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Chemistry: A-level > Module 2: Foundations In Chemistry > Flashcards

Module 2: Foundations In Chemistry Flashcards

1
Q

What three subatomic particles are atoms made up of?

A
  • Protons
  • Neutrons
  • Electrons
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2
Q

What is the mass and relative charge of a proton?

A

Mass: 1
Charge: +1

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3
Q

What is the mass and relative charge of neutrons?

A

Mass: 1
Charge: 0 (neutral)

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4
Q

What is the mass and relative charge of electrons?

A

Mass: 1/1836
Charge: -1

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5
Q

What is the mass number?

A

The larger number on the periodic table. The number of protons + neutrons in an atom.

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6
Q

What is the atomic number?

A

The smaller number on the periodic table. The number of protons.

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7
Q

What are Isotopes?

A

atoms of the same element with different numbers of neutrons.

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8
Q

Do isotopes of the same atom have the same chemical properties and the same physical properties? Explain why.

A

They will have the same chemical properties as they have the same electronic structure.
However may have different physical properties due to different masses.

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9
Q

What are Ions?

A

Atoms that have lost or gained electrons.

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10
Q

What is meant by Relative Atomic Mass?

A

The mean mass of an atom compared with 1/12 the mass of one atom of carbon-12

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11
Q

What is meant by Relative Isotopic Mass?

A

The mass of one atom of an isotope compared with 1/12 the mass of one atom of carbon-12.

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12
Q

How to calculate relative atomic mass from isotopic abundances?

A

Multiply each relative atomic mass by its % relative isotopic abundance and add the results, then divide by 100.

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13
Q

What is the name and the charge of the molecule NH4?

A

ammonium forms a +1 molecular ion

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14
Q

What is the name and the charge of the molecule OH?

A

Hydroxides form -1 molecular ions.

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15
Q

What is the name and the charge of the molecule NO3?

A

Nitrates form -1 molecular ions.

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16
Q

What is the name and the charge of the molecule CO3?

A

Carbonates form -2 molecular ions.

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17
Q

What is the name and the charge of the molecule SO4?

A

Sulphates form -2 molecular ions.

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18
Q

How do you work out an ionic formula from the ionic charges?

A

In an ionic compound, the overall charge is zero. Therefore, the ions in the compound must balance to zero, so you can add more + or - ions to balance the charges.

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19
Q

What do the state symbols (s), (l), (g), (aq) represent?

A

(s): solids
(l): liquids
(g): gas
(aq): aqueous

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20
Q

What is the rule for significant figures?

A

Round the result to the lowest amount of significant figures in the question.

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21
Q

How do you convert between Celcius and Kelvins?

A

Celsius + 273 = Kelvins

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22
Q

How to convert between dm3 and cm3?

A

cm3 / 1000 = dm3

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23
Q

How to convert between cm3 and m3?

A

cm3 / 1 000 000 = m3

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24
Q

How to convert between dm3 and m3?

A

dm3 / 1000 = m3

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25
Q

What is the definition of ‘a mole?’

A

A mole is the amount of a substance that contains the same number of particles as the number of carbon-12 atoms in 12g.

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26
Q

What does Avogadro’s number represent?

A

The number of particles in a mole.

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27
Q

What is the equation for moles?

A

Moles = Mass (g) / Molar Mass (g mol -1)

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28
Q

How to figure out the number of particles in a substance?

A

No. of particles = No. moles x Avogadro’s constant.

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29
Q

What is the Molar Gas Volume at room temperature and pressure?

A

24 dm3

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30
Q

What is the equation for ideal gas law?

A

PV = nRT
pressure (Pa) x volume (m3) = mols x gas constant (J/kmol) x temperature (K)

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31
Q

Equations for the concentration of solutions?

A

concentration (g dm-3) = mass (g) / volume (dm3)

concentration (mol dm-3) = mol / volume (dm3)

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32
Q

What is meant by the term ‘Molecular Formula?’

A

The actual number of atoms in an element.

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33
Q

What is meant by the term ‘Empirical Formula?’

A

The simplest whole number ratio of atoms in a compound.

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34
Q

What does it mean if a crystalline structure is ‘hydrated?’

A

Contains water molecules.

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35
Q

What does it mean if a substance structure is ‘anhydrous?’

A

The substance contains no water molecules.

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36
Q

How do you calculate percentage yield?

A

% Yield = ( Amount of product, in mol / theoretical amount of product, in mol) x 100

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37
Q

What are some of the reasons for a lower than predicted yield in reactions?

A
  • The reaction may not go to completion.
  • Unpredicted side reactions may occur.
  • The reactants may not be pure.
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38
Q

How to calculate Atom Economy?

A

% Atom Economy = (molar mass of useful products / sum of molar mass of all products) x 100

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39
Q

How can Atom Economy be improved?

A
  • Alternative reaction pathways
  • Finding uses for waste products.
40
Q

What is the formula and ions in Hydrochloric Acid?

A

HCl (g)
H+(aq) + Cl-(aq)

41
Q

What is the formula and ions of Sulphuric Acid?

A

H2SO4 (l)
H+(aq) + HSO4-(aq)

42
Q

What is the formula and ions of Nitric Acid?

A

HNO3
H+(aq) + NO3-(aq)

43
Q

What is the formula and ions of Sodium Hydroxide?

A

NaOH
Na+ and OH-

44
Q

What is the formula and ions of Potassium Hydroxide?

A

KOH
K+ and OH-

45
Q

How can you tell if an acid is strong or weak?

A

Strong acids fully dissociate, to release H+ ions, in aqueous solutions. Weak acids partially dissociate.

46
Q

What is the Neutralisation Equation?

A

H+ + OH- (reversible reaction arrow) H2O

47
Q

Metal + Acid -> ?? + ??

A

Metal + Acid -> Salt + Hydrogen

48
Q

Metal Oxide + Acid -> ?? + ??

A

Metal Oxide + Acid -> Salt + Water

49
Q

Metal Hydroxide + Acid -> ?? + ??

A

Metal Hydroxide + Acid -> Salt + Water

50
Q

Metal Carbonate + Acid -> ?? + ?? + ??

A

Metal Carbonate + Acid -> Salt + Water + Carbon Dioxide

51
Q

What is the oxidation state of all uncombined elements?

52
Q

TRUE OR FALSE
The sum of oxidation states in a neutral compound add to +1?

A

FALSE
The sum of oxidation states in a neutral compound add to 0.

53
Q

TRUE OR FALSE
The sum of oxidation states in an ion is equal to the charge on that ion?

54
Q

What is the usual oxidation state of Oxygen and what are the exceptions?

A

Usually -2, except in peroxides when it is -1 and when bonded with Fluorine when it is +2.

55
Q

What is the usual oxidation state of hydrogen and what are the exceptions?

A

Usually +1, except in metal hydrides where it is then -1.

56
Q

What is the oxidation number in combined Fluorine?

57
Q

What is the usual oxidation state for combined Group 1 elements?

58
Q

What is the usual oxidation state for combined Group 2 elements?

59
Q

What is Oxidation?

A

The loss of electrons.

60
Q

What is Reduction?

A

The gain of electrons.

61
Q

What does an increase in oxidation states tell us?

A

Gain of electrons, reduced. (Vice vera for decrease in oxidation number)

62
Q

What is a disproportionation reaction (oxidation and redox)?

A

The oxidation and reduction of the same element in a redox reaction.

63
Q

What do you balance oxygen/hydrogens/charges with in redox reaction half equations?

A

Oxygen: H2O
Hydrogen: H+
Charges: e-

64
Q

What sub-shell is in the first shell?

65
Q

What sub-shells are in the second shell?

66
Q

What sub-shells are in the third shell?

A

3s, 3p and 3d

67
Q

What sub-shells are in the fourth shell?

A

4s, 4p, 4d and 4f

68
Q

What is the order of sub-shells 1-4 in order of increasing energy?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

69
Q

How many electrons can ‘s’ sub-shells hold?

70
Q

How many electrons can ‘p’ sub-shells hold?

71
Q

How many electrons can ‘d’ sub-shells hold?

72
Q

What is Ionic Bonding?

A

The electrostatic attraction between a metal and a non-metal: resulting in a positive, metal ion and negative, non-metal ion.

73
Q

Properties of Giant Ionic Structures and why?
(hints: m.p. b.p., s, e)

A
  • High melting and boiling points due to strong electrostatic attractions.
  • Soluble in polar solvents (such as water) as the to polar molecules attract the charged ions in the lattice, disrupting it.
  • Conduct electricity when molten/dissolved as the solid lattice breaks down and ions are free to move, allowing them to carry electrical charges.
74
Q

What is Covalent Bonding?

A

The sharing of electrons between two non-metals.

75
Q

What is Dative Covalent Bonding? How are these bonds represented?

A

When one element gives both electrons for the covalent bond. Represented using an arrow pointing from the element giving the bonding pair.

76
Q

What is the name of the shape and bond angle in a molecule with one/two bonding pairs and no lone pairs?

A

Name: Linear Molecule
Angle: 180*

77
Q

What is the name of the shape and bond angle in a molecule with three bonding pairs and no lone pairs?

A

Name: Trigonal Planar
Bond Angle: 120*

78
Q

What is the name of the shape and bond angle in a molecule with four bonding pairs and no lone pairs?

A

Name: Tetrahedral
Bond Angle: 109.5*

79
Q

What is the name of the shape and bond angles in a molecule with five bonding pairs and no lone pairs?

A

Name: Trigonal Bioyramidal?
Bond Angles: of 120* and 90*

80
Q

What elements always obey the octet rule?

A

Carbon, Nitrogen, Oxygen and Fluorine.

81
Q

What is the name of the shape and bond angles in molecule with six bonding pairs and no lone pairs?

A

Name: Octahedral
Angles: 3D 90*

82
Q

What is Electron-Pair Repulsion Theory?

A

Electrons repel each other, so the electrons on the outer shell arrange each other so they are as far apart as possible.

83
Q

What is Electronegativity?

A

The power of an atom to attract the two electrons in a covalent bond.

84
Q

Does electronegativity increase or decrease across a period? Why?

A

Electronegativity increases as the number of protons increases, the number of shells remains the same and so the atomic radius is decreasing as the electrons are pulled in.

85
Q

Does electronegativity increase or decrease down a group? Why?

A

Electronegativity decreases as there is an increase in electron shielding.

86
Q

What is a Permanent Dipole?

A

A small charge difference across a bond that results from a difference in the electronegativities of the bonded atoms.

87
Q

What are Intermolecular Forces?

A

Forces between molecules.

88
Q

What is a Permanent Dipole-Induced Dipole Interaction?

A

When a molecule with a permanent-dipole goes near a non-polar molecule, it causes the electrons in the shell of the non-polar molecule to shift slightly and become slightly polar. Attraction then occurred, the polar molecule having induced the non-polar molecule.

89
Q

What is a Polar Covalent Bond?

A

A covalent bond that has a permanent dipole.

90
Q

What are London (dispersion) Forces?

A

Very weak attractive forces between induced dipoles.

91
Q

What is an Instantaneous Dipole?

A

A dipole that will occur at a random moment in a usually non-polar molecule because electrons are constantly moving.

92
Q

When does Hydrogen Bonding occur?

A

When hydrogen is bonded to Nitrogen, Oxygen or Fluorine (the highly electronegative elements)

93
Q

What is Hydrogen Bonding?

A

A strong permanent dipole-permanent dipole attraction between:
- An electron deficient hydrogen atom bonded with another atom
…and…
- A lone pair of electrons on a highly electronegative atom bonded with another atom.

94
Q

What are the effects of hydrogen bonding on the properties of water and why?

A
  • Ice is less dense than water as hydrogen bonds pushes molecules further apart.
  • Water has a higher than expected melting point and boiling point as hydrogen bonds are stronger than most intermolecular forces.
95
Q

Properties of Simple Molecular Structures?

A
  • Low melting and boiling points
  • Insoluble in water
  • Does not conduct electricity

Chemistry: A-level (5 decks)

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