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GCSE Chemistry > C1 - Atomic Structure & the Periodic Table > Flashcards

C1 - Atomic Structure & the Periodic Table Flashcards

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1
Q

What is the definition of an element? Give an example

A

Elements are made from one type of atom (eg. sodium)

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2
Q

What is the definition of a compound? Give an example

A

Compounds are two or more different elements chemically bonded (eg. sodium chloride)

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3
Q

What is the definition of a mixture? Give an example

A

Mixtures are two or more elements or compounds not chemically bonded (eg. air)

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4
Q

Name the 3 subatomic particles

A

Protons, electrons and neutrons

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5
Q

Which particles are found in the nucleus?

A

Protons and neutrons

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6
Q

Which particle orbits the nucleus on shells?

A

Electrons

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7
Q

What is the charge of an electron?

A

-1, negative

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8
Q

What is the charge of a neutron?

A

0, neutral

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9
Q

What is the charge of a proton?

A

+1, positive

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10
Q

What is the overall charge of an atom? Why?

A

Neutral - there are the same number of electrons and protons (they cancel each other out)

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11
Q

How do you work out the number of protons in an atom?

A

Use the atomic number

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12
Q

How do you work out the number of electrons in an atom?

A

Use the atomic number

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13
Q

How do you work out the number of neutrons in an atom?

A

Atomic mass number - atomic number

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14
Q

Which subatomic particle decides what the element is?

A

The number of protons

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15
Q

What did the plum pudding model suggest the atom was like?

A

A ball of positive charge with negatively charged electrons scattered in it

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16
Q

What did Rutherford (nuclear) model think the atom looked like?

A

Positively charged nucleus (small) with electrons surrounding the nucleus

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17
Q

What did Bohr discovery about the atom?

A

He discovered that electrons went around in shells

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18
Q

State the size of the nucleus

A

1 x 10-14m

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19
Q

State the size of the atom

A

1 x 10-10m

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20
Q

What is an isotope?

A

An element with the same number of protons and electrons but different numbers of neutrons

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21
Q

How are elements organised in the modern periodic table?

A

According to atomic / proton number

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22
Q

How were elements ordered in the early periodic table?

A

According to atomic weight

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23
Q

What were the main issues with the early periodic table?

A
  • It had no gaps for undiscovered elements

- Some elements were placed in groups with elements that had very different properties

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24
Q

How are elements in the modern periodic table grouped?

A

They are grouped according to their properties

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25
Q

Where are metals found on the periodic table?

A

Left and bottom

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26
Q

What is group 1 called?

A

Alkali metals

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27
Q

What is group 7 called?

A

Halogens

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28
Q

What is group 0/8 called?

A

Noble gases

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29
Q

How many electrons can go on each shell?

A

1st - 2, 2nd - 8, 3rd - 8, 4th- 18

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30
Q

How are group number and outer electrons number related?

A

Group number = number of electron on outer shell

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31
Q

How are period number and shell number related?

A

Period number = number of shells

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32
Q

What are the properties of group 1 metals?

A
  • Conduct electricity, can be cut
  • React with water and oxygen
  • Low density
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33
Q

How does the reactivity of group 1 elements change down the group?

A

They get more reactive

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34
Q

Complete the equation:

metal + water > … + …

A

Metal + water = metal hydroxide + hydrogen

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35
Q

What is the test for hydrogen?

A

Squeaky pop

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36
Q

How do I test that metal hydroxides are alkali?

A

Add universal indicator - it will turn purple

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37
Q

Complete the equation:

Metal + oxygen > ……

A

Metal + oxygen > metal oxide

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38
Q

What is the formula for chlorine?

A

Cl2 (diatomic)

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39
Q

What is the formula for fluorine?

A

F2 (diatomic)

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40
Q

What is the colour and state of fluorine at room temperature?

A

Yellow / green gas

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41
Q

What is the colour and state of chlorine at room temperature?

A

Green gas

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42
Q

What is the colour and state of bromine at room temperature?

A

Brown liquid

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43
Q

What is the state symbol for a gas, liquid and a solid?

A

(s), (l), (g)

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44
Q

Name some properties of noble gases

A

Low melting points, unreactive

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45
Q

Why are noble gases unreactive?

A

They have a full outer shell

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46
Q

What is the easiest way to identify whether something is a element?

A

Look in the periodic table

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47
Q

What is a molecule?

A

Two or more atoms chemically bonded

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48
Q

What is the mass of a proton?

A

1

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49
Q

What is the mass of a neutron?

A

1

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50
Q

What is the mass of a electron?

A

0

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51
Q

What did Thompson (plum pudding model) suggest the atom was like?

A

A ball of positive charge with negative electrons scattered in it

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52
Q

Rutherford realised that the plum pudding model was wrong.

What was Rutherford’s model of the atom like?

What did he call his model of the atom?

A

Rutherford thought the atom was made of a tiny positively charged nucleus, surrounded by scattered electrons.

He called it ‘the nuclear model’

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53
Q

How did Bohr improve Rutherford’s nuclear model of the atom?

A

Bohr realised that electrons orbit the nucleus on shells

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54
Q

How big is the nucleus of an atom compared to the whole atom?

A

The nucleus of the atom is 1/10,000 of the size of an atom

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55
Q

Where is most of the mass of the atom found?

A

In the nucleus, because this is where protons and neutrons are found

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56
Q

How did Mendeleev organise the elements in his periodic table?

A

Mainly order of relative atomic mass but he changed the order slightly to put elements with similar properties in the same group

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57
Q

What unusual step did Mendeleev take to make this possible?

A

He left gaps in his periodic table

58
Q

How did the gaps in Mendeleev’s periodic table hep to prove his ideas?

A

Elements were discovered that fitted the gaps and had similar properties to the other elements in the same group
(For example, gallium fitted in a gap in group 3 and it had similar properties to aluminium)

59
Q

What do elements in the same group of the periodic table have in common?

A
  • The same number of electrons in their outer shell (e.g group 1 elements have 1 electron in their outer shell)
  • Similar properties (e.g all group 1 elements react with water producing hydrogen and metal hydroxides)
60
Q

Where are nonmetals found in the periodic table?

A

On the right and towards the top

61
Q

What properties do most metals have in common?

A
  • They are strong
  • They can be bent or hammered into shape (Malleable)
  • They are good at conducting heat and electricity.
  • They have high melting and boiling points
62
Q

What properties do most nonmetals have in common?

A
  • They are brittle
  • Dull looking
  • Often liquids or gases at room temperature.
  • Poor conductors
63
Q

What do we call the metals in the middle of the periodic table?

A

Transition metals

64
Q

Is hydrogen a metal or nonmetal?

A

Nonmetal

65
Q

What do we call atoms that have lost or gained electrons and so have a positive or negative charge?

A

Ions

66
Q

If two elements are in the same group in the periodic table, what does this tell you about them?

A
  • That they have the same number of electrons in their outer shell (e.g all group 1 element have 1 electron in their outer shell).
  • That they have similar properties
67
Q

What charge do group 1 ions have (e.g sodium ions)?

A

+1

68
Q

What happens when group 1 elements react with water?

A
  • They float and move around on the surface fizzing.
  • Hydrogen (which may burst into flame) is made.
    Metal hydroxides (e.g sodium hydroxide) are made. These dissolve in the water making alkaline solutions.
  • sodium + water → sodium hydroxide + hydrogen
  • 2Na (s) + 2H2O (l) →2NaOH (aq) + H2 (g)
69
Q

Lithium + oxygen→……

A

Lithium + oxygen →lithium oxide

70
Q

Sodium + oxygen →……

A

Sodium + oxygen →sodium oxide (+ sodium peroxide)

71
Q

Sodium + chlorine →……

A

Sodium + chlorine→sodium chloride

72
Q

Potassium + chlorine →……

A

Potassium + chlorine→potassium chloride

73
Q

Explain why the group 1 elements (alkali metals) get more reactive as you go down the group.

A

The outer electron is further from the nucleus
There is a weaker (electrostatic) attraction between the nucleus and outer electron.
It is easier to lose the outer electron

74
Q

Compare the properties of group 1 metals and transition metals

A
  • Group 1 are more reactive
  • Alkali metals have much lower boiling points
  • Transition metals are more dense, strong and hard
  • Transition metals have much higher melting point
75
Q

Fluorine and chlorine are both in group 7 of the periodic table, how many electrons do they have in their outer shell?

A

7

76
Q

Group 7 elements form diatomic molecules. What are diatomic molecules?

A

Molecules made from two atoms bonded together

77
Q

How does the reactivity of group 7 elements (halogens) change as you go down the group?

A

They get less reactive as you go down the group

78
Q

Complete the equation below and explain why this happens:

Chlorine + sodium bromide→……

A

Chlorine displaces bromine, because it is more reactive than bromine.

Chlorine + sodium bromide →sodium chloride and bromine

79
Q

What is oxidised (loses electrons) and what is reduced (gains electrons) in this reaction.

A
  • Chlorine atoms are reduced (gain electrons) forming chloride ions.
  • Bromide ions are oxidised (lose electrons) forming bromine molecules
80
Q

Explain why the group 7 elements (halogens) get less reactive as you go down the group.

A
  • The outer shell is further from the nucleus.
  • The (electrostatic) attraction between the nucleus and outer shell/electrons is weaker.
  • It’s more difficult to attract electrons to the outer shell
81
Q

How many electrons do the group 0 elements (noble gases) have in their outer shell?

A

8, except helium which has 2 electrons in its outer shell

82
Q

Why are group 0 elements unreactive?

A
  • They have full outer shells of electrons.

- They have a stable arrangement of electrons so they don’t need to gain/lose or borrow electrons from other atoms

83
Q

What is the state and colour of group 0 elements at room temperature?

A

They are all colourless gases

84
Q

Why does the boiling point of group 1 elements increase as you go down the group?

A
  • The atoms have more electrons.
  • There are stronger forces of (electrostatic) attraction between the atoms.
  • It takes more energy to break these forces
85
Q

Which elements form diatomic molecules (molecules made from two atoms joined together)?

A
  • Oxygen (O2)
  • Hydrogen (H2)
  • Nitrogen (N2)
  • Group 7 elements:
    • Fluorine F2
    • Chlorine Cl2
    • Bromine Br2
    • Iodine I2
86
Q

A nanometer is one ___th of a metre

A

Billionth

87
Q

What is one nanometer in standard form?

A

1x10-9m

88
Q

How many different elements are there?

A

118

89
Q

What is relative atomic mass and why is it used?

A

Because many elements can exist as different isotopes, RAM is used as a mass number. It is an average, taking into account the different masses and their abundance

90
Q

What formula is used to calculate the RAM of an element?

A

Sum of all the elements’ masses in a compound, mixture etc.

91
Q

Chromatography is used to separate different dyes within an ink. Explain how a chromatography experiment is set up.

A
  1. Draw a line near the bottom of a sheet of filter paper in pencil (since it’s insoluble).
  2. Add a spot of ink to this line.
  3. Place the sheet in a beaker of solvent (e.g. water - depends on what’s being tested), making sure the ink doesn’t touch the solvent.
  4. Put a lid on the container to stop the solvent evaporating
92
Q

Explain what happens during chromatography

A
  • The mobile phase (solvent) seeps up the paper, carrying the ink with it.
  • Each dye within the ink moves up the paper at a different rate; the dyes separate out into spots.
    If any of the dyes are insoluble, they stay on the stationary phase (pencil line).
  • When the solvent has nearly reached the top, the paper is taken out and allowed to dry; the result is a chromatogram
93
Q

What are the mobile and stationary phases in chromatography?

A
  • Mobile phase = solvent (because it moves up the paper).

- Stationary phase = pencil line (it doesn’t move)

94
Q

How would you separate an insoluble solid from a liquid?

A

Filtration

95
Q

How would you separate a soluble salt from a liquid? Give 2 ways.

A
  • Evaporation: heating the solution, in an evaporating dish, over a bunsen burner. However, this can only be used if the salt doesn’t thermally decompose. If it does:
  • Crystallisation: gently heating the solution until crystals start to form, then leaving it to cool. More crystals will form, which are filtered out of the solution and left in a warm place to dry
96
Q

How would you separate out rock salt (a mixture of salt and sand)?

A
  • Grind it to make the salt crystals small (so that they dissolve easily).
  • Put it in water and heat; salt will dissolve but sand won’t.
  • Filter out the sand.
  • Evaporate the water from the salt, so that it forms dry crystals
97
Q

What is simple distillation used for?

A

Separating out a liquid from a solution

98
Q

How does simple distillation separate out a liquid from a solution?

A
  • Solution is heated.
  • The part of the solution with the lowest boiling point evaporates first.
  • The vapour is cooled in a condenser, then collected in a beaker at the end of the condenser.
  • The rest of the solution remains in the flask
99
Q

What is the problem with simple distillation?

A

You can only use it to separate liquids with very different boiling points

100
Q

What is fractional distillation used for?

A

Separating out liquids, from a mixture, with close boiling points

101
Q

How does fractional distillation work?

A
  1. Heat the mixture in the flask
  2. The liquid with the lowest boiling point evaporates first.
  3. When the thermometer’s temperature reaches the boiling point of this liquid, it reaches the top of the column, condenses, and is collected.
  4. Other liquids may start to evaporate, but will condense part of the way up the column (bc it’s cooler towards the top).
  5. Once the first liquid’s collected, the temperature is raised until the next one reaches the top of the column. Etc.
102
Q

What was Dalton’s model of the atom?

A

Atoms were solid, indivisible spheres. Each element was made of a different type of sphere

103
Q

What was Thomson’s model of the atom?

A

He discovered electrons which could be removed from atoms, disproving Dalton’s theory of indivisibility.

Thomson suggested atoms were spheres of positive charge, with negative electrons scattered throughout (plum pudding model)

104
Q

What was Rutherford’s model of the atom?

A

Positively charged nucleus, surrounded by a cloud of negative electrons. This was the first nuclear model of the atom

105
Q

What discovery did Bohr make? How was he proven right?

A

Electrons orbit the nucleus in fixed positions, called energy levels.

His theoretical calculations agreed with experimental data

106
Q

What discovery was made after electron orbits?

A

Experiments showed that the nucleus’ positive charge was subdivided between a group of particles, called protons

107
Q

What discovery did Chadwick make? Why did this make sense?

A

He proved the existence of neutrons in the nucleus. This explained the imbalance between the atomic and mass numbers

108
Q

Who theorised that atoms were solid, indivisible spheres?

A

Dalton

109
Q

Who theorised the plum pudding model?

A

Thompson

110
Q

Whose discoveries led to the nuclear model of the atom?

A

Rutherford

111
Q

Who calculated that electrons had fixed orbits?

A

Bohr

112
Q

Who discovered neutrons?

A

Chadwick

113
Q

How was Rutherford’s alpha scattering experiment conducted?

A

Scientists in Rutherford’s lab fired a beam of alpha particles at a thin sheet of gold foil

114
Q

In Rutherford’s alpha scattering experiment, what did the scientists expect to happen and why?

A

From the plum pudding model, they expected particles to pass straight through the gold foil, or only be slightly deflected

115
Q

What were the results of Rutherford’s alpha scattering experiment? Why wasn’t this as expected?

A

The majority of the alpha particles went straight through the gold sheet. However, some were deflected more than expected and a few were deflected right back the way they came. The plum pudding model couldn’t explain this

116
Q

In Rutherford’s alpha scattering experiment, a few alpha particles were deflected back. What did the scientists deduce from this?

A

They realised that most of the atom’s mass was concentrated in a central nucleus, which must also have a positive charge, since it repelled the positive alpha particles

117
Q

In Rutherford’s alpha scattering experiment, most of the alpha particles passed straight through. What did the scientists deduce from this?

A

They realised that most of the atom is just empty space, and that the nucleus is very small relative to this space

118
Q

What was deduced from the results of Rutherford’s alpha scattering experiment?

A
  1. Because a few alpha particles were deflected back, they realised that most of the atom’s mass was concentrated in a central nucleus, which must also have a positive charge, since it repelled the positive alpha particles.
  2. Because most of the alpha particles passed straight through, they realised that most of the atom is just empty space, and that the nucleus is very small relative to this space
119
Q

How were elements organised in the early periodic table, and what was the problem with this?

A

Elements were organised according to their relative atomic masses.

The table wasn’t complete, and some elements were placed in the wrong group because the ordering didn’t take into account their properties

120
Q

How was Mendeleev’s periodic table different to - and better than - the early versions of the periodic table?

A
  • He organised elements mainly in order of atomic mass, but also taking into account their properties.
  • The discovery of isotopes showed he was right to do this.
  • He left gaps which predicted the existence of undiscovered elements, allowing him to predict their properties.
  • The discovery of certain elements that fit this pattern confirmed his ideas, for example, he left a gap for an element with similar properties to aluminium.
  • Gallium was discovered, and it showed these properties
121
Q

Transition metals are found in the middle of the periodic table. Give 7 properties of transition metals

A
  • Good conductors of electricity + heat
  • Strong
  • Dense
  • Shiny
  • Can form more than one ion
  • Colourful
  • Transition metal compounds often make good catalysts
122
Q

What are the trends in properties of alkali metals as you go down the group?

A

Increasing reactivity
Decreasing melting + boiling points
Increasing relative atomic mass

123
Q

Alkali metals only ever form ionic compounds. What are these compounds usually like?

A

White solids that dissolve in water to form colourless solutions

124
Q

Explain why sodium is more reactive than lithium

A
  • Sodium is further down group 1 than lithium.
  • Group 1 ions lose one electron to become stable.
  • Sodium’s outer electron is further from the nucleus that lithium’s.
  • So the electrostatic attraction between the electron and nucleus is weaker
  • So the electron is more easily released.
  • So sodium is more reactive
125
Q

Compare the reactions of lithium and potassium with water. Describe what would be observed in each reaction and explain the similarities and differences [6 marks]

A
  • A vigorous reaction would be observed with both metals
  • The reaction with potassium would be more vigorous as it is more reactive than lithium
  • This is because the outer electron of potassium is further from its nucleus and therefore more easily lost
  • Both metals would float, as they are less dense than water.
  • Both metals will fizz, producing bubbles of hydrogen gas
  • The reaction with potassium would produce more hydrogen bubbles, as potassium is more reactive
  • The reaction with potassium would release more energy, to the extent that the potassium and hydrogen could ignite
126
Q

Compare the properties of transition and alkali metals

A

Alkali metals are much more reactive.
Transition metals are stronger, denser and harder.
Alkali metals have much lower boiling points.

127
Q

How do alkali metals react with water?

A
  • They react vigorously to produce hydrogen gas and metal hydroxides
  • The violence of reactivity (and energy released from the reaction) increases as you go down the group
128
Q

How do alkali metals react with chlorine?

A
  • When heated in chlorine gas, they form white metal chloride salts
  • The reaction gets more vigorous as you go down the group
129
Q

How do alkali metals react with oxygen?

A

They react with oxygen to form metal oxides (which is why they tarnish in the air), different types of oxide forming depending on the metal:

Lithium → lithium oxide.

Sodium → mixture of sodium oxide and sodium peroxide

Potassium → mixture of potassium peroxide and potassium superoxide

130
Q

What are the trends in properties of the halogens as you go down the group?

A
  • Decreasing reactivity
  • Increasing melting + boiling points
  • Increasing relative atomic masses
131
Q

Give 3 properties of the halogens

A

Form diatomic molecules
Bond ionically w/ metals
Bond covalently to form simple molecular compounds
A more reactive halogen will displace a less reactive one from an aqueous solution of its salt

132
Q

What is the colour and state of fluorine at room temperature? Give 2 of its properties

A
  • Very reactive and poisonous

- Yellow gas

133
Q

What is the colour and state of chlorine at room temperature? Give 3 of its properties

A
  • Green gas.

- Fairly reactive, poisonous, dense

134
Q

What is the colour and state of bromine at room temperature? Give 3 of its properties

A
  • Red-brown liquid

- Poisonous, dense, volatile

135
Q

What is the colour and state of iodine at room temperature? What does its gas look like?

A
  • Dark grey, crystalline solid

- Purple gas

136
Q

A chemist places a glass lid over a beaker containing aqueous iodine. She places an ice cube on top of the lid and then gently heats the solution. Predict and explain her observations

A

A purple vapour will form over the solution as the iodine evaporates.

Purple/grey crystals will form on the lid as the iodine deposits/crystallises/solidifies

137
Q

Give 4 properties of noble gases

A

Inert
Non-flammable (bc of above)
Exist as monatomic gases
All colourless gases at room temperature

138
Q

Why are the noble gases inert?

A

They have full outer shells which are energetically stable. They therefore don’t need to react to gain or lose electrons to become stable

139
Q

What are the trends in properties of the noble gases as you go down the group?

A
  • Increasing boiling points (greater no. electrons in each element leads to greater intermolecular forces
  • Increasing relative atomic mass
140
Q

If isotopes have different numbers of neutrons how else must they also be different?

A

They must have different atomic masses.

GCSE Chemistry (10 decks)

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