide
- Diffuses into blood meaning less oxygen is transported around the body
- Oxygen deprivation
- Colorless and oader less
Nitrogen oxides and sulfur dioxides cause acid rain
Formation of Ions (Metal + Non Metal)
Ion
- is a charged particle, can be a single or group amount of atoms
Ions are formed by gaining or losing an electron(s)
do this to get a complete outer shell
Periodic table, group number is the amount of atoms in the outer shell
To become stable they must lose or gain electrons
Requires energy to change amount of electrons so is easier for the ones who only require 1 or 2
Group 1,2,6,7 all easiest to become an Ion as they only need 1 or 2
group 3,4,5,6 hardest
Na -> Na+ + e-
e = Electron with a negative charge
Ionic Bonding
Ions
- Formed when atoms lose or gain electrons
Transfer electrons to stabiles both atoms e.g. gain 1 and lose one
Two ions have opposite charges so will be attracted as opposites attract
Ionic Bond is formed
Using a dot and Cross Diagram
- Use to know which electrons belong to which atom
- Box with square brackets and + or - to show charge
Ionic Compounds
Metals normally form ions which have a positive charge
Non metals usually have a negative charge
From in a Lattice structure
Properties
- Very high melting and boiling points
- can conduct electricity only when melted and dissolved in water
- Strong ionic Bonds
- Requires a lot of energy to break apart
- When solid everything is fixed
- when they can move about depends on the charge allows the conductivity of electricity
Formula
Chemical symbol together
if one has 2 electrons would be xy2
Covalent Bonding (Non metals)
Covalent bonding
Bonds when two Non Metal elements with the same requirements e.g. need electrons
They can share electrons so will share 1 each, both get 1 extra
Dot and Cross diagram and a space to overlap
Cl-Cl
Shares what they need so become bonded together
Types of Covalent Structures
Simple Molecular Substances
- Small easy formations
Giant Covalent structures
- Forms things like diamonds
Covalent bonds are very strong
Simple
Just need to break the intermolecular forces so wont require a lot of energy
More intermolecular forces require more energy to break them a part
Do not conduct electricity and have no electric charge
Giant
- Huge numbers of non metal atoms
- regular repeating lattices
- Diamond, graphite and silicon dioxide
- very strong
- high melting and boiling points
- dont conduct electricity
Diamond and Graphite
Only made of Carbon Giant Covalent bond
Allotrope
- Different structural forms of the same element in the same physical state
Regular lattice, making them very strong
Diamond
- Each carbon atom is covalently bonded to 4 other carbon atoms
- recognisable
- very strong
- doesn’t conduct electricity
graphite
- Bonded to only 3 carbon atoms
- Arranged into hexogen’s to arrange several layers
- Each layer held together weakly
- relatively soft
- high melting point
- Layers can slide
- One spare electron which becomes delocalized
- Can conduct electricity and heat due to it
- One layer known as graphene
Graphene and Fullerenes
Allotropes of Carbon
Fullerenes and graphene
Each layer of graphite is known as graphene very strong bond.
Graphene can conduct electricity
Graphene is completely natural
Graphene tubes and spheres are known as fullerenes
fullerenes
- act as spheres around other molecules like a cage to get specific dosages to specific areas
- Can also be used as an industrial catalyst
- Large surface area
- Nanotubes for Nano technology
- Very long and thin
first ever known as
Buckminsterfullerene
- Hallow sphere
- 60 carbon atoms
C60
- Nanotechnology
Metallic Bonding ( Two Metals)
Bonding between metal atoms
Giant structure of atoms arranged in a regular pattern
Share with all the other atoms in the metal
All become positive atoms and the delocalized electrons run in between
strong electrostatic attraction between delocalized electrons and the positive metal atoms
high melting and boiling point
good conductor of electricity and heat
metals are malleable
due to the different layers being able to slide over each other
Alloys
- Contain 2 or more different elements picked with different sized atoms
- disrupts the metals regular pattern
- the different layers cannot slide over each other as easily
Examples: Steel made from Iron, carbon and sometimes chromium manganese and vanadium
Relative formula Mass
Mass numbers = the larger number of the symbol
Relative atomic mass (Ar)
- Average mass of all the isotopes of that element
Relative formula mass
- Add together the relative atomic masses of all the atoms in that compounds molecular formula.
Example
H2SO4
4 O = 4x16 = 64
2 H = 2x1 = 2
1 S = 1x32 = 32
32+2+64= 98
Mr=98
Calculate the percentage mass of a particular mass in a compound.
Example
H2SO4 Percentage mass of S
(Ar of S x Number of S) / Mr of the H2SO4 x 100
32 x 1 / 98 =0.327
0.327 x 100 = 32.7%
Moles and Mass
Formula
Number of moles in a sample = Mass of that element / Mr
Mole
- Unit to measure the amount of a chemical we have
1 Mole of any substance that contains 6.02 x 10^23
Known as Avogadro’s constant
Will be the same as relative formula mass in grams equal to the Mole of a substance
Conservation of Mass
Conservation of mass is that mass is always conserved
No atoms are created or destroyed, only the bonds change
Amount of atoms doesn’t change based on the equitation
If scales used mass should stay the same
Mass may appear to change if it dispersed into surroundings
chemical change
- Rearrangements of the atoms in the reactants to form the products
Limiting Reactants
Limiting reactants
- It all reacts and limits how much product can be formed once all of the reactant has been used up
Refers to the reactant that is fully used up known as the limiting reactant
Whatever other reactant is known as being in excess
Limiting factor changes the amount of product
Concentration Calculation
Measured in grams per decimeter cubed
= 1000cm
Concentration = Mass/Volume
Mass = G
Volume = dm^3
Acids and Bases
Ph Measure of how acidic or alkaline a solution is
Ph Scale
0- Red Highly acidic
1- Red Highly Acidic
2- Red Highly Acidic
3- Orange/red Strong Acidic
4 - Orange Acidic
5 - Orange/Yellow Less acidic
6 - Yellow Weak Acidic
7 - Green Neutral
8 - Turquoise Weak Alkaline
9 - Blue Less Alkaline
10 - Purple/Blue Alkaline
11 - Purple Strong Alkaline
12 - Purple Highly Alkaline
13 - Purple Highly Alkaline
14 - Violet Highly Alkaline
Pure water is 7 (Neutral)
pH measured using an indicator, chemical dyes based on pH level.
Wide range indicators
Universal Indicator makes the colour of the normal pH scale
pH probe
- Electronically measure the pH
- Much more perceive and there is no room for misinterpretation of the colour
Acid
- Any substance that forms an aqueous solution with a pH less than 7
Bases
- Any substances with a pH greater than 7
Alkalis- A base that dissolves in water to form a substance of pH 8 or higher
Neutralization reaction
Acid and Base combine to form a neutral substance making the pH roughly 7 (Neutral)
Acids
- Hydrochloric acid
HCl
- sulfuric acid
H2SO4
- Nitric Acid
HNO3
Base
- Sodium Hydroxide
NaOH
- Calcium Carbonate
CaCo3
Strong Acids and Weak Acids
Acid
- Substance that forms aqueous solutions with a pH less than 7
- Ionize with the aqueous solution
Strong Acid
- All of the acids particles will Disassociate
Example
Hydrochloric acid
Hcl
Nitric Acid
HNO3
Sulfuric acid
Weak Acid
- Dont fully ionize
- dissociation of weak acids is reversible reaction
- Equilibrium Shift to the left
Examples
- Citric acid
- Ethanoic Acid
- Carbonic Acid
**Concentration **
- How much acid there is in a certain volume.
Strength
- How much an acid dissociates
Higher Hydrogen the pH gets lower
each pH is equivalent to a factor of 10
Neutralization Reactions
Acid reacting with, Metal oxides or metal hydroxides or metal carbonates all class as a neutralization reaction
Metal Oxides/ Hydroxides
whenever reacted with acid would produce salt and water
Metal Oxides/hydroxides + Acid -> Salt + water
Take negative from the acid and positive from the base
Metal Carbonates
Produces salt, water and Carbon dioxide
Acid + Metal Carbonates -> Salt, H20+ CO2
Soluble salts
Sodium Chloride
NaCl
Potassium Sulfate
K2SO4
Calcium Nitrate
Ca(NO3)2
Insoluble base
- Place acid dilute in a beaker and gently heat over budson burner
- Add insoluble base a little bit at a time
- When it stops disappearing the base has become an excess
- Acid has been neutralized
- Filter any excess
Should give solution
To get crystals
- Water bath at a gentle temp
- When first crystals stop heating and more will form
- Filter crystals and dab dry leaving in a warm place
Separating Metals from Metal Oxides
Oxidation
- gaining oxygen
Reduction
- Lose of oxygen
Oil Rig Acronym
Oxidation
Iis
LLoss
Reduction
Is
Gain
Metals naturally Oxide as they are reactive
Unreactive metals like gold are too unreactive to react with oxygen.
Cheapest method (Reduction)
- Metal Oxides reacted with carbon, forming carbon dioxide and leaving behind a pure metal
- Only works on the ones less reactive than carbon
- (Zinc, Iron and Copper) from oxides
Any more reactive use Electrolysis but is very expensive
Ores
- Metal rich compounds
Redox Reactions
Redox reactions
- Reduction and oxidation take place at the same time
Oxidation (Positive)
- Loss of Electrons
Reduction (Negative)
- Gain of electrons
Oil Rig Acronym
Oxidation
Iis
LLoss
Reduction
Is
Gain
Mg + H2 -> Mg 2+ atoms +H2
Displacement Reaction
- A more reactive metal displacing a less reactive one
Calcium + Iron Sulfate -> Calcium Sulfate + Iron
Calcium Displaces Iron
It will lose 2 electrons to bond with the negatively charged sulfate
Half Equation
Ca -> Ca+2 + 2e-
Ca - Calcium
E - Electron
Electrolysis
Equipment
- Beaker
- Electrolyte (Free to move Ions)
- Electrodes
- Wire
- Power supply like a battery
Soluble
- Dissolve it in water
Insoluble
- Melt it to get a molten liquid
- only way to ensure they can move around
Electrodes
- Solid Conductors
- Right is the positive electrode known as the anode
- Left is the negative electrode known as the cathode
- Usually inert carbon meaning it is unreactive and will not take part in the reaction
Electrolysis
- Splitting up with electricity
Negative Ions attract to Anode while positive ions attract to Cathode
Once there they will discharge and become normal atoms then pairing up forming a gas or solid
Anode - Oxidized - Positive - Negatives attracted
Cathode - Reduced - Negative - Positives attracted
Electrolysis Separating reactive metals from oxides
Works on all reactivity series however is only economical for the ones more reactive than carbon
Solid (Ions are fixed) so isn’t an electrode
- Purify raw ore
- Mix the aluminum with cryolite which lowers the melting point
- Now set up the equipment
Purify, then mix with cryolite then electrolyze
Electrolysis in Aqueous substances
Soluble
- Dissolve them in water
Insoluble
- Melt into molten
Always Present
H20 -> H+ + OH-
Rules
- Cathode will attract negative Ions
- Will only discharge one of the ions
- Ion of the least reactive element will be discharged
- Positive Anode if Halide is present than that will be discharged
- If no Halide is present it will be the Hydrogen that gets discharged
Group 7 and Group 0
Group 7 known as Halogen
Group 0 Known as Nobel gasses
Halogens
- Fluorine - Poisonous yellow gas
- Chlorine - Less reactive but poisonous
- Bromine - Red poisonous
- Iodine - Poisonous purple vapors
- Diatomic molecules (2 atoms)
Boiling point and melting point increases further down you go same as Nobel gasses
reactivity decreases further down you go as the outer most shell is further away from the nucleolus so the attractive force is weaker
Ionic Bonds with metals known as Halide
Fluoride, chloride ,bromide ,iodide
Rate of Reactions
Rate of reaction
- Refers to the speed that reactants turn into products
Rates can be very different, some within seconds other over years
To measure the rate of reactants we need to measure, how fast products are formed or how fast the reactants are used up
Rate of reaction = Quaintly of reactants used / Time taken
Rate of Reaction = Quantity of products formed / Time Taken
Grams and Cms
Factors Affecting rate of reaction
Factor
- Temperature
- Surface area
- concertation/pressure
- catalyst
Collison theory
- For particles to react they collide with each other with sufficient energy (Activation Energy)
1- Amount of energy the particles have
2- Frequency of the collisions
Higher tempreture = quicker rate of reaction
More surface area = Quicker rate of reaction
More pressure - quicker rate of reaction
Catalyst = Increases rate of reaction
Reversible reactions and Dynamic Equilibriums
Concentration of reactants and products wont change any more
- is constant doesn’t mean that they are the same
Reversible reaction can both take place at different rates
Must be in a closed system as to not lose and products or reactants
Equilibrium
- Both reactions still happening
- Canceling out each other
When there is more products the equilibrium lies to the right
when more reactants the equilibrium would lie to the left
Adding heat would move it to the right
if cooled would go back to the left
Le Chateliers Principles
Position of equilibrium during a reversible reaction
position will try to counter act the change
Negative shows it is exothermic forward reaction
Endothermic is backwards if negative
**Decreasing temperature **would move it in the exothermic direction
If more products than reactants the position will move to the right
If added pressure the position will move to side with least molecules
If added concentration then the equilibrium would shift to the opposite side
Negative Joules = Exothermic
Positive Joules - Endothermic
Purity and Formulations
Chemical Analysis
- The instruments and methods used to identify, separate and quantity results
Purity
- Something that contains only one type of compound or element
- e.g pure water (H2O)
Has specific boiling points used to identify unknown substances
Physical test
- Testing physical qualities
Chemical tests
- Using another chemicals
Impure
- Melt and boil over a range of temperature dependent on what is in the mixture
- Usually lowers the melting point but increases the boiling point
Formulations
- Mixtures that have been prepared using a specific formula
- Precise amounts of different components
- Have a particular function
- Fuels, medicines, cleaning agents, food and drink, alloys and fertilizer
Tests for gasses
Tests for
Chlorine
oxygen
hydrogen
carbon dioxide
chlorine
- Requires test tube and sample gas
- Blue litmus paper and dampen it
- **If chlorine is present will turn from blue to white **
- Wear mask or do in a fume cupboard
Oxygen Glowing splint
- Glowing splint without a flame
- If gas is present will relight the splint
Hydrogen Lighted splint in mouth of test tube
- Squeaky Pop
- Splint should be burning
- **A squeaky pop should be heard if hydrogen is present **
Carbon dioxide Bubble through limewater
- Limewater test bumbling the gas sample
- if gas is present the water will go cloudy
Life Cycle Assessment
LCA
Analyze the different stages in a products life cycle to asses the effects on the environments
- Extracting and processing the raw materials
- Manufacturing and packaging the product
3.using your product
4.disposing of product
Stage One: Extracting
- Can damage trees and wildlife to achieve the item
- or during fracing
- a lot of pollutants released
Packaging and manufacturing
- energy used
- waste products
- pollutuion
Usage of product
- How much damage it does in its life
- How long it is used for and how many times it will be used
How its disposed
- Using landfill may have chemicals or take up land for the environment
- May release pollutants
- Take into account energy used
Limitations of LCA
- Involves a lot of steps and is hard to quantify all of them
- Asses the harm of each step in comparison to others
- Companies can manipulate to look more favorable
Potable water
Potable water
- Water that is safe to drink
- Doesn’t mean it is pure water
- pH of 6.5 and 8.5
- No microorganisms
- Levels of dissolved substances must be fairly low
Fresh water
Countries with access to fresh water comes from rainwater usually
Collects as surface or ground water
Surface water
- Lakes
- Rivers
Ground water
- Aquifers
Requires treatment
- Pass it through wire mesh
- Pass it through bed of sand and gravel to remove smaller bits
- sterilize it to kill and microbes
- chlorine gas
- UV light
- Ozone
Desalination
- Extract potable water from sea water
Distillation and reverse osmosis are very expensive
distillation
- same process as simple distillation
Reverse osmosis
- Passed through a membrane only allowing the water molecules to get through
Waste water management
** 3 main sources**
- Domestic - goes to sewers
- Agricultural systems
- Industrials
All has to be treated to make it safe before its disposed off.
Remove any Organic matter and harmful microbes
How sewage treatment actually works
- Screening the sewage, removing anything large
- Let it sit in a settlement tank to undergo sedimentation
Sludge will since and go to a separate tank while the effluent (Top) will go to its own tank
effluent is under aerobic conditions
Sludge is under anaerobic conditions
3. Break down all organic matter, Relying on biological breakdown.
effluent water is now safe
Anaerobic digestion used as a fertilizer as well as a gas source
