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Chemistry > Module 2: Foundations in chemistry > Flashcards

Module 2: Foundations in chemistry Flashcards

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1
Q

Describe the basic structure of an atom

A

Made from subatomic particles - protons, neutrons and electrons

Protons and neutrons are in the nucleus but electrons are in a region outside in shells

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2
Q

Describe the relative masses and charges of the three subatomic particles

A

Proton: mass 1, charge +1
Neutron: mass 1, charge 0
Electron: mass 1/1835 (negligible), charge -1

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3
Q

What is an isotope?

A

Atoms of the same element with different numbers of neutrons and different masses

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4
Q

What is atomic number? What is it represented by?

What is mass number? What is it represented by?

A

Number of protons -Z

Sum of protons + neutrons - A

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5
Q

Describe the similarities / differences in the properties of isotopes of the same element

A
  • Same chemical reactions as an atom’s chemistry is determined by the behaviour of its electrons and isotopes have the same number + configuration of electrons
  • Slight variations in physical properties (e.g boiling points) as they have different masses
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6
Q

What is relative isotopic mass?

A

The mass of an atom of an isotope compared to 1/12 the mass of an atom of carbon-12

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7
Q

What is relative atomic mass?

A

Ar - the weighted mean mass of an atom of an element compared to 1/12 the mass of carbon-12

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8
Q

How is Ar determined?

A

Mass spectrometry - a mass spectrum is produced showing the isotopes present in a sample of an element and their relative abundances.

Positive ions of isotopes are shown as a mass/charge ratio (m/z)

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9
Q

What is the formula for relative atomic mass?

A

((isotopic mass1 x abundance1) + (isotopic mass2 x abundance2))/100

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10
Q

What is relative molecular mass?

A

Used for simple molecular compounds. The sum of the relative atomic masses of all atoms in the compound.

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11
Q

What is relative formula mass?

A

Used for giant ionic or giant covalent compounds. Sum of the relative atomic masses of all atoms in the formula

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12
Q

How do you show ionic formulae? Why?

A

You can’t write a molecular formula for an ionic compound (no simple molecules)

An empirical formula is used instead - ratio of cations (positive ions) to anions (negative ions) present in the ionic lattice.
Worked out from balancing the charges of ions (as the compound is neutral)

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13
Q

What is the formula of a nitrate ion?
Carbonate?
Sulfate?
Hydroxide?

A

NO3 -
CO3 2-
SO4 2-
OH-

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14
Q

What is the formula of an ammonium ion?
Zinc?
Silver?

A

NH4 +
Zn 2+
Ag+

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15
Q

What is the ionic formula of barium chloride?

A

Ba 2+ Cl-
Ba 2+ 2Cl-
BaCl2

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16
Q

Balance this equation:
Na + Cl2 -> NaCl

A

2Na + Cl2 -> 2NaCl

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17
Q

Write the ionic equation for this reaction:
KI (aq) + AgNO3 (aq) = AgI (s) + KNO3 (aq)

A

KI (aq) + AgNO3 (aq) = AgI (s) + KNO3 (aq)

(write out all aqueous ions)

K+(aq) + I-(aq) + Ag+(aq) + NO3-(aq) = AgI(s) + K+(aq) + NO3-(aq)

(cancel out spectator ions)

I-(aq) + Ag+(aq) = AgI(s)

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18
Q

What is the mole?
What is the Avogadro constant?

A

The unit for amount of substance. One mole contains the same number of particles as there are atoms in 12g of Carbon-12. This number is the Avogadro constant (6.02 x 10^23)

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19
Q

What is molar mass?
How is it calculated?

A

The mass in grams of one mole of a substance = Mr

Moles = mass/Molar mass
n = m/Mr

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20
Q

How do you find the number of particles in a substance?

A

Calculate the number of moles by the Avogadro constant

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21
Q

What is empirical formula?
What is molecular formula?

A
  • Simplest integer ratio of atoms of each element present in a compound
  • Number and type of atoms of each element in a molecule
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22
Q

How would you find the empirical formula from given percentage composition or masses of elements?

A
  • Divide the % or mass by Ar to find moles
  • Divide by the smallest number to find an integer ratio
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23
Q

How would you find the molecular formula from a given empirical formula and molecular mass?

A
  • Find the mass of the empirical formula
  • Divide molecular mass by this
  • Multiply the number of each atom in the empirical formula by this value
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24
Q

What is a hydrated compound?
What is an anhydrous compound?

A
  • Crystalline and contains water molecules
  • Contains no water molecules
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25
Q

What is water of crystallisation?
How is it represented?

A

Water molecules that are bonded to the crystalline structure of a hydrated compound.

Amount is shown after a dot in the formula e.g CuSO4 . 5H20

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26
Q

How do you find the formula of a hydrated salt?

A
  • Calculate the moles of the anhydrous salt and the moles of water
  • Find the ratio between them (will be 1:xH20)
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27
Q

How would you check that a salt is anhydrous in a practical?

A

Check that water of crystallisation has been removed by heating to a constant mass

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28
Q

What is concentration?
How is it calculated?

A

The amount (in moles or grams) of a dissolved substance in 1dm^3 of a solution.

moles or mass = concentration x volume

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29
Q

What is molar gas volume?
How does it change with temperature and pressure?

A

The volume per mole of a gas at a stated temperature and pressure in dm^3 mol^-1

As temperature increases, it increases
As pressure increases, it decreases

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30
Q

What is the volume of 1 mole of any gas at RTP?

A

24dm^3

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31
Q

What is the ideal gas equation? Give units

A

pV = nRT
pressure = kPa / Pa
volume = dm^3 / m^3
n = moles
R = gas constant = 8.314
temperature = K (C+273)

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32
Q

How would you determine relative molecular mass of a liquid?

A
  • Evaporate the liquid
  • Use ideal gas equation to find moles
  • Use mass/moles = Mr
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33
Q

What is stoichiometry?

A

The ratio of moles in a reaction (shown by the balanced equation)

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34
Q

How would you work out an unknown quantity from the balanced equation?

A
  • Find the moles of the known value
  • Use the ratio to find the moles of the unknown substance
  • Calculate mass by moles x Mr
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35
Q

What is atom economy?

A

The percentage proportion of reactants that are converted into useful products

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36
Q

How do you calculate atom economy?

A

(sum of molar masses of desired products / sum of molar masses of all products) x 100

Take balancing numbers into account

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37
Q

What does a high atom economy suggest?

A

The reaction is efficient and sustainable (produces less waste and uses fewer raw materials)

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38
Q

What reactions have 100% atom economy?

A

Reactions with only one product

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39
Q

What is percentage yield and how is it calculated?

A

The actual yield shown as a percentage of theoretical yield

(actual yield / theoretical yield) x 100

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40
Q

What is an acid?
Give common examples

A

A substance that releases H+ ions when dissolved in water.
Hydrochloric acid (HCl), Sulfuric acid (H2SO4), Nitric acid (HNO3), Ethanoic acid (CH3COOH)

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41
Q

What is a base?
Give some examples

A

Substances that accept H+ ions
Metal oxides, metal hydroxides, metal carbonates, alkalis

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42
Q

What is the difference between an acid and a base in reference to H+ ions?

A

A H+ ion is a proton
Acids are referred to as proton donors
Bases are referred to as proton acceptors

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43
Q

What is a strong acid?
Give some examples
Give an example of an ionic equation

A

Fully dissociate when dissolved in water
HCl, HNO3, H2SO4

HCl + aq —> H+ + Cl-

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44
Q

What is a weak acid?
Give an example
Give an example of an ionic equation

A

Partially dissociate when dissolved in water
CH3COOH
CH3COOH (reversible arrow) CH3COO- + H+

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45
Q

What are the differences between strong and weak acids?

A

A strong acid produces more H+ ions and has a lower pH than a weak acid with the same concentration.

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46
Q

What is an alkali?
Give some common examples
Give an example of an ionic equation

A

A base that dissolves in water, releasing OH- ions

Sodium hydroxide (NaOH), Potassium hydroxide (KOH), ammonia (NH3)

NaOH + aq –> Na+ + OH-

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47
Q

What is a neutralisation reaction?
How do the ions behave?

A

A reaction between an acid and a base to form a salt

Salt is formed when the acid’s H+ ion is replaced by a metal / ammonium ion

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48
Q

Give the general neutralisation reactions
Explain how ions behave when applicable

A

acid + carbonate –> salt + water + carbon dioxide
acid + metal oxide –> salt + water
acid + alkali –> salt + water (H+ + OH- –> H2O)
acid + ammonia solution –> ammonium salt
(NH3 accepts H+ ions from acids, forming ammonium salts with the ammonium ion NH4+)

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49
Q

What are acid-base titrations used for?
What do they require?

A

Used to find information about a substance in solution
Needs:
- Solution of known concentration (standard solution) reacts with a solution of unknown concentration
- Very precise apparatus (burettes, pipettes) used to measure volumes
- Indicator used to show end point (when exact neutralisation has occurred)

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50
Q

How do you prepare a standard solution?

A

1) Work out mass of solute that needs to be weighed out for the known concentration and volume
2) Weigh out that mass and add to a beaker
3) Dissolve it in distilled water using a stirring rod and pour into a volumetric flask of the volume required
4) Rinse beaker with distilled water using a stirring rod and wash rinsings into the volumetric flask
5) Add distilled water until the bottom of the meniscus is exactly on the graduation line
6) Place the stopper on and invert the flask several times to mix the solution

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51
Q

How do you do acid-base titration calculations?

A
  • Write balanced equation to find molar ratio
  • Set up a table with volume, moles and concentration . Fill in given values
  • Work out the moles of known concentration solution. Use ratio to find the moles of the other solution
  • Calculate the unknown concentration (n = c x v)
52
Q

How do you carry out an acid-base titration?

A

1) Use pipette to add 25cm^3 of unknown concentration solution into a conical flask
2) Place flask on white tile and add several drops of indicator
3) Add the other solution to burette and record initial burette reading to nearest 0.05cm^3
4) Add solution from burette to the flask, swirling it to mix the contents
5) Colour of indicator changes at the end point. Record final burette reading to nearest 0.05cm^3
6) Calculate the total volume of solution added. This is trial titre (rough idea of volume needed to neutralise)
7) Repeat but add burette solution dropwise near the end point for an accurate titre
8) Repeat until you get concordant results (2 titres within 0.1cm^3)
9) Average the concordant titres for a mean titre

53
Q

What is an oxidation number?

A

Shows how many electrons that atom has lost/gained in a reaction.
Has a + or - followed by a number

54
Q

What are the rules when assigning oxidation numbers?

A
  • In a pure element, it is 0
  • In monatomic ions, oxidation number = ionic charge
  • In polyatomic compounds, sum of oxidation numbers = overall charge
  • In covalent compounds, more electronegative element keeps the negative oxidation number
55
Q

What are the oxidation numbers in CaSO4?

A

Ca = +2 (from periodic table)
O = -2 (from periodic table)
S = +6 (SO4 = 2-, O(-2) x 4 = -8, -2–8 = +6)

56
Q

What is a redox reaction?

A

A reaction with reduction and oxidation occurring together

57
Q

Describe the transfer of electrons in a redox reaction

A

Oxidation = loss of electrons
Reduction = gain of electrons
Total number of electrons lost in oxidation = total number of electrons gained in reduction

58
Q

Describe how oxidation numbers change in a redox reaction

A

Oxidation = increase in oxidation number
Reduction = decrease in oxidation number
Total increase in oxidation number in oxidation = total decrease in oxidation number in reduction

59
Q

What is disproportionation?

A

The same element is both oxidised and reduced in a reaction.

60
Q

What is oxidised and what is reduced in this equation?
Zn + 2HCl -> ZnCl2 + H2

A

Zn: oxidation number of 0 -> oxidation number of +2 = oxidation

H: oxidation number of +1 -> oxidation number of 0 = reduction

61
Q

What is oxidised and what is reduced in this equation?
Cl2 + 2NaOH -> NaClO + NaCl + H2O

A

Cl2 -> NaCl: Cl oxidation number from 0 to -1 = reduction

Cl2 -> NaClO: Cl oxidation number from 0 to +1 = oxidation

Disproportionation reaction

62
Q

What do half equations show?
Give the two half equations of this reaction:
Mg + CuSO4 -> MgSO4 + Cu

A

Show the electron transfer in redox reactions

Mg -> Mg(2+) + 2e(-) = oxidation
Cu(2+) + 2e(-) -> Cu = reduction

63
Q

Describe atomic shells

A

Electrons are contained in shells surrounding the nucleus.
Shells are known as energy levels where the energy increases with shell number.
Shell number = principal quantum number (n)
In each shell you can fit a max of 2n^2 electrons

64
Q

What is an atomic orbital?

A

A region in space where there is a high probability of finding an electron.
Each orbital can hold 2 electrons with opposite spins. There are different types with different shapes.

65
Q

What are the theoretical properties of an electron that relate to orbitals?

A

Can think of an electron as a negative-charge cloud with the shape of the orbital referred to as an electron cloud.
Wave-particle duality - have properties of both a wave and a particle.

66
Q

Describe s orbitals

A

Every shell from n=1 has an s orbital.
Spherical shape
Greater shell number = greater radius of orbital

67
Q

Describe p orbitals

A

3 separate orbitals at right angles to each other (px, py, pz)
Shaped like an infinity sign along each axis.
Greater shell number = further from nucleus

68
Q

Describe d and f orbitals

A

Each shell from n=3 contains 5 d orbitals

Each shell from n=4 contains 7 f orbitals

69
Q

What is a sub-shell?

What sub-shells and how many electrons are present in the first 4 shells?

A

Within a shell, orbitals of the same type are grouped as sub-shells

Shell 1: 1s, 2 electrons
Shell 2: 2s + 2p, 8 electrons
Shell 3: 3s + 3p + 3d, 18 electrons
Shell 4: 4s + 4p + 4d + 4f, 32 electrons

70
Q

What order are orbitals filled in?

What is the electron configuration of Calcium (20)?

What is the shorthand electron configuration of Potassium (19)?

A

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

1s2 2s2 2p6 3s2 3p6 4s2

[Ar] 4s1

71
Q

How do you find the shorthand electron configuration of an atom?

Why is it useful?

A

Expressed as the previous noble gas + the outer shell electrons

Emphasises similarities in electron configurations of outer shells, quicker to write

72
Q

How are orbitals filled with electrons? Why?

A
  • Each orbital can hold 2 electrons with opposite spins. Electrons repel each other (negatively charged). Have opposite spins to counteract repulsion
  • Orbitals with the same energy are occupied singly first. One electron occupies each orbital before pairing begins to prevent repulsion between paired electrons.
73
Q

What are the two types of ion?

How is the periodic table divided in terms of sub-shells?

A

Cations (+), Anions (-)

Divided into blocks depending on their highest energy sub-shell. E.g s block = highest energy electrons in the s sub-shell.

74
Q

How do you write the electronic configuration of ions?

What is the configuration of an Ni 2+ ion?
O 2-?

A

When forming ions, highest energy sub-shells lose/gain electrons. The 4s sub-shell fills + empties first (because once filled, 3d falls below the 4s energy level)

Ni 2+: 1s2 2s2 2p6 3s2 3p6 3p8

O 2-: 1s2 2s2 2p6

75
Q

What is ionic bonding?

Describe the structure of an ionic compound

A

The electrostatic attraction between positive and negative ions.

Each ion attracts oppositely charged ions in all directions, forming a giant ionic lattice structure, with the number of ions determined by the size of the crystal.

76
Q

What are the properties of an ionic compound?

A
  • High melting/boiling points (solid at room temp)
  • Most are soluble
  • Non-conductors (electrical + heat) when solid
  • Conductors when molten or aqueous
77
Q

Why do ionic compounds have high melting points?

How do melting points differ?

A

High temps are needed to provide the large amount of energy needed to overcome the strong electrostatic attraction between ions

Lattices with ions with greater ionic charges have higher melting points (stronger attraction)
A smaller ionic radius = higher melting point because the nucleus is closer to the oppositely charged ion, so a stronger attraction

78
Q

Explain the solubility of ionic compounds

A

Most dissolve in polar solvents (e.g water).
Polar molecules break down the lattice and surround each ion (attracted to the ionic charges)
If ions have large charges, the ionic attraction may be too strong for water to break it down, making the compound less soluble

79
Q

Explain the conductivity of ionic compounds

A
  • Non-conductors when solid - ions in a fixed position in the giant ionic lattice so there are no mobile charge carriers
  • Conductors when molten or aqueous - solid ionic lattice has broken down, ions are free to move as mobile charge carriers
80
Q

What is covalent bonding?

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Overlap of atomic orbitals, each with 1 electron to give a shared pair of electrons

Bonded atoms have outer shells with the same electron structure as the nearest noble gas

81
Q

What do covalent bonds occur in?

What structures do they form?

A

-Non-metallic elements, compounds of non-metallic elements, polyatomic ions

  • Small molecules (e.g H2), giant covalent structures (e.g SiO2), charged polyatomic ions (e.g NH4 +)
82
Q

Between how many atoms does a covalent bond occur?

What are lone pairs of electrons?

A

They are localised (only between the shared electron pair and nuclei of the bonded atoms)

Paired electrons that are not shared

83
Q

What is the octet rule?

What are the exceptions?

A

Says that bonded atoms must have 8 outer shell electrons

  • Electron deficient - B and Be can have too few electrons because they have <4 electrons on the 2nd energy level so can’t reach a full outer shell
  • Electron rich with an expanded octet - P, S, Cl can have too many electrons because in the 3rd energy level, outer shell can hold 18 electrons
84
Q

What other types of covalent bond are there other than coordinate bonds?

A
  • Double bonds (electrostatic attraction between 2 shared electron pairs + nuclei of bonding atoms)
  • Triple bonds (same but with 3)
  • Dative (coordinate) bonds - shared pair of electrons has been supplied by only one atom. Usually from an original lone pair
85
Q

How do you draw displayed formulae of covalent bonds?

A
  • Single bonds connect atoms with one line, double bonds have 2, triple bonds have 3.
  • Dative bonds are shown by an arrow going from the atom that supplied the electrons to the other
  • Lone pairs are usually shown as dots next to that atom
86
Q

How do you draw dot-and-cross diagrams of covalent bonds?

A
  • Overlapping outer shells between atoms
  • Shared electrons in the overlap
  • One atom’s electrons represented by dots, other by crosses
87
Q

How do you draw dot-and-cross diagrams of ionic bonds?

A
  • Draw ions separately inside square brackets
  • Show outer shells only
  • One atom’s electrons represented by crosses, other by dots
  • Cations should have an empty outer shell, anions should have a full outer shell with a combination of dots and crosses
88
Q

What is average bond enthalpy (covalent bonding)?

A

A measurement of covalent bond strength
A larger value = stronger covalent bond

89
Q

What is the electron-pair repulsion theory?

A

Electrons have a negative charge so electron pairs repel one another

  • Electron pairs surrounding a central atom determines the shape of the molecule/ion
  • Electron pairs repel each other so they are arranged as far apart as possible
  • Arrangement of electron pairs minimises repulsion and holds bonded atoms in a definite shape
  • Different number of pairs = a different shape
90
Q

What three symbols do you use to display the 3D shape of molecules on paper?

A
  • Solid line = a bond on the plane of the paper
  • Solid wedge = a bond that comes out the plane of the paper
  • Dotted wedge = a bond that goes into the plane of the paper
91
Q

What is special about lone pairs in terms of the shape of molecules?

How do we consider double/triple bonds when determining shape?

A
  • Lone pairs are slightly closer to the central atom and occupy more space so they repel more strongly than a bonding pair

Multiple bonds are treated as a single bonding region

92
Q

Name the 7 different shapes of molecules

A
  • Linear
  • Non-linear
  • Trigonal planar
  • Pyramidal
  • Tetrahedral
  • Trigonal bipyramidal
  • Octahedral
93
Q

Describe the linear shape of molecules

A
  • Two bonding regions repel each other as far apart as possible.
  • Bond angle = 180 degrees

e.g O=C=O

94
Q

Describe the non-linear shape of molecules

A
  • Two bonding regions repel each other as far apart as possible but lone pair(s) reduce the bond angle further.
  • Bond angle = 104.5 degrees

e.g F
B
F F

95
Q

Describe the pyramidal shape of molecules

A
  • Three bonding regions repel each other as far apart as possible. Lone pairs repel more than bonding pairs, reducing bond angles by about 2.5 degrees
  • Bond angle = 107 degrees
    ..
    e.g N
    H H H
96
Q

Describe the tetrahedral shape of molecules

A
  • Four bonding regions repel each other as far apart as possible
  • Bond angle = 109.5 degrees

e.g H
C H
H H

97
Q

Describe the trigonal bipyramidal shape of molecules

A
  • Five bonding regions repel each other as far apart as possible
  • Bond angles = 120 degrees, 90 degrees

e.g F
F P F
F F

98
Q

Describe the octahedral shape of molecules

A
  • Six bonding regions repel each other equally as far apart as possible
  • Bond angle = 90 degrees

e.g F
F F
S
F F
F

99
Q

How do you work out the shape of ions?

A

Exactly the same as molecules but with square brackets and its charge

100
Q

What is electronegativity?

A

The ability of an atom to attract the shared pair of electrons in a covalent bond

101
Q

When does unequal attraction of the shared pair of electrons in a covalent bond occur?

A
  • Nuclear charges are different
  • Atoms are different sizes
  • Electron pair is closer to one nucleus than the other
102
Q

How is electronegativity measured?
What is the trend?

A
  • Measured by the Pauling scale, where a large Pauling electronegativity value means the atom is very electronegative
  • Electronegativity increases across and up the periodic table, where Fluorine is the most electronegative element
103
Q

Explain the trend in electronegativity on the periodic table

A

Increases as you go across and up because:
- As you go across, nuclear charge increases (greater attraction to shared pair)
- As you go across, atomic radius decreases (shared pair is closer to nucleus)
- As you go up, electron shielding decreases (greater attraction to shared pair)

104
Q

What makes a bond ionic, covalent or polar covalent?

A

Ionic bonds have a large difference in electronegativity (>1.8)
Covalent bonds have no difference in electronegativity (0)
Polar covalent bonds have a small difference in electronegativity (0 to 1.8)

105
Q

Describe what happens in non-polar bonds and when they occur

A

Bonded pair of electrons is shared equally between bonded atoms
Occurs when:
- bonded atoms are the same element
- Atoms have similar electronegativity values

= a pure covalent bond

106
Q

Describe what happens in polar bonds and when they occur

A

Bonded pair is shared unequally between bonded atoms.
Occurs when bonded atoms are different and have different electronegativity values

Have a partially positive and a partially negative atom

107
Q

What is a dipole?

A

The separation of opposite charges.
A dipole in a covalent bond doesn’t change so is called a permanent dipole

108
Q

What is a polar molecule?
When will molecules not be polar?

A

Polar molecules have a permanent dipole in one direction over the whole molecule

Symmetrical molecules are non-polar because bond dipoles cancel each other out

109
Q

Describe how polar solvents break down ionic lattices

A
  • Each polar end of the solvent molecule attracts oppositely charged ions
  • The ionic lattice breaks down as it dissolves
  • Solvent molecules surround the ions
110
Q

What are intermolecular forces?

A

Weak interactions between dipoles of different molecules.
Are largely responsible for physical properties like melting/boiling points while covalent bonds determine the identity and chemical reactions of molecules

111
Q

What are the three different types of intermolecular forces?

A
  • London forces (induced dipole-dipole interactions)
  • Permanent dipole-dipole interactions
  • Hydrogen bonds
112
Q

What are London forces?

A

Weak intermolecular forces that exist between all molecules, whether polar or non-polar. They are between induced dipoles of molecules. Are only temporary and are the weakest intermolecular force

113
Q

How do London forces occur?

A
  • Electrons constantly move randomly = a changing dipole in the molecule
  • At any instant, electrons can be more on one side than the other, creating an instantaneous dipole
  • Instantaneous dipole induces a dipole on a neighbouring molecule.
  • Induced dipole further induces dipoles on neighbouring molecules that attract each other.
114
Q

How does the number of electrons affect London forces?

A

More electrons =
- Larger instantaneous and induced dipoles
- Greater induced dipole-dipole interactions
- Stronger attractive forces between molecules

115
Q

What is a permanent dipole-dipole interaction?

A

Intermolecular forces that act between permanent dipoles in different polar molecules.
Stronger than London forces.
Substances with permanent dipole-dipole interactions also have London forces

116
Q

What is a hydrogen bond?

A

A special type of permanent dipole-dipole interaction. The strongest type of intermolecular force.
Between molecules containing a hydrogen atom bonded to an electronegative atom with a lone pair (O,N,F).
Act between a lone pair of electrons on the electronegative atom of one molecule and the hydrogen atom of another.

117
Q

What are water’s anomalous properties?

A
  • Solid ice is less dense than liquid water
  • Has a high melting / boiling point
  • Has relatively high surface tension and viscosity
118
Q

Explain why solid ice is less dense than liquid water

A
  • Each molecule has 2 lone pairs around oxygen and 2 hydrogen atoms so can form 4 hydrogen bonds.
  • These extend outwards, holding molecules apart and forming an open tetrahedral lattice full of holes. Bond angle around the H atom is almost 180 degrees
  • This decreases density when solid
119
Q

Explain why water has a high melting/boiling point

A
  • Water has hydrogen bonds as well as London forces between molecules
  • So a large amount of energy is needed to break the hydrogen bonds in the ice lattice
120
Q

What is a simple molecular substance?

A

Made up from simple molecules (small units containing a definite number of atoms with a definite molecular formula)

121
Q

How are simple molecules arranged when solid?

A

Form a simple molecular lattice
- Molecules held in place by weak molecular forces
- Atoms within each molecule bonded strongly by covalent bonds

122
Q

What are the properties of a simple molecular substance?

A
  • Low melting / boiling point
  • Non-polar molecules are usually soluble in non-polar solvents
  • Non-polar molecules are usually insoluble in polar solvents
  • Polar covalent substances are sometimes soluble in polar solvents
  • Don’t conduct electricity
123
Q

Explain why simple molecular substances have a low melting/boiling point

A
  • When the lattice is broken down, only the weak intermolecular forces break, not the strong covalent bonds
  • It takes little energy to overcome them
124
Q

Why are non-polar simple molecular substances usually soluble in non-polar solvents?

A
  • Intermolecular forces form between the molecules and the solvent
  • Interactions weaken intermolecular forces in the simple molecular lattice.
  • Intermolecular forces break and the compound dissolves
125
Q

Why are non-polar simple molecular substances usually insoluble in polar solvents?

A
  • There is little interaction between molecules in the lattice and solvent molecules
  • Intermolecular bonding within the polar solvent is too strong to be broken
126
Q

Describe the solubility of polar simple molecular substances

A
  • May dissolve in polar solvents
  • Polar solute molecules and polar solvent molecules can attract each other
  • Solubility depends on the strength of the dipole and can be hard to predict
  • Some compounds contain both polar and non-polar parts so can dissolve in both types of solvent
127
Q

Why do simple molecular substances not conduct electricity?

A
  • No mobile charged particles present
  • Nothing to complete an electrical circuit

Chemistry (4 decks)

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