Chapter 2 - "Atoms, Elements and Compounds" - Long Questions
"Atoms, Elements and Compounds" - Long Questions:-
Elements, Compound and Mixtures
Q1
(a)
Define an element.
(b) Explain why an element cannot be broken down by chemical methods.
Answer:
(a) An element is a pure substance that contains only one type of atom.
(b) An element cannot be broken down by
chemical methods because it consists of only one type of atom and there are no
simpler substances chemically bonded together to separate.
(a) Define a compound.
(b) Explain how compounds differ from mixtures.
Answer:
(a) A compound is a pure substance formed when two or more different elements are chemically combined in fixed proportions.
(b) In a compound, elements are chemically bonded and can only be separated by chemical reactions. In a mixture, substances are not chemically bonded and can be separated by physical methods.
Q3
State two differences between a compound and a mixture.
Answer:
Compounds have fixed compositions, while mixtures can have variable compositions.
Compounds can only be separated chemically, while mixtures can be separated physically.
Q4
Air is described as a mixture. Explain why.
Answer:
Air is a mixture because it contains different gases that are not chemically
bonded. Each gas retains its own properties and
the gases can be separated by physical methods such as fractional distillation.
Q5
Explain why water has different properties from hydrogen and oxygen.
Answer:
Water is a compound in which hydrogen and oxygen are chemically bonded. Chemical
bonding changes the arrangement of electrons, producing new properties that are
different from those of the individual elements.
Q6
Name
a suitable method to separate:
(a) Sand and water
(b) Salt from salt solution
Answer:
(a) Filtration
(b) Evaporation or crystallization
Q7
Explain why filtration cannot be used to separate a solution.
Answer:
Filtration cannot separate a solution because the dissolved particles are too
small to be trapped by the filter paper and pass through with the solvent.
Q8
Define a pure substance.
Answer:
A pure substance contains only one element or one compound and has a fixed
composition throughout.
Q9
Explain why seawater is not a pure substance.
Answer:
Seawater contains water mixed with dissolved salts and gases. These
substances are not chemically bonded and can be separated physically.
Q10
Explain why compounds have fixed melting points but mixtures do not.
Answer:
Compounds have fixed melting points because they have a fixed composition and
regular bonding. Mixtures have variable
compositions, so they melt over a range of temperatures.
Q1
Define an atom.
Answer:
An atom is the smallest unit of an element that can take part in chemical
reactions and still retain the chemical properties of that element.
Q2
State the location, charge, and relative mass of:
(a) Proton
(b) Neutron
(c) Electron
Answer:
| Particle | Location | Charge | Relative Mass |
|---|---|---|---|
| Proton | Nucleus | +1 | 1 |
| Neutron | Nucleus | 0 | 1 |
| Electron | Shells | −1 | 1/1836 |
Q3
Explain why almost all the mass of an atom is in the nucleus.
Answer:
Protons and neutrons have much greater mass than electrons and are located in
the nucleus, while electrons have negligible mass.
Q4
Define proton number.
Answer:
The proton number is the number of protons in the nucleus of an atom.
Q5
Explain why proton number identifies an element.
Answer:
Each element has a unique number of protons. Changing
the proton number changes the element itself.
Q6
Define mass number.
Answer:
The mass number is the total number of protons and neutrons in the nucleus of
an atom.
Q7
Calculate the number of neutrons in an atom of sodium-23.
Answer:
Protons = 11
Neutrons = 23 − 11 = 12
Q8
Define isotopes.
Answer:
Isotopes are atoms of the same element with the same number of protons but
different numbers of neutrons.
Q9
Explain why isotopes have the same chemical properties.
Answer:
Chemical properties depend on electron arrangement. Isotopes
have the same number of electrons and the same electronic configuration.
Q10
Explain why isotopes have different physical properties.
Answer:
Isotopes have different masses due to different numbers of neutrons, affecting
properties such as density and rate of diffusion.
Q1
Write
the electronic configuration of:
(a) Oxygen
(b) Sodium
(c) Magnesium
Answer:
(a) Oxygen: 2,6
(b) Sodium: 2,8,1
(c) Magnesium: 2,8,2
Q2
Explain why noble gases are unreactive.
Answer:
Noble gases have full outer electron shells, making them stable and unlikely to
gain, lose, or share electrons.
Q3
Explain the link between group number and outer shell electrons.
Answer:
For Groups I–VII, the group number equals the number of electrons in the outer
shell.
Q4
Explain the link between period number and shells.
Answer:
The period number shows the number of occupied electron shells in the atom.
Q5
The
diagram shows an atom with configuration 2,8,1.
(a) Name the element.
(b) State its group and period.
(c) Predict the ion it forms.
Answer:
(a) Sodium
(b) Group I, Period 3
(c) Na⁺ formed by losing one electron
Ionic Bonding
Q1
Define an ion.
Answer:
An ion is a charged particle formed when an atom gains or loses electrons.
Q2
Explain how a sodium ion is formed.
Answer:
A sodium atom loses one outer shell electron, forming a Na⁺ ion with a full
outer shell.
Q3
Define ionic bonding.
Answer:
Ionic bonding is the electrostatic attraction between oppositely charged ions.
Q4
Explain why ionic compounds have high melting points.
Answer:
Strong electrostatic attractions between ions in a giant lattice require large
amounts of energy to overcome.
Q5
Explain why solid ionic compounds do not conduct electricity.
Answer:
The ions are fixed in position and cannot move to carry charge.
Q6
The
diagram shows a giant ionic lattice.
Explain how this structure affects melting point and conductivity.
Answer:
The lattice contains strong electrostatic attractions, giving a high melting
point. When molten or dissolved, ions are free to move and conduct electricity.
Covalent Bonding & Simple Molecules
Q1
(a) Define
a covalent bond.
(b) State the type of elements that usually form covalent bonds.
Answer:
(a) A covalent bond is a bond formed when a pair of electrons is shared between
two atoms.
(b) Covalent bonds usually form between non-metal atoms.
Q2
Explain why covalent bonding involves sharing electrons rather than transfer.
Answer:
Non-metals have similar attraction for electrons, so neither atom loses
electrons easily. Sharing allows both atoms to achieve a full outer shell and
become more stable.
Q3
Hydrogen
forms H₂.
(a) Explain why each hydrogen atom forms one covalent bond.
(b) Describe what is shared in the bond.
Answer:
(a) Hydrogen has 1 electron and needs 1 more to fill the first shell (2
electrons).
(b) One pair of electrons is shared, one electron from each hydrogen atom.
Q4
Chlorine forms Cl₂. Explain why chlorine forms a single covalent bond in Cl₂.
Answer:
Each chlorine atom has 7 outer electrons and needs 1 more to complete the outer
shell. Two chlorine atoms share one pair of electrons, so each achieves 8 outer
electrons.
Q5
Dot-and-cross
diagram question (H₂)
Draw a dot-and-cross diagram for H₂.
Answer
(text model):
H(·) shares with H(×) to form one shared pair (·×) between the atoms. Each H
now has 2 electrons in the first shell.
Q6
Dot-and-cross
diagram question (Cl₂)
Explain what a dot-and-cross diagram for Cl₂ must show.
Answer:
It must show one shared pair of electrons between the two chlorine atoms and
three lone pairs around each chlorine atom to complete each outer shell.
Q7
Hydrogen
chloride is HCl.
(a) State whether the bonding is ionic or covalent.
(b) Explain why.
Answer:
(a) Covalent.
(b) Both hydrogen and chlorine are non-metals, so they share electrons rather
than transfer them fully.
Q8
Water
is H₂O.
(a) State how many covalent bonds oxygen forms in water.
(b) Explain using outer-shell electrons.
Answer:
(a) Two covalent bonds.
(b) Oxygen has 6 outer electrons and needs 2 more to complete its outer shell,
so it shares one electron with each hydrogen atom.
Q9
Methane is CH₄. Explain why carbon forms four covalent bonds.
Answer:
Carbon has 4 outer electrons and needs 4 more to reach 8. It shares one
electron with each of four hydrogen atoms, forming four shared pairs.
Q10
Ammonia
is NH₃.
(a) State how many covalent bonds nitrogen forms.
(b) State how many lone pairs nitrogen has.
Answer:
(a) Three covalent bonds.
(b) One lone pair (because nitrogen has 5 outer electrons and uses 3 for
bonding, leaving 2 as a lone pair).
Q11
Oxygen forms O₂ with a double bond. Explain why a double bond forms.
Answer:
Each oxygen atom has 6 outer electrons and needs 2 more. Two oxygen atoms share
two pairs of electrons, so both achieve 8 outer electrons.
Q12
Nitrogen forms N₂ with a triple bond. Explain why a triple bond forms.
Answer:
Each nitrogen atom has 5 outer electrons and needs 3 more. Two nitrogen atoms
share three pairs of electrons so both reach 8 outer electrons.
Q13
Carbon
dioxide is CO₂.
(a) State the type of bonding in CO₂.
(b) State how many double bonds it contains.
Answer:
(a) Covalent bonding.
(b) Two double bonds, one between carbon and each oxygen.
Q14
Explain why simple molecular substances have low melting and boiling points.
Answer:
The forces between molecules (intermolecular forces) are weak and require
little energy to overcome. The strong covalent bonds inside molecules are not
broken when melting or boiling.
Q15
Explain why most simple molecular substances do not conduct electricity.
Answer:
They contain neutral molecules and have no mobile ions or delocalized electrons
to carry charge.
Q16
Compare conductivity of an ionic compound and a simple molecular substance when molten.
Answer:
Molten ionic compounds conduct because ions are free to move. Molten simple
molecular substances usually do not conduct because they remain as neutral
molecules.
Q17
Explain why chlorine (Cl₂) is a gas at room temperature but sodium chloride is a solid.
Answer:
Cl₂ consists of small molecules with weak intermolecular forces, so it boils
easily. Sodium chloride has a giant ionic lattice with strong attractions, so
it stays solid and has a high melting point.
Q18
A student says “covalent bonds are weak because molecular substances melt easily.” Explain why this is incorrect.
Answer:
The covalent bonds inside molecules are strong. Molecular substances melt
easily because the forces between molecules are weak, not the covalent bonds.
Q19
A substance has low melting point and does not conduct electricity in any state. Predict its structure and bonding.
Answer:
It is likely a simple molecular substance with covalent bonding and weak
intermolecular forces.
Q20
Explain, using bonding and structure, why carbon dioxide is a gas and silicon dioxide is a solid with a high melting point.
Answer:
CO₂ is made of small molecules. The covalent bonds inside each molecule are
strong, but the forces between molecules are weak, so little energy is needed
to separate molecules and CO₂ is a gas. Silicon dioxide forms a giant covalent
lattice with many strong covalent bonds throughout the structure, so a lot of
energy is needed to break bonds and it has a very high melting point and is
solid.
Q1
Define a giant covalent structure.
Answer:
A giant covalent structure is a structure in which a very large number of atoms
are joined together by strong covalent bonds in a continuous network.
Q2
State one similarity and one difference between a giant covalent structure and a simple molecular structure.
Answer:
Similarity: both have covalent bonds.
Difference: giant covalent structures have covalent bonds throughout the whole
lattice, while simple molecular substances have separate molecules with weak
forces between them.
Q3
Diamond and graphite are allotropes of carbon.
Define an allotrope.
Answer:
Allotropes are different forms of the same element in the same physical state, with different structures and different properties.
Q4
Describe the structure of diamond.
Answer:
Each carbon atom forms four covalent bonds to four other carbon atoms in a
tetrahedral arrangement, forming a rigid 3D lattice.
Q5
Explain why diamond is very hard.
Answer:
Diamond has a strong 3D network of covalent bonds. To scratch or cut it, many
strong covalent bonds must be broken, requiring a lot of force.
Q6
Explain why diamond has a very high melting point.
Answer:
To melt diamond, many strong covalent bonds throughout the lattice must be
broken. This requires a large amount of energy.
Q7
Explain why diamond does not conduct electricity.
Answer:
All four outer electrons of each carbon atom are used in covalent bonding, so
there are no delocalized electrons
to carry charge.
Q8
Describe the structure of graphite.
Answer:
Each carbon forms three covalent bonds, creating layers of hexagonal rings. One electron per carbon is delocalized. Layers are held together by weak forces.
Q9
Explain why graphite is soft and slippery.
Answer:
The layers are held together by weak forces, so the layers can slide over one
another easily.
Q10
Explain why graphite conducts electricity.
Answer:
Each carbon has one delocalized electron
that can move through the layers and carry charge.
Q11
Explain why graphite is used as a lubricant.
Answer:
Its layers slide easily, reducing friction between surfaces.
Q12
Explain why graphite is used as an electrode.
Answer:
It conducts electricity due to delocalized electrons
and is chemically stable at high temperatures.
Q13
Silicon(IV) oxide is SiO₂. Describe its structure.
Answer:
It forms a giant covalent lattice where each silicon is bonded to four oxygens
and each oxygen is bonded to two silicons,
forming a 3D network.
Q14
Explain why silicon(IV) oxide has a high melting point.
Answer:
Many strong covalent bonds extend throughout the structure. A lot of energy is
needed to break these bonds.
Q15
Explain why silicon(IV) oxide does not conduct electricity.
Answer:
It has no mobile ions and no delocalized electrons. All electrons are held in covalent bonds.
Q16
State two similarities between diamond and silicon(IV) oxide.
Answer:
Both have giant covalent lattices.
Both have high melting points and do not conduct electricity.
Q17
State one key structural difference between diamond and graphite.
Answer:
Diamond has a 3D network with each carbon bonded to 4 others; graphite has layered sheets with each carbon bonded to 3 others.
Q18
Explain why graphite is less dense than diamond.
Answer:
Graphite has a layered structure with spaces between layers, while diamond is a
tightly packed 3D lattice.
Q19
A student says “graphite has weak bonds so it is soft.” Explain why this statement is not fully correct.
Answer:
The covalent bonds within each layer are strong. Graphite is soft because the
forces between layers are weak, allowing layers to slide.
Q20
Compare diamond and graphite in terms of structure, bonding, electrical conductivity, and one use for each.
Answer:
Diamond: each carbon forms four covalent bonds in a rigid 3D lattice; no
delocalized electrons so it does not conduct electricity; very hard so used in
cutting tools.
Graphite: each carbon forms three covalent bonds in layers; one delocalized
electron per carbon so it conducts electricity; layers slide so it is soft and
used as a lubricant or electrode.
Mettalic Bonding
Q1
Define metallic bonding.
Answer:
Metallic bonding is the electrostatic attraction between a lattice of positive
metal ions and a sea of delocalized electrons.
Q2
Describe the structure of a metal.
Answer:
Metals consist of a giant lattice of positive ions surrounded by delocalized
electrons that can move through the structure.
Q3
Explain how the positive metal ions are formed.
Answer:
Metal atoms lose one or more outer-shell electrons, becoming positive ions. The
lost electrons become delocalized.
Q4
Explain why metals conduct electricity.
Answer:
Delocalized electrons are free to move. When a voltage is applied, electrons
flow and carry charge.
Q5
Explain why metals conduct heat well.
Answer:
Delocalized electrons transfer kinetic energy quickly through the lattice,
spreading thermal energy efficiently.
Q6
Define malleability.
Answer:
Malleability is the ability of a metal to be hammered or rolled into thin
sheets.
Q7
Explain why metals are malleable.
Answer:
Layers of metal ions can slide over each other, while delocalized electrons
maintain attraction and hold the structure together.
Q8
Define ductility.
Answer:
Ductility is the ability of a metal to be drawn into wires.
Q9
Explain why metals are ductile.
Answer:
As the metal is stretched, ion layers slide but metallic bonding continues due
to the electron sea, preventing the structure from breaking.
Q10
Metals are usually strong. Explain this in terms of bonding.
Answer:
Strong electrostatic attraction between positive ions and delocalized electrons
holds the lattice together firmly.
Q11
Explain why metals have high melting points (in general).
Answer:
A lot of energy is required to overcome strong metallic bonds between ions and
delocalized electrons.
Q12
Why is sodium softer and has a lower melting point than aluminum? (structure and bonding)
Answer:
Sodium forms Na⁺ and contributes one delocalized electron per atom, giving
weaker metallic bonding. Aluminum forms Al³⁺ and contributes three delocalized
electrons, producing stronger attraction and higher melting point.
Q13
Compare electrical conduction in metals and ionic compounds.
Answer:
Metals conduct as solids because electrons are delocalized and mobile. Ionic
compounds do not conduct as solids because ions are fixed, but conduct when
molten or in solution.
Q14
Compare electrical conduction in metals and simple molecular substances.
Answer:
Metals conduct due to delocalized electrons. Simple molecular substances do not
conduct because they have no mobile ions or delocalized electrons.
Q15
Explain why metals are not brittle, unlike many ionic crystals.
Answer:
In metals, electrons keep ions bonded even after layers shift. In ionic
crystals, shifting layers can bring like charges together, causing repulsion
and shattering.
Q16
A metal is an element. Explain why it can still form ions inside a lattice.
Answer:
In metallic bonding, metal atoms release outer electrons into a shared electron
sea. The atoms become positive ions, but the whole structure remains
electrically neutral due to the electrons.
Q17
Describe what happens to the delocalized electrons when a metal is bent.
Answer:
The electrons move with the ions and continue to attract the positive ions, so
the bonding remains and the metal does not break.
Q18
A student says “metallic bonding is like ionic bonding because it involves attraction.” State one similarity and one difference.
Answer:
Similarity: both involve electrostatic attraction.
Difference: ionic bonding is between oppositely charged ions, while metallic
bonding is between positive ions and delocalized electrons
Q19
Explain why copper is used for electrical wiring.
Answer:
Copper has mobile delocalized electrons, so it conducts electricity well. It is
also ductile so it can be drawn into wires
Q20
Compare ionic bonding, covalent bonding, and metallic bonding in terms of particles involved, type of structure formed, and electrical conductivity.
Answer:
Ionic bonding: attraction between positive and negative ions, forms a giant
ionic lattice; does not conduct when solid, conducts when molten or aqueous.
Covalent bonding (simple molecules): atoms share electrons to form molecules; weak
forces between molecules; does not conduct in any state because no mobile
charged particles.
Metallic bonding: positive ions in a lattice with delocalized electrons;
conducts electricity as solid and liquid because electrons are free to move.
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