Inorganic common questions

Reactivity of halogens down the group:

• Decreases

• Electronegativity decreases

• Less attractive to electron addition

• Weaker oxidising agents

• Greater nuclear attraction outweighs increased shielding

Boiling point of hydrogen halides down the group:

• Increases

• The van der Waals (VdW) / London dispersion forces increase.

• Due to the increasing number of electrons.

• H-F does not follow this trend because:

  • It is the most electronegative.

  • Contains hydrogen bonding.

  • Has the strongest intermolecular force (IMF).

  • More energy is needed to overcome the forces.

Thermal Stability of hydrogen halides down the group:

• Thermal stability decreases down the group.

• The larger the halogen atom, the more shielding it has.

• Bond energies decrease, and less energy is needed to break the H-X bond.

Volatility of halogens down the group:

• Decreases

• Increasing number of electrons.

• Increasing strength of van der Waals forces.

Oxidising agent strength of halogens down the group:

• Oxidising power decreases down the group.

• Ability to accept electrons decreases because the outer shell experiences more shielding and increased distance from nucleus to outer shell.

Melting points of Halogens:

• Increasing melting point down the group

• Because van der Waals forces, intermolecular forces (IMFs), and London dispersion forces increase

Reaction of HCl and HClO added to dilute and cold NaOH. One of the products acts as a bleach:

2𝑁𝑎𝑂𝐻+𝐻𝐶𝑙+𝐻𝐶𝑙𝑂→𝑁𝑎𝐶𝑙𝑂+𝑁𝑎𝐶𝑙+2𝐻2𝑂

𝑁𝑎𝑂𝐻+𝐻𝐶𝑙→𝑁𝑎𝐶𝑙+𝐻2𝑂

𝑁𝑎𝑂𝐻+𝐻𝐶𝑙𝑂→𝑁𝑎𝐶𝑙𝑂+𝐻2𝑂

Reaction of Cold vs Hot NaOH

Explain the differences in the reactions of concentrated sulfuric acid with NaI and with NaCl. Your answer should refer to the role of sulfuric acid in each reaction:

NaCl+H2SO4→NaHSO4+HCl      

2NaI+H2SO4→SO2+I2+2NaOH

1. Sulfuric acid acts as an acid with NaCl and NaB

2. Sulfuric acid acts as an oxidising agent with NaI.

3. NaI is a reducing agent.

4. Iodine (I) is more powerful than chlorine (Cl).

5. Sulfuric acid can oxidise iodine but not chlorine.

Describe and explain the relative reactivity of halogens down the group when they react with hydrogen to form HCl, HBr, and HI:

• Reactivity decreases down the group

• Atomic radius increases, increased leangth of H-X bond therefore less heat energy is required to break the bond down the group

• Electronegativity decreases down the group.

• Less attractive to electron addition.

• Weaker oxidising agents.

• Greater nuclear charge outweighs increased shielding.

What makes Magnesium Oxide suitable for the manufacture of heat-resistant bricks for furnace linings?

• High melting and boiling points

• Strong forces of attraction between oppositely charged ions

• Strong ionic bonding

When calcium metal is placed in dilute sulfuric acid, it reacts vigorously at first. After a short time, a crust of calcium sulphate forms on the calcium metal, and the reaction stops. Some of the calcium metal and dilute sulfuric acid remain unchanged. Why?

• The crust of calcium sulphate prevents hydrogen bubbles from rising.

• This stops/slows down the reaction.

• CaSO4 does not react with the excess unreacted sulfuric acid.

State and explain the pattern of change of oxidation number which occurs to both chlorine and the different Period 3 elements when they react together:

• Oxidation number of the elements increases by one.

• Due to the addition of an electron to the valence shell.

• Oxidation number of chlorine stays the same.

Melting point of group 2 metals:

• Decreases

• Increased metallic / cationic radius / size.

• Decreasing cation charge density.

• Decreased attraction of ions for delocalized / outer electrons.

First Ionization energy of group 2 metals:

• Decreases

• Because the distance of outer electrons from the nucleus increases.

• Shielding increases.

• decrease in attraction between nucleus and valence electrons.

• Hydroxide and Nitrates Solubility increases down the group.

• Sulphates solubility decreases down the group.

Physical properties of ceramics:

• High melting point.

• Electrical insulator.

• Hard.

• Retains strength at high pressures and temperatures due to a giant lattice structure, requiring a lot of energy to break strong bonds.

Nitrogen oxides are involved in the formation of photochemical smog through several processes:

• They react with unburnt hydrocarbons in the presence of sunlight.

• This reaction leads to the formation of peroxyacetyl nitrate (PAN), which is a key component of photochemical smog.

• PAN contributes to the formation of secondary pollutants and enhances the harmful effects of smog, including respiratory issues and environmental damage.

Explain the general increase in first ionisation energy with atomic numbers 11-18 (Na-->Ar):

• Increasing nuclear attraction

• Increasing nuclear charge

• Increasing number of protons and similar shielding

Explain the decrease in first ionisation energies between Magnesium and Aluminium:

• Aluminium has an outer electron orbital of 3p, which is at a higher energy level than Magnesium.

• Magnesium has an outer electron orbital of 3s.

• 3p is more shielded.

• Ionisation energy of Mg > ionisation energy of Al

Explain the decrease in first ionisation energies between Phosphorus and Sulphur:

• The electrons in the 3p phosphorus subshell are unpaired.

• There is greater electron pair repulsion, so more energy is required to remove that electron.

• Sulphur has a pair of electrons in 3p orbital.

• Phosphorous has half-filled stability (3p3)

Explain the trend in the first ionisation energies of the elements with atomic numbers 2, 10, 18 (Helium, Sodium, Argon):

• Decreasing ionisation energy down due to increasing nuclear attraction, increasing shielding, increase in number of shells, increase in distance of outer electrons from nucleus.

• (the distance between the nucleus and the outer electron increases 

• the shielding by complete inner shells increases

• These two factors outweigh the increased nuclear charge. So first IE decreases)

Increase in melting point from sodium to magnesium:

• More energy is needed to break metallic bonding due to increasing strength between delocalized electrons and cation as mg has greater charge; the number of delocalized electrons increases.

Why is the increase from magnesium to aluminium much smaller than from sodium to magnesium:

• Attraction for electrons is too strong to fully delocalize all 3 electrons in aluminium.

• Difference in size between Magnesium and Aluminium is less than the size of Sodium and Magnesium.

• Magnitude of increase in charge is less from 2+ to 3+ than 1+ to 2+.

Explain the trend and pattern of melting points from phosphorus to argon:

• S8 > P4 > Cl2 > Ar

• Melting points depend on the van der Waals forces.

• The greater the number of electrons, the greater the melting point.

Why does silicon have the greatest melting point:

• Has a giant covalent structure.

• Has many strong bonds.

Why does the atomic radius decrease across  the period :

• Increase in nuclear charge, same shielding; electrons are added to the same shell, stronger nuclear attraction with valence electrons and nucleus.

Physical properties of ceramics:

• High melting point.

• Electrical insulator.

• Hard.

• Retains strength at high pressures and temperatures.

Why are calcium hydroxide and calcium carbonate suitable for agriculture?

• Reduce soil acidity / increase soil pH.

• Both are basic.

In terms of electronegativity, why is the bonding in NaCl different from the bonding in SiCl4:

• Sodium is less electronegative than Silicon.

• NaCl includes the transfer of electrons.

• SiCl4 has a shared pair of electrons.

What is meant by amphoteric oxide?

• Reacts with both acid and base.

  

State the trends in acid/base behaviour of the period 3 elements:

• More acidic and less basic from Na to S.

Explain the variation in ionic radius for the elements with atomic number 14 to 17:

• Increasing proton number.

• Same number of electrons.

• Increasing attraction of nucleus from outer electrons.

Explain why the solution produced after aluminium chloride is added to water has a pH of 1-2:

• HCl dissolved in water to make acid.

Which of the group 2 metals, calcium, magnesium or radium requires the highest temperature to decompose:

• Radium.

• The thermal stability of group 2 metals increases down the group.

• The more thermally stable a metal is, the more energy is needed to decompose.

Group 2 nitrates and carbonates stability:

• More stable down the group.

• More decomposing energy is needed because the group 2 ion has a lower charge density and because the atom increases in size.

Cold water reacts slowly with a piece of magnesium to produce bubbles of H2.

Cold water reacts rapidly with burning Mg to produce H2 in an explosive mixture. Explain why the rate of reaction of cold water with burning magnesium is greater:

• Heat released from burning Mg provides more particles with energy.

• Frequency of collisions is greater.

Nitrogen molecules contain 2 atoms attracted to each other by a triple covalent bond. Describe how the triple bond forms in the N2 molecule. Refer to the orbital overlap and hybridization in your answer:

• 2  pi bonds and 1  sigma bond.

• Contains sp hybridization in each N atom (in general has sp3 hybridization).

• Sigma bond causes a head-to-head overlap.

• Pi bond causes a sideways overlap.

Outline how oxides of nitrogen are involved in the formation of photochemical smog:

• Reacts with unburnt hydrocarbons.

• Forming PAN (peroxyacetyl nitrate).

Identify a reagent and the conditions required to make NH3 from NH4Cl:

• Any group 1 hydroxide.

Construct an equation to show the behaviour of NH3 as a weak Bronsted-Lowry base when dissolved in water:

• NH3 + H2O <--> NH4+ + OH-

An excess of Cl- is added to 1cm^3 of Br2, describe what is observed:

• No visible reaction, orange colour.

• Bromine cannot oxidise chloride.

• Bromine is not a strong oxidising agent.

Explain why the second ionisation energy is higher than the first ionisation energy:

• The nuclear attraction increases.

• The 2nd electron is being removed from a smaller ion.

Describe the reaction of phosphorus oxide with water:

• Vigorous and exothermic.

What is needed to calculate the Ar of a substance?

• Percentage Abundance of each isotope and Isotope number.

Attraction types:

• Between atoms: ionic / covalent.

• Between molecules: van der Waals.

Nitrogen and Oxygen do not react at normal atm why?

• Nitrogen is nonpolar (it’s diatomic) and has a triple bond.

• Not enough activation energy for the reaction to take place.

What is meant by Bronsted-Lowry Acid?

• Proton donor.

What is meant by Bronsted-Lowry Base?

• Proton Acceptor.

Two Reasons why Ammonia is added to some soils:

• Reduces acidity of soil 

• because it is a base.

Volatile meaning:

• Easily vaporised.

Identify the species that are electrostatically attracted to one another in:

• An ionic bond: cations and anions.

• A covalent bond: shared electrons of the bond and 2 nuclei.

To identify the bonding and structure of a compound/element from its reactions and properties:

• Reacts with water = Covalent because it is hydrolyzed.

• Solid at room temp = simple molecular due to low melting point.

• Forms acidic solution = Covalent bond.

• High melting point and it's a metal = metallic bonding.

• High electrical conductivity = metallic bonding due to free delocalized sea of electrons.

• Conducts electricity when liquid or molten = ionic bonding due to free-flowing ions in that state.

Suggest why reactions may occur at different rates:

• Different activation energies.

Suggest why reacting burning magnesium with cold water is faster than cold water and a piece of Mg:

• Heat released from burning Mg provides more particles with energy.

• Frequency of collisions is greater.

Outline how oxides of nitrogen are involved in the formation of a photochemical smog:

• Reacts with unburnt hydrocarbons.

• Forming PAN (peroxyacetyl Nitrate).

The first ionisation energies of elements in the first row of the d block (Sc-Cu) are very similar. Why? 

• Electrons are added to the same last orbital / d orbital 

• increased shielding 

• similar nuclear attraction

Shielding proportions 

• Indirectly proportional to ionisation energy 

• Directly proportional to radius size

• Indirectly proportional to element electronegativity 

• Shielding is indirectly proportional to nuclear attraction of valence electron 

Write an equation to show the acid base behaviour of sulfuric acid with water

• H2SO4 (aq) + H2O(l) → H3O+ (aq) + HSO4- (aq) 

What is meant by high vapour pressure 

• gas particles in equilibrium 

• greater proportion of gas particles than liquid 

9701/21/o/n/18 idk if its in the syllabus with us 

Give the meaning of term enhanced greenhouse effect 

• Human activity creates thick layer in greenhouses gases 

Explain how CHClF2 may contribute this effect 

• Traps more heat 

• Leads to greater global warming 

Suggest another environmental problem associated with the use of CHClF2

• Ozone depletion / thinning 

Explain why more energy is required to remove the outer electrons in magnesium than in calcium 

• fewer inner shells of electrons / less shielding 

• smaller distance of the outer electrons 

• stronger attraction to the outer electrons 

Ionisation energy 

• It is the energy needed to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of + ions, also in its gaseous state

Ionisation energy in period 2 (topic 1) 

• Positive charge in nucleus increases 

• as the number of protons increases 

• increasing the attraction between the nuclei and electrons 

• atomic radius decreases across the period 

• outer electron are more attracted to the Nucleus 

• this causes the 1st ionisation energy to increase 

• shielding effect is similar 

Exceptions Explanation (these ideas are similar to Mg vs Al and S vs P in period 3)

1. Beryllium vs Boron

• The 2p (in boron) subshell has a higher energy than the 2s (in beryllium) 

• so less energy is required to remove the outer electron 

2. Nitrogen vs Oxygen 

• In the p subshell in oxygen there is a paired pair of electrons, these electrons enforce spin pair repulsion 

• so less energy is needed to remove the electron in oxygen than in nitrogen

• Despite Ammonia having 3 hydrogen bonds: only one is actually “hydrogen bonded - in terms of intermolecular forces” which is the one with the lone pair 

Explain the trends in the melting points of the oxides down group 2 

• Decreases 

• Lower forces of attraction between cations and anions 

• Larger cations and constant charge 

• decreasing charge density of cation down the group 

describe the trend in the colour of the group 17 elements down the group 

• they become darker