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OCR AS/A Level Chemistry A > Module 2.1 Atoms and Reactions > Flashcards

Module 2.1 Atoms and Reactions Flashcards

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1
Q

15th C - The Greek atom

A

Greek philosopher Democritus developed the first idea of the atom

Suggested you could divide matter only a certain number of times - “átomos” meaning “indivisible”

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2
Q

Early 1800s - Dalton’s atomic theory

A

Atoms are tiny particles that make up elements

Atoms cannot be divided

All atoms of a given element are the same

Atoms of one element are different from those of every other element

Dalton used his own symbols to represent atoms

Dalton developed the first table of atomic masses

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3
Q

1897-1906 - Joseph John Thomson

A

He discovers electrons

Discovered “cathode rays” had a negative charge, could be deflected by a magnet and an electric field and had a v v small mass

Cathode rays were electrons

The idea that an atom could not be split any further (proposed by the ancient Greeks and Dalton) had been disproved

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4
Q

1897-1906 - Joseph John Thomson

A

Proposed that atoms were made up of negative electrons moving around a “sea” of positive charge - “plum pudding model”

He thought that the overall negative charge = the overall positive charge - an atom is neutral with no overall charge

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5
Q

1909-1911 Ernest Rutherford

A

Gold leaf experiment:
Directed a-particles towards a sheet of very thin gold foil. They measured any deflection of the particles.

What the results showed:
Most particles weren’t deflected at all

A small percentage of particles deflected at large angles

Few particles were deflected back towards the source

His new model (1911):
The positive charge of an atom and most of the mass are concentrated in the nucleus, in the centre

Negative electrons orbit the nucleus - just like how planets orbit the sun

The + and - charges must balance

Rutherford had disproved the plum pudding model - the expected result was that the a-particles would pass through undisturbed but the observed results were that a few particles were deflected - indicating a small concentrated positive charge (the nucleus)

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6
Q

1913 - Niels Bohr and Henry Moseley

A

Bohr thought electrons could only follow certain paths otherwise they would spiral into the nucleus - “planetary atom” - electrons orbited the central nuclear “sun” in “shells”

Bohr’s model helped explain some periodic properties e.g. spectra lines seen in the emission spectra, the energy of electrons at different distances from the nucleus

Henry Moseley discovered the link between x-ray frequencies and an element’s atomic number - however, he couldn’t explain this at the time

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7
Q

1918- Rutherford

A

Discovers the proton

Could explain Moseley’s finding that an atom’s atomic number was linked to x-ray frequencies

We now know that atomic no. = no. of protons in the atom

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8
Q

1923-1926

A

Louis de Broglie suggested that particles could have the nature of both a wave and a particle

Erwin Schrödinger suggested that an electron had wave-like properties in an atom. He also introduced the idea of atomic orbitals

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9
Q

1932- James Chadwick

A

Discovers neutrons

Observed a new type of radiation emitted from some elements

Uncharged particles - approx. same mass as a proton

These were neutrons

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10
Q

Modern day

A

Now thought that protons and neutrons are made up of even smaller particles called quarks

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11
Q

Current model of the atom

A

Protons and neutrons in the nucleus

Nucleus = centre of atom

Electrons orbit the nucleus in “shells”

Nucleus = absolutely tiny

Nucleus = extremely dense, accounts for almost all of the atom’s mass

Most of an atom is empty space

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12
Q

Relative mass and charge of the sub-atomic particles

A

Proton:
Mass = 1
Charge = +1

Neutron:
Mass = 1
Charge = 0

Electron:
Mass = 1/2000
Charge = -1

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13
Q

Isotopes

A

Different masses

Same no. of protons

Same no. of electrons

Different number of neutrons

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14
Q

Atomic number

A

Proton number

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15
Q

Mass number

A

Protons + neutrons

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16
Q

Relative molecular mass, Mr

A

Add together the relative atomic masses of each atom making up a molecule

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17
Q

Relative formula mass

A

Add together the relative atomic masses of each atom making up the formula unit

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18
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 about molecules

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19
Q

A mass spectrum shows

A

Relative or percentage abundance on the y-axis and mass-to-charge ratio on the x-axis

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20
Q

How is mass-to-charge ratio in all mass spectra shown as

A

m/z

m= mass
z= the charge of the ion (usually 1)
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21
Q

Atoms of metals in groups 1-13

A

Lose electrons

Form positive ions with the electron configuration of the previous noble gas in the periodic table

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22
Q

Atoms of non-metals in groups 15-17

A

Gain electrons

Form negative ions with the electron configuration of the next noble gas in the periodic table

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23
Q

Why do atoms of Be, B, C and Si not usually form ions?

A

It requires too much energy to transfer the outer shell electrons to form ions

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24
Q

Ammonium

A

NH4+

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25
Q

Hydroxide

A

OH-

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26
Q

Nitrate

A

NO3-

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27
Q

Carbonate

A

CO3^2-

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28
Q

Sulphate

A

SO4^2-

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29
Q

Avogadro’s constant

A

6.02 times 10^23 mol^-1

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30
Q

Number of moles

A

Mass divided by molar mass

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31
Q

Empirical formula

A

Write out the atoms present in the compound

Divide the amount of each element present by its molar mass (this gives you the molar ratio)

Divide these values by the smallest value to simplify the ratio

Ensure your answer is the simplest whole number ratio

Write out the empirical formula

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32
Q

At RTP, one mole of gas molecules occupies approximately

A

24.0 dm^3

24,000 cm^3

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33
Q

n=

A

v DIVIDED BY 24.0 (dm^3)

For gases

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34
Q

n=

A

v DIVIDED BY 24,000 (cm^3)

For gases

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35
Q

Gases are assumed to behave in an ideal way:

A

In continuous motion

Don’t experience any intermolecular forces

Exert pressure when they collide with each other and the container walls

All collisions are elastic - no K.E. lost

K.E. inc. with inc. temp.

Gas molecules are so small so any differences in sizes of different gas molecules can be ignored

36
Q

Ideal gas equation:

pV = nRT

A

p= pressure (Pascal, Pa. 1 atm =101325 Pa)

V= volume (Cubic metres, m^3. 1 m^3 = 1000 dm^3)

n= number of moles

R= gas constant (8.314 J mol^-1 K^-1)

T= temperature (Kelvin, K. 0*C = 273 K)

37
Q

Concentration tells you

A

How much solute is dissolved in a given amount of solvent

Unit: mol dm^-3

38
Q

n=

A

c TIMES BY v (in dm^3)

For solutions

39
Q

n=

A

c TIMES BY V DIVIDED BY 1000 (in cm^3)

For solutions

40
Q

In order to make a standard solution, you will need to..

A

Know the concentration and volume of the solution you need to make

Work out the amount, in mol, of the solute needed

Convert this amount of solute into a mass, in g, so you know how much to weigh out

41
Q

Step-by-step of how you make up a standard solution

A

Step 1: Accurately weigh the required mass of solute. Use a beaker/weigh boat, a balance and a spatula.

Step 2: Transfer the solute into a beaker. Thoroughly wash the weigh boat with distilled water. Add enough distilled water to the beaker to dissolve the solute. Use a glass rod to mix the solution until all the solute dissolves.

Step 3: Use distilled water to wash out the contents of the beaker and pour all the contents into a volumetric flask. Rinse the glass rod with distilled water too and add that to the volumetric flask.

Step 4: Fill up the volumetric flask with distilled water. Stop just before the line on the flask.

Step 5: Use a pipette to get up to the required volume. The bottom of the meniscus needs to touch the line on the flask. If you go over the line then you have to start again!

Step 6: Put the stopper in the flask. Invert the flask multiple times.

42
Q

Mass concentration are measured in the units

A

g dm^-3

43
Q

M refers to a solution with a concentration in

A

moles per cubic decimetre

44
Q

Giant structures

A

Ionic compounds

Giant covalent structures

All metals

Some non-metals e.g. carbon, silicon and boron

45
Q

Percentage yield

A

Actual amount in mol of product DIVIDED BY theoretical amount in mol of product

TIMES BY 100

46
Q

Atom economy

A

Molecular mass of the desired product DIVIDED BY the sum of the molecular masses of all the products

TIMES BY 100

47
Q

Sulfuric acid

A

H2SO4

48
Q

Hydrochloric acid

A

HCl

49
Q

Nitric acid

A

HNO3

50
Q

Ethanoic (acetic) acid

A

CH3COOH

51
Q

An acid is

A

A proton donor

52
Q

Strong acids

A

V good at giving up H+ ions

Dissociate fully or almost fully

53
Q

Weak acids

A

Not v good at giving up H+ ions

Only partially dissociate

54
Q

Common bases

A

Metal oxides and hydroxides

E.g. MgO, CuO, NaOH, Mg(OH)2

55
Q

Ammonia (base)

A

Ammonia = NH3

Amines = CH3NH2

56
Q

Magnesium hydroxide, Mg(OH)2 use

A

“Milk of magnesia”

Used for treating acid indigestion

57
Q

Calcium hydroxide, Ca(OH)2 use

A

Also known as lime

Used for reducing the acidity of acid soils

58
Q

A base

A

Is a proton, H+ acceptor

Bases neutralise acids

59
Q

Alkalis

A

Release OH-(aq) ions when they dissolve in water

60
Q

Amphoteric substances

A

Behave as acids and bases

61
Q

Salts

A

Ionic compounds

Has a cation, usually a metal or ammonium ion

Has an anion, derived from the acid

The formula of the salt is the same as that of the parent acid, except that the H+ ion has been replaced by the positive ion

62
Q

Salts can be produced by neutralising acids with bases (proton acceptors) e.g.

A

Carbonates

Metal oxides

Alkalis

63
Q

BASHO

A

Base + acid –> salt + water

64
Q

MASH

A

Metal + acid –> salt + hydrogen

65
Q

CASHOCO

A

Carbonate + acid –> salt + water + carbon dioxide

66
Q

Hydrous

A

The crystalline form containing water

67
Q

Anhydrous

A

The crystalline form containing no water

68
Q

Dot formula

A

Gives the ratio between the number of compound molecules and the number of water molecules within the crystalline structure

69
Q

In a titration you MUST know

A

The concentration of one of your solutions

70
Q

The unknown info in a titration may be

A

The conc. of the solution

A molar mass

A formula

The number of molecules of water of crystallisation

71
Q

How to carry out a titration

A
  1. Using a pipette, add a measured volume of 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 (end point). Measure the volume of solution added from the burette
  4. You now know the volume of one solution that exactly reacts with the volume of the other solution
72
Q

Methyl orange

A

Red in acid
Yellow in base
Orange is end point colour

73
Q

Bromothymol blue

A

Yellow in acid
Blue in base
Green is end point colour

74
Q

Phenolphthalein

A

Colourless in acid
Pink in base
Pale pink is end point colour*

*assuming that the aq base has been added from the burette to the aq acid. If the acid is added to the base, the titration is complete when the solution goes colourless.

75
Q

Oxidation number of uncombined element

A

0

76
Q

Oxidation number of oxygen in peroxides

A

-1

77
Q

Oxidation number of combined hydrogen

A

+1

78
Q

Oxidation number of combined hydrogen in metal hydrides

A

-1

79
Q

Oxidation number of simple ions

A

The charge on the ion

80
Q

Oxidation number of combined fluorine

A

-1

81
Q

Oxidation numbers in compounds

A

The sum of the oxidation numbers must equal the overall charge of 0

82
Q

Oxidation numbers in molecular ions

A

The sum of the oxidation numbers must equal the overall charge

83
Q

Transition elements form ions with

A

Different oxidation numbers

84
Q

Oxyanions

A

Negative ions that contain an element along with oxygen

85
Q

Oxidation

A

Gain of oxygen

Loss of electrons

An increase in oxidation number

86
Q

Reduction is

A

Loss of oxygen

Gain of electrons

A decrease in oxidation number

87
Q

In redox reactions of metals with acids

A

The metal is oxidised, forming positive metal ions

The hydrogen in the acid is reduced, forming the element hydrogen as a gas (H2)

OCR AS/A Level Chemistry A (7 decks)

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