Group 2 Chemistry Concepts

• Redox reactions involve both reduction and oxidation.

Oxidation

• Adding oxygen indicates oxidation has occurred.
• Removing hydrogen indicates oxidation has occurred.
• Losing electrons (increase in the oxidation number) indicates oxidation has occurred.

Reduction

• Removing oxygen indicates reduction has occurred.
• Adding hydrogen indicates reduction has occurred.
• Gaining electrons (decrease in the oxidation number) indicates reduction has occurred.

Oxidizing Agent (Oxidant)

• A substance that oxidizes another substance.
• It is itself reduced in the process.
• Examples include acidified KMnO4, acidified K2Cr2O7, chlorine, oxygen, and hydrogen peroxide.

Reducing Agent (Reductant)

• A substance that reduces another substance.
• It is itself oxidized in the process.
• Examples include hydrogen, carbon, carbon monoxide, potassium iodide, and reactive metals.

Reduction of Hematite (Iron Oxide to Iron)

• Fe2O3 is reduced as oxygen is removed, hence it is the oxidizing agent in the reaction Fe2O3 + 3CO → 2Fe + 3CO2.
• CO is oxidized as oxygen is added, hence it is the reducing agent in the reaction Fe2O3 + 3CO → 2Fe + 3CO2.

Reaction Between Chlorine and Hydrogen Sulfide

• Cl2 is reduced as hydrogen is added, making it the oxidizing agent in the reaction Cl2 + H2S → 2HCl + S.
• H2S is oxidized as hydrogen is removed, making it the reducing agent in the reaction Cl2 + H2S → 2HCl + S.

Combustion of Magnesium

• Mg is oxidized as it loses electrons, acting as the reducing agent in the reaction 2 Mg + O2 → 2 MgO.
• Mg → Mg2+ + 2e−
• O2 is reduced as it gains electrons, acting as the oxidizing agent in the reaction 2 Mg + O2 → 2 MgO.
• O2 + 2e− → O2−

Oxidation Number

• A number assigned to each atom or ion in a compound.
• This number indicates its degree of oxidation.
• Oxidation numbers can be positive, negative, or zero.
• The oxidation number of any uncombined element is zero.
• For a monoatomic ion, the oxidation number is the same as the charge of the ion.
• Cu2+ has an oxidation number of +2.
• Cr3+ has an oxidation number of +3.
• Cl− has an oxidation number of -1.
• Group I elements in a compound always have an oxidation number of +1.
• Group II elements in a compound always have an oxidation number of +2.
• Fluorine in a compound always has an oxidation number of -1.
• Oxygen in most compounds has an oxidation number of -2.
• Hydrogen in most compounds has an oxidation number of +1.
• The more electronegative element in a substance is assigned a negative oxidation number.
• Transition metals have variable oxidation numbers in compounds, denoted by roman numerals.

Sum of Oxidation Numbers

• The sum of oxidation numbers in a neutral compound is zero.
• The sum of oxidation numbers in an ion is equal to the charge of the ion.
• Metals generally form positive ions by losing electrons, which increases the oxidation number.
• Non-metals generally form negative ions by gaining electrons, which decreases the oxidation number.

Deducing Oxidation Numbers

• Deduce the oxidation number of chlorine in NaCl:
  • Na has an oxidation number of +1, so chlorine (Cl) must be -1 to balance the charge.
• Deduce the oxidation number of chlorine in NaClO:
  • Na has an oxidation number of +1 and O has -2, so Chlorine (Cl) must be +1 to balance the charge.
• Deduce the oxidation number of chlorine in NaClO3:
  • Na is +1, O is -2, therefore Chlorine (Cl) is: +1 + X + (3 x -2) = 0 and X = +5
• Deduce the oxidation number of nitrogen in NH3:
  • H is +1, therefore N is more electronegative and N is: X + (3 x +1) = 0 and X = -3
• Deduce the oxidation number of nitrogen in NO2-:
  • O is -2, and O is more electronegative, therefore N is: X + (2 x -2) = -1 and X = +3
• Deduce the oxidation number of nitrogen in NO3-:
• O is -2, and O is more electronegative, therefore N is: X + (3 x -2) = -1 and X = +5

Using Oxidation Numbers to Classify Reactions

• In the reaction Cl2 + H2S → 2HCl + S, chlorine's oxidation number decreases from 0 to -1, thus chlorine is reduced.
• In the reaction Cl2 + H2S → 2HCl + S, sulfur's oxidation number increases from -2 to 0, thus sulfur is oxidized.
• This is a redox reaction since both reduction and oxidation occur.
• In the reaction NaOH + HCl → NaCl + H2O, there is no change in oxidation numbers.
• Therefore this is not a redox reaction.
• In the reaction 2 NaOH + Cl2 → NaCl + NaClO + H2O, the oxidation number of chlorine increases from 0 to +1 in NaClO and decreases from 0 to -1 in NaCl.
• This indicates a redox reaction and disproportionation reaction.

Disproportionation Reaction

• Involves simultaneous oxidation and reduction of an element in a single species.

Reaction of Copper (I) Oxide and Sulfuric Acid

• In the reaction Cu2O + H2SO4 → CuSO4 + Cu + H2O, one Cu+ ion in copper (I) oxide loses electrons to form Cu2+, while another Cu+ ion gains electrons to form Cu.
• Copper is involved in both oxidation and reduction processes.

Constructing Equations Using Oxidation Numbers

• Half-ionic equations show what happens to the electrons in reactions where atoms, molecules, or ions are either gaining or losing them.
• They are called half equations because each half equation represents only half of what is happening in a reaction that involves electron transfer.
• To construct a full ionic equation for a reaction, add the two half-ionic equations together in a way that cancels out the electrons.

Reaction Between Zinc and Copper (II) Sulfate

• The half equations are:
  • Zn(s) → Zn2+(aq) + 2e−
  • Cu2+(aq) + 2e− → Cu(s)
• Full ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Reaction Between Chlorine Gas and Iron (II) Chloride

• The half equations are:
  • Fe2+(aq) → Fe3+(aq) + e−
  • Cl2(g) + 2e− → 2Cl−(aq)
• Full ionic equation: 2Fe2+(aq) + Cl2(g) → 2Fe3+(aq) + 2Cl−(aq)
• Write the unbalanced equation to show the species that undergoes reduction or oxidation
• Balance the atoms being oxidized or reduced, if necessary.
• Add H2O to balance O atoms.
• Add H+ to balance H atoms.
• Add e− to balance the charge.

Example: Dichromate Ions and Iron (II) Ions in Acidic Conditions

• The unbalanced equations are:
  • Fe2+ → Fe3+ + e-
  • Cr2O72- → 2Cr3+
• The balanced equations are:
  • 6 Fe2+ → 6 Fe3+ + 6e-
  • Cr2O72- + 14 H+ + 6e- → 2Cr3+ + 7 H2O
• Full ionic equation: Cr2O72- + 14 H+ + 6 Fe2+ → 6 Fe3+ + 2Cr3+ + 7 H2O

Balancing Equations Using Oxidation Numbers

• Identify the elements whose oxidation numbers change.
• Changes from +4 to +6, this is a '2 electron' change occurs for S.
• Ag changes from +1 to 0, this is a ‘I electron' change.
• With SO2(g) + ....H2O(l) + 2Ag+(aq) → SO42-(aq) + ....H+(aq) + 2 Ag(s) the ratio of SO2 to Ag+ is 1 : 2.
• With SO2(g) + 2 H2O(l) + 2Ag+(aq) → SO42-(aq) + 4H+(aq) + 2 Ag(s) balance the H and O atoms to receive = SO2(g) Group I & II Elements

Trends in Ionization Energies

• Ionization energy depends on the electronic structure of an element and affects its physical and chemical properties
• Energy needed for ionization is used to overcome the electrostatic attraction between the electron being removed and the protons in the nucleus.

First Ionization Energy

• It is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions (M(g) → M+(g) + e−).

Second Ionization Energy

• It is the energy required to form one mole of gaseous +l ions to form one mole of gaseous 2+ ions (M+(g) → M2+(g) + e−).
• Effective nuclear charge, atomic radius and shielding affect ionization.
• 1st ionization energy decreases as we go down both Groups 1 and 2.
• Because even though the nuclear charge increases down the group, atomic radius increases so the outer electron is further from the nucleus and is attracted less strongly.
• The number of filled inner shells increases, so the repulsion between filled inner shells and the electron being removed increases.

Trends in Reactivity of Group 1 and II Elements

• Down groups 1 and 2 reactivity increases.
• Can be explained by decrease in removal energy electrons from the outermost shell

Reaction with Oxygen (Group I Metals)

• Group I metals are only shiny when they are kept away from air and tarnish when exposed forming dull dark layer on surface
• When oxygen is in the air, the general reaction is
• 4M(s) + O2(g)--> 2M2O(s)
• The products are oxides containing M+ and O2- ions
• Lithium forms lithium oxide only
• Sodium forms both sodium oxide and sodium peroxide
• This is because sodium has larger ionic radius than lithium

Reaction with oxygen (Group 2 metals)

• Group 2 metals burn when heated in air forming metal oxide
• Even with heating there is slow reaction with O2 forming metal oxide coating which prevents metal from reaction
• products are oxides containing M+2 and O -2 ions
• General equation from Group 2 metals in air 2M(s) + O2(g) --> 2MO(s)

Reaction with Chlorine (Group 1 & Group 2 Metals)

• Group 1 & 2 metals combine with chlorine when chlorine gas when heated
• With reactions getting more vigorous down both groups
• General equation from Group I metal and chlorine: 2M (s) + Cl2 (g) → 2 MCI (s)
• General equation from Group II metal and chlorine: M (s) + Cl2 (g) → MCI2 (s)

Reaction with Water (Group 1 Metals)

• Group 1 metals react vigorously w/ water
• General equation of Group 1 metals: 2M (s) + 2H2O () → 2 MOH (aq) + H2(g)
• The products are hydrogen gas and M+ and OH- ions
• When lithium is added to water, reaction forms colorless LiOH and fizzs forming H2 gas
• Reactions get more vigorous as the group goes down

Reaction With Water (Group 2 Metals)

• Group 2 metals react w/ water with Mg reacting slowly and with Ca, Sr & Ba reacting more vigorously
• General Equation: M(s) + 2H20(l) --> M(OH)2 (aq) + H2(g)
• Products are hydroxides and Hydrogen-gas
• With Magnesium reacting differently when heated w/ steam to rapidly form Magnesium Oxide and Hydrogen gas, from formula : Mg (s) + H2) (g) --> MgO (s) + H2 (g) gas
• The Hydrogen formed is burned as it's leaves tube to prevent highly flammable gas inside lab

Reactions of Oxides and Hydroxides of Group 1 & 2

• Reactions with water:
• The oxides of group I and II are basic and forms Alkaline solutions
• Equations of the groups can be simplified because there is no change of M+ ions, 02-+ h2o ---> 2OH
• Reactions with Acids:
• The Group 1 & 2 both oxides & hydroxides forms salt & water known as Neutralization Reaction
• The example equations are
• Na2O+ H2SO4---> Na2SO4 + H2O
• MgO + 2HNO3 ----> Mg (NO3)+ H2O
• KOH + HCI ---> KCI + H2O

Making Use Of Lime

• Agriculture = Lime increases crop yield in agriculture by controlling the PH level acidity used in soil Formula: Ca (OH)2 + 2 HNO ---> Ca(NO3)2 + 2H20
• Testing For Carbon Dioxide: Formula: , CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(1)
• Milk of Magnesia: Used to treat indigestion/antacid from formula Mg(OH)2 + 2HCI ---> MgCl2 2H2O

Trends in Sollubility of Group 2 Hydroxides and Sulfates

• For hydroxides, as you move down the group, sollubility increases as well as alkaline
• As you move group Sulfates sollubility increases as you move down the group
• Barium meal contains barium, and is used in X-rays on the digestion track to see all parts of the inside body as the soft tissues clear

Trends on Thermal Stability

• Thermal stability is a measure of heat & compound decomposition
• Group 1 & 2 nitrates And carbonates help increase size of ions as more power is used to thermally decompose them.

Thermal decomposition

• Carbonates: Thermal Decomposition is the breakdown of a compound into two or more substances through heating.
• Group 1: Usually white solids who decompose with metal & carbon dioxide
• Group 2: Almost the same process as group 1 who decompose when heated with oxide & dioxide
• For Carbon Dioxide: used when testing in liquid or turns milky.

Thermal decomposition (Nitrites)

• Solids are white, and usually decompose with metallic and nitrite oxide with oxygen/
• Tested with nitrogen dioxide or oxygen.

Flame Test

• Metal Ions produces a colour with the flame test to identify colour of substance in the periodic table
• The heat causes the Electrons, the more amount colours the more electrons emit.
• Magnesium does not produces color as it's does.

Test for Ions:

• With Ammonium release of sodium with hydroxide
• With Carbonate with hydrochloric and dilute acid form water & carbon

Titration Experiments

• Titration is used to determine the concentration of a solution
• Example is where planning to find HCI with NaOH & apparatus in lab
• Steps: 1 Rincing till condensation 2 Use of pipette with hydrochloric 3 Use to transfer 4 Add of drops 5 Rince 6 Fill burette 7 Record burette 8 Add slowly and swirl 9 Each titration must be recorded down on graph.

Important points

1 Burette must be washed for previous uses 2 Burette washed with water only 3 Flask with solid tip is filled

Measurement Uncertainty

• Uncertainty is for measurement purposes, and digital balances are use is measurement, as well as calculating a percentage Group Seven (Halogens)

Group 7 (Halogens)

• Otherwise known as Halogens, meaning to produce Salt
• There is a variation of color as it changes states in the periodic table
• They usually from Halides when together

Melting and Boiling points

• Halogens exist as diatomic & molecules, and the intermolecular forces depend on these factors
• Their molecules don't fluctuate the same
• The more molecules exists, the strength of the inter-forces increases

The trends

• As the electro-negativity increases down the group, so does shield and radius while decreasing.
• As more reacting occurs, the higher number more oxidize agent

Reactions of group I + 2 metals

• Has the same reaction of metals & forms halo-lines in ionic compounds

Halide

• Is like displace reactions as usually a reactive is usually displacement by another
• If cyclohexane, then usually an upper organic area is made.

Disproportion reactions

• The element in these reactions undergoes both oxidation and reduction simultaneously.
• The types of processes are chlorine + water, cold alkali, & Hot kali
• When added, it leads to reactions which leads to salts and oxide

Reactions of Halides

• Halides reaction and oxidation is the main process
• And there’s hydrogen sulfide
• The reaction of a Halide and concentrated Sulfur is mainly a reduction by itself

Key aspects with sodium cholride

• Sulfuric acid behaves like it’sonly an acid, chloride is with it is very powered reduced. Bromidee
• There is also two parts, where it’s sulphuric and reduced

Reactions

• To note is Sodium Iodide when it is reduced & strong reduced.
• There will be a mix and the results may vary.

Group Seven (Reactions)

• The aq halides with a silver nitrate have a nitrate solution in the aq.
• The solution must contain a lot of solids.
• The aq ammonium can test these when used.
• the colors usually test
• chloride- white solid and dilute
• bromide- cream color and soluble ammonia Iodide - yellow and insolubel

Explore the chemistry of Group 2 elements, covering topics such as magnesium hydroxide's safety, solubility trends, oxidation numbers, barium sulfate's use in X-rays, and thermal stability factors.