Unit 1: States of Matter
1.1 Solids, Liquids and Gases
1. State the distinguishing properties of a solid in terms of its shape and volume.
A solid has a fixed shape and a fixed volume.
2. How does the shape and volume of a liquid differ from a solid?
A liquid has a fixed volume but no fixed shape; it takes the shape of the container it is in.
3. Describe the properties of a gas regarding shape and volume.
A gas has no fixed shape and no fixed volume.
4. Describe the separation of particles in a solid compared to a gas.
Particles in a solid are very close together (touching), whereas in a gas, they are far apart.
5. What is the arrangement of particles in a solid?
They are in a regular arrangement (or lattice).
6. Describe the arrangement of particles in a liquid.
They are in a random (or irregular) arrangement.
7. What is the specific motion of particles in a solid?
They vibrate about fixed positions.
8. Describe the motion of particles in a liquid.
They move around and slide over each other.
9. Describe the motion of particles in a gas.
They move quickly and randomly in all directions.
10. Define the change of state from solid to liquid and liquid to solid.
Solid to liquid is melting; liquid to solid is freezing.
11. Define the changes of state from liquid to gas.
This can occur through boiling or evaporating.
12. What is the name for the process of a gas turning into a liquid?
Condensing.
13. What is the effect of increasing temperature on the volume of a gas?
It causes the volume to increase.
14. What is the effect of increasing pressure on the volume of a gas?
It causes the volume to decrease.
15. Explain melting in terms of kinetic particle theory.
As a solid is heated, particles gain kinetic energy and vibrate faster until they have enough energy to overcome the attractive forces holding them in a regular arrangement.
16. Explain boiling in terms of kinetic particle theory.
As a liquid is heated, particles gain enough energy to completely overcome the forces of attraction between them, allowing them to move far apart as a gas.
17. On a heating curve, what is happening to the temperature during a change of state?
The temperature remains constant (shown as a horizontal line) because the energy is being used to break the forces between particles.
18. Explain, using kinetic particle theory, why increasing temperature increases gas volume.
Heating gives particles more kinetic energy, so they move faster and collide with the container walls more frequently and with more force, pushing them outward.
19. Explain, using kinetic particle theory, why increasing pressure reduces gas volume.
Increased pressure forces the particles closer together, reducing the amount of empty space between them.
1.2 Diffusion
20. What is diffusion?
The movement of particles from a region of higher concentration to a region of lower concentration.
21. Explain diffusion in terms of kinetic particle theory.
It occurs because particles are in constant random motion, which causes them to spread out and mix.
22. Why does diffusion occur faster in gases than in liquids?
Gas particles have more kinetic energy and move much faster than liquid particles, with more space between them to move through.
23. How does relative molecular mass affect the rate of diffusion of a gas?
Gases with a lower relative molecular mass diffuse faster than gases with a higher relative molecular mass.
24. Explain why a "lighter" gas diffuses faster than a "heavier" gas.
At any given temperature, particles with lower mass move faster on average than heavier particles.
2.1 Elements, Compounds and Mixtures
1. What is the definition of an element?
A substance that consists of only one type of atom.
2. How is a compound defined in chemistry?
A substance formed from two or more different elements that are chemically combined.
3. What is a mixture?
Two or more substances that are physically mixed together but not chemically combined.
4. Describe a primary difference between a compound and a mixture regarding how they are joined.
In a compound, the elements are chemically bonded in fixed proportions, whereas in a mixture, the components are physically combined and can be separated by physical means.
2.2 Atomic Structure and the Periodic Table
5. Describe the general structure of an atom.
An atom consists of a central nucleus containing protons and neutrons, surrounded by electrons arranged in shells.
6. What particles are found in the nucleus of an atom?
Protons and neutrons.
7. Where are electrons located in an atom?
In shells surrounding the central nucleus.
8. State the relative charge of a proton, a neutron, and an electron.
A proton has a charge of +1, a neutron has a charge of 0 (neutral), and an electron has a charge of -1.
9. State the relative mass of a proton, a neutron, and an electron.
A proton has a mass of 1, a neutron has a mass of 1, and an electron has a mass that is negligible (approx. 1/1840).
10. Define the term "proton number" (also known as atomic number).
The number of protons in the nucleus of an atom.
11. Define the term "mass number" (also known as nucleon number).
The total number of protons and neutrons in the nucleus of an atom.
12. What is the electronic configuration of an atom with a proton number of 13?
2, 8, 3.
13. What is unique about the electronic configuration of Group VIII noble gases?
They have a full outer electron shell.
14. For elements in Groups I to VII, what is the relationship between the group number and the electronic configuration?
The number of outer shell electrons is equal to the group number.
15. How is the period number of an element related to its electron shells?
The number of occupied electron shells is equal to the period number.
2.3 Isotopes
16. Define "isotopes."
Different atoms of the same element that have the same number of protons but different numbers of neutrons.
17. In the symbol 126C, what does the lower number (6) represent?
The proton number (atomic number).
18. In the symbol 3517Cl-, what does the negative superscript indicate?
It indicates that the species is an anion (a negative ion).
19. Why do isotopes of the same element have the same chemical properties?
Because they have the same number of electrons and therefore the same electronic configuration.
20. How is the relative atomic mass of an element calculated?
It is calculated from the relative masses and the percentage abundances of all its isotopes.
2.4 Ions and Ionic Bonds
21. How is a positive ion (cation) formed?
By an atom losing one or more electrons.
22. How is a negative ion (anion) formed?
By an atom gaining one or more electrons.
23. Define an "ionic bond."
A strong electrostatic attraction between oppositely charged ions.
24. Between which types of elements do ionic bonds typically form?
Between metallic and non-metallic elements.
25. Describe the typical melting and boiling points of ionic compounds.
They have high melting points and high boiling points.
26. Describe the electrical conductivity of ionic compounds in different states.
They have good electrical conductivity when aqueous or molten, but poor electrical conductivity when solid.
27. Describe the "giant lattice structure" of ionic compounds.
A regular arrangement of alternating positive and negative ions.
28. Explain why ionic compounds have high melting and boiling points in terms of their structure.
They have a giant lattice structure with strong electrostatic forces of attraction between oppositely charged ions, which require significant thermal energy to overcome.
29. Explain why ionic compounds conduct electricity when molten but not when solid.
In the solid state, ions are fixed in a lattice and cannot move; when molten or in solution, the ions are free to move and carry the electric charge.
30. What diagram is used to show the formation of ionic bonds?
Dot-and-cross diagrams.
2.5 Simple Molecules and Covalent Bonds
31. How is a covalent bond formed?
When a pair of electrons is shared between two atoms, leading to noble gas electronic configurations.
32. Which simple molecules must you be able to describe using dot-and-cross diagrams?
H2, Cl2, H2O, CH4, NH3, HCl, CH3OH, C2H4, O2, CO2, and N2.
33. Describe the melting and boiling point characteristics of simple molecular compounds.
They have low melting points and boiling points.
34. Describe the electrical conductivity of simple molecular compounds.
They have poor electrical conductivity.
35. Explain the low melting and boiling points of simple molecular compounds in terms of bonding.
They have weak intermolecular forces between the molecules which require very little energy to overcome.
36. Explain why simple molecular compounds are poor conductors of electricity.
They do not have delocalised electrons or mobile ions to carry a charge.
2.6 Giant Covalent Structures
37. What are the two giant covalent structures of carbon described in the syllabus?
Graphite and diamond.
38. Relate the structure of graphite to its use as a lubricant.
Graphite is arranged in layers that can slide over each other due to weak forces between them, making it slippery.
39. Why is graphite used as an electrode?
It has delocalised electrons between its layers that allow it to conduct electricity.
40. Why is diamond used in cutting tools?
It is extremely hard because every carbon atom is bonded to four others in a rigid giant covalent lattice.
41. Describe the giant covalent structure of silicon(IV) oxide (SiO2).
It is a giant lattice where each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms.
42. Why does silicon(IV) oxide share similar properties with diamond?
Both have giant covalent structures with strong covalent bonds throughout the entire lattice, resulting in high hardness and high melting points.
2.7 Metallic Bonding
43. Define "metallic bonding."
The electrostatic attraction between positive ions in a giant metallic lattice and a ‘sea’ of delocalised electrons.
44. Explain the good electrical conductivity of metals.
The delocalised electrons are free to move throughout the giant metallic lattice and carry an electric charge.
45. Explain why metals are malleable and ductile.
The layers of positive ions can slide over each other without breaking the metallic bond because the delocalised electrons maintain the attraction.