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physical chemistry 1 Flashcards

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1
Q

what was the order of discovery for the history of the atom? (9 people) (atomic structure)

A
  • Robert Boyle (1661)
  • John Dalton (1803)
  • Henri Becquerel (1896)
  • J.J. Thompson (1897)
  • Ernest Rutherford (1911)
  • Neils Bohr (1913)
  • Erwin Schrödinger (1926)
  • James Chadwick (1932)
  • Gilbert Lewis (1932)
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2
Q

what was Robert’s Boyle’s contribution to the history of the atom? (atomic structure)

A
  • 1661
  • proposed that there were some substances that couldn’t be made simpler
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3
Q

what was John Dalton’s contribution to the history of the atom? (atomic structure)

A
  • 1803
  • suggested the elements were made of ‘invisible’ atoms & that they couldn’t be broken down
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4
Q

what was Henri Becquerel’s contribution to the history of the atom? (atomic structure)

A
  • 1896
  • discoverer radioactivity showing particles could come from inside an atom
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5
Q

what was J.J. Thompson’s contribution to the history of the atom? (atomic structure)

A
  • 1897
  • plum pudding model
  • discovered the electrons were negative in a positive mass
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6
Q

what was Ernest Rutherford’s contribution to the history of the atom? (atomic structure)

A
  • 1911
  • alpha particle scattering experiment
  • showed there was a positive nucleus
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7
Q

what was Neils Bohr’s contribution to the history of the atom? (atomic structure)

A
  • 1913
  • proposes that there were electrons that orbited shells at fixed levels
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8
Q

what was Erwin Schrödinger’s contribution to the history of the atom? (atomic structure)

A
  • 1926
  • quantum mechanics
  • electrons had properties of waves as well as properties of particles
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9
Q

what was James Chadwick’s contribution to the history of the atom? (atomic structure)

A
  • 1932
  • discovered the neutron
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10
Q

what was Gilbert Lewis’ contribution to the history of the atom? (atomic structure)

A
  • 1932
  • proposes the inertness of novel gases is due to their full outer shell
  • atoms bond by sharing electrons (to get a full outer shell)
  • ions are formed by atoms gaining/losing electrons to get a full outer shell
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11
Q

what is the relative mass & charge of a proton? (atomic structure)

A
  • relative mass = 1
  • relative charge = +1
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12
Q

what is the relative mass & charge of a neutron? (atomic structure)

A
  • relative mass = 1
  • relative charge = 0
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13
Q

what is the relative mass & charge of an electron? (atomic structure)

A
  • relative mass = 1/1840
  • relative charge = -1
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14
Q

what are nucleons & where are they found? (atomic structure)

A
  • protons & neutrons
  • found in the nucleus
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15
Q

what are the forces in the atomic structure? (atomic structure)

A
  • strong nuclear forces between protons & neutrons
  • electrostatic forces of attractions hold electrons in orbit
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16
Q

why do we assume that there is the same number of protons & electrons in an atom? (atomic structure)

A
  • because atoms have no relative charge
  • opposing charges of electrons & protons cancel out so there must be the same number of each
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17
Q

what is an atomic number & what symbol is it represented by? (atomic structure)

A
  • the number of protons in an element
  • represented by symbol Z
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18
Q

what is the mass number of an element & what symbol is it represented by? (atomic structure)

A
  • number of protons & neutrons
  • always an integer
  • represented by symbol A
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19
Q

what is the definition of relative atomic mass? (atomic structure)

A
  • the weighted mass of an atom of an element relative to 1/12 the mass of a carbon-12 atom
  • measured on a scale where the mass of carbon-12 is exactly 12
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20
Q

what 3 things is the relative atomic mass also relative to? (atomic structure)

A
  • relative isotopic mass
  • relative formula mass
  • relative molecular mass
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21
Q

why does mass spectrometry need to happen in a vacuum? (atomic structure)

A
  • to prevent ions that are formed from colliding with molecules from the air
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22
Q

what are the stages of time-of-flight mass spectrometry? (atomic structure)

A
  • ionisation (electron impact or electrospray)
  • acceleration
  • flight tube
  • detection
  • data analysis
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23
Q

what is the process of electron impact ionisation? (5 things) (atomic structure)

A
  • sample is vaporised & high energy electrons are fired at it from an electron gun
  • electron gun = hot wire with a filament with a current running through it that emits a beam of high energy electrons
  • this knocks off an electron from each particle
  • forms a 1+ ion
  • traditional method is still used for elements & low Mr compounds
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24
Q

what is the equation for electron impact ionisation? (atomic structure)

A
  • X(g) + e- —> X+(g) + 2e-
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25
Q

what is the process of electrospray ionisation? (6 things) (atomic structure)

A
  • sample is dissolved in a volatile solution
  • fired through a fine hypodermic needle that is connected to the positive terminal of a high voltage power supply
  • produces tiny positively charged droplets that have lost electrons to the positive terminal
  • solvent evaporates from droplets into vacuum & droplets get smaller until they contain only a positively charged ion
  • used for high Mr compounds (e.g. proteins)
  • gains a proton
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26
Q

what is the equation for electrospray ionisation? (atomic structure)

A
  • X(g) + H+ —> XH+(g)
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27
Q

what is the acceleration process in TOF mass spectrometry? (atomic structure)

A
  • positive ions are attracted to a negatively charged plate & are accelerated towards it
  • ions are accelerated to the same kinetic energy
  • lighter & more highly charged particles achieve a higher speed
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28
Q

what is the process of ion drift (flight tube stage) in TOF mass spectrometry? (atomic structure)

A
  • ions pass through a slit in the negatively charged plate
  • this forms a beam
  • ions travel along the flight tube to the detector
  • lighter particles move faster & have a shorter time of flight
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29
Q

what is the process of detection in TOF mass spec? (atomic structure)

A
  • positive ions hit the negative plate & produce a current
  • size of current is proportional to the number of ions
  • lighter ions arrive first due to higher velocities
  • flight time is recorded
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30
Q

what is the data analysis process in TOF mass spec? (atomic structure)

A
  • the signal from the secretory is passed to a computer
  • this generates a mass spectrum
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31
Q

what are the 6 basic principles (equipment) of TOF mass spec? (atomic structure)

A
  • sample inlet
  • ionisation chamber
  • accelerating
  • drift region
  • receptor
  • mass spectrogram
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32
Q

how is the relative abundance of ions detected in TOF mass spec? (atomic structure)

A
  • the abundance is proportional to the size of the current produced
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33
Q

what is the formula for kinetic energy, with units? (atomic structure)

A
  • Ke = 1/2 x m x v squared
  • Ke = kinetic energy (J)
  • m = mass (kg)
  • v = velocity (m/s-1)
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34
Q

what is the formula that links velocity, distance & time, with i=units? (atomic structure)

A
  • v=d/t
  • v = velocity (m/s-1)
  • d = distance (m)
  • t = time (s)
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35
Q

what are the 6 main formulas for isotope ions with the same kinetic energy travelling down the same flight tube? (atomic structure)

A
  • KE (1) = KE (2)
  • 1/2 m (1) v (squared) (1) = 1/2 m (2) v (squared) (2)
  • 1/2 m (1) d (squared)/t (squared) (1) = 1/2 m (2) d (squared)/t (squared) (2)
  • m (1)/t (squared (1) = m (2)/ t (squared) (2)
  • t (squared) (2) = m (2)/ m(1) x t (squared) (2)
  • t (2) = square root - m (2)/m (1) x t (1)
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36
Q

how do you calculate relative atomic mass? (atomic structure)

A
  • (abundance 1 x isotope 1) + (abundance 2 + isotope 2)/100
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37
Q

what are the different electron orbitals? (atomic structure)

A
  • s
  • p
  • d
  • f
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38
Q

what is the order in electron orbitals up to 6s? (atomic structure)

A
  • 1s - 3d
  • 2s - 4p
  • 2p - 5s
  • 3s - 4d
  • 3p - 5p
  • 4s - 6s
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39
Q

how many electrons can an s-orbital hold? (atomic structure)

A
  • 2 electrons
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40
Q

how many electrons can a p-orbital hold? (atomic structure)

A
  • 6 electrons
  • 2 for each but travel in groups of 3
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41
Q

how many electrons can a d-orbital hold? (atomic structure)

A
  • 10 electrons
  • 2 for each but travel in groups of 5
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42
Q

how many electrons can an f-orbital hold? (atomic structure)

A
  • 14 electrons
  • 2 for each but travel in groups of 7
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43
Q

what symbol represents electron configuration? (atomic structure)

A
  • arrows
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44
Q

define the term first ionisation energy (atomic structure)

A
  • the energy required to remove one mole of electrons from one mole of atoms in a gaseous state
  • stated in kJ mol-1
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45
Q

what is the trend for first ionisation energy across a period & why? (atomic structure)

A
  • first ionisation increases
  • increases due to increased nuclear attractions by the increasingly positive nucleus
  • nuclear charge increases & this makes it more difficult to remove an electron
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46
Q

are there any exceptions for increase in 1st ionisation energy across a period? (atomic structure)

A
  • Be & B - indicates evidence for a difference in energy levels in s & p orbitals
  • N & O - indicates the half-filled nature of p orbitals in nitrogen & the relative instability of pairing electrons
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47
Q

why does the first electron removed require the leat amount of energy? (atomic structure)

A
  • because it is being removed from a neutral atom
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48
Q

define the term successive ionisation energies (atomic structure)

A
  • the energy required to remove one mole of electrons from the previous ionisation state in their gaseous state
  • applies to second ionisation energy onwards
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49
Q

what are the first three ionisation energies for sodium (?) (atomic structure)

A
  • 1st IE = Na (g) —> Na+ (g) + e- (+496 kJ mol-1)
  • 2nd IE = Na+ (g) —> Na 2+ (g) + e- (+ 4563 kJ mol-1)
  • 3rd IE + Na 2+ (g) —> Na 3+ (g) + e- (+6913 kJ mol-1)
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50
Q

why would you expect a significant increase of ionisation energy between the fourth ionisation energy? (atomic structure)

A
  • because it requires a full set of orbitals to be broken apart so the electrons can be removed
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51
Q

what happens to the ionisation energy between Mg & Al (?) (atomic structure)

A
  • decreases despite an increase in nuclear charge
  • because the outer electron in Al is inactive 3p orbital which is of a slightly higher energy than a 3s orbital
  • therefore it needs less energy to remove it
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52
Q

why is there a decrease in ionisation energy down group 2? (atomic structure)

A
  • because the outer electron is in a main level that gets further from the nucleus in each case
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53
Q

what happens to nuclear charge going down a group? (atomic structure)

A
  • increases
  • the actual positive charge ‘felt’ by an electron in the outer shell is less than the full nuclear charge
  • this is because of the effect of the inner electrons shielding the nuclear charge
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54
Q

what is relative atomic mass & how do you calculate it? (amount of substance)

A
  • the weighted average mass of an element relative to 1/12 the mass of a carbon-12 atom
  • average mass of one atom 1/12 the mass of one 12C atom
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55
Q

what is relative molecular/formula mass & how do you calculate it? (amount of substance)

A
  • average mass of one molecule compared to 1/12 the Ar of one carbon-12 atom
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56
Q

when is the term ‘relative formula mass’ used & why? (amount of substance)

A
  • used for ionic compounds because they don’t exist as molecules
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57
Q

what is the value of avogadro’s constant? (amount of substance)

A
  • 6.022 x 10 23
  • same number of atoms in 12g of carbon-12
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58
Q

what is a mole? (amount of substance)

A
  • the amount of substance that contains 6.022 x 10 23 particles
  • Ar of any element in g contains one mole of atoms
  • Mr of any substance in g contains one mole of entities
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59
Q

what is an entity? (amount of substance)

A
  • a general word for a particle
  • can refer to an atom, molecules, ion or electron
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60
Q

what do balanced equations tell us? (amount of substance)

A
  • how much of a substance reacts
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61
Q

How many cm cubed are in 1 dm cubed? (amount of substance)

A
  • 1000 cm cubed
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62
Q

how many dm cubed are in 1 m cubed? (amount of substance)

A
  • 1000 dm cubed
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63
Q

how many Pa are in 1 ATM? (amount of substance)

A
  • 100 000 Pa
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64
Q

what is the conversion number between ˚C & ˚K? (amount of substance)

A
  • 273
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65
Q

what are the units for STP (standard temperature & pressure)? (amount of substance)

A
  • 298K
  • 102 kPa (102 000 Pa)
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66
Q

what is the theory behind Boyle’s law? (amount of substance)

A
  • pressure & volume of a gas are linked
  • as you decrease the volume a gas is occupying you increase the pressure
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67
Q

what is the equation for Boyle’s law? (amount of substance)

A
  • pressure x volume = constant
  • P x V = constant
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68
Q

what is the theory behind Charles’ Law? (amount of substance)

A
  • so long as pressure remains constant, the volume of gas is proportional to the temperature
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69
Q

what is the equation for Charles’ law? (amount of substance)

A
  • volume/temperature = constant
  • V/T = constant
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70
Q

what is the constant in the ideal gas equation? (amount of substance)

A
  • given by symbol R
  • has a value of 8.31 JK-1 mol-1
71
Q

state the ideal gas equation (amount of substance)

A
  • PV = nRT
72
Q

what are the units for each element of the ideal gas equation? (amount of substance)

A
  • pressure (P) = Pa
  • volume (V) = m cubed
  • moles (n) = moles
  • constant (R) = JK-1 mol-1
  • temperature (T) = K
73
Q

define empirical formula (amount of substance)

A
  • the empirical formula is the formula that represents the simplest whole number ratio of the atoms of each element in a compound
74
Q

define molecular formula (amount of substance)

A
  • gives the actual number of atoms of each element in one molecule of the compound
  • it only applies to substances that exist as molecules
75
Q

what is stoichiometry? (amount of substance)

A
  • the whole number ratio in which reactants combine & produce products
76
Q

what is the calculation on for atom economy? (amount of substance)

A
  • mass of desired products/total mass of products x 100
77
Q

what does atom economy tell us? (amount of substance)

A
  • atom economy is theoretical & will never be achieved
  • it tells us in theory how many atoms must be wasted in a reaction
78
Q

what does percentage yield tell us? (amount of substance)

A
  • tells us about the practical efficiency of the process, how much is lost by the practical process of obtaining a product & as a result of reactions that do not go to completion
79
Q

give two ways you can click late percentage yield (amount of substance)

A
  • (no. mol of specified products/theoretical no. mol of products) x 100
  • (mass of theoretical product/theoretical mass of product) x 100
80
Q

what is an ionic equation? (atomic structure)

A
  • an equation that is written only using the ions in the equation
  • ions that exist on both sides of the equation are called spectator ions & are not involved in the reaction so cancel each other out
81
Q

what is ionic bonding? (bonding)

A
  • occurs between metals & non-metals
  • electrons are transferred from the metal atom to the non-metal atom
  • positive & negative ions are formed
82
Q

why do metals from positive ions? (bonding)

A
  • they lose a negative electron
83
Q

why do non-metal atoms from negative ions? (bonding)

A
  • because they gain a negative electron
84
Q

why are the ions in ionic bonds attracted to each other? (bonding)

A
  • they are oppositely charged so there are electrostatic forces of attraction
85
Q

what are the properties of ionic compounds? (5) (bonding)

A
  • always solids at room temperature
  • have giant structures & therefore have high melting points
  • conduct electricity when molten/aqueous
  • brittle & shatter easily
  • mostly soluble in water
86
Q

is the melting point of ionic compounds high or low & why? (bonding)

A
  • high
  • there are strong electrostatic forces of attraction between oppositely charged ions that require lots of energy to overcome
87
Q

what is the electrical conductivity of ionic compounds? (bonding)

A
  • only when molten/aqueous
  • when solid the ions can’t move
  • when liquid ions can move & carry a current
88
Q

what is the strength of ionic compounds? (bonding)

A
  • brittle
  • if you apply a force, like charges come together/align & repel, so it breaks into pieces
89
Q

what is the solubility of ionic compounds? (bonding)

A
  • mostly soluble in water
  • both ionic compounds & water molecules are polar (partially charged)
  • partial charges of water break apart the lattice by pulling oppositely charged ions apart
90
Q

what are some examples of ionic compounds? (bonding)

A
  • magnesium oxide (MgO)
  • sodium chloride (NaCl)
91
Q

what is covalent bonding? (bonding)

A
  • forms between pairs of non-metal atoms
  • atoms share outer electrons so both have a stable noble gas structure
  • is a shared pair of electrons
92
Q

how are the atoms in a covalent bond held together? (bonding)

A
  • by strong electrostatic forces of attraction (between oppositely charged ions) between the nuclei & shared electrons
93
Q

what is a double covalent bond? (bonding)

A
  • 4 electrons (2 pairs) are shared
  • e.g. O=O
94
Q

what are the properties of simple covalent bonds? (4) (bonding)

A
  • solids, liquids or gases with low melting points
  • poor conductors of electricity
  • mostly insoluble in water
  • brittle
95
Q

is the melting point of simple covalent molecules high or low & why? (bonding)

A
  • low
  • forces between the molecules are weak & not much energy is required to overcome them
96
Q

what is the electrical conductivity of simple covalent molecules? (bonding)

A
  • poor conductors of electricity
  • molecules are neutral
  • there are no charged particles to carry current
97
Q

what is the solubility of simple covalent molecules? (bonding)

A
  • mostly insoluble in water
  • due to the ability to form hydrogen bonds
  • if dissolve in water & remain as molecules, the solutions don’t conduct electricity
  • some exceptions (e.g. alcohol, ammonia & sugar)
98
Q

what is the strength of simple covalent molecules? (bonding)

A
  • brittle
  • weak forces between molecules are easy to break
99
Q

what are some examples of covalent molecules? (bonding)

A
  • hydrogen chloride (HCl)
  • hydrogen (H2)
  • water (H2O)
  • ammonia (NH3)
100
Q

what is metallic bonding? (bonding)

A
  • between two metals
  • metals can’t transfer electrons unless there is a non-metal to receive them
  • outer main levels merge so outer/delocalised electrons aren’t associated with a particular element
  • number of delocalised electrons depends on how many electrons have been lost by each atom
  • metallic bonding spreads throughout structures = giant structure
101
Q

what are the properties of metallic bonds? (4) (bonding)

A
  • high melting points
  • good conductors of heat & electricity
  • malleable & ductile
  • insoluble in water & break apart
102
Q

are the melting points of metallic substances high or low & why? (bonding)

A
  • high
  • strong electrostatic attractions between between positive metal ions & negative delocalised electrons that need lots of energy to overcome
103
Q

what is the electrical conductivity of metallic substances? (bonding)

A
  • good conductors of heat & electricity
  • delocalised electrons are free to carry current throughout the structure
104
Q

what is the strength of metallic compounds? (bonding)

A
  • malleable
  • regulate structure allows layers to slide over each other
  • ductile
105
Q

what is the solubility of metallic substances? (bonding)

A
  • insoluble in water
  • the electrostatic attractions between ions in the compound are much stronger than that of ions in the water
106
Q

what is a coordinate/dative covalent bond? (bonding)

A
  • a covalent bond in which both electrons come from the same atom
107
Q

give an example of a molecule with a coordinate bond (bonding)

A
  • carbon monoxide
  • nitric acid
108
Q

give the written formula for the coordinate bond in NH4 (bonding)

A
  • H
    I
    H — N —> H
    I
    H
  • coordinate bond is represented by the arrow
109
Q

what are the bonding pairs & bond angled if a molecule with a linear shape? (bonding)

A
  • bonding pairs = 2
  • bond angle = 180˚
110
Q

what are the bonding pairs & bond angle for a molecule with a trigonal planar shaped molecule? (bonding)

A
  • bonding pairs = 3
  • bond angles = 120˚
111
Q

what are the bonding pairs & bond angles for a tetrahedral shaped molecule? (bonding)

A
  • bonding pairs = 4
  • bond angles = 109.5˚
112
Q

what are the bonding pairs & bond angles for a molecule with a trigonal planar shape? (bonding)

A
  • bonding pairs = 5
  • bond angles = 90˚ & 120˚
113
Q

what are the bonding pairs & bond angles for a molecule with an octahedral shape? (bonding)

A
  • bonding pairs = 6
  • bond angles = 90˚
114
Q

what is a lone pair of electrons? (bonding)

A
  • a pair of electrons what is not involved in a (covalent) bond
115
Q

what is the process of working out the shape of a molecule? (bonding)

A
  • (a) find out the number of outer shell electrons on the central atom & adjust for charge
  • (b) find out the number of atoms bonded to the central atom (bonding pairs)
  • find out the number of lone Paris (1/2 the difference of a + b)
116
Q

what is electron pair repulsion theory? (bonding)

A
  • each pair of electrons around atom will repel all other electron pairs
  • the pairs of electron will therefore take up positions as far apart os possible to minimise repulsion
117
Q

define electronegativity (bonding)

A
  • the power of an atom to attract the electrons in a covalent bond
118
Q

define electron density (bonding)

A
  • the probability of electrons being found in a particular volume in space
119
Q

define ‘polar’ (bonding)

A
  • describes a molecule in which the charge in not symmetrically distributed so that one area is slightly more positively charged & another slightly negatively charged
120
Q

what 3 things does electronegativity depend on? (bonding)

A
  • the nuclear charge
  • the distance between the nucleus & the outer shell electrons
  • the shielding of the nuclear charge by electrons inner shells
121
Q

what is the electronegativity in smaller atoms & why? (bonding)

A
  • the smaller the atom, the closer the nucleus is to the shared outer main level electrons & the greater its electronegativity
122
Q

what is the electronegativity of atoms with a large nuclear charge? (bonding)

A
  • the larger the nuclear charge (for a given shielding effect), the greater the electronegativity
123
Q

what happens to electronegativity going up a group & why? (bonding)

A
  • electronegativity increases
  • the atoms get smaller & there is less shielding by electrons in inner shells
124
Q

what happens to electronegativity across a period & why? (bonding)

A
  • electronegativity increases
  • the nuclear charge increases, the summer of inner main levels remain the same & the number of atoms becomes smaller
125
Q

where are the most electronegative atoms found on the periodic table? (bonding)

A
  • top right hand corner
  • e.g. fluorine, oxygen, nitrogen
126
Q

what type of bond does no electronegativity difference lead to? (bonding)

A
  • a pure non-polar covalent bond
127
Q

what type of bond does a small electronegativity difference lead to? (bonding)

A
  • a polar covalent bond
128
Q

what type of bond does a large electronegativity difference lead to? (bonding)

A
  • an ionic bond
129
Q

what happens to electronegativity when two atoms are the same? (bonding)

A
  • the bond must be shared equally between the atoms
  • they have the same electronegativity
  • the bond is non-polar
130
Q

what happens to electronegativity when atoms are different? (bonding)

A
  • the atoms will not have the same electronegativity
  • the electrons in the bond will to be shared equally between the atoms
131
Q

what symbol shows if atoms are more or less electronegative? (bonding)

A
  • less electronegative = δ+
  • more electronegative = δ-
132
Q

what are intermolecular forces? (bonding)

A
  • the forces between molecules
133
Q

name the three types of intermolecular force (bonding)

A
  • van der Waals
    -dipole-dipole
  • hydrogen bonds
134
Q

what types of molecules do van der Waal forces occur in & are they found on their own? (bonding)

A
  • act between all atoms & molecules
  • present in all molecular substances
  • in additions to any other intermolecular
135
Q

What are van der Waal forces & how do they occur? (bonding)

A
  • electrons are constantly moving around & there will be an uneven distribution at any given time
  • this causes a temporary dipole within the molecules
  • this temporary molecules induces a temporary dipole in a neighbouring molecule
  • the bigger the molecule the stronger the van der Waals forces
136
Q

what can van der Waals forces help to explain about boiling points? (bonding)

A
  • why the boiling points in noble gases increase down the down group
  • why the boiling points of hydrocarbons increases with chain length
137
Q

when do dipole-dipole forces occur in molecules? (bonding)

A
  • present between polar molecules
  • act between molecules that have permanent dipoles
  • in addition to van der Waals forces
138
Q

what are dipole-dipole forces? (bonding)

A
  • whatever their starting points, the molecules will ‘flip’ to give an arrangement where the two molecules attract
  • in molecules with more than one polar bond, the effects of each bond may cancel & leave it no dipole moment
  • the effects may also add up & reinforce each other
139
Q

what is hydrogen bonding? (bonding)

A
  • when hydrogen is bonded between nitrogen, oxygen or fluorine
  • the polar bond between the hydrogen atom & the N/O/F atom leaves the hydrogen nucleus exposed as hydrogen only has one electron
  • therefore there is a strong attraction from the lone pair
140
Q

what are the 2 conditions for a hydrogen bond? (bonding)

A
  • there must be a hydrogen atom that is bonded to a very electronegative atom. This will produce a strong partial positive charge in the hydrogen atom
  • there must be a very electronegative atom with alone pair of electrons. These will be attracted to the partially charged hydrogen atom in another molecule & form the bond
141
Q

what happens to hydrogen bonds in the structure of ice? (bonding)

A
  • in the water, the hydrogen bonds break & reform easily as the molecules are moving around
  • when water freezes, the molecules are no longer free to move & the hydrogen bonds hold the molecules in fixed positions
  • the resulting 3D structure is similar to the structure of diamond
142
Q

why is ice less dense than water? (bonding)

A
  • to fit into its diamond like structure, the molecules are slightly less closely packed than in liquid water
  • this means that ice is less dense & forms on the tops on ponds (helps to insulate them & keep fish alive in winter)
  • ice has a lattice structure
  • water molecules are further apart & so have more space between space
143
Q

define thermochemistry (energetics)

A
  • the study of heat changes during chemical reactions
144
Q

what is an exothermic reaction? Give an example (energetics)

A
  • a reaction where energy is given out at the end of it
  • the neutralising of an acid with an alkali
145
Q

what is an endothermic reaction? Give an example (energetics)

A
  • a reaction where energy has been taken in at the end of it
  • e.g. the breakdown of limestone (calcium carbonate) into lime (calcium oxide)
146
Q

what happens to mass in a chemical reaction? (energetics)

A
  • mass in conserved in chemical reactions
147
Q

what happens to atoms & bonds when a chemical reaction takes place? (3) (energetics)

A
  • atoms are rearranged
  • to do this, bonds must be broken, which requires the activation energy
  • then new bonds can be formed, which releases energy
148
Q

what does the amount if heat given out/taken in during a chemical reaction depend on? (energetics)

A
  • the quantity of reactants
149
Q

what is the symbol for enthalpy change? (energetics)

A
  • ΔH
150
Q

define ‘enthalpy change’ & state its units (energetics)

A
  • the amount of energy (kJ) that is released or absorbed when one mole of a given substance reacts
  • units = kJmol-1
151
Q

what 2 things do you have to give when stating an enthalpy change? (energetics)

A
  • a balanced equation
  • state symbols (in the balanced equation)
152
Q

what are the standard conditions for measuring standard enthalpies? (energetics)

A
  • 298K (room temperature)
  • 100kPa
  • the state at which that substance would be in these conditions (solid/liquid/gas)
153
Q

at what point is a standard reaction thought to be over? (energetics)

A
  • when the products have cooled back to 298K (room temperature)
154
Q

what is the name given to the heat given out while a reaction mixture cools? (bonding)

A
  • the enthalpy change
155
Q

when are reactions that are carried out in aqueous solutions. Considered to be over & how does it reach this point? (energetics)

A
  • when the products have warmed up to the temperature at which they started
  • the solution takes in heat from the surroundings in order to become warmer
156
Q

how can pressure affect the amount of heat energy given out by reactions that involve gases? (2) (energetics)

A
  • the greater the atmospheric pressure, the more energy is used to push away the atmosphere
  • this means that less energy is left to be given out as heat by the reaction
157
Q

can the physical state of reactants/products affect enthalpy change? (energetics)

A
  • yes
158
Q

what do enthalpy level diagrams show? (energetics)

A
  • the relative enthalpy levels of the reactants & products in a reaction
159
Q

draw the enthalpy level diagram for an exothermic reaction (energetics)

A
  • y axis should be labelled ‘enthalpy’
  • x axis should be labelled ‘extent of reaction’
  • reactants should have more energy than products
  • ΔH should be negative
  • no curve/Ea
160
Q

draw the enthalpy level diagram for an endothermic reaction (energetics)

A
  • y axis should be labelled ‘enthalpy’
  • x axis should be labels ‘extent of reaction’
  • products should have more energy than reactants
  • ΔH should be positive
  • no curve/Ea
161
Q

define ‘standard enthalpy of combustion’ & state the symbol used for it (energetics)

A
  • the enthalpy change when one mole of substance is completely burnt with oxygen under standard conditions (all reactants & products are in their standard states)
  • ΔcH˚ (˚ = theta)
162
Q

define ‘standard enthalpy of formation’ & state the symbol used for it (energetics)

A
  • the enthalpy change when one mole of substance is formed from its constituent elements under standard conditions (all reactants & products are in their standard states)
  • ΔfH˚(˚ = theta)
163
Q

define ‘standard enthalpy of formation’ & state the symbol used for it (energetics)

A
  • the enthalpy change when one mole of substance is formed from its constituent elements under standard conditions (all reactants & products are in their standard states)
  • ΔfH˚(˚ = theta)
164
Q

should you sue fractions when balancing equations for standard enthalpies? Why/why not? (energetics)

A
  • yes
  • you can’t change the substance that you’re measuring the standard enthalpy of
165
Q

is there an enthalpy charge for incomplete combustion? (energetics)

A
  • no
166
Q

what 3 things do we need to know to measure enthalpy change? (energetics)

A
  • the mass of the substance that is being heated up/cooled down
  • the temperature change
  • the specific heat capacity of the substance
167
Q

define ‘specific heat capacity’ & state its units (energetics)

A
  • the amount of heat needed to raise the temperature of 1g of substance by 1K
  • units = Jg-1K-1
168
Q

what is the difference between heat & temperature? (energetics)

A
  • temperature is the average kinetic energy of the particles in a system
  • heat is the total energy of all particles present in a given amount of substance
169
Q

what is the equation for calculating enthalpy change? Give the units for each component (energetics)

A
  • q = McΔT
  • q = energy (J)
  • M = mass (g)
  • c = specific heat capacity (4.18 Jg-1K-1)
  • ΔT = temperature change (K)
170
Q

What do simple calorimeters do? Outline how to use one (energetics)

A
  • find the approximate enthalpy change when a fuel burns
  • can be used to compare the ΔcH values of a series of similar reactions as the errors for each experiment will be similar
  • burn the fuel to heat a known mass of water & then measure the temperature rise of the water
    -assume that all the fuel goes into the water
171
Q

what do flame calorimeters do & how are they designed to minimise heat loss? (energetics)

A
  • used for measuring the enthalpy changes of combustion
  • has a spiral chimney made of copper, an enclosed flame & fuel is burned in pure oxygen - reducing heat loss
172
Q

why are expanded polystyrene beakers used when measuring enthalpy changes in solution)

A
  • they are good insulators (reduces heat loss through sides) & have a low heat capacity (absorbs very little heat)
  • heat is generated in solutions & only has to be kept in the calorimeter
173
Q

how’d o you find the enthaply change in neutralisation reactions? (energetics)

A
  • reaction is exothermic so gives out heat
  • to find enthaply change, use the quantities in moles given by the balanced equation
  • use the total volume of mixture
174
Q

how can you measure enthalpy change in a displacement reaction? (energetics)

A
  • if a compound produced dissolves in water, the reaction can be investigated using a polystyrene beaker

chemistry 🧪 (12 decks)

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