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Chemistry > 1. Key Concepts In Chemistry > Flashcards

1. Key Concepts In Chemistry Flashcards

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1
Q

Why has the Dalton model of an atom changed over time?

A

Because of the discovery of subatomic particles.

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2
Q

What is 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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3
Q

Relative charge of proton

A

+1

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4
Q

Relative mass of proton

A

1

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5
Q

Relative charge of neutron

A

0

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6
Q

Relative mass of neutron

A

1

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7
Q

Relative charge of electron

A

-1

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8
Q

Relative mass of electron

A

1/2000

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9
Q

Why do atoms contain equal number of protons and electrons?

A

Atoms are neutral as the charges of a proton is +1 and an electron is -1, therefore the amount of protons = amount of electrons, so the charges cancel each other out.

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10
Q

TRUE OR FALSE: The nucleus of an atom is very small compared to the overall size of an atom.

A

True

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11
Q

Where is the most mass of an atom concentrated?

A

In the nucleus.

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12
Q

What is mass number?

A

Number of protons + neutrons

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13
Q

TRUE OR FALSE: The atoms of the same element have different number of protons in the nucleus.

A

False - they have the same number of protons and this number is unique to that element.

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14
Q

What are isotopes?

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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15
Q

What is atomic number?

A

Number of protons (= number of electrons if it’s an atom, because atoms are neutral)

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16
Q

Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

A
  • The relative atomic mass is calculated using the abundance of different isotopes and because it is an average it can lead to the relative atomic mass not being a whole number
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17
Q

How to calculate relative atomic mass

A

R.A.M = (mass of isotope-A x % of isotope-A) + (mass of isotope-B x % of isotope-B) / 100

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18
Q

A sample of chlorine gas is a mixture of 2 isotopes, chlorine-35 and chlorine-37. These isotopes occur in specific proportions in the sample i.e. 75% chlorine-35 and 25% chlorine-37. Calculate the R.A.M of chlorine in the sample.

A

R.A.M = (35 x 75) + (37 x 25) / 100

R.A.M = 3550/100
R.A.M = 35.5

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19
Q

How did Mendeleev arrange the elements, known at that time, in a 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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20
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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21
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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22
Q

What are vertical columns called in the periodic table?

A

Groups

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23
Q

What are horizontal rows called in the periodic table?

A

Periods

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24
Q

What do elements in the same group have in common?

A

They have the same amount of electrons in their outer shell, which gives them similar chemical properties. The number of outer shell electrons determines how an atom reacts.

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25
Q

How are elements arranged in the periodic table?

A

In order of increasing atomic number in rows called periods.

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26
Q

What are metals

A

Elements that react to form positive ions by losing electrons.

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27
Q

Where in the periodic table are metals found.

A

To the left and towards the bottom of the periodic table - most elements are metals.

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28
Q

What are non-metals

A

Elements that react to form negative ions by gaining electrons.

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29
Q

What is electronic configuration of an element and what does it tell us?

A

Tells you how many electrons are in each shell around the nucleus.
- For example, sodium has 11 electrons: 2 in its most inner shell, then 8, then 1 in its outermost shell. 2.8.1

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30
Q

How does group number relate to electronic configuration?

A

The group an electron is in tells you how many electrons are in its outermost shell, e.g. group 1 elements have 1 electron in their outer shell.

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31
Q

How does period number relate to electronic configuration?

A

The period an electron is in tells you which number shell an element’s outermost
electron is found in, e.g. period 3 elements have their outermost electrons in
shell 3.

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32
Q

How are ionic bonds formed?

A

By the transfer of electrons between atoms to produce cations and anions.

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33
Q

What happens when there is ionic bonding between a metal and a non-metal?

A

Electrons in the outer shell of the metal atom are transferred.
- Metal atoms lose electrons to become a positively charged ion (cation).
- Non-metal atoms gain electrons to become a negatively charged ion (anion).

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34
Q

How is electron transfer during the formation of an ionic compound represented?

A

Dot and cross diagram.

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35
Q

What is an ion?

A

An atom or group of atoms with a positive or negative charge.

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36
Q

How to calculate number of electrons in an ion (as ions have different number of electrons to protons)

A
  • Work out how many electrons an atom of the element would have (same as proton number).
  • Work out how many electrons have been gained or lost.
  • Calculate number of electrons in atom plus electrons gained or minus electrons lost.
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37
Q

What group are noble gases found?

A

Group 0

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38
Q

How many electrons will metals in group 1 gain/lose and what type of ions will they form?

A

Lose 1 electron , form +1 ions

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39
Q

How many electrons will metals in group 2 gain/lose and what type of ions will they form?

A

Lose 2 electrons , form +2 ions

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40
Q

How many electrons will nonmetals in group 6 gain/lose and what type of ions will they form?

A

Gain 2 electrons , form 2- ions

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41
Q

How many electrons will nonmetals in group 7 gain/lose and what type of ions will they form?

A

Gain 1 electron , form 1- ions

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42
Q

What overall charge will an ionic compound have?

A

An overall charge of 0 because you need to balance out the + and - charges.

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43
Q

Explain the use of the –ide ending in the names of compounds

A

-ide means the compound contains 2 elements (one is the nonmetal negative ion)

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44
Q

Explain the use of the –ate ending in the names of compounds

A

-ate​ ​means​ ​the​ ​compound contains​ ​at​ ​least​ ​3​ ​elements,​ ​one​ ​of​ ​which​ ​is​ ​oxygen.

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45
Q

What charge are ions ending in -ide or -ate in the names of compounds?

A

Negative

46
Q

How do you deduce the formulae of ionic compounds?

A

You​ ​need​ ​to​ ​balance​ ​out​ ​the​ ​+​ ​and​ ​-​ ​charges to​ ​make​ ​the​ ​overall​ ​charge​ ​0.​ ​You​ ​do​ ​this​ ​by​ ​writing​ ​a​ ​little​ ​number​ ​below​ ​the​ ​element e.g.​ ​Cl​₃​​ ​or​ ​for​ ​ions​ ​with​ ​more​ ​than​ ​one​ ​element​ ​you​ ​draw​ ​a​ ​bracket​ ​round​ ​first​ ​e.g.​ ​(SO​₄)​₂

47
Q

Explain the structure of an ionic compound

A
  • A giant structure of ions = ionic compound.
  • Held together by strong electrostatic forces of attraction between oppositely charged ions.
  • The forces act in all directions in the lattice, which is called ionic bonding.
  • The lattice has regular arrangement of ions.
48
Q

How are covalent bonds formed?

A

When there is a pair of shared electrons between two non-metal atoms.

49
Q

Covelently bonded substances can be…

A
  • Small molecules (e.g. water)
  • Large molecules (e.g. polymers)
  • Giant covalent structures (e.g. diamond)
50
Q

Are covalent bonds weak or strong?

A

Very strong

51
Q

How can covalent bonds be presented?

A

Dot and cross diagrams

52
Q

In covalent bonding, the atoms need to obtain electrons which give them a full outer shell.

A

For example: chlorine (Cl₂)
- Each chlorine atom begin with 7 electrons in its outer shell.
- By sharing 1 pair of shared electrons in a single covalent bond, each Cl atom obtains a full outer shell with 8 electrons.

53
Q

Types of covalent bonds

A

Single - 1 shared pair of electrons.
Double - 2 shared pairs of electrons.
Triple - 3 shared pairs of electrons.

54
Q

Properties of ionic compounds

A
  • Made up of a metal and a nonmetal.
  • Regular structures (giant ionic lattices) with strong electrostatic forces of attraction in all directions between oppositely charged ions.
  • High melting and boiling points, as a lot of energy is required to break the strong bonds.
  • Conduct electricity when melted or dissolved in water because the ions are free to move and carry current.
  • Do not conduct when solid because the ions are fixed and are not able to move, carrying charge with them.
  • Often dissolve in water to form an aqueous solution.
55
Q

Properties of simple molecular compounds

A
  • Usually gases or liquids.
  • Low melting and boiling points.
  • Made up of nonmetal elements.
  • Weak intermolecular forces between molecules, which are broken in boiling or melting, not the covalent bonds.
  • The intermolecular forces increase with the size of molecules, so larger molecules have higher melting and boiling points.
  • Don’t conduct electricity because small molecules do not have an overall electric charge, although some break down in water to form ions which can conduct electricity.
  • Many are insoluble in water, but some are soluble because they can form intermolecular forces with water which are stronger than those between water molecules or their own molecules already (e.g. CO₂ and NH₃ are soluble)
56
Q

Properties of giant covalent structures

A
  • Made up of nonmetal elements.
  • Solids with very high melting points.
  • All atoms in these structures are linked to other atoms by strong covalent bonds, and these bonds must be overcome to melt or boil these substances.
  • Some can conduct electricity, whereas others can’t.
57
Q

Properties of metals

A
  • Consist of giant structures of atoms arranged in a regular pattern.
  • Electrons in outer shell are delocalised and so are free to move through the whole structure.
  • The sharing of delocalised electrons gives rise to strong metallic bonds.
  • Most have high melting and boiling points.
  • Conduct heat and electricity because of the delocalised electrons, depends on the ability for electrons to move throughout the metal.
  • Layers of atoms in metals are able to slide over each other, so metals can be bent and shaped.
  • Insoluble in water, but some will react with it instead.
  • High density.
58
Q

What are graphite and diamond an example of?

A

Giant covalent structures

59
Q

Structure of diamond

A

Each carbon is joined to 4 other carbons covalently.
- It’s very hard, has a very high melting point and does not conduct electricity.

60
Q

Structure of graphite

A

Each carbon is covalently bonded to 3 other carbons, forming layers of hexagonal rings, which have no covalent bonds between the layers.
- The layers can slide over each other due to no covalent bonds between layers, but weak intermolecular forces, meaning graphite is soft and slippery.

One electron from each carbon atom is delocalised.
- This makes graphite similar to metals, because of its delocalised electrons.
- It can conduct electricity, unlike diamond.

61
Q

Why is graphite used to make electrodes?

A

Because graphite can conduct electricity.

62
Q

Why is graphite used to act as a lubricant?

A

Graphite has weak intermolecular forces and no covalent bonds between the layers, therefore it is soft and slippery.

63
Q

Why is diamond used in cutting tools?

A

Because diamond is hard due to its rigid structure.

64
Q

What is C60 known as?

A

Buckminsterfullerene

65
Q

Properties of graphene

A
  • Single layer of graphite.
  • Very high melting point.
  • Very strong because of its large regular arrangement of carbon atoms joined by covalent bonds.
  • 2D hexagonal shape.
  • Like graphite, conducts electricity due to delocalised electrons that move freely.
66
Q

Define fullerene

A

A fullerene is a molecular form of the element carbon with hollow shapes.

67
Q

Examples of fullerenes

A
  • Nanotubes
  • Buckyballs (buckminsterfullerene)
68
Q

Properties of nanotubes

A
  • Cylindrical fullerenes with very high length to diameter ratios.
  • A nanotube resembles a layer of graphene, rolled into a tube shape.
  • High tensile strength, so they are strong in tension and resist being stretched.
  • Strong.
  • Conduct electricity due to delocalised electrons.
69
Q

Properties of buckyballs (buckminsterfullerene)

A
  • Spheres or squashed spheres of carbon atoms.
  • Made up of large molecules so not classed as giant covalent networks.
  • Weak intermolecular forces between buckyballs.
  • These need little energy to overcome, so substances consisting of buckyballs are slippery and have lower melting points than graphite or diamond.
70
Q

What do simple polymers consist of?

A

Consist of large molecules containing chains of carbon atoms.

71
Q

Atoms in polymer molecules are linked to other atoms by what type of bonds?

A

Strong covalent bonds.

72
Q

Why are polymer molecule substances solid at room temperature?

A

Because intermolecular forces between polymer molecules are relatively strong.

73
Q

What does it mean when metals are described as being malleable?

A

Malleable = the layers of atoms in metals are able to slide over each other.

74
Q

Why can metals conduct electricity?

A

Because the delocalised electrons can move.

75
Q

What is relative formula mass?

A

The relative formula mass of a compound is equal to the sum of the relative atomic masses of all the atoms in the compound.

76
Q

What is relative atomic mass?

A

The relative atomic mass of an atom is the average mass of the atom.

77
Q

Calculate relative formula mass (Mr) of hydrochloric acid (HCl)

A

Identify the relative atomic mass (Ar) of each element in the compound:
1) Ar of hydrogen = 1
2) Ar of chlorine = 35.5
3) 35.5 + 1 = 36.5
Therefore, the total Mr of HCl is 36.5

78
Q

What are the steps when calculating the relative formula mass of a compound?

A

1) Identify the Ar of each element.
2) Multiply the Ar by the number of atoms.
3) Add up the totals for each element.

79
Q

What is empirical formula?

A

The empirical formula of an ionic compound is the simplest ratio of ions possible.

80
Q

Work out empirical formula of: Ca²⁺ and F⁻

A

In order for the charges to balance out and for the overall formula to be neutral, there must be 2 F⁻ ions per Ca²⁺ ion.

The empirical formula is CaF₂

81
Q

Work out empirical formula of: Na⁺ and SO₄²⁻

A

In order for the charges to balance out and for the overall formula to be neutral, there must be 2 Na⁺ ions per SO₄²⁻ ion.

The empirical formula is Na₂SO₄-

82
Q

Molecular vs Empirical Formula

A

The molecular formula shows the actual amount of atoms which make up a molecule. The empirical formula shows the simplest ratio of atoms which make up a molecule.
Molecular formula:
e.g. glucose = C₆H₁₂O₆
Empirical formula:
e.g. glucose = CH₂O

83
Q

How to calculate empirical formula from reacting masses

A

1) Work out moles of each using moles = mass / molar mass.
2) Work out the ratio of moles.
3) Multiply the ratio so that you get the smallest whole numbers possible.
4) Find the formula by multiplying each element by their number in the ratio
(remember to use little numbers not a big number at the front).

84
Q

Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide

A

1) Measure the mass of the solid that is going to be reacted (e.g. magnesium ribbon).
2) React the solid with the oxygen in the air by heating in a crucible.
3) Find out the mass of the product (magnesium oxide).
4) Calculate the mass of oxygen gained (final mass - initial mass).
5) Use the mass of magnesium and oxygen to calculate the empirical formula.

85
Q

Deduce: empirical formula of a compound from the formula of its molecule

A

If you have a common multiple e.g. Fe₂ and O₄, the empirical formula is the simplest whole number ratio, which would be FeO₂.
If there is no common multiple, you already have the empirical formula.

86
Q

Deduce: molecular formula of a compound from its empirical formula and its relative molecular mass

A

1) Find relative molecular mass of the empirical formula.
2) Divide relative molecular mass of compound by that of the empirical formula.
3) Multiply the number of each type of atom in the empirical formula by this number.
e.g. if answer was 2 and the empirical formula was Fe₂O₃ then the molecular formula would be empirical formula x2 = Fe₄O₆

87
Q

Law of conservation of mass

A

No atoms are lost or made during a chemical reaction so the mass of the products = mass of the reactants

88
Q

Explain the law of conservation of mass applied to a closed system
including a precipitation reaction in a closed flask

A

Precipitate that forms is insoluble and is a solid,
as all the reactants and products remain in the sealed reaction container then it is easy to show that the total mass is unchanged.

89
Q

Explain the law of conservation of mass applied to a non-enclosed
system including a reaction in an open flask that takes in or gives out a gas

A

Does not hold for a reaction in an open flask that takes in or gives out a gas, since mass will change from what it was at the start of the reaction as some mass is lost when the gas is given off.

90
Q

Calculate masses of reactants and products from balanced equations, given the mass of one substance

A

1) Find moles of that one substance: moles = mass / molar mass
2) Use balancing numbers to find the moles of desired
reactant or product (e.g. if you had the equation:
2NaOH + Mg -> Mg(OH)₂ + 2Na, if you had 2 moles of
Mg, you would form 2x2=4 moles of Na)
3) Mass = moles x molar mass(of the reactant/product) to
find mass

91
Q

How to calculate concentration of a solution

A

concentration (g dm³) = mass of solute (g) / volume (dm³)

92
Q

How to calculate mass of solute in a given volume of a known concentration

A

mass = concentration x volume i.e. g = g/dm³ x dm³

93
Q

When can substances in chemical reactions seem to gain or lose mass?

A

If one of the reactants or products is a gas, as they can escape into the atmosphere.

94
Q

One mole of particles of a substance is defined as :

A

The Avogadro constant number of particles (6.02 x 10²³ atoms, molecules, formulae or ions) of that substance and a mass of ‘relative particle mass’ g

95
Q

What is the mass of 1 mole of carbon

A

12g , which is the same as the relative atomic mass of carbon.

96
Q

How to calculate number of moles

A

mass / relative formula mass

97
Q

How many moles are there in 24g carbon

A

24 / 12
= 2 moles

98
Q

TRUE OR FALSE: The mass of one mole of a substance in grams is numerically equal to its relative formula mass.

A

True
For example, the Ar of Iron is 56, so one mole of iron weighs 56g

99
Q

What is a limiting reactant 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.

100
Q

Limiting reactant when neutralising an acid

A

When neutralising an acid, it is important that there’s no acid left after the reaction. To make sure this happens, we use an excess amount of the other reactant.

101
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.
102
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.

103
Q

Describe the atomic model proposed by JJ Thomson

A

Plum pudding model.

Negatively charged electrons scattered through a positively charged material.

104
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.

105
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.
106
Q

When are atoms most stable?

A

When they have full electron shells.

107
Q

What is an ionic bond?

A

A bond between a metal and non-metal involving the transfer of electrons.

108
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

109
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

110
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

111
Q

How many atoms are in 3 moles of copper?

A

Number of atoms = Avogadro’s constant x Moles
= 6.02 x 10²³ x 3
= 1.81 x 10 to power of 24

112
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

Chemistry (7 decks)

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