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Chemistry Paper 1 > Bonding, Structure And Properties Of Matter > Flashcards

Bonding, Structure And Properties Of Matter Flashcards

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1
Q

What are the melting and Boiling Points of a Substance

A

-Melting point is the temperature at which a substance melts or freezes. Below this point, a substance is a solid.

-Boling point is the temperature at which a substance boils or condenses. Above this point, a substance is a gas.

-Between these points, a substance is a liquid.

-The more energy needed to break the bonds in a substance, the higher the melting/ boiling point.

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2
Q

Describe the Properties of Solids

A

-Cannot be compressed (no spaces between the particles).

-Cannot flow and take shape of container (particles cannot move).

-The particles are in neat rows, very close together and cannot move. The bonds are very strong.

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3
Q

Describe the Properties of Liquids

A

-Cannot be compressed (very few spaces between particles).

-Can flow and take shape of container (the particles can move).

-The particles are touching but able to move and are mixed up. The bonds are not very strong.

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4
Q

Describe the Properties of Gases

A

-Can be compressed (there are large spaces between the particles).

-Can flow and take shape of container (the particles can move).

-The particles are spread out and move very fast. There are no bonds.

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5
Q

Give Limitations of the Particle Model

A

-It assumes that particles are solid spheres but we know they are mostly empty space.

-it assumes that particles are stationary but we know particles vibrate and move around

-It assumes that there are no forces between particles but we know there are different types and strengths of attractions between particles.

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6
Q

Explain in Terms of Electrons, Why Elements React

A

-Elements react to achieve a full outer electron shell.

-Elements in group 0 (noble gases) do not react because they already have a full outer electron shell.

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7
Q

What are Ions

A

-An ion is an atom or molecule that has lost or gained an electron(s) to form a stable electronic structure.

-A positive ion is formed when a metal loses an electron from its outer shell.

-A negative ion is formed when a non-metal gains an electron.

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8
Q

What is Ionic Bonding

A

-An ionic bond is a strong electromagnetic attraction between oppositely charged ions.

-Ionic bonding is between metal and non-metal atoms.

-Metals lose an electron(s).

-Non-metals gain an electron(s).

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9
Q

What Happens during Ionic Bonding (Between Group 1 and Group 7)

A

-When a group 1 and a group 7 element react, ionic bonding takes place.

-The group 1 element loses one electron to form an ion with a charge of +1.

-The group 7 element gains one electron to form an ion with a charge of -1.

-Both ions have a full outer electron shell just like noble gases.

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10
Q

What Happens during Ionic Bonding (Between Group 2 and Group 6)

A

-When a group 2 and a group 6 element react, ionic bonding takes place.

-The group 2 element loses two electrons to form an ion with a charge of +2.

-The group 6 element gains two electrons to form an ion with a charge of -2.

-Both ions have a full outer electron shell just like noble gases.

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11
Q

Why are atoms uncharged before they react

A

-Before it reacts, an atom has the same number of protons and electrons.

-The negative electrons cancel out the positive protons, making the atoms uncharged.

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12
Q

Why does the Charge of Lithium and Fluorine change when they React

A

-When it reacts, the lithium atom loses one electron from its outer electron shell. Now it has two electrons (two negative charges) but it still has three protons (three positive charges).

-The two electrons can cancel out the charges on two of the protons. Therefore the lithium ion has one overall positive charge.

-When it reacts, the fluorine atom gains one electron into its outer energy level. Now it has ten electrons (ten negative charges) but it still has nine protons (nine positive charges).

-The nine protons can cancel out the charges on nine of the electrons. Therefore the fluorine ion has one overall negative charge

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13
Q

How does a Group 2 Element React with a Group 7 Element

A

-Each group 2 atom will lose 2 electrons from the outer energy level.

-However, each group 7 atom can only gain 1 electron (as they already have 7 electrons in their outer energy level).

-So one group 2 atom gives 1 electron each to two group 7 atoms.

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14
Q

What are the Advantages and Disadvantages of the Ball and Stick Model when Drawing Lattices

A

Advantage: The electrostatic forces of attraction (ionic bonds) are shown

Disadvantage: The sizes of each ion are not shown correctly

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15
Q

What are the Advantages and Disadvantages of the Space Filling Model Model when Drawing Lattices

A

Advantage: The sizes of each ion are shown correctly

Disadvantage: Electrostatic forces of attraction (ionic bonds) are shown

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16
Q

Describe what Happens to Particles when an Ionic Solid is Heated

A

-When we heat an ionic solid, the ions vibrate.

-As the strength of the vibration increases, there comes a point where the ions are moving with so much energy that the electrostatic forces between the ions weaken or break.

-Now the ions are free to move and the ionic compound has melted.

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17
Q

Explain why Ionic Compounds have High Melting and Boiling Points

A

-Ionic compounds have high melting and boiling points because the electrostatic attraction between positive, metal ions and negative, non-metal ions is strong.

-This takes a lot a lot of energy to weaken (melting point) or overcome (boiling point).

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18
Q

Explain why Ionic Compounds do not Conduct Electricity when Solid

A

-When solid, ionic compounds do not conduct electricity because the ions are fixed in a lattice by the strong electrostatic forces of attraction..

-This means they are not free to move and carry charge.

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19
Q

Explain why Ionic Compounds Conduct Electricity when Melted or Dissolved in Water

A

-Melting an ionic compound or dissolving it in water breaks the electrostatic forces of attraction between the ions.

-This means the ions are free to move and can carry charge.

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20
Q

Explain why Ionic Compounds are Brittle.

A

-Brittle means the compound shatters when hit.

-Ionic compounds are brittle because the layers in a lattice cannot slide over each other without repelling.

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21
Q

What is Covalent Bonding

A

-When a non-metal reacts with a non-metal, covalent bonding takes place.

-A pair of outer shell electrons is shared between two atoms.

-During covalent bonding, the atoms get a full outer shell, so become stable.

-Covalent bonding can be shown as a dot and cross diagram or a stick diagram.

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22
Q

Explain why only the Outer Electron Shell is shown when Drawing Covalent Bonds

A

Only the outer electron shells are involved in covalent bonding so we only show these on diagrams.

-Internal electron shells are already full so we don’t need to show these

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23
Q

Explain why one Atom of Oxygen must form Covalent Bonds with two Atoms of Hydrogen

A

-An oxygen atom has 6 electrons in its outer energy level.

-That means that 2 more electrons are required for oxygen to have a full outer energy level.

-Each hydrogen atom has 1 electron, so one oxygen atom has to form covalent bonds with two atoms of hydrogen.

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24
Q

Explain why one Atom of Nitrogen must form Covalent Bonds with three Atoms of Hydrogen

A

-A nitrogen atom has 5 electrons in its outer energy level.

-That means that 3 more electrons are required for nitrogen to have a full outer energy level.

-Each hydrogen atom has 1 electron, so one nitrogen atom has to form covalent bonds with three atoms of hydrogen.

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25
Q

Explain why one Atom of Carbon must form Covalent Bonds with four Atoms of Hydrogen

A

-A carbon atom has 4 electrons in its outer energy level.

-That means that 4 more electrons are required for carbon to have a full outer energy level.

-Each hydrogen atom has 1 electron, so one carbon atom has to form covalent bonds with four atoms of hydrogen.

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26
Q

Explain how Oxygen Atoms get Full Outer Electron Shells when they bond to form a Molecule

A

-Oxygen atoms have 6 electrons in their outer energy level which means that 2 more electrons are required for a full outer energy level.

-In the case of the double covalent bond, each oxygen
atom puts 2 electrons into the bond.

-So both atoms now have 8 electrons in their outer electron shell (a full outer electron shell).

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27
Q

Explain how Nitrogen Atoms get Full Outer Electron Shells when they bond to form a Molecule

A

-Nitrogen atoms have 5 electrons in their outer energy level which means that 3 more electrons are required for a full outer energy level.

-In the case of the triple covalent bond, each nitrogen
atom puts 3 electrons into the bond.

-So both atoms now have 8 electrons in their outer electron shell (a full outer electron shell).

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28
Q

What is a Small Molecule

A

-Small molecules are substances made up of non-metal atoms.

-A molecule is a particle made from atoms joined together by covalent bonds (between atoms).

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29
Q

Explain why Small Molecules have Low Melting/ Boiling Points

A

-Small molecules have low melting/ boiling points.

-This is because of weak intermolecular forces of attraction between the molecules requiring little energy to break them.

-The strong covalent bonds are not affected.

-Small molecules are often gases at room temperature.

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30
Q

What happens to the Melting/ Boiling Point as the Size of Covalent Molecules Increases.

A

-As the size of covalent molecules increases, the melting/ boiling points increase as well.

-this is because the strength of the intermolecular forces of attraction increases.

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31
Q

Explain why Small Molecules do not Conduct Electricity

A

-Small molecules do not conduct electricity because molecules are not electrically charged.

-They are neutral.

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32
Q

Explain why Large Covalent Structures are Strong

A

-Large covalent structures are strong (hard except graphite) because the atoms are held together by many strong covalent bonds.

-These take a lot of energy to break.

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33
Q

Explain why Large Covalent Structures have High Melting/ Boiling points

A

-Large covalent structures have high melting/ boiling points because the atoms are held together by many strong covalent bonds.

-These take a lot of energy to weaken the melting point and overcome the boiling point.

-They are always solids at room temperature.

34
Q

Explain why Large Covalent Structures do not Conduct Electricity

A

-Large covalent structures (except graphite) do not conduct electricity because there are no charged particles (electrons/ ions) free to move.

35
Q

Describe the Structure and Bonding of a Diamond Molecule

A

-Diamond is made from the element carbon.

-Each carbon atom forms four covalent bonds to four other carbon atoms.

-That means that a single diamond molecule is a giant molecule containing many millions of carbon atoms joined by covalent bonds.

36
Q

Explain why the Melting Point of Diamond is Greater than 3700℃

A

-When we melt diamond, we need to break the covalent bonds between the carbon atoms.

-Covalent bonds are extremely strong and take a great deal of energy to break.

-So to provide this energy, we need to heat diamond to a very high temperature.

37
Q

Explain why Diamond cannot Conduct Electricity

A

-Diamond cannot conduct electricity.

-This is because every electron in the outer energy level of the carbon atoms are involved in a covalent bond.

-There are no free electrons to move and carry charge.

38
Q

Explain why Diamond is Hard

A

-Diamond is hard because it forms a rigid lattice.

-Many strong covalent bonds have to be broken to break the lattice which requires a lot of energy to do.

39
Q

What are the Uses of Diamond

A

-The rigid network of carbon atoms, held together by strong covalent bonds, makes diamond very hard.

-This makes it useful for cutting tools, such as diamond-tipped glass cutters and oil rig drills.

40
Q

Which two Elements are Present in Silicon Dioxide (Silica)

A

-One element is silicon. We can tell that because silicon is in group 4 of the periodic table, just like carbon. And just like carbon, silicon forms 4 covalent bonds.

-The other element is oxygen. We can tell that because oxygen is in group 6 of the periodic table. That means that it requires 2 electrons to have a full outer energy level, so it forms 2 covalent bonds.

41
Q

How is the Structure of Silicon Dioxide (Silica) Similar to the Structure of Diamond

A

-The structure of silicon dioxide is similar to diamond because they are both giant covalent molecules containing many millions of atoms joined by covalent bonds.

-Also, in silicon dioxide, each silicon atom forms four covalent bonds (just like the carbon atoms in diamond).

42
Q

How is the Structure of Silicon Dioxide (Silica) Different from the Structure of Diamond

A

-The structure of silicon dioxide is different to diamond as each silicon atom is not bonded to other silicon atoms.

-Instead, they are bonded to oxygen atoms, which then bond to silicon atoms

43
Q

Why is the Melting Point of Silicon Dioxide over 1700℃

A

-Similarly to diamond, when we melt silicon dioxide we have to break a very large number of covalent bonds.

-These bonds take a lot of energy to break.

-That means that silicon dioxide has a very high melting point.

44
Q

Why can Silicon Dioxide not Conduct Electricity

A

-In silicon dioxide, every electron in the outer energy levels of both the silicon and oxygen atoms are in covalent bonds.

-This means, again just like diamond, that silicon dioxide has no free electrons to move and carry an electric current.

45
Q

Describe the Structure and Bonding of Graphite

A

-Each carbon atom forms three covalent bonds with other carbon atoms.

-The carbon atoms form layers of hexagonal rings.

-There are no covalent bonds between the layers.

-There is one delocalised electron from each atom

46
Q

Why does Graphite have a Melting Point of 4000℃

A

-Graphite has a high melting point because it has millions of strong covalent bonds.

-These take a lot of energy to break.

47
Q

Explain why Graphite Feels Slippery to Touch

A

-Graphite consists of layers which are not strongly bonded to each other (there are forces but
they’re not strong).

-When we apply a force to graphite (for example rubbing it), the layers can slide off each other.

-This is why graphite feels slippery to touch (and also explains why we can use graphite to draw with for example in pencil and charcoal).

48
Q

Explain why Graphite can Conduct both Heat and Electricity

A

-The fourth electron on the outer shell of carbon is not used for bonding. It is delocalised so free to carry charge through the structure. This makes graphite good conductor of electricity.

-The moving delocalised electrons can also carry thermal energy, which makes graphite a good conductor of heat.

49
Q

What are the Uses of Graphite

A

-Graphite can conduct electricity. This makes graphite useful for electrodes in batteries and for electrolysis.

-Graphite is also slippery, so it is useful as a lubricant for machines and in pencils.

50
Q

Describe the Structure and Uses of Graphene

A

-Graphene is a single layer of graphite.

-It is a layer of carbon atoms arranged into hexagonal rings. Each carbon atom is only bonded to three other carbon atoms.

-This means the fourth electron is delocalised so can move and carry charge. This makes it a good conductor of electricity so is used in electronics.

-Graphene is very strong and has a high melting point which also makes it useful for new materials.

51
Q

Describe the Structure of Fullerenes

A

-Fullerenes are molecules of carbon atoms with hollow shapes.

-Their structures are based on hexagonal rings of carbon atoms joined by covalent bonds.

-Some fullerenes include rings with five or seven carbon atoms.

-Two examples of fullerenes are buckminsterfullerene and nanotubes.

52
Q

Describe the Structure and Uses of Buckminsterfullerene

A

Buckminsterfullerene was the first fullerene to be discovered. Its molecules are made up of 60 carbon atoms joined together by strong covalent bonds. Molecules of C60 are spherical.

There are weak intermolecular forces between molecules of buckminsterfullerene. These need little energy to overcome, so buckminsterfullerene is slippery and has a low melting point.

-They have large surface area to volume ratios.

-These properties make them useful as lubricants, catalysts and for delivering pharmaceuticals.

53
Q

Describe the Structure and Uses of Carbon Nanotubes

A

-A nanotube is like a layer of graphene, rolled into a cylinder. The length of a nanotube is very long compared to its width, so nanotubes have high length to diameter ratios.

-Nanotubes have high tensile strength, so they are strong in tension and resist being stretched. Like graphene, nanotubes are strong and conduct electricity because they have delocalised electrons.

-These properties make nanotubes useful for nanotechnology, electronics, specialised materials and strengthening other materials.

54
Q

What are Polymers

A

-Polymers are long chains of many monomers (small molecules) joined together.

-The bonding between atoms is covalent because they are all non-metals.

55
Q

What are the Properties of Polymers

A

-Polymers are very long chains.

-These chains are held to other chains by a lot of intermolecular forces.

-These intermolecular forces are strong and require a lot of energy to weaken or break.

-This means that polymers have high melting points and are usually solids at room temperature.

56
Q

How is Bonding between carbon Atoms Different in Monomers Compared to Polymers

A

-In the monomer (which is an alkene), the carbon atoms are joined by a double covalent bond.

-However, in the polymer, the carbon atoms are joined by single covalent bonds.

57
Q

Describe the Three Key Features of any Repeating Unit

A

-The bond between the carbon atoms must a single covalent bond not a double covalent bond.

-There must be a bond extending out from each carbon atom to show that this is only a tiny part of the polymer.

-There must be a lower case “n” to the right to show that the polymer consists of many of these repeating units joined together.

58
Q

What is Metallic Bonding

A

-A metallic bond is a strong electrostatic attraction between positive metal ions and nearby delocalised electrons.

-Metal atoms always lose outer shell electrons to form positive ions.

-Metallic bonds are non-directional. The attraction is not between a specific metal ion and a specific electron.

59
Q

Describe the Structure of Metals

A

-Metals consist of a giant structure.

-Atoms are arranged in layers.

-Metals contain delocalised electrons which can move

60
Q

What are the Limitations of the Model of Metals

A

-2d and not 3d

-Can’t explain all properties

-Doesn’t show empty space

61
Q

Explain why Metals are Good Conductors of Electricity

A

-Metals are good conductors of electricity.

-This is because their outer shell electron(s) are delocalised.

-This means that they are free to move and carry charge through the structure.

62
Q

Explain why Metals are Good Conductors of Heat

A

-Metals are good conductors of heat.

-This is because their outer shell electron(s) are delocalised.

-This means that they are free to move and carry thermal energy through the structure.

63
Q

Explain why Metals have High Melting and Boiling Points

A

-Metals have high melting and boiling points.

-This is because the electrostatic attraction between positive metal ions and negative delocalised electrons is strong (this is also known as a metallic bond).

-These require a lot of energy to weaken the melting point and overcome the boiling point.

64
Q

Explain why Pure Metals are Malleable

A

-In pure metals, all of the atoms are the same (therefore the same size).

-These atoms are arranged in layers.

-This means the layers of atoms can slide over each other to from different shapes.

65
Q

What is Meant by an Alloy

A

An alloy is a mixture of a metal with another metal.

66
Q

Explain why Alloys are harder than Pure Metals

A

-Alloys are a mixture of different sized atoms.

-This means the layers become distorted.

-This means the layers can’t slide over each other.

-Therefore, alloys are harder than pure metals.

67
Q

What is Nanoscience

A

-Nanoscience is the study of structures that are between 1 and 100 nanometres (nm) in size.

-Most nanoparticles are made up of a few hundred atoms.

68
Q

Give the Name, Symbol and Standard Form Power of Different Measurements

A

-Terametre, T, 10 to the power of 12

-Gigametre, G, 10 to the power of 9

-Megametre, M, 10 to the power of 6

-Kilometre, k, 10 to the power of 3

-Millimetre, m, 10 to the power of -3

-Micrometre, μ, 10 to the power of -6

-Nanometre, n, 10 to the power of -9

69
Q

What are Coarse Particles

A

-Coarse particles are also known as dust.

-They have a particle matter number (PM) of 10.

-They have a size range of 10 µm to 2.5 µm.

-They contain many thousand of atoms.

70
Q

What are Fine Particles

A

-They have a particle matter number (PM) of 2.5.

-They have a size range of 2.5 µm to 0.1 µm.

-They contain several thousand of atoms.

71
Q

What are Nano Particles

A

-They have a particle matter number (PM) of 0.1.

-They have a size range of 0.1 µm to 1 nm.

-They contain a few hundred atoms.

72
Q

What are Nanoparticulate Materials

A

-A substance that consists of nanoparticles is described as being nanoparticulate.

-Some properties of nanoparticulate materials are different from the properties of the same material in bulk, as powders, lumps or sheets.

-This difference in properties is the result of two things:

-The tiny size of nanoparticles compared to the same material in bulk.

-The large surface area to volume ratios of nanoparticulate materials compared to the same material in bulk

73
Q

What are the Uses of Nanoparticles

A

-Medical treatments

-Cosmetics, deodorants and sunscreens

-Electronics

-Catalysts

74
Q

What are the Advantages of Nanoparticles

A

-Higher surface area to volume ratio means a smaller quantity of a substance can be used for the same effect.

-Antimicrobial effects.

75
Q

What are the Disadvantages of Nanoparticles

A

-Nanoparticles leaking into the atmosphere may cause environmental damage.

-Scientists are unsure of the long term effects of nanoparticles on humans such as through ingestion or absorption through skin.

-This is why some are cautious around using nanoparticles for products such as deodorants or sun creams.

76
Q

What are the Advantages and Disadvantages of Dot and Cross Diagrams

A

-Advantage: Because we use dots to represent the electrons from one atom and crosses to represent the electrons from another atom, it is very clear where the
electrons are coming from.

Disadvantage: They do not tell us about the shape of the molecule.

77
Q

What are the Disadvantages of 2d Stick Diagrams

A

-Because the covalent bond is shown by a stick,
we cannot tell which electron in the bond came from which atom.

-Stick diagrams also do not give us any information about electrons that are not in bonds.

-2D stick diagrams do not give us accurate information about the shape of the molecule.

78
Q

What is the Advantage of 3d Stick Diagrams

A

-The advantage of a 3D stick diagram over a 2D stick diagram is that a 3D stick diagram shows us the shape of the molecule.

79
Q

What are the Advantages and Disadvantages of Ball and Stick Diagrams

A

Advantage: allows us to clearly see the ions in 3 dimensions.

Disadvantage: the ions are shown as widely spaced but in reality, the ions are closely packed together in a crystal lattice.

-Also, a ball and stick diagram gives us no idea of the relative sizes of the ions.

-Only show a tiny part of the giant crystal lattice. In fact, a crystal lattice is a giant structure.

80
Q

What are the Advantages and Disadvantages of Space Filling Diagrams

A

Advantage: Gives us a more realistic idea of how closely packed the ions are in the crystal lattice.

-Also a space filling diagram gives us an idea of the relative sizes of the ions.

Disadvantage: Because the ions are shown packed together, it can be difficult to see the 3D packing with a space filling diagram.

-Only show a tiny part of the giant crystal lattice. In fact, a crystal lattice is a giant structure.

Chemistry Paper 1 (5 decks)

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