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Chemistry > Module 2.1 - Atoms and Reactions > Flashcards

Module 2.1 - Atoms and Reactions Flashcards

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1
Q

What were the ideas about the atom in:

  • Greek times?
  • Early 1800s?
  • 1897-1906?
  • 1909-1911?
  • 1913?
  • 1918?
  • 1923-1926?
  • 1932?
  • Modern day?
A
  • Greek: Democritus first idea of atom, suggesting could divide matter only a certain number of times until you got to one that couldn’t split any further: atomos = indivisible
  • Early 1800s: John Dalton: atomic theory (atoms=tiny particles making up elements, can’t divide, all atoms of an element are the same and different from other elements)
  • 1897-1906: Joseph John Thompson: discovered electrons: negative charge, deflected both by magnets and electric fields, very small mass), disproved that atoms couldn’t be split further
  • 1909-1911: Ernest Rutherford: directed alpha particles to sheet of thin gold foil: most not deflected, some were at large angles, few deflected towards source. +ve charge + mass in nucleus, which is v small, negative orbits nucleus like plants + sun, +ve + -ve balance
  • 1913: Niels Bohr + Henry Moseley: Bohr: electrons follow certain path or would spiral into nucleus. Bohr’s model explained some period properties, e.g. Spectral lines seen in emission spectra, energy of electrons at different distances from nucleus. Moseley: link between x ray frequency + atomic number
  • 1918: Rutherford discovered protons, could explain why atomic number was linked to x ray frequency
  • 1923-1926: Louis de Broglie: particles could have nature of wave + particle. Erwin Schrodinger: electron had wave like properties. Introduced idea of atomic orbitals
  • 1932: James Chadwick: observed a new type of radiation, made up of uncharged particles with same mass as a proton: neutrons
  • Modern day: protons + neutrons made up of smaller particles: quarks.
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2
Q

What is the current model of the atom?

A
  • Protons (positive) + neutrons found in the nucleus, in the centre of the atom
  • Electrons (negative) orbit nucleus in shells
  • Nucleus is tiny compared to total volume of atom
  • Nucleus is extremely dense + accounts for almost all the atom’s mass
  • Most of an atom consists of empty space between the tiny nucleus + electron shells
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3
Q

What does an atom have the same number of?

A

Protons and electrons

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4
Q

What is the relative charge and mass of a proton, neutron and electron?

A

Proton: mass = 1, charge = +1
Neutron: mass = 1, charge = 0
Electron: mass = 1/2000, charge = -1

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5
Q

How do isotopes of the same element compare to each other?

A
  • Different masses (different mass number)
  • Same number of protons + electrons
  • Different number of neutrons
  • Same atomic number
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6
Q

What does atomic number, Z, show?

A

The number of protons in the nucleus

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7
Q

What does the mass number, A, show?

A

The number of protons and neutrons in the nucleus

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8
Q

How do the reactions of isotopes compare to the reactions of atoms?

A

React in the same way as they have the same number of electrons which are what are involved in the reaction (neutrons do not affect the reaction)

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9
Q

What is the international standard for the measurement of relative mass?

A

Carbon-12

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10
Q

What are atomic masses measured in?

A

Unified atomic mass unit, u (1 u = 1.660540210 x 10^-27 kg)

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11
Q

What is the mass of a carbon-12 atom in unified atomic mass unit?

A

12 u

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12
Q

What is the relative isotopic mass for oxygen-16?

A

16

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13
Q

What is the relative isotopic mass for sodium-23?

A

23

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14
Q

What assumptions have to be made when calculating relative mass?

A
  • Neglected the tiny mass of the electrons

- Assumed the mass of both protons and neutrons as 1u

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15
Q

What factors influence the atomic mass of an element?

A
  • Relative mass of the isotope

- Percentage abundance of the isotope

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16
Q

When do you use relative formula mass and relative molecular mass?

A

Relative molecular mass: simple molecules

Relative formula mass: giant molecules

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17
Q

What is mass spectrometry used for?

A
  • Identify an unknown compound
  • Find the relative abundance of each isotope of an element
  • Determine structural information
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18
Q

How does mass spectrometry measure the mass of a molecule or isotope?

A

Measures the mass to charge ratio

  • Causes substances to become positive ions
  • Positive ions then passed through apparatus and separated according to their mass and charge
  • Computer within the mass spetrometer analyses data about the ions present + produces a mass spectrum
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19
Q

How do you interpret a mass spectrum?

A
  • y axis shows the relative abundances

- x axis shows the mass to charge ratio (m/z)

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20
Q

How do you interpret a mass spectrum with no values on the y axis?

A

Measure the length of each line, divide each length by the length of all the lines then x100 to get a percentage

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21
Q

How do you work out relative atomic mass?

A

Multiply each relative abundance by its corresponding relative isotopic mass. Add them all together and divide by the sum of all the relative abundances

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22
Q

How do metals in groups 1-13 form ions?

A
  • Lose electrons

- Form positive ions (cations) with the electron configuration of the previous noble gas in the periodic table

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23
Q

How do non-metals in groups 15-17 form ions?

A
  • Gain electrons

- Form negative ions (anions) with the electron configuration of the next noble gas in the period table

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24
Q

How does Be, B, C and Si form ions?

A

Don’t usually form ions as it requires too much energy to transfer the outer shell electrons to form an ion

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25
Q

What is a molecular ion?

A

Groups of covalently bonded atoms that gain or lose electrons to form ions

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26
Q

What is the overall charge of an ionic compound?

A

0

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27
Q

What is the value of Avogadro’s constant?

A

6.02 x 10 ^23

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28
Q

What is the amount of substance measured in?

A

Moles, mol

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29
Q

What is molar mass measured in?

A

g mol^-1

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30
Q

What is the equation linking moles, molar mass and mass?

A

moles (mol) = mass (g) / molar mass (g mol^-1)

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31
Q

How do you calculate the empirical formula of a compound?

A
  • Divide the amount of each element by its molar mass to give a molar ratio
  • Divide each of the answers by the smallest molar ratio, to ensure the ratio is 1:x
  • If necessary, multiply the ratio up to get it into a whole number ratio
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32
Q

What compounds is an empirical formula mostly used for?

A

Giant structures e.g. Ionic compounds like NaCl or giant covalent compounds like SiO2

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33
Q

What compounds are molecular formulae used for?

A

Compounds that exist as simple molecules

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34
Q

What is the empirical formula and molecular formula for butane?

A

Empirical formula: C2H5

Molecular formula: C4H10

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35
Q

How do you work out the molecular formula of a compound?

A
  • Calculate the empirical formula
  • Calculate how many times the empirical formula’s molar mass goes into the Mr
  • Multiply by the empirical formula
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36
Q

What is Avogadro’s law?

A

One mole of any gas occupies 24.0dm^3 (24000cm^3) at room temperature and pressure
The volume per mole of gas molecules is 24.0dm^3 mol^-1

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37
Q

What is the volume per mole of gas molecules also known as?

A

Molar gas volume

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38
Q

Why does air consist largely of oxygen and nitrogen?

A

Have similar relative molecular masses, e.g. H2 would be much lighter than air. CO2 would be much easier than air

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39
Q

How do you convert between moles of a gas and its volume?

A 
n = V (dm^3) / 24.0 
n = V (cm^3) / 24000
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40
Q

What is the ideal gas equation?

A

pV=nRT

p=pressure (Pa), V=volume (m^3), n=moles (mol), R=gas constant (8.314 J mol^-1 K^-1), T= temperature (K)

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41
Q

What is the ideal way in which gases are assumed to behave?

A
  • In continuous motion and don’t experience intermolecular forces
  • Exert pressure when they collide with each other + walls of container
  • Collisions between gas molecules and other gas molecules/container walls are elastic (do not cause kinetic energy to be lost)
  • KE of gases increases with increasing temperature
  • Gas molecules are so small compared to the size of the container they’re in that different sizes of different gas molecules can be ignored
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42
Q

What is 1 atm in Pa?

A

101325 Pa

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43
Q

What is 1m^3 in dm^3?

A

1000 dm^3

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44
Q

What is 0ºC in K?

A

273K

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45
Q

What is the equation linking concentration, moles and volume?

A 
n = c x (V in cm^3 / 1000)
n = c x (V in dm^3)
46
Q

What is a standard solution?

A

A solution of known concentration. Normally used in titration to determine unknown information about another substance

47
Q

How do you work out the mass of solvent, in g, needed to make the standard solution?

A
  • Use a known concentration and volume of the standard solution you need to make to work out the mol of solute needed
  • Convert mol to mass using the Mr
48
Q

How do you make a standard solution?

A
  1. Weight out the mass of the solute (mass of weighing about and solute - mass of weighing boat = mass of solute OR set to 0 when weight boat put on)
  2. Completely dissolve the solute in solvent in a beaker. Transfer solution to the flask + rinse the beaker repeatedly, using more solvent, adding the rinsing to the flask
  3. Add solvent to the flask, but don’t fill up to the graduation line
  4. Carefully add solvent drop by drop up to the line on the flask, until the bottom of the meniscus sits exactly on the graduation mark on the flask. If the solution goes over the meniscus line, you must throw it away and start again
  5. Finally, you must mix your solution thoroughly, by inverting the flask several times
49
Q

What does it mean if a solution is concentrated?

A

A large amount of solute per dm^3 (e.g. Greater than 10 mol dm^-3)

50
Q

What does it mean if a solution is dilute?

A

Small amount of solute per dm^3 (e.g. Normal bench solutions at 1 or 2 mol dm^3)

51
Q

What is the unit Molar, M, used to refer to?

A

-Solution with a concentration in mol dm^-3 (moles per cubic decimetre)
2 mol dm^-3 = 2M
Should quote concentrations in mol dm^-3

52
Q

What is the species in a reaction?

A

Type of particle that takes place in a reaction e.g. Atom, ion, molecule, empirical formula, electron

53
Q

Give some examples of metals with giant structures.

A
  • All metals

- Some non-metals e.g. Carbon, silicon, boron

54
Q

How are compounds with giant structures shown in an equation?

A

Empirical formula

55
Q

How are elements with giant structures shown in an equation?

A

Just by its symbol (effectively the empirical formula)

56
Q

Why is the percentage yield rarely 100%?

A
  • Reaction may be at equilibrium + may not go to completion
  • Side reactions may occur, leading to by-products
  • May be impurities in reactants
  • Some of the reactants/products may be left behind in the apparatus used in the experiment
  • Separation + purification may result in loss of some of the product
57
Q

How do you calculate percentage yield?

A

Actual amount of product (mol) / theoretical amount of product (mol) x100

58
Q

In terms of waste, if a reaction has a high atom economy, what does this mean?

A

Little is wasted

59
Q

How do you work out atom economy?

A

Molecular mass of the desired product / su of molecular masses of all products x 100

60
Q

What is an addition reaction and what is its atom economy?

A

When two reactants join together to make one product e.g. Propene + bromine –> 1,2-dibromopropane
Addition reactions always have an atom economy of 100%

61
Q

What kind of reactions have an atom economy of less than 100%?

A

Reactions involving substitution or elimination

62
Q

What pH do acids give in water?

A

Less than 7.0 (as H+ ions no longer bonded)

63
Q

What are the three main acids and their formulae?

A
  • Sulfuric acid, H2SO4
  • Nitric acid, HNO3
  • Hydrochloric acid, HCl
64
Q

What is the formula for ethanoic acid?

A

CH2COOH

65
Q

Using equations as an example, explain what happens when an acid is added to water.

A
  • Acid releases H+ ions (protons) into the solution
  • HCl(g) –> H+(aq) + Cl-(aq)
  • H2SO4(aq) –> H+(aq) + HSO4-(aq)
  • H+ ion is responsible for all acid reactions
  • Acid is a proton donor
66
Q

How do weak and strong acids differ in their ability to give up and accept H+ ions?

A
  • Strong: very good at giving up. Fully or almost fully dissociate. Poor at gaining H+ ions
  • Weak: poor at giving up H+ ions. Very good at accepting protons. Only partially dissociate. Once H+ ions released, they’re quickly taken back
67
Q

What are some examples of a base?

A
  • Metal oxides, MgO, CuO
  • Metal hydroxides, Mg(OH)2, NaOH
  • Ammonia, NH3, and its organic compounds, amines, CH3NH2
68
Q

What are some examples of bases in everyday life?

A
  • Mg(OH)2, milk of magnesia, treating acid indigestion

- Ca(OH)2, lime, reducing acidity of acid soils

69
Q

What is a base?

A
  • Proton H+ acceptor

- Base neutralises acid

70
Q

What is an alkali?

A
  • A substance that gives a pH greater than 7.0 when dissolved in water
  • Releases OH- ions when dissolved in water
71
Q

What are some common alkalis and their formulae?

A
  • Sodium hydroxide, NaOH
  • Potassium hydroxide, KOH
  • Ammonia, NH3
72
Q

Which are more corrosive, acids or alkalis?

A

Alkalis

73
Q

What is the ionic equation for the reaction of an alkali in water?

A

H+(aq) + OH-(aq) –> H2O (l)

74
Q

Give an equation to show that ammonia acts as a base.

A

NH3(aq) + H2O(l) (equilibrium sign) NH4+(aq) + OH-(aq)

Ammonia is a weak base as only small proportion of the dissolved NH3 reacts with water, shown by the equilibrium sign

75
Q

What is an amphoteric substance?

A

Substance that behaves like both an acid and a base

76
Q

Give an example of an amphoteric substance.

A
  • Amino acid molecule
  • e.g. glycine (has a carboxyl group, COOH, which is able to donate a proton, and an amino group, NH2, which could accept a proton)
77
Q

What are the features of a salt?

A
  • Positive ion (cation) in a salt usually a metal or ammonium ion (NH4+)
  • Negative ion (anion) in a salt derived from an acid
  • Formula of a salt the same as the parent acid, except H+ ion has been replaced by a positive ion
78
Q

What is a diprotic acid?

A

Has 2 H+ ions that can be replaced by a different positive ion

79
Q

Describe the formation of an acid salt, and describe how it behaves.

A
  • H2SO4–>NaHSO4 (sodium hydrogensulfate)

- Behaves like an acid as the other H+ ion can be replaced by a cation

80
Q

How do alkalis react with acids?

A

alkali + acid –> salt + water

81
Q

What happens when ammonia reacts with acids?

A
  • Ammonium salts formed

- e.g. NH3(aq) + HNO3(aq) –> NH4NO3(aq)

82
Q

What is a hydrated compound?

A

The crystalline form containing water molecules

83
Q

What is an anhydrous compound?

A

The form of a substance that contains no water molecules

84
Q

How does a compound get water in the form water of crystallisation?

A

If a compound crystallises within water, the water becomes part of the resulting crystalline structure

85
Q

What do you need to know to carry out an acid-base titration?

A

The concentration of one of the solutions (usually a standard solution)

86
Q

What can you be asked to find out from an acid-base titration?

A
  • Concentration of the unknown solution
  • A molar mass
  • A formula
  • Number of molecules of water of crystallisation
87
Q

How do you carry out a titration?

A
  1. Using a pipettes, add a measured volume f one solution to a conical flask. Add a suitable indicator.
  2. Place the other solution in a burette.
  3. Add the solution in the burette to the solution in the conical flask until the reaction has just been completed - the end point of the titration. Measure the volume of the solution added from the burette.
  4. You now know the volume of one solution that exactly react with the volume of the other solution.
88
Q

What indicators can be used in an acid-base titration?

A
  • Methyl orange
  • Bromothymol blue
  • Phenolphthalein
89
Q

How does the colour of methyl orange change in an acid-base titration?

A

Acid: red
Base: yellow
End point: orange

90
Q

How does the colour of bromothymol blue change in an acid-base titration?

A

Acid: yellow
Base: blue
End point: green

91
Q

How does the colour of phenolphthalein change in an acid-base titration?

A

Acid: colourless
Base: pink
End point: pale pink (base added from the burette to the acid, if acid added from burette to base the titration if complete when the solution goes colourless)

92
Q

What is always the first step when calculating unknowns in an acid-base titration?

A

Write the balanced equation (to work out the ratio of moles)

93
Q

Describe an investigation to determine the dot formula of a hydrated salt.

A
  • Weigh the mass of the hydrated salt (either mass of both boat and salt or difference of boat and boat and salt)
  • Heat on a Bunsen burner
  • Weigh and then heat further, reweighing until there is no further change in mass after heating - this is when all the water of crystallisation has evaporated
  • Weigh the mass of the anhydrous salt, and then calculate the change of mass from the hydrated salt to calculate the mass of water
  • Calculate the moles of the anhydrous salt and the water, and multiply up into a simple ratio
94
Q

What is the oxidation number of uncombined elements, e.g. O2, C?

A

0

95
Q

What is the oxidation number of combined oxygen, e.g. H2O?

A

-2

96
Q

What is the oxidation number of combined oxygen in peroxides, e.g. H2O2?

A

-1

97
Q

What is the oxidation number of combined hydrogen, e.g. NH3?

A

+1

98
Q

What is the oxidation number of combined hydrogen in metal hydrides, e.g. LiH?

A

-1

99
Q

What is the oxidation number of a simple ion, e.g. Mg2+, Cl-?

A

The charge on the ion

100
Q

What is the oxidation number of combined fluorine, e.g. NaF?

A

-1

101
Q

What is the sum of the oxidation numbers in a compound?

A

0

102
Q

What is the sum of the oxidation numbers in a molecular ion?

A

Must equal the overall charge of the ion

103
Q

When might you see oxidation numbers in chemical names?

A
  • Transition metals

- Oxyanions

104
Q

What is the formula and oxidation numbers of iron (III) chloride and copper (I) oxide?

A
  • FeCl3, Fe: +3, Cl: -1

- Cu2O, Cu: +1, O: -2

105
Q

What is an oxyanion?

A

Negative ions hat contain an element along with oxygen

106
Q

What is the formula and oxidation numbers of the following oxyanions:

  • nitrate (III)-?
  • nitrate (V)-?
  • sulfate (IV)2-?
  • sulfate (VI)2-?
A
  • NO2-; N: +3, O: -2
  • NO3-; N: +5, O: -2
  • SO3 2-; S: +4, O: -2
  • SO4 2-; S: +6, O: -2
107
Q

What is oxidation?

A
  • Gain of oxygen
  • Increase in oxidation number
  • Loss of electrons
108
Q

What is reduction?

A
  • Loss of oxygen
  • Decrease of oxidation number
  • Gain of electrons
109
Q

What is a redox reaction?

A

A reaction in which both oxidation and reduction occurs (metal tends to be oxidised by losing electrons and the non metals tend to be reduced by gaining electrons)

110
Q

Describe the redox reaction of magnesium and chlorine.

A

Mg + Cl2 –> MgCl2
Mg has been oxidised as the oxidation number has increased from 0 to +2
Chlorine has been reduces as the oxidation number has decreased from 0 to -1
Mg –> Mg2+ + 2e-
Cl2 + 2e- –> 2Cl-

111
Q

Describe the redox reaction of metal with an acid.

A

-Metal is oxidised, forming positive metal ions
-Hydrogen in the acid is reduced, forming hydrogen gas
-metal + acid –> salt + hydrogen
E.g. Mg + 2H+ –> Mg2+ + H2

Chemistry (13 decks)

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