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Chemistry AS > AS Chemistry Term 1 > Flashcards

AS Chemistry Term 1 Flashcards

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1
Q

Define relative atomic mass

A

The relative atomic mass (Ar) is the mass of an atom compared to 1/12 of the mass of an atom of carbon-12.

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2
Q

Define relative molecular mass

A

The relative molecular mass (Mr) is the mass of a molecule compared to 1/12 of the mass of an atom of carbon-12.

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3
Q

Define relative formula mass

A

The relative formula mass is the mass of a compound compared to 1/12 of the mass of an atom of carbon-12.

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4
Q

Define relative isotopic mass

A

The relative isotopic mass is the mass of an isotope compared to 1/12 of the mass of an atom of carbon-12.

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5
Q

What is Avagadro’s number/constant?

A

It is the number of atoms of carbon in 12 g of carbon-12. It has a value of 6.02 x 10^23. It is essentially the number of particles (atoms, molecules) in a mole.

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6
Q

Define empirical formula

A

The simplest whole number ratio of the elements present in one molecule or formula unit of a compound.

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7
Q

Define molecular formula

A

The actual number of each of the different atoms present in a molecule.

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8
Q

List the three ways in which moles can be calculated

A
  1. By Mass
    Amount = mass/molar mass
    n = m/M
  2. By Volume
    Amount = Volume/24
    n = V/24
  3. By Concentration
    Amount = Concentration x Volume
    n = c x v
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9
Q

State Avagadro’s Law

A

Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of particles.

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10
Q

Define an isotope

A

Atoms of the same element with different nucleon numbers

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11
Q

Define an orbital

A

An orbital is a region in space where there is a high probability (90-95%) of finding an electron. It is a three-dimensional region. The position of the electron is not certain.

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12
Q

Describe an s orbital

A

It is a spherical shape. There is one s orbital per energy level. It can hold a maximum of two electrons.

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13
Q

Describe a p orbital

A

It has a dumbbell shape with two lobes extending out into three-dimensional space. There are 3 p-orbitals per energy level with each orbital extending out on a different axis. As each orbital can hold two electrons, the p subshell holds 6 electrons in total.

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14
Q

Describe d and f orbitals

A

d: hold 5 orbitals, meaning can hold 10 electrons
f: hold 7 orbitals

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15
Q

State the rules for filling up electrons

A

Aufbau Principle:
Enter lowest energy orbital available, those of lowest energy being filled first. Energy levels are not entered until those below them are filled.

Pauli’s Excursion Principle:
Each orbital can hold a maximum of two electrons, provided they have opposite spin.

Hund’s Rule:
Orbitals of the same energy remain singly occupied before pairing up. This is due to repulsion between electron pairs.

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16
Q

What are the exceptions to the Aufbau Principle?

A

Chromium: ends in 4s1 3d5 rather than the expected 4s2 3d4. This is because of the relative stability of a half full set of d orbitals.

Copper: ends in 4s1 3d10 rather than expected 4s2 3d9 due to stability of full d sub orbital and the closeness in energies of 3 and 4.

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17
Q

Define 1st ionisation energy

A

The energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous cations.

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18
Q

Describe the factors that influence ionisation energies

A
  1. Size of nuclear charge: As atomic number increases the positive nuclear charge increases. This means greater attractive forces between nucleus and electrons. This means ionisation energies increase across the period.
  2. Distance of outer electrons from nucleus: Decreases as distance increases.
  3. Shielding effect of inner electrons: Electrons repel each other so electrons in full inner shells repel electrons in outer shells. Both distance and shielding mean that ionisation energy decreases down a group as distance and number of electrons increase.
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19
Q

List types and differences of chemical and physical bonds

A 
Chemical bonds: Strong
- Ionic/Electrovalent
- Covalent
- Dative covalent/co-ordinate
- Metallic
Physical Bonds: Weak
- Van der Waals forces - weakest
- Dipole-Dipole interaction
- Hydrogen bonds - Strongest
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20
Q

Describe the formation of ions

A

Positive ions: Electron is removed from the atom and are therefore, smaller than the original atom. The energy required to remove the electron is the ionisation energy.

Negative ions: Larger than the original atom due to electron repulsion in outer shell. Negative ions are formed when electrons are added to the atom. The energy that is released by the pulling of the electron is the electron affinity.

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21
Q

Describe the formation of ionic bonds

A
  • Formed between a cation and anion.

- Electrons are transferred from one atom to another

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22
Q

Define the terms Monoatomic ion, Polyatomic ion and Valency

A

Monoatomic Ion: A single atom with an overall charge
Polyatomic Ion: A stable group of atoms with an overall charge
Valency: The size of the charge on an ion (integer value)

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23
Q

Describe the structure of ionic compounds and the properties this gives the compound

A

It is a giant ionic lattice that has oppositely charged ions held in a regular 3-dimensional lattice by electrostatic attraction.
Properties:
- High melting and boiling points: Large amounts of energy required to overcome strong electrostatic attractions. Charge influences melting and boiling points, higher charge means higher melting and boiling points.

  • Solid does not conduct electricity as ions are in fixed lattice structure and cannot carry charge. Can conduct in molten or aqueous state as organised structure is broken and ions are free to move and can carry charge.
  • Are hard due to strong electrostatic forces of attraction holding crystal together
  • Dissolve in water to some extent as water molecules can form attractions to positive and negative ions, breaking up the lattice structure and keeping the ions apart.
  • Brittle, ions are arranged in planes which can be distorted under a strong force. This causes ions of the same charge to be pushed together which repel each other and break the crystal.
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24
Q

Describe the structure of metallic bonding

A

3 dimensional lattice of metal ions surrounded by valence electrons. The electrons move from ion to ion in a ‘sea of mobile electrons’. The electrons are delocalised. The strength of the bond depends on the ability of metal ions to pack efficiently and the number of valence electrons the metal ion can contribute to the ‘sea of delocalised electrons’.

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25
Q

Describe the properties of metallic solids

A
  • Conduct electricity due to mobile electrons that can move through the lattice carrying charge.
  • High melting and boiling points: Due to strong attraction between the metal ions and delocalised electrons.
  • Shiny: close pack structure that reflects light from its surface.
  • High Densities: Due to close packed nature of their lattice structure
  • Insoluble in solvents
  • Malleable and Ductile: Layers in the metal can reform without breaking. This is because delocalised electrons can move to prevent repulsive forces from breaking the structure.
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26
Q

Define electronegativity

A

The measure of attractive force of the nucleus of an atom on a shared electron pair in a covalent bond.

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27
Q

Describe a non-polar covalent bond

A

Formed between two atoms with the same electronegativity and are usually of the same atom. This means electrons are distributed evenly

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28
Q

Describe a polar covalent bond

A

When there is more than one element involved, electrons can be shared unevenly between the unlike atoms. This is because of differing electronegativities.

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29
Q

List properties of a non-polar covalent bond

A
  • Low melting and boiling points: Due to weak intermolecular forces.
  • Do not conduct electricity: no mobile electrons or ions
  • Soluble in non-polar solvents
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30
Q

List properties of a polar covalent molecule

A
  • Low melting and boiling points: Due to weak intermolecular forces.
  • Do not conduct electricity: no mobile electrons or ions
  • Soluble in polar solvents such as water: Attractions form between charged ends of the polar molecules and the charged ends of the solvent molecules.
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31
Q

Explain the terms bond length and bond disassociation energy

A

Bond length: Due to the attractive forces between the nuclei of one atom and the electron pair of another, a balance between the two forces must be found. This balance between the attractive and repulsive forces will determine the bond length.

Bond (disassociation) energy: The amount of energy required to break a bond in the gaseous state. The bond energy indicates the strength of the bond. If the bond energy is higher, the bond is harder to break and is thus, stronger.

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32
Q

Describe co-ordinate/dative covalent bonding

A

A dative covalent bond has all the same properties of a normal covalent bond except for the fact that one atom provides both the electrons required for the covalent bond. For dative covalent bonding we need:

  1. One atom having a pair of lone electrons
  2. A second atom with an unfilled orbital to accept the lone pair i.e. an electron deficient compound.
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33
Q

State properties of a network covalent/macromolecular solid

A
  • Have strong covalent bonds throughout the entire structure.
  • Very hard: due to strong 3-dimensional bonds throughout structure
  • Do not conduct electricity: All electrons are localised to covalent bonds
  • High melting and boiling points: Due to strong 3-dimensional bonds throughout structure
    -Insoluble in all solvents
    Graphite is an exception as it is soft and conducts electricity. This is because it is only bonded in 2 dimensions rather than 3 making a layered structure held together by delocalised electrons (Van Der Waals).
34
Q

List and describe the different shapes of molecules

A 
Linear:
- 2 areas of negative charge about the central atom
- 2 occupied sites/bonded pairs
- Bond angle: 180°
Trigonal Pyramidal:
- 3 areas of negative charge
- 3 bonded pairs
- Bond angle: 120°
Tetrahedral:
- 4 areas of negative charge
- 4 bonded pairs
- Bond angle: 109.5°
Trigonal Pyramidal:
- 4 areas of negative charge
- 3 bonded pairs, 1 lone pair
- Bond angle: 107°
V-shaped:
- 4 areas of negative charge
- 2 bonded pairs, 2 lone pairs
- Bond angle: 104.5°
Trigonal Bypyrimidal:
- 5 areas of negative charge 
- 5 bonded pairs
- Bond angle: 90° on top and bottom, 120° on plane
Octahedral:
- 6 areas of negative charge 
- 6 bonded pairs 
- Bond angle: 90°
35
Q

Describe Van Der Waals forces

A

Electron clouds in non polar molecules are constantly moving meaning more of the charge cloud could be one side side than the other. This means that one end of the molecule temporarily has more of a negative charge. This results in a temporary dipole which can induce dipoles on neighbouring molecules. This can create forces of attraction between the molecules. The strength of the forces increase with:

  1. increasing the number of electrons and protons in the molecule
  2. increasing the number of contact points between the molecules.
36
Q

Describe Permanent dipole-dipole forces

A

The uneven distribution of charge in some molecules create a permanent dipole. The positive end of one molecule’s dipole can attract the negative end of another.

37
Q

Describe hydrogen bonding

A

Strongest intermolecular force. Requires:
1. One molecule having a hydrogen atom covalently bonded to F, O or N
2. A second molecule with an F,O or N atom with an available lone pair of electrons
A hydrogen bond is a stronger dipole-dipole force that results from a hydrogen atom bonded to a very electronegative atom, creating a highly polarised bond.

38
Q

Describe the unusual properties of water

A
  1. Enthalpy change of vaporisation and boiling point: Much higher than expected due to extensive hydrogen bonding. This means water is a liquid at room temperature when it would be expected to be a gas.
  2. Surface Tension and Viscosity: High surface tension and viscosity as hydrogen bonding reduces the ability of water molecules to slide over each other. Hydrogen bonds also exert a downward force at the surface of the liquid.
  3. Ice is less dense than water: Due to hydrogen bonds being quite long the rigid lattice of water molecules are in a ‘more open’ arrangement. When the ice melts, the molecules collapse in on each other making water more dense than ice.
39
Q

State the kinetic theory of matter

A
  • All matter is composed of tiny particles that may be of different sizes
  • The arrangement of particles in different states is different
  • Attractions (bonds) only exist between particles in the solid and liquid states.
  • Melting and boiling points are determined by the strength of the intermolecular forces attracting the particles together.
40
Q

State the assumptions of the kinetic theory of gases

A
  • Gases consist of very small particles whose sizes are negligibly small compared to the distances in between them.
  • Particles in the gas are in constant random and rapid motion.
  • The particles in the gas collide constantly with each other and the walls of the container.
  • The collisions are perfectly elastic - no energy loss
  • Particles have no effect on each other except during collisions
  • The kinetic energy of the particles is proportional to the temperature of the gas.
41
Q

What are the limitations of ideal gas laws?

A

At high pressures and low temperatures:

  • molecules are close together
  • the volume of the molecules is not negligible compared with the volume of the container
  • there are Van der Waals or dipole-dipole forces of attraction between the molecules
  • attractive forces pull the molecules towards each other and away from the walls of the container
42
Q

State the gas laws

A

Boyle’s Law: At constant temperature, and for a given amount of gas, volume is inversely proportional to pressure.

Charles’ Law: The volume of a given amount of gas at constant pressure is proportional to its absolute temperature.

General gas equation: Incorporates Avagadro’s expression. PV = nRT. P in N/m2, V in m3, T in K, n in mol. R is the gas constant and is approximately 8.31 when using these units.

43
Q

Describe vapour pressure

A

In a sealed container with a liquid, the liquid may evaporate and will form a pressure above the liquid. A balance point will eventually be reached where the rate of condensation of the liquid will equal the rate of evaporation. This equilibrium point is called the saturated vapour pressure. This points increases with temperature. The temperature at which vapour pressure is equal to atmospheric pressure is the boiling point of the liquid.

44
Q

Describe a giant molecular structure

A

Three dimensional network of covalent bonds throughout the whole structure.

45
Q

List properties of graphite and diamond

A

Graphite: Carbon atoms are arranged in planar layers. Within the layers the carbon atoms are arranged hexagonally. Each carbon atom is covalently bonded to 3 others with the the 4th electron delocalised above and below the carbon layers. There are weak Van der Waals between each layer.

  • High melting and boiling points: strong covalent bonds within layers
  • Soft: attraction between layers are weak and layers can slide over each other when force is applied.
  • Good conductor: due to delocalised electrons

Diamond: Each carbon atom covalently bonded with four others and are arranged tetrahedrally.

  • High m.p and b.p: strong covalent bonds throughout structure
  • Hard: Difficult to break three dimensional network of covalent bonds
  • Do not conduct: no free electrons as all atoms are involved in covalent bonding.
46
Q

List properties of silicon dioxide

A

Silicon Dioxide: Each Si atom is bonded to 4 oxygen atoms and each oxygen atom is bonded to 2 Si atoms.
- Properties are similar to diamond.

47
Q

List properties of fullerenes

A

Fullerenes: Allotropes of carbon in the form of hollow tubes or spheres.
Buckminsterfullerene - Sphere
- Low sublimation point: only weak Van der Waals between fullerene molecules
- Soft, weak intermolecular forces
- Poor conductor of electricity, extent of electron delocalisation is low
- Slightly soluble
- More reactive than graphite or diamond due to relatively high electron density in certain parts of the molecule.
Nanotubes - Cylinder
- High electrical conductivity, some electrons are delocalised
- High tensile strength
- High melting point as there is strong covalent bonding throughout the structure.

48
Q

List properties of graphene

A

A single isolated layer of graphite, not rigid and can be distorted

  • Most chemically reactive form of carbon.
  • Extremely strong for its mass
  • Conducts heat and electricity better than graphite.
49
Q

Why should we recycle materials?

A
  • Saves money: keeps prices lower as it is cheaper to reuse than to extract new materials
  • Saves natural resources for the future as well as associated energy inputs e.g. fossil fuels
  • Saves environment: by recycling we do less mining or drilling into the Earth.
  • Reduces problem of waste disposal: recycling prevents litter.
50
Q

Define organic chemistry and an organic compound

A

Organic Chemistry: The study of a large class of chemical compounds containing carbon.

Organic Compound: A compound consisting of carbon and any other elements

51
Q

Define functional group and homologous series

A

Functional Group: The site on an organic molecule where reactions take place. Properties of an organic molecule are hence, predominantly determined by the properties of the functional group.

Homologous Series: Group or family of compounds all having the same functional group and represented by the same general formula.

52
Q

Define a saturated and unsaturated compound

A

Saturated: A compound with no double or triple bonding. It is incapable of undergoing addition reactions.

Unsaturated: A compound containing multiple bonding and is capable of undergoing addition reactions.

53
Q

Define catenation

A

The ability to form bonds between atoms of the same element. Carbon catenates to form chains and rings.

54
Q

Describe the types of formula used for organic molecules

A

General Formula: Represents any member of the homologous series e.g. CnH2n

Empirical Formula: Shows simplest whole number ratio of atoms of each constituent element.

Molecular Formula: Shows number of each type of atom in the molecule

Structural Formula: Two-dimensional arrangement of atoms in the molecule giving an unambiguous structure using minimal detail and conventional groups.

Displayed Formula: Shows relative placing of atoms and number of bonds between them.

Skeletal Formula: Used to show a simplified organic formula by removing hydrogen atoms leaving just the carbon skeleton and functional groups.

55
Q

Define isomerism

A

Occurs when compounds have the same molecular formula but different structural formulae.

56
Q

Describe and list types of structural isomers

A

Arises where compounds have different structural formula - same number and type of atoms arranged in a different way.

  1. Chain Isomers: Compounds with different carbon skeletons.
  2. Positional Isomers: Molecules that have a substituent in different positions on the same carbon skeleton.
  3. Functional Group Isomers: Molecules belong to different homologous series.
57
Q

Describe and list types of stereoisomerism

A

Arises due to different spatial arrangements of atoms in a molecule.
Geometric Isomerism: Occurs in compounds in which free rotation is prevented by the presence of a double bonds or any other means. They are called -cis and -trans isomers and have identical atoms or groups on the same carbon atom but have different spatial arrangements.

Optical Isomers: Occurs when four different atoms are bonded to a central, chiral, carbon atom. The isomers are called enantiomers. The two isomers are unable to superimpose onto each other. They have identical physical and chemical properties but are distinguished by their ability to rotate the plane of polarised light. Both isomers rotate polarised light the same amount but they do so in opposite directions. The isomer that rotates light to the right is the dextrorotary or ‘d’ enantiomer and the one that rotates light to the left is the laevorotatory or ‘l’ enantiomer.

58
Q

Describe addition reactions

A

A reaction in which atoms or groups of atoms are introduced into an unsaturated molecule.

59
Q

Describe elimination reactions

A

Reactions in which new multiple bonds are made between carbon atoms with the elimination of a small molecule such as water or hydrogen chloride.

60
Q

Describe substitution reactions

A

A reaction in which one atom is exchanged for another. An example is when alkanes react with halogens which requires light or high temperatures. These specific substitution reactions are known as halogenation reactions and if chlorine is involved, it is called chlorination.

61
Q

Describe a condensation reaction

A

Two large molecules combine, discarding a smaller molecule - often water or HCl.

62
Q

Describe a redox reaction

A

Reactions involving oxidation and reduction half reactions.

63
Q

Describe hydrolysis reactions

A

Decomposition or alteration of a chemical substance by water.

64
Q

Describe homolytic fission

A

The breaking of a covalent bond leaving both atoms involved in the bond with one electron each. These atoms are called free radicals. A half curly arrow is used to indicate the movement of a single electron.

65
Q

Describe heterolytic fission

A

When one atom takes both the bonding electrons when the covalent bond is broken. Ionic intermediates are formed from the unequal splitting. The resulting particles of heterolytic fission are ions. Curly arrows are used to indicate the movement of a pair of electrons.

66
Q

Describe a nucleophilic reaction

A

A nucleophile is a reagent which seeks out the slightly positive end of a polar bond. This is because nucleophiles have a lone pair of electrons which are attracted to the positive charge. A reaction in which a molecule is attacked by a nucleophile will be a nucleophilic reaction.

67
Q

Define atomic radius and describe its trend

A

Definition: It is half the average distance between the nuclei of two covalently bonded atoms or metal-bonded atoms.

Trend: Atomic Radius decreases across periods as across the period the number of protons increase with each successive element, meaning the size of the nuclear charge increases. The extra electrons occupy the same energy level so shielding is fairly constant, so the nuclear charge pulls the valence electrons closer to the nucleus, hence, decreasing atomic radii.

68
Q

Describe the trend of ionic radii across a period

A

Positive ions are formed by atoms that have lost electrons, and so have lost their outer energy level. This decreases their radii significantly. The radii also decreases down from Na+ to Si 4+ due to increasing nuclear charge. Negative ions increase in size as, as extra electrons are added onto the same energy level, the repulsion between electrons increase, pushing the electrons further from the nucleus.

69
Q

Describe the trend of ionisation energy

A

Generally increases across the period as nuclear charge increases. Shielding and distance are similar across the period. However there are small dips in the trend due to the electron being removed from s1, p1 or p4.

70
Q

Describe the trend of electronegativity

A

Increases across a period due to the increase in nuclear charge with the increase in atomic number. The increased nuclear charge means a shared electron pair will be attracted more strongly.

71
Q

Describe the trend of melting point

A

Based on bonds and structures in molecule. Metallic bonding in metals means there is a relatively high melting point. The melting point increases from Na to Al as Al donates more electrons to the sea of delocalised electrons meaning there are greater electrostatic forced of attraction. Si has the highest melting point as it is a giant covalent structure meaning covalent bonds must be broken for it to melt. P, S, Cl and Ar have low melting points as for them to melt, the weak intermolecular forces must be overcome. P and S are slightly higher due to the fact that they are found as P4 and S8, meaning their melting points are higher.

72
Q

Describe the trend of electrical conductivity

A

Metals conduct electricity. Al is the best conductor from Na, Mg and Al because each atom donates three electrons to the sea of delocalised electrons meaning there are a greater number of delocalised electrons, thus increasing conductivity. Silicon is a semi-conductor and practically has no conductivity, while non-metals do not conduct.

73
Q

List the reactions of elements with oxygen

A
  • 4Na + O2 → 2Na2O : Reacts with air
  • 2Mg + O2 → 2MgO : Burns when heated in air
  • 4Al + 3O2 → 2Al2O3 : Will burn when heated strongly in powder form as it has a protective oxide layer.
  • Si + O2 → SiO2
  • 4P + 5O2 → P4O10 : Ignites spontaneously in air
  • S + O2 → SO2 : Burns when heated in air
  • Chlorine and Argon do not react with Oxygen
74
Q

List the reactions of elements with chlorine

A
  • 2Na + Cl2 → 2NaCl
  • Mg + Cl2 → MgCl2
  • 2Al + 3Cl2 → Al2Cl6
  • Si + 2Cl2 → 4SiCl4 (l)
  • 2P + 5Cl2 → 2PCl5
75
Q

List the reactions of elements with water

A
  • 2Na + 2H20 → 2NaOH + H2 : Vigorous reaction with cold water
  • Mg + 2H2O(l) →Mg(OH)2 + H2 : Slow reaction with water
  • Mg + H20 (g) → MgO + H2 : Reacts vigorously with steam.
76
Q

Describe the trend in oxidation number

A

Increases across the period as number of valence electrons increase.

77
Q

List the reactions of the oxides in water

A
  • Na2O + H20 → 2NaOH : Strongly alkaline
  • MgO + H2O → Mg(OH)2 : Weakly alkaline
  • SO2 + H2O → H2SO3
78
Q

Explain the amphoteric nature of aliminium

A

Amphoteric means it reacts with both an acid and a base. Aliminium is acidic as the high 3+ charge draws electrons out of the O - H bond as water molecules approach: [Al(H2O)6]3+ → [Al(H2O)5OH]2+ + H+

However it also the same character as all other metal hydroxides: Al(OH)3 + 3H+ → Al3+ + 3H2O

79
Q

List the reactions of chlorides with water

A
  • NaCl → Na+ + Cl-
  • MgCl2 + 6H2O → [Mg(H2O)6]2+ + 2Cl - : 2+ charge attracts water molecules and forms dative bonds using lone pairs in oxygen. Mg2+ attracts electrons in O-H resulting in hydrogens being more positive. This gives it a slight acidity:
  • [Mg(H20)6]2+ + H2O → [Mg(H2O)5OH]+ + H30+
  • Al2Cl6 + 3H2O → Al2O3 + 6HCl : Displays some molecular character
  • SiCl4 + 2H2O → SiO2 + 4HCl
  • PCl3 + 3H2O → H3PO3 + 3HCl
  • PCl5 + 4H2O → H3PO4 + 5HCl
80
Q

Describe the term ceramics

A

Range of materials that consist of a three-dimensional network of covalent bonds. They are hard, strong and brittle. They do not conduct electricity and have very high melting points.
Types:
Magnesium Oxide - used in refractory bricks in furnace lining for high temperatures

Aliminium Oxide - Electrical Insulator

Silicon Dioxide - In glass products

Chemistry AS (3 decks)

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