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Key Concepts Flashcards

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1
Q

What are the three ideas in John

Dalton’s theory about the atom?

A
  • Atoms cannot be created, divided or destroyed.
  • Atoms of the same element are exactly the same
    and atoms of different elements are different.
  • Atoms join with other atoms to make new
    substances.
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2
Q

What discovery caused the original Dalton model of

an atom to change?

A

The discovery of subatomic particles.

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3
Q

How did JJ Thomson discover the electron?

A 
Thomson experimented with a cathode ray
tube.
The beam moved towards the positively
charged plate so he knew that the particles
must have a negative charge.
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4
Q

Describe the atomic model proposed by JJ Thomson

A

Plum pudding model.
Negatively charged electrons scattered
through a positively charged material.

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5
Q

What did Ernest Rutherford discover from his gold

foil experiment?

A

He shot a beam of positively charged particles at sheet of gold
foil.
- Most of the particles passed straight through suggesting that
atoms were mostly empty space.
- A few particles were deflected and a few bounced directly
back showing that there must be a tiny, dense and
positively-charged nucleus.

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6
Q

Describe Rutherford’s new model of the atom

A 
- Mass is concentrated in the central
nucleus.
- Mostly empty space.
- Electrons travel in random paths
around the nucleus.
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7
Q

Describe the structure of an atom

A

Small central nucleus made up of
protons and neutrons.

Electrons orbit (move around) the
nucleus in shells.
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8
Q

What is the radius of the nucleus? How large is it

compared to the radius atom?

A

The radius of the nucleus is 1x10-14 m.

This is 1/10000 of the atomic radius.

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9
Q

What are the relative masses of protons, neutrons

and electrons?

A

Proton: 1
Neutron: 1
Electron: 1/1836

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10
Q

What are the relative charges of protons, neutrons

and electrons?

A

Proton: +1
Neutron: 0
Electron: -1

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11
Q

Why do atoms contain equal numbers of protons and

electrons?

A

Atoms are stable with no overall charge.
Protons are positively charged and electrons
are negatively charged. For the charges to
balance, the number of protons and electrons
must be equal.

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12
Q

Where is the mass of an atom concentrated?

A

In the nucleus

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13
Q

What does the atomic number of an atom represent?

A

The number of protons.

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14
Q

What does the mass number of an atom

mean?

A

The mass number is the number of
protons and neutrons found in the
nucleus of an atom.

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15
Q

Fill in the blank: ‘Atoms of the same element have
the same number _____ of in the nucleus and this
is unique to that element’

A

Protons

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16
Q

What is an isotope?

A

Isotopes are atoms with the same number
of protons (so they are the same element)
but a different number of neutrons.
Isotopes of an element have the same
atomic number but different mass numbers.

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17
Q

Boron has the atomic number 5 and the mass
number 11. How many protons, electrons and
neutrons does Boron have?

A

5 protons
5 electrons
6 neutrons

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18
Q

Why is the relative atomic mass not always a whole

number?

A

Different isotopes of the same element
have different mass numbers. The
relative atomic mass is an average of the
masses of all these isotopes.

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19
Q

What two values would be required to calculate the

relative atomic mass of chlorine?

A

Mass numbers and relative abundances

of all the isotopes of chlorine.

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20
Q

How did Mendeleev arrange the elements in his

periodic table?

A

Elements arranged with increasing atomic masses.
Elements with similar properties put into groups (due to
periodic trends in chemical properties).
Switched the position of some elements.
Gaps left for undiscovered elements.

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21
Q

How was Mendeleev able to predict the properties of

new elements?

A

Mendeleev left gaps in his periodic table.
He used the properties of elements next
to these gaps to predict the properties of
undiscovered elements.

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22
Q

Mendeleev’s table lacked some amount
of accuracy in the way he’d ordered his
elements. Why was this?

A

Isotopes were poorly understood at the time.
Protons and neutrons had not yet been
discovered.

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23
Q

How are elements arranged in the modern periodic

table?

A

In order of increasing atomic number.

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24
Q

Fill in the blank: ‘Elements in the same group

(column) have similar _________ ’

A

Chemical properties

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25
Q

Why do elements in the same column have similar

chemical properties?

A

Same number of outer shell electrons.
Number of outer shell electrons
determines how an atom reacts.

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26
Q

What does the period (row) number tell you about all

the elements in that period?

A

Elements in the same period have the
same number of electron shells.
e.g. all elements in period 4 have 4
shells of electrons.

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27
Q

What does group (column) number tell you about all

the elements in that group?

A

All elements in the same group have the
same number of outer electrons.
e.g. all elements in group 2 have 2
electrons in their outer shell

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28
Q

On which side of the periodic table are

the metals positioned?

A

Left hand side

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29
Q

What determines whether an element is a metal or

non-metal?

A

Atomic structures of the elements.

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30
Q

What is the maximum number of electrons allowed in

each of the first 3 shells?

A

1st shell: 2
2nd shell: 8
3rd shell: 8

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31
Q

When are atoms most stable?

A

When they have full electron shells.

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32
Q

The atomic number of Na is 11.

What is the electron configuration of Na?

A

2, 8, 1

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33
Q

How is the electron configuration of an element

related to its position in the periodic table?

A

Diagram:
The number of rings is the period.
The number of electrons in the outer ring is the group.
The total number of electrons is the atomic number.
Number (e.g. sodium, 2.8.1):
Sodium has 3 shells as there are 3 numbers.
The last number is the number of outer shell electrons / the group.
The sum of the digits is the atomic number.

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34
Q

What is an ionic bond?

A

Abond between a metal and non-metal

involving the transfer of electrons.

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35
Q

In terms of electrons, describe what happens to the

metal and non-metal when an ionic bond forms

A

The metal atom loses electrons to
become a positively charged ion (cation).
The non-metal gains electrons to become
a negatively charged ion (anion).

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36
Q

What is an ion?

A

An ion is an atom or group of atoms with

a positive or negative charge.

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37
Q

If an ion is positively charged, has it lost or gained

electrons?

A

It has lost electrons. There are fewer
negatively charged electrons to cancel out
the charge of the positive protons. This
means the overall charge becomes positive.

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38
Q

Na+ has the atomic number 11 and the mass number
23. How many protons, neutrons and electrons are
in this ion?

A

Protons: 11
Electrons: 10
Neutrons: 12

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39
Q

O2- has the atomic number 8 and the mass number
16. How many protons, neutrons and electrons are
in this ion?

A

Protons: 8
Electrons: 10
Neutrons: 8

40
Q

Why do elements in groups 1, 2, 6 and 7

readily form ions?

A 
So they become more stable and
achieve the electron structure of the
noble gases (group 8).
41
Q

What type of ions do elements in group 1 and 2

form?

A

Cations (positive)

  • Group 1 metals will form 1+ ions
  • Group 2 metals will form 2+ ions
42
Q

What type of ions do elements in groups 6 and 7

form?

A

They are non-metals so form anions
(negative)
- Group 6 will form 2- ions
- Group 7 will form 1- ions

43
Q

What does it mean if an ionic compound ends in

-ide?

A

The compound contains 2 elements

44
Q

What does it mean if an ionic compound ends in

-ate?

A

The compound contains at least three

elements, one of which is oxygen.

45
Q

Describe the structure of an ionic compound

A

Lattice structure.
Regular arrangement of ions.
Ions held together by strong electrostatic
forces between ions with opposite charges.

46
Q

What is the chemical formula of sodium oxide,

formed from Na+ and O2-?

A

Na2O
Charges must balance so 2 sodium ions
are required.

47
Q

What is the chemical formula of magnesium

hydroxide, formed from Mg2+ and OH-?

A

Mg(OH)2
Charges must balance so 2 hydroxide
ions are required

48
Q

What is a covalent bond?

A

A bond formed when an electron pair is

shared between two atoms

49
Q

What forms as a result of covalent bonding?

A

A molecule

50
Q

Draw a dot and cross diagram for the formation of
methane (CH4
)

A

https://www.physicsandmathstutor.com/pdf-pages/?pdf=https%3A%2F%2Fpmt.physicsandmathstutor.com%2Fdownload%2FChemistry%2FGCSE%2FNotes%2FEdexcel%2F1-Key-Concepts-in-Chemistry%2FFlashcards%2520-%2520Topic%25201%2520Key%2520Concepts%2520in%2520Chemistry%2520-%2520Edexcel%2520Chemistry%2520GCSE.pdf

page 101

51
Q

True or false?

‘Covalent bonds are weak’

A

FALSE

52
Q

Which is smaller, an atom or a molecule?

A

An atom

Simple molecules consist of atoms
joined by strong covalent bonds within
the molecule.

53
Q

Why do ionic compounds have high melting points?

A

Strong electrostatic attraction between
the positive and negative ions requires a
lot of energy to overcome

54
Q

When do ionic compounds electricity? Why?

A

When molten or aqueous (dissolved in
water) because the ions are charged and
free to move. When solid, the ions are fixed
in an ionic lattice so can’t move.

55
Q

Why do simple molecular compounds have low

melting and boiling points?

A

They have weak intermolecular forces
(forces between molecules) which only
require a little energy to overcome.

56
Q

Do simple molecular compounds conduct electricity?

Why / why not?

A

No because there are no charged

particles.

57
Q

Do giant covalent structures have a high melting

point? Explain your answer

A

Yes because they have lots of strong
covalent bonds which require a lot of
energy to break

58
Q

How do metals conduct electricity and heat?

A

The positive ions are fixed in a sea of
delocalised electrons. These electrons
are free to move and carry charge /
energy.

59
Q

True or false?

‘Metals are insoluble in water’

A

TRUE

60
Q

Name two giant covalent structures formed from

carbon atoms

A

Graphite

Diamond

61
Q

Describe the structure and properties relating to

graphite

A

Each carbon atom bonded to 3 other carbon atoms.
Layers of hexagonal rings of carbon atoms.
Weak intermolecular forces between layers.
One delocalised electron per carbon atom.

62
Q

Describe and explain the properties of graphite

A

Graphite is soft / slippery because there are only weak
intermolecular forces between layers which allow the
layers to slide over one another.
Graphite conducts electricity because there is one
delocalised electron per carbon atom. The delocalised
electrons are mobile charges.

63
Q

Describe the structure of diamond

A

All carbon atom are covalently bonded to
four other carbon atoms.
No delocalised electrons.

64
Q

Describe the properties of diamond

A

Very hard.
Very high melting point.
Doesn’t conduct electricity as there are
no charged particles.

65
Q

What are the uses of graphite? Why?

A

Electrodes because graphite conducts
electricity and has a high melting point.
Lubricant because it’s slippery (the layers
in graphite can slide over each other).

66
Q

Why is diamond used in cutting tools?

A

It’s very hard

67
Q

What is a fullerene?

A

A fullerene is a molecule made of
carbon, shaped like a closed tube or
hollow ball

68
Q

Name two fullerenes

A

Graphene

C60 (buckminsterfullerene)

69
Q

What are the properties of the fullerene

C60?

A
  • Slippery due to weak intermolecular forces.
  • Low melting point.
  • Spherical.
  • Strong covalent bonds between carbon atoms
    in a molecule.
  • Large surface area
70
Q

What are the properties of graphene?

A
  • High melting point due to covalent
    bonding between carbon atoms.
  • Conducts electricity because it has
    delocalised electrons.
71
Q

Why is graphene useful in electronics?

A

It is extremely strong and has delocalised
electrons which are free to move and carry
charge.
It is only one atom thick as it is a single layer of
graphite

72
Q

What is a polymer? Name a polymer

A

Long chain molecules formed from many
monomers.
E.g. poly(ethene)

73
Q

Draw the structure of poly(ethene)

A

https://www.physicsandmathstutor.com/pdf-pages/?pdf=https%3A%2F%2Fpmt.physicsandmathstutor.com%2Fdownload%2FChemistry%2FGCSE%2FNotes%2FEdexcel%2F1-Key-Concepts-in-Chemistry%2FFlashcards%2520-%2520Topic%25201%2520Key%2520Concepts%2520in%2520Chemistry%2520-%2520Edexcel%2520Chemistry%2520GCSE.pdf

Page 147

74
Q

What bond is formed between carbon atoms in

polymer molecules?

A

Covalent bonds

75
Q

What are the properties of metals?

A
  • High melting point.
  • High density.
  • Good conductors of electricity.
  • Malleable and ductile.
  • Generally shiny.
76
Q

Explain why metals are malleable

A

The atoms are arranged in uniform rows

which can slide over one another

77
Q

Explain why metals can conduct electricity

A

The electrons in the metal are charges

that can move.

78
Q

What are the general properties of non-metals?

A
  • Low boiling points.
  • Poor conductors of electricity.
  • Brittle when solid.
79
Q

List the limitations of the following models when
representing ionic compounds: dot and cross, 2D
diagrams and 3D diagrams

A

● Dot and cross - no lattice structure or ionic bonds.
● 2D diagrams - only shows one layer, doesn’t show
formation of ions.
● 3D diagrams - shows spaces between the ions,
doesn’t show charges.

80
Q

List the limitations of the following models when
representing covalent molecules: dot and cross and
ball and stick

A

● Dot and cross - doesn’t show relative sizes of
atoms or intermolecular forces.
● Ball and stick - bonds shown as sticks rather
than forces, doesn’t show how covalent bonds
form.

81
Q

How do you calculate the relative formula mass of a

compound?

A

Add together all the relative atomic

masses of the atoms in the compound.

82
Q

What is the empirical formula? What 2 values could
be used to calculate the empirical formula of a
simple compound?

A

● The empirical formula is the smallest whole
number ratio of the atoms of each element in a
compound.
● Reacting masses or percentage composition
can be used to calculate the empirical formula.

83
Q

What is the empirical formula for Fe2O4

?

A

FeO2

84
Q

What is the molecular formula?

A

Actual number of atoms of each element

in a compound.

85
Q

Describe an experiment to work out the empirical

formula of magnesium oxide

A

● Weigh a sample of magnesium.
● Heat the sample in a crucible.
● Weigh the mass of magnesium oxide at the end.
● Calculate the mass of oxygen (this is the increase of mass).
● Calculate the moles of magnesium and oxygen using the
experimental mass and relative atomic mass.
● Work out the whole number ratio of the number of moles of
magnesium to oxygen.

86
Q

What is the law of conservation of mass?

A

No matter is lost or gained during a

chemical reaction

87
Q

If a reaction is carried out in a closed system, what
can you say about the total mass of the reaction
throughout the experiment?

A

Mass stays constant.

88
Q

If a reaction is carried out in an open flask and a gas
is produced, what can you say about the total mass
of the reaction throughout the experiment?

A

Mass decreases as the gas escapes

89
Q

52g of calcium reacts with oxygen to form 79g of
calcium oxide. Using the law of conservation of
mass, what mass of oxygen is needed?

A

79 - 52 = 27

Mass of oxygen = 27g

90
Q

What equation links mass, moles and relative atomic

mass? (higher only)

A

Mass (g) = Moles x Relative atomic mass (Mr)

91
Q

How can you calculate concentration in

g/dm3?

A

Concentration(g/dm^3

) = Mass (g) /Volume (dm^3)

92
Q

What is the Avogadro constant?

higher only

A

The number of atoms, molecules or ions in
one mole of a given substance.
The value of the constant is 6.02 x 10^23
.

93
Q

What is the mass of 20 moles of calcium carbonate,
CaCO3
?

A

Mass (g) = Moles x Relative atomic mass (Mr)
Mr = 100
20 x 100 = 2000 g

94
Q

What formula links the Avogadro constant, moles
and number of particles?
(higher only)

A

Number of particles =Avogadro constant x Moles

95
Q

How many atoms are in 3 moles of copper?

A

Number of atoms =Avogadro’s constant x Moles
= 6.02 x 10^23 x 3
= 1.81 x 10^24

96
Q

What is a limiting reagent in a chemical reaction?

A

The chemical that is used up first in a reaction,
preventing the formation of more product.

Typically, an excess of one of reactants is used
to ensure that the other reactant is completely
used up.

Chemistry (30 decks)

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