OCR A-level chemistry - module 2

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Last updated 9:24 AM on 2/5/26
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106 Terms

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nuclear symbols

top/bigger no = RELATIVE ATOMIC MASS (-AN=neutrons)

bottom/smaller no = ATOMIC NUMBER (total proton/electrons)

<p>top/bigger no = RELATIVE ATOMIC MASS (-AN=neutrons)</p><p>bottom/smaller no = ATOMIC NUMBER (total proton/electrons)</p>
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ions

- diff number of protons and electrons

- negative ion = more electrons

- positive ion = less electrons

e.g F = 9 protons

so F^- = 10 electrons

overall charge = 9 - 10 = -1

Mg = 12 protons

so Mg^2+ = 10 electrons

overall charge 12 - 10 = +2

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define 'isotopes'

isotopes of an element are atoms with the same number of protons but different numbers of neutrons

(same atomic number, different mass number)

<p>isotopes of an element are atoms with the same number of protons but different numbers of neutrons</p><p>(same atomic number, different mass number)</p>
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why do Isotopes have similar chemical properties?

they have the same configuration of electrons

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why do isotopes have slightly different physical properties? e.g densities and rates of diffusion

physical properties depend more on the mass of the atom

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define 'relative atomic mass'

Ar is the weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12

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define 'relative isotopic mass'

the mass of an atom of an isotope of an element compared to 1/12th of the mass of an atom of carbon-12

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how do you calculate relative atomic mass Ar?

sum of (isotope abundance x isotope mass number) / 100 (sum of abundances of all the isotopes)

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3 applications of a mass spectrometer

1. drug analysis

2. forensic analysis

3. carbon dating

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how to read a mass spectrum

y-axis = abundance of ions (%), peak of each element = relative isotopic abundance

x-axis = relative isotopic mass

- sum of multiplying each mass by % abundance

- divide by 100 (should be sum of abundances)

= RELATIVE ATOMIC MASS

e.g (35 x 75) + (37 x 25) / 100

= 35.5 (RAM of chlorine)

<p>y-axis = abundance of ions (%), peak of each element = relative isotopic abundance</p><p>x-axis = relative isotopic mass</p><p>- sum of multiplying each mass by % abundance</p><p>- divide by 100 (should be sum of abundances)</p><p>= RELATIVE ATOMIC MASS</p><p>e.g (35 x 75) + (37 x 25) / 100</p><p>= 35.5 (RAM of chlorine)</p>
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what if the relative abundances are not as a %

- sum of multiplying each mass by its abundance

- divide by the SUM of relative abundances

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how to calculate % abundance in a table

(relative abundance of isotope / sum of all relative abundance of isotopes) x 100

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one mole of a substance is

6.02 × 10^23 particles (Avogadro's constant)

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formula for finding the moles from the number of atoms or molecules

number of moles = number of particles/6.02 x 10^23

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what is molar mass?

the mass of one mole of a substance, has the same numerical value as the Mr, just add g mol-1

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formula incl moles, mass and Mr

moles = mass/Mr

<p>moles = mass/Mr</p>
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how do you find the number of molecules, particles or atoms in the sample from the moles?

moles x (6.02x10^23) = number of molecules

then x by how many atoms for atoms

e.g H2S = 3 atoms

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m3 to dm3

dm3 to cm3

m3 to cm3

1. x1000

2. x1000

3. x1000,000

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what is room temperature and pressure?

298K and 100kPa

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formula for number of moles at room temp and pressure

moles = volume in dm3 / 24

or vol in cm3 / 24 000

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what is the ideal gas equation?

pV = nRT

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pV = nRT units

p = pressure (Pa)

V = volume (m3)

n = moles

R = the gas constant (8.314 JK-1 mol-1)

T = temperature (K)

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degrees celcius to kelvin

+273

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kPa to Pa

x1000

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cm3 to m3

divide by 1,000,000

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dm3 to m3

divide by 1000

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what is the relationship between pressure and volume?

pressure is inversely proportional to volume P ∝ 1/V

as P decreases V increases, and as P increases V decreases

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what is concentration and its units?

how many moles/grams of something are dissolved per 1dm3 of solution

mol dm-3 (when w moles)

or g dm-3 (when w mass)

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formula icl moles, concentration and vol in dm3

moles = concentration mol dm-3 x vol dm3

<p>moles = concentration mol dm-3 x vol dm3</p>
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formula icl mass, concentration and vol in dm3

mass = concentration g dm-3 x vol dm3

<p>mass = concentration g dm-3 x vol dm3</p>
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difference between empirical and molecular formula e.g butane

molecular formulas tell you how many atoms of each element actually are in a compound, and empirical formulas tell you the simplest or most reduced ratio of elements in a compound.

e.g butane

mf = C4H10

ef = C2H5, so ratio of carbon to hydrogen atoms is 2:5

<p>molecular formulas tell you how many atoms of each element actually are in a compound, and empirical formulas tell you the simplest or most reduced ratio of elements in a compound.</p><p>e.g butane</p><p>mf = C4H10</p><p>ef = C2H5, so ratio of carbon to hydrogen atoms is 2:5</p>
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how do you calculate the molecular formula of a compound when you know the empirical formula and molecular mass?

1. find empirical mass (Mr)

2. divide molecular mass (sometimes given if diff) by empirical mass = how many empirical units in the molecule

3. molecular formula = empirical formula X that number

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how do you calculate empirical formulas from experimental data (mass) e.g hydrocarbon from CO2 and H2O

1. use n=g/Mr to find how many moles of each product has been made

2. use moles to figure out how many moles of each atom u started with

e.g for hydrocarbon

CO2 = 0.10 moles so 0.10 moles of C

H20 = 0.10 moles so

2 moles of H = 0.10 x 2 = 0.20 moles

3. write this as a ratio

C:H = 0.10:0.20

4. divide both sides by the smallest number (0.10) = 1:2

5. so empirical formula of this hydrocarbon is CH2

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the mass of the products =

the mass of the reactants

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how do you calculate empirical formulas from percentage compositions

1. assume u have 100 g of the compound, change all %s into g

2. use n=g/Mr to find the moles of each element

3. divide each of the moles by the smallest number of moles

4. this tells you the ratio and empirical formula

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how do you calculate the molecular formula when you ONLY know the molar mass and mass

1. find the empirical formula using the steps from before (using mass)

2. find the MR of the empirical formula e.g C2H6O

(12x2)+(6)+16 = 46 g

3. divide the molar mass of the alcohol (what ur finding) by this

e.g 92 g mol-1 / 46 = 2

4. multiply empirical formula by this to get the molecular formula

e.g C2H6O x 2 = C4H12O2

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formula of ionic compounds e.g potassium sulfate

balance charges

e.g K group one = K+ = +1

SO42- = -2

so for every sulfate ion, you need 2 K ions to = 0

K2SO4

<p>balance charges</p><p>e.g K group one = K+ = +1</p><p>SO42- = -2</p><p>so for every sulfate ion, you need 2 K ions to = 0</p><p>K2SO4</p>
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how do you work out ionic equations

1. make sure the equation is BALANCED

2. break up aqueous compounds (ionic) into their ions with charges (swap and drop)

3. cross out spectator ions

4. check charges on each side are balanced

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how do you calculate masses of products (or reactants)

1. write out balanced equation

2. work out how many moles you have from the given mass

3. molar ratio of given : what you're working out

4. work out the moles of what you're working out

5. find the mass using n x rfm

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how do you calculate gas volumes

1. work out moles using n = m/mr

2. use molar ratio

3. number of moles x 24 = vol in dm3

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how do u determine the charge on ions

their group on the periodic table e.g Na group 1 so Na+

S group 6 so S2-

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nitrate ion

knowt flashcard image
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carbonate ion

knowt flashcard image
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sulfate ion

knowt flashcard image
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hydroxide ion

knowt flashcard image
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ammonium ion

knowt flashcard image
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zinc ion

knowt flashcard image
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silver ion

knowt flashcard image
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phosphate ion

knowt flashcard image
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what is theoretical yield

the mass of the product that should be formed in a chemical reaction assuming no chemicals are 'lost' in the process

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how do you calculate theoretical yield

1. work out how many moles of limiting reactant you have using n=g/r

2. ratio this to work out how many moles of the product u expect (usually same)

3. calculate mass of product using g=nxr

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how do you calculate % yield

actual yield/theoretical yield X 100

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what is atom economy

a measure of the proportion of reactant atoms that become part of the desired product in the balanced chemical equation

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how do you calculate % atom economy

RFM of desired product / RFM of all products X 100 (make sure equation is balanced)

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define acids

proton donors

release H+ ions when dissolved in water

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define bases

proton acceptors

releases OH- ions when dissolved in water

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define alkalis

soluble bases

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examples of strong acids

hydrochloric acid HCl

sulfuric acis H2SO4

nitric acid HNO3-

phosphoric acid H3PO4

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examples of weak acids

ethanoic acid C2H5OH

carbonic acid H2CO3

citric acid

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examples of bases

most oxides, all hydroxides and all carbonates AND ammonia

sodium hydroxide NaOH

ammonia NH3

potassium hydroxide KOH

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example of an alkali

sodium carbonate

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the salts

-chloride

-sulfate

-nitrate

-citrate

-phospate

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metal + acid

salt + hydrogen

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metal oxide + acid

salt + water

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metal hydroxide + acid

salt + water

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metal carbonate + acid

salt + carbon dioxide + water

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ammonia + acid

salt + water

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what is water of crystallisation

the water contained in a crystal lattice

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what is a hydrated salt?

a solid salt containing water of crystallisation

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what is an anhydrous salt?

a salt which does not contain water of crystallisation

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one mole of a particular hydrated salt always has the same number of ...

moles of water of crystallisation, it's formula shows how many

e.g: CuSO4 . 5 H2O

hydrated copper sulfur has 5 moles of water for every mole of the salt

(the dot shows they care not joined by a covalent bond)

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how to find the formula of hydrated salts from the mass of the salt when hydrated and anhydrous

1. first find the mass of water lost by taking the mass of the anhydrous salt away from the mass of the hydrated

2. find the number of moles of water lost using that mass and n=g/r

3. find the number of moles of anhydrous salt that is produced using given mass then n=g/r

4. work out the ratio of moles of anhydrous salt to moles of water

5. scale up or down so the ratio is in the form 1 : n then round up and boom!

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what if water is given as %

assume 100g of hydrated salt and convert all %s into masses (g)

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why do we perform titrations

to find out exactly how much acid is needed to neutralise a quantity of alkali

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how do we perform a titration

- first, do a rough titration to get a rough idea of where the end point is

- to do this take an initial reading to see how much alkali is in the burette to begin with

- add the acid to the alkali, giving the flask a regular swirl

- stop when the indicator shows a permanent colour change, END POINT

- record final reading from the burette

- now do an accurate titration

- run the acid in to within 2cm3 of the end point, then add an acid drop wise using the burette

- stop when the indicator shows a permanent colour change, END POINT

- work out the volume of acid used to neutralise the alkali by subtracting the initial reading from the final reading

- this volume is the TITRE

- repeat the titration 3 times

- readings should be concordant (within 0.1cm3 of eachother)

- wash the conical flask between each titration

- then calculate a mean, ignoring anomalous results

<p>- first, do a rough titration to get a rough idea of where the end point is</p><p>- to do this take an initial reading to see how much alkali is in the burette to begin with</p><p>- add the acid to the alkali, giving the flask a regular swirl</p><p>- stop when the indicator shows a permanent colour change, END POINT</p><p>- record final reading from the burette</p><p>- now do an accurate titration</p><p>- run the acid in to within 2cm3 of the end point, then add an acid drop wise using the burette</p><p>- stop when the indicator shows a permanent colour change, END POINT</p><p>- work out the volume of acid used to neutralise the alkali by subtracting the initial reading from the final reading</p><p>- this volume is the TITRE</p><p>- repeat the titration 3 times</p><p>- readings should be concordant (within 0.1cm3 of eachother)</p><p>- wash the conical flask between each titration</p><p>- then calculate a mean, ignoring anomalous results</p>
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why do we wash the conical flask between each titration

to remove any acid or alkali left in it

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methyl orange

turns from YELLOW to RED when adding acid to an alkali

<p>turns from YELLOW to RED when adding acid to an alkali</p>
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phenolphthalein

turns from PINK to COLOURLESS when adding acid to alkali

<p>turns from PINK to COLOURLESS when adding acid to alkali</p>
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when taking reading from a burette you should take readings from ....

the bottom of the meniscus

<p>the bottom of the meniscus</p>
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why are measuring volumes for satndard solutions or initial vols, using pipette and volumetric flasks more accurate rather than a burette

they only measure a fixed volume e.g 10cm3

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when filling these

be at EYE-LEVEL with the GRADUATION LINE and fill exactly up to it

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what is a standard solution

a solution that has a precisely known concentration

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how to prepare a standard solution

1. using a precise balance, carefully weigh out the required mass of solid onto a watch glass

2. transfer this solid to a clean, dry beaker and use some water to wash any bits of solid from the glass watch into the beaker

3. add water to the beaker to completely dissolve to completely dissolve the solid and use a glass rod to stir the solution to help the solid dissolve

4. once the solid has dissolved, transfer the solution into a volumetric flask. (use a volumetric flask that's the same size as the vol of solution you want to make up) rinse the beaker and glass rod with water, transferring this water into the volumetric flask

5. use water to fill the volumetric flask up to the graduation line. use a pipette to add the final few drops to make sure you don't add too much water and overshoot the gline

6. insert the stopper and invert several times to mix thoroughly

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what equation do you use to calculate concentrations from standard solutions

moles = conc x volume in cm3(/100)

vol needs to be in DM3

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how to calculate how much distilled water is needed to dilute to make further standard solutions of different concs

- divide the conc of the standard solution you want by the conc of the standard solution you have

- multiply by the volume of new standard solution you want = vol of concentrated solution to use

- subtract this volume from the total volume of dilute standard solution you desire = amount of distilled water to use

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know how to do titration calculations.

use table and balance

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what are oxidation numbers/states

of an element tells you the total number of electrons it has donated or accepted to form an ion or part of a compound

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oxidation number of ELEMENTAL substances

e.g Mg, H2, P4

ZERO

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oxidation number of SIMPLE monatomic IONS

e.g NA+, Mg2+

its CHARGE

Na+ = +1

Mg2+ = +2

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oxidation number NEUTRAL COMPOUNDS

e.g MgCl2

ZERO

Mg = +2

Cl = -1 (x 2) = -2

-2 + 2 = 0

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oxidation number of MOLECULAR IONS

e.g SO4^2- ion

SUM of all the elements o numbers which is equal to the IONS OVERALL CHARGE

O = -2 (x 4) = -8

S = +6 because overall charge (-2)

so -8 + 6 = -2

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the oxidation number of oxygen is always __ EXCEPT...

-2

in peroxides (O2^2-), where it is -2

molecular oxygen (O2), where it is 0

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systematic names (year 13?)

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oxidation

gain of oxygen

loss of electrons

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reduction

gain of electrons

loss of oxygen

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what is a redox reaction

a chemical reaction in which both oxidation and reduction happen simultaneously, there is a transfer of electrons

<p>a chemical reaction in which both oxidation and reduction happen simultaneously, there is a transfer of electrons</p>
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what is an oxidising agent

accepts (gains) electrons and gets REDUCED

<p>accepts (gains) electrons and gets REDUCED</p>
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what is a reducing agent

donates (loses) electrons and gets OXIDISED

<p>donates (loses) electrons and gets OXIDISED</p>
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as an electron is lost (oxidised)

the oxidation number for an atom will INCREASE by 1 for every e- lost

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as an electron is gained (reduced)

the oxidation number will DECREASE by 1 for every e- gained