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GCSE chemistry > Paper 1 > Flashcards

Paper 1 Flashcards

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1
Q

Why do isotopes have the same chemical properties?

A

This is because they contain the same number of electrons. Electrons determine chemical properties of atoms.

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2
Q

John Dalton

A

John Dalton described atoms as tiny spheres that could not be divided

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3
Q

JJ Thomspson

A
  • J.J. Thomson discovered the plum pudding model
  • It is a ball of positive charge with electrons embedded in the ball of positive charge.
  • No empty space in the atom
  • Mass spread throughout
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4
Q

Ernest Rutherford:

A
  • Ernest Rutherford’s alpha scattering experiment led to the plum pudding model being replaced by nuclear model.
  • Positive charge concentrated at the centre
  • Most of the atom is empty space
  • Mass is concentrated at the centre
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5
Q

Niels Bohr

A
  • Niels Bohr adapted the nuclear model by suggesting that the electrons orbit the nucleus at specific distances.
  • The electrons are on electron shells/energy level.
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6
Q

Later experiments:

A
  • Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge.
  • The name proton was given to these particles.
  • James Chadwick provided the evidence for the existence of neutrons in the nucleus.
  • This was about 20 years after the nucleus became an accepted scientific idea.
  • This could help explain the existence of isotopes.
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7
Q

Modern periodic table

A
  • Arranged in order of atomic number, elements with similar properties are in columns, known as groups.
  • Periodic table- siimilar properties occur at regular intervals.
  • Elements in same group - same number of electrons in their outer shell- gives them similar chemical properties.
  • Elements in the same period have the same number of electron shells
  • As you go down the group, the number of shells increases.
  • Metals are found on the left-hand side of the periodic table whilst non-metals on the right-hand side
  • Most of the elements found on the period table are metals
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8
Q

Properties of metals:

A
  • High density
  • High melting point
  • Shiny
  • Malleable
  • Conducts electricity and heat
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9
Q

Properties of non-metals:

A
  • Low density
  • Low melting point
  • Dull
  • Brittle
  • Does not conduct electricity nor heat
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10
Q

What were early periodic tables arranged in order of?

A

In order of atomic weight.

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11
Q

Newland’s periodic table:

A
  • Have more than one element in a box
  • Missing Noble gases
  • Arranged elements in order of atomic weight

DISADVANTAGES:
* He did not leave gaps for undiscovered elements
* He had many dissimilar elements in a column

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12
Q

Mendeleev’s periodic table:

A
  • Have more than one element in a box
  • Missing Noble gases
  • Arranged elements in order of atomic weight

Advantages of Mendeleev
* He left gaps for undiscovered elements
* He changed the order of some elements

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13
Q

Why was Mendeleev’s periodic table accepted?

A
  • He predicted properties of missing elements
  • Elements with properties predicted by Mendeleev were discovered
  • Elements were discovered which fitted the gaps
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14
Q

Group 0 (Nole gases)

A
  • Noble gases exist as single atoms (monoatomic) as they have a full outer shell, which means they do not lose or gain electrons.
  • They are unreactive because they have a full outer shell and they do not need to lose
    or gain electrons.
  • The boiling points of the noble gases increase with increasing relative atomic mass (going down the group).
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15
Q

Group 1 metals:

A
  • They react with non-metals to form ionic compounds
  • The compounds are white solids which dissolve in water to form colourless solutions.
  • The alkali metals all have low densities.
  • The first three alkali metals (Li, Na & K) are all less dense than water – this is why they float on the surface.
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16
Q

Observation of the first three alkali metals with water

A
  • The alkali metals react with water to form a metal hydroxide and hydrogen gas.
  • When universal indicator is added to the metal hydroxide solution, it turns purple as it contains OH-(hydroxide) ions.
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17
Q

Lithium reaction with water:

A
  • Fizzing/ bubbles
  • Moves on the surface of the water
  • Gets smaller in size
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18
Q

Sodium reaction with water:

A
  • Fizzing/ bubbles
  • Moves quickly on the surface of water
  • Gets smaller in size
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19
Q

Potassium reaction with water:

A
  • Fizzing/ bubbles
  • Moves very quickly on the surface of water
  • Gets smaller in size
  • Lilac flame
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20
Q

Why does the reactivity of group 1 metals increases down the group?

A
  • More electron shells down the group
  • Weaker attraction between the nucleus and the outer shell electron
  • Easy to lose the outer electron
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21
Q

Halogens:

A
  • All have 7 electrons on their outer shell
  • They have coloured vapours
  • Consist of molecules which are made up of pairs of atoms (diatomic) For example, Cl2, F2, Br2
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22
Q

Why does the melting point and boiling point increase down group 7?

A
  • The melting and boiling point increases as you go down group 7 because the relative molecular mass increases.
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23
Q

Why does reactivity decrease down group 7?

A
  • More electron shells
  • Weaker attraction between the nucleus and the outer shell electron
  • Harder to gain an electron
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24
Q

Halogen- displacement reaction:

A

A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.

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25
Q

Compared with group 1 elements, the transition metals:

A
  • Have higher melting points
  • Have higher densities
  • Are stronger
  • Are harder
  • Are muss less reactive- so do not react as vigorously with water and oxygen
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26
Q

Properties of transition metals:

A
  • Have ions with different charges
  • Form coloured compounds
  • Are useful as catalysts
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27
Q

Atoms:

A
  • All substances are made of atoms.
  • An atom is the smallest part of an element that can exist.
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28
Q

Element:

A

A substance made form one type of atom

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29
Q

Compounds:

A
  • Contains atoms of two or more different elements chemically bonded together.
  • Is difficult to separate.
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30
Q

Mixture:

A
  • A mixture is when two or more different elements or compounds that are not chemically bonded together.
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31
Q

Key definitions- seperation techniques:

A
  • Soluble – it dissolves.
  • Insoluble – it will not dissolve. Examples of insoluble substances include mud and sand.
  • Solute – a solid being dissolved. Example include salt and sugar.
  • Solvent – a liquid a solute can be dissolved into.
  • Solution – a mixture of a solute and solvent that does not separate.
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32
Q

Filtration:

A

Used to seperate an insoluble solid from a liquid

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33
Q

Simple distillation:

A
  • Used to seperate liquids with very diffferent boiling points
  • Used to seperate a liquid from a soluble solid
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34
Q

Fractional distillation:

A

Used to seperate solutions with similar boiling points

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35
Q

Chromatography:

A

Chromatography is used to separate mixtures of coloured compounds.

36
Q

Chromatography practical:

A
  1. Use a pencil and ruler to draw startline and solvent front on a chromatography paper
  2. Fill a beaker with the solvent and ensure that the startline is above the solvent.
  3. Place dots of the samples to be tested on the startline
  4. Place the chromatography paper in the beaker and add lid.
  5. Allow the solvent to move through the paper
  6. Take out the paper and allow it to dry.
  7. Analyse the separation.
37
Q

Two phases of chromatography:

A
  1. A stationary phase
  2. A mobile phase
    * The paper is the stationary phase
    * The solvent is the mobile phase
38
Q

Chromatography key points:

A
  • Start line drawn in pencil to prevent it from dissolving in the solvent
  • The spots should be above the solvent to prevent it from dissolving in the solvent
  • A lid is used to prevent solvent from evaporating
  • The most soluble colour will travel the furthest and the least soluble will travel the least.
  • If a colour is insoluble, it will not move therefore the solvent has to be changed.
  • A pure sample only has one dot in a vertical line.
  • The sample is a mixture if there is more than one dot in a vertical line.
39
Q

How the different dyes separated by paper chromatography? 4 marks

A
  • Solvent moves through paper
  • Different dyes have different solubilities in solvent
  • Different dyes have different attractions for the paper
  • So are carried different distances.
40
Q

Rf value:

A

Distance moveed by substance/ distance moved by solvent

41
Q

Structure of ionic compounds:

A
  • A giant ionic lattice
42
Q

Ionic lattice:

A
  • The strong ionic bonds require lots of energy to break
  • Ionic compounds are mostly white solids. Just like salt (NaCl).
  • Ionic compounds easily dissolve in water to form aqueous solutions.
  • As an ionic compound dissolves, its ionic lattice breaks up, and the ions are free to move.
43
Q

Why can solid ionic compounds not conduct electricity?

A
  • Ions are fixed in position
  • Ions can only vibrate (they can’t move)
  • Ions can’t carry electrical charge
44
Q

Why can molten/ aqeous ionic compounds conduct electricity?

A
  • Ions are free to move
  • Ions can carry an electrical charge
45
Q

Explain why sodium chloride has a high mp?

A
  • Giant ionic lattice
  • SEABOCI
  • A lot of energy is needed to overcome the forces
46
Q

What is a compound ion?

A

Groups of elements that have a charge

47
Q

Small covalent molecule:

A
  • Small covalent molecules are gases at room temperature
  • Weak intermolecular forces between molecules which require little energy to break.
48
Q

Why is ammonia a gas at room temperature?

A
  • Ammonia is a small covalent molecule
  • Ammonia is made up of small molecules
  • There are weak intermolecular forces beetween bonds
  • Little energy is needed to break bonds
49
Q

Graphite (giant covalent structure):

A
  • A giant covalent structure
  • Only made from carbon atoms
  • Arranged in hexagonal layers
  • There is weak intermolecular forces between layers
  • Each carbon atom is bonded to three other carbon atoms by STRONG covalent bonds
50
Q

Uses of graphite:

A
  • Lubricants
  • In pencils (lead)
  • Electrodes
51
Q

Why is graphite soft?

A
  • There are only weak intermolecular forces between layers.
  • The layers (of carbon atoms) in graphite can slide.
52
Q

Why does graphite conduct electricity?

A
  • Each carbon atom is bonded to three other carbon atoms.
  • One delocalised electron per carbon atom which moves throughout the structure and carry electrical charge.
53
Q

Why does graphite conduct thermal energy?

A
  • Each carbon is bonded to **three other carbon atoms. **
  • One delocalised electron per carbon atom which moves throughout the structure and can transfer thermal energy.
54
Q

Why does graphite have high melting point?

A
  • Giant covalent structure
  • Each carbon atom is bonded to three other carbon atoms by strong covalent bonds.
  • Need a lot of energy to break the bonds.
55
Q

Diamond:

A

Structure and bonding
* A giant covalent structure.
* Made only from carbon atoms.
* Each carbon atom is bonded to four other carbon atoms.

56
Q

Uses of diamond:

A
  • Laser beam
  • Cutting tools and drills
  • Jewellery
57
Q

Silicon dioxide:

A

Structure and bonding
* Giant covalent structure.
* Made from silicon and oxygen atoms.
* Each oxy atom- bonded to 2 silicon atoms by strong covalent bonds.
* Each silicon atom- bonded to four other oxygen atoms by strong covalent bonds.
* Has a formula SiO2.

58
Q

Uses of silicon dioxide:

A

Lining for furnaces- due to high mp.

59
Q

Graphene:

A
  • Giant covalent structure
  • It is made of a single layer of graphite.
  • Each carbon atom forms three strong covalent bonds.
  • It has a low density
  • Useful in elctronic devices- it is transparent
60
Q

Fullerenes:

A

HOLLOW SHAPED MOLECULES THAT ARE ONLY MADE FROM CARBON ATOMS
* Drug delivery
* Lubricants- spherical, so can roll
* Catalysts- large SA to V ratio

61
Q

Carbon nanotube:

A
  • Cylindrical fullerenes made from ONLY carbon atoms.
  • They are strong and have a low density.
  • Each carbon atom in carbon nanotubes forms three strong covalent bonds.
  • This means there are delocalised electrons- carry an electrical charge
  • They also have a high tensile strength (don’t break when stretched).
  • Nanotechnology
  • Electronics
  • Materials as it has a low density.
62
Q

Metals:

A
  • Giant structure of positively charged ions in a lattice.
  • Delocalised electrons move throughout the whole structure.
  • Strong electrostatic attraction between positive ions and negative electrons.
63
Q

Why are metals malleable?

A

Layers of atoms can slide

64
Q

Why do metals have high mp and bp?

A
  • Giant lattice.
  • Strong electrostatic force of attraction between positive ion and negative electrons.
  • A lot of energy is required to break the force.
65
Q

Nanoparticles:

A

Very high SA to V ratio

USES:
* Huge surface area to volume ratio, so they could make new catalysts
* Drug delivery
* Some nanoparticles conduct electricity, so they can be used in tiny electric circuits for computer chips
* They are also used in cosmetics. Used to improve moisturisers without making them
really oily.

66
Q

Risk of nanoparticles:

A
  • Some nanoparticles could be toxic
  • Cell damage to the body
  • Un-discovered harmful side effects
67
Q

Coarse particles diameter:

A

2.5 * 10 ^-6 m

68
Q

Fine particles diameter:

A

1 * 10^-7 m

69
Q

Nanoparticles diameter:

A

1 * 10^-9 m.

70
Q

Law of conservation of mass:

A
  • No atoms are lost or made during chemical reaction
  • Mass of the products equals the mass of reactants
71
Q

Thermal decomposition of copper carbonate:

A

Copper carbonate → Copper oxide + carbon dioxide

72
Q

When is atom economy 100%?

A

When only one product is formed

73
Q

Atom economy:

A

(Total mass of desired product atoms/ total mass of reactant atoms) * 100

74
Q

Mass of a sinle atom in grams:

A

Mass of a single atom in g = Relative atomic mass / (6.02 × 1023)

75
Q

Percentage yield:

A

(Mass of product actually made/ maximum theoretical mass) * 100

76
Q

Moles of GAS:

A

Volume of gas dm^3 / 24 dm^3

77
Q

Cells:

A
  • A cell is a system that contains chemicals which react to produce electricity.
  • A simple cell can be made by connecting two different metals in contact with an electrolyte.
78
Q

Alkaline batteries:

A
  • Alkaline batteries contain hydroxide ions
  • They are used in remote controls
  • Over time, the ions in the electrolyte and the metal ions on the electrode gets used up and turns into products
  • Once any one of the reactants get used up, the reaction can’t happen and so no electricity is produced
  • Reactions in an alkaline battery is irreversible therefore they are non-rechargeable.
79
Q

Hydrogen fuel cells:

A

Hydrogen + oxygen = water + energy

80
Q

Anode- hydrogen fuel cell:

A

OXIDATION:
H2→2H+ + 2e-

81
Q

Cathode- hydroge fuel cells:

A

REDUCTION:
2H2 + O2 → 2H2O

82
Q

Hydrogen fuel celss + and -:

A

Advantages
* No pollutants-Produces only water
* Hydrogen can be made renewable if made by electrolysis using renewable energy.

Disadvantages
* Hydrogen is highly flammable
* Hydrogen not renewable if produced using fossil fuels
* Not many hydrogen filling stations

83
Q

Rechargeable batteries + and -:

A

Advantages:
* Charging points more widely available
* Rechargable

Disadvantages:
* May release toxic chemicals on disposal
* Rechargeable batteries have finite life-time
* Can catch fire

84
Q

Bond energies:

A
  • Energy supplied to break chemical bonds so bond breaking is an endothermic process
  • Making new bonds transfers energy to the surrounding therefore is an exothermic process.
85
Q

Endothermic reaction:

A

Energy needed to break existing bonds is greater than the energy released from forming new bonds.

86
Q

Exothermic reaction:

A

Energy released from forming new bonds is greater than the energy needed to break existing bonds.

87
Q

Metals and extraction:

A
  • Metals that are less reactive than carbon can be extracted from their oxide by reduction with carbon.
  • Metals more reactive than carbon are extracted by electrolysis.

GCSE chemistry (7 decks)

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