RX + PT Flashcards
Define oxidation (3)
- Addition of oxygen
- Loss of hydrogen
- Loss of electrons
Define reduction (3)
- Loss of oxygen
- Addition of hydrogen
- Gain of electrons
Define oxidation number
The oxidation state/charge of an element or ionic substance
Oxidation number of uncombined elements
0
Oxidation number of a monoatomic ion
The charge of the ion
Oxidation number of molecular ions
Sum of oxidation numbers is equal to overall charge
Oxidation number of a neutral compound
Overall charge is 0, sum of oxidation numbers is 0
Oxidation number of hydrogen
+1
Oxidation number of oxygen
-2
Oxidation number of halogens
-1
Oxidation number of group 1 metals
+1
Oxidation number of group 2 metals
+2
Oxidation number of oxygen in peroxides
-1
Oxidation number of hydrogen in metal hydrides (MHx)
-1
When are Roman numerals used?
For transition metals with variable oxidation numbers
Oxidation state of S block elements (Group 1 and 2)
- Lose electrons
- Oxidised
- Positive oxidation numbers
Oxidation state of P block elements
- Gain electrons
- Reduced
- Negative oxidation number
What is an oxidising agent?
- Gains electrons
- Reduced
- Increase in oxidation number
What is a reducing agent?
- Donates/loses electrons
- Oxidised
- Decrease in oxidation number
Define redox reaction
A reaction where reduction and oxidation occurs simultaneously
Define disproportionation reaction
A reaction in which an element in a single species is simultaneously oxidised and reduced
Trend in oxidation of metals
- Lose electrons
- Increase in oxidation number
- Oxidised
- Positive ions formed
Trend in oxidation of non metals
- Gain electrons
- Decrease in oxidation number
- Negative ions formed
Steps for ionic half equations and full ionic equations
- Write half equation for first element
- Balance all species excluding oxygen and hydrogen
- Balance oxygen by adding H20
- Balance hydrogen by adding H+ ions
- Balance charges using e- electrons
- Write half equation for second element
- Balance
- Ensure both equations have the same number of electrons
- Combine
- Cross off any common substances
What is the trend in ionisation energy going down groups 1 and 2?
Decrease in ionisation energy
1. More electrons going down the group
2. More quantum shells
3. Increased atomic radius
4. Distance between outermost electron and nucleus is greater
5. Weaker electrostatic attraction between nucleus and outermost electron
6. Less energy is required for the outermost electron to be removed
What is the reactivity trend going down groups 1 and 2?
Increase in reactivity
1. More electrons
2. More quantum shells
3. Larger atomic radius
4. Greater distance between outermost electron and nucleus
5. Weaker electrostatic attraction between outermost electron and nucleus
6. Less energy is required to remove the outermost electron
Reaction between group 2 metal with water
M(s) + 2H20(l) —> M(OH)2 (aq) + H2(g)
Group 2 metal + water —> Metal hydroxide + hydrogen
Reaction between group 2 metal and oxygen
2M(s) + O2(g) —> 2MO(s)
Group 2 metal + oxygen —> metal oxide
Reaction of group 2 metals (Sr and Ba) with excess oxygen
M(s) + O2(g) —> MO2
Reactions between group 2 metal and chlorine
M(s) + Cl2(g) —> MCl2 (s)
Group 2 metal + chlorine —> metal chloride
Reaction between group 1 metal and oxygen
4M + O2 —> 2M2O
Group 1 metal + oxygen —> metal oxide
Reaction between group 1 metal and excess oxygen
2M + O2 —> M2O2
Group 1 metal + excess oxygen —> metal peroxide
Reactions of group 1 metals and water
Group 1 metal + water —> metal hydroxide + H2
Reactions of group 1 metals and chlorine
Group 1 metal + chlorine —> metal chloride
Reaction between group 1/2 oxide and water
Group 1/2 oxide + water —> metal hydroxide
Group 1: M2O + H2O —> 2MOH
Group 2: MO + H2O —> M(OH)2
Reaction between group 1/2 oxides and dilute acid
Group 1/2 oxide + hydrochloric acid —> metal chloride + water
Group 1: M2O + 2HCl —> 2MCl + H2O
Group 2: MO + 2HCl —> MCl2 + H2O
Reactions between group1/2 hydroxides and dilute acid
Group 1/2 hydroxide + hydrochloric acid —> metal chloride + water
Group 1: MOH + HCl —> MCl + H20
Group 2: M(OH)2 + 2HCl —> MCl2 + 2H2O
Solubility trend of Group 2 hydroxides
Increases down the group
Solubility trend of Group 2 sulfates
Decrease down the group
Thermal stability of Group 1 and 2 compounds (carbonates and nitrates)
- Increases down the group
- Atomic radius increases
- Cations are bigger and have a smaller charge
- Lower charge density
- Lower polarising power
- Less distortion of carbonate/nitrate ions
- More stable compounds formed
Why are group 2 compounds less thermally stable then group 1 compounds?
- Group 2 cations have a higher charge density
- More polarising
- Less stable
Thermal decomposition of Group 1 carbonates
- All group 1 carbonates are thermally stable
- Decompose at higher temperatures
- Only lithium carbonate decompose
- Li2CO3(s) —> Li2O(s) + CO2(g)
Thermal decomposition of group 1 nitrates
- Group 1 nitrates decompose to form nitrite and oxygen
- 2MNO3 —> 2MNO2 + O2
- Only LiNO3 decomposes to form Li2O + NO2 + O2
Thermal decomposition of group 2 carbonates
All group 2 carbonates decompose to form oxide and carbon dioxide
MCO3 —> MO + CO2
Thermal decomposition of group 2 nitrates
All nitrates decompose to form oxide, nitrogen dioxide and oxygen
2M(NO3)2 —> 2MO + 4NO2 + O2
Test for nitrogen dioxide
- Brown gas
- Acidic pH when dissolved in water
Test for carbon dioxide
Bubble the gas through limewater, if CO2 is present limewater turns cloudy
Test for oxygen
A glowing splint relights in the presence of oxygen
Suggest how characteristic flame colours are formed during a flame test
- Energy absorbed from the flame causes electrons to become “excited”, causing them to move to a higher energy level
- Electron is unstable in this energy level
- Drops back down to original energy level
- Energy is emitted in the form of visible light wavelength which allows for a colour to be observed
Describe the method for a flame test
- Take a nichrome wire and dip it into HCl
- Heat the wire in the flame
- Take the clean ncihrome wire and dip it into HCl
- Dip into sample
- Place wire with sample in flame and observe colour change
Lithium colour in flame test
Red
Sodium colour in flame test
Orange/yellow
Potassium colour in flame test
Lilac
Rubidium colour in flame test
Red
Caesium colour in flame test
Blue
Calcium colour in flame test
Brick red
Strontium colour in flame test
Crimson
Barium colour in flame test
Green
Test for carbonate ions
- CO2 is produced
- Bubble the gas through limewater
- Limewater turns cloudy in the presence of carbon dioxide
Test for hydrogencarbonate ions
- CO2 is produced
- Bubble the gas through limewater
- If carbon dioxide is present, limewater turns cloudy
Test for sulfate ions
- Add acidified barium chloride
- A white precipitate of barium sulfate forms
Test for hydroxide ions
Hydroxide ions turn damp red litmus paper blue
Test for ammonium ions
- Add sodium hydroxide
- Gently warm
- Damp red litmus paper turns blue
Colour and state of fluorine at room temperature
Pale yellow, gas
Colour and state of chlorine at room temperature
Pale green, gas
Colour and state of bromine at room temperature
Red/brown, liquid
Colour and state of iodine at room temperature
Grey/black, solid
Colour of chlorine in aqueous solution
Pale yellow
Colour of bromine in aqueous solution
Orange/red
Colour of iodine in aqueous solution
Brown
Colour of chlorine in organic solvent
Pale yellow
Colour of bromine in organic solution
Red
Colour of iodine in organic solvent
Purple
Trend in melting point of halogens
Down the group
1. Increase in quantum shells/shielding
2. Increase in number of electrons
3. Stronger London dispersion forces
4. More energy is needed to overcome strong intermolecular forces
5. Melting point and boiling point increases
Trend in electronegativity of halogens
Down the group
1. Atomic radius increases
2. Outer electrons are further away from the nucleus
3. Incoming electron experiences more shielding from the attraction of the positive nuclear charge (nucleus)
4. More difficult to attract an electron
5. Electronegativity decreases
Trend in reactivity of halogens
Down the group
1. Atomic radius increases
2. Outer electrons are further away from the nucleus
3. Increased shielding
4. Harder to attract an electron
5. Decreased reactivity
Group 1 oxide + halogen, change in oxidation number
Group 1 element is oxidised: 0 —> +1
Halogen is reduced: 0 —> -1
Group 2 oxide + halogen, change in oxidation number
Group 2 element is oxidised: 0 —> +2
Halogen is reduced: 0 —> -2
Disproportionation reaction of chlorine with water equation
Cl2(g) + H2O(l) —> HCl(aq) + HClO(aq)
Chlorine + water —> hydrochloric acid and hydrochlorous acid
HClO(aq) +H2O(l) —> ClO- (aq) + H3O+(aq)
Hydrochlorous acid + water —> chlorite ions
What is chlorine used for?
Water treatment
Disproportionation of chlorine with cold alkalis equation
2NaOH(aq) + Cl2(g) —> NaClO(aq) + NaCl(aq) + H2O(l)
Sodium hydroxide + chlorine —> Sodium chlorate + sodium chloride + water
Oxidation of chlorine: 0 —> +1 in NaClO (sodium chlorate/bleach)
Reduction of chlorine: 0 —> -1 in NaCl (sodium chloride)
Disproportionation of chlorine with hot alkalis
3Cl2 + 6NaOH —> NaClO3 + 5NaCl + 3H2O
Oxidation of chlorine: 0 —> +5
Reduction of chlorine: 0 —> -1
Trend in reducing power of halides
Down the group
1. Ionic radius increases
2. Electrons are further away from the nucleus
3. Greater shielding effect
4. Easier to remove and electron
Reaction of potassium fluoride and sulfuric acid
KF(s) + H2SO4(l) —> KHSO4(s) + HF(g)
Steamy fumes produced
NOT REDOX
Reaction of potassium chloride and sulfuric acid
KCl(s) + H2SO4(l) —> KHSO4(s) + HCl(g)
Steamy fumes produced
NO REDOX
Reaction of potassium bromide with sulfuric acid
KBr(s)+H2SO4(l) —> KHSO4(s) +HBr
2HBr(aq) + H2SO4(l) —>Br2(g) + SO2(g) + 2H2O(l)
Change in oxidation state of sulfur and bromine
Oxidation state of S: +6 —> +4= reduction
Oxidation state of Br: -1 —> 0 = oxidation
Reaction of potassium iodide and sulfuric acid
KI + H2SO4 —> KHSO4 + HI
2HI + H2SO4 —> I2 + SO2 + 2H2O
6HI + SO2 —> 2H2S + 3I2 + 2H2O
