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GCSE Chemistry P2 > Topic 1 - Key concepts of Chemistry > Flashcards

Topic 1 - Key concepts of Chemistry Flashcards

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1
Q

What was Dalton’s model of the atom like?

A
  • atoms cannot be broken down into anything simpler
  • the atoms of a given element are identical to each other
  • the atoms of different elements are different from one another
    during chemical reactions atoms rearrange to make different substances.
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2
Q

Why has Daltons’s model of the atom changed?

A

The Dalton model has changed over time because of the discovery of subatomic particles.

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3
Q

Where are the subatomic particles located?

A

Neutron - Nucleus
Proton - Nucleus
Electron - Shells orbiting the nucleus

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4
Q

What is the mass and charge of a proton?

A

Mass - 1

Charge - 1

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5
Q

What is the mass and charge of a neutron?

A

Mass - 1

Charge - 0

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6
Q

What is the mass and charge of a Electron?

A

Mass - 0.0005

Charge - (-1)

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7
Q

Which two subatomic particles are of equal number in a neutral atom?and why?

A

Protons and Electrons. This is so that the atom is in a neutral state as it has a charge of 0.

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8
Q

Where is most of the mass of an atom concentrated?

A

The mass of electrons is very small compared to protons and neutrons. Since a nucleus contains protons and neutrons, most of the mass of an atom is concentrated in its nucleus.

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9
Q

What is the mass number?

A

The mass number is the number above the element symbol on the periodic table. It tells you the total number of neutrons and protons in an atom.

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10
Q

What is the atomic number?

A

The stomic number is the number below the element symbol on the periodic table. It tells you the total number of protons in an atom.

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11
Q

What must every atom of the same element have in common?

A

The same number of protons, the number of protons is unique to each element.

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12
Q

What is the relative atomic mass?

A

The relative atomic mass of an element is the weighted average of the relative atomic masses of the isotopes in the element. This means that some relative atomic masses are not whole number

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13
Q

What are isotopes?

A

Isotopes are different forms of the same element, which have the same number of protons but a different number of neutrons. They have the same atomic number but different mass numbers.

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14
Q

What are the main differences between Mendeleev’s periodic table and the modern periodic table?(4)

A
  • The current periodic table has its elements ordered in increasing atomic number whereas Mendeleev’s periodic table has its elements ordered in increasing atomic weight.
  • More elements have been discovered and added in the current periodic table: such as the noble gases group 0.
  • There are no gaps in the current periodic table. The gaps in Mendeleev’s table have been filled by the elements that were discovered.
  • Transition metals have been listed in a separate block whereas Mendeleev included them with other elements.
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15
Q

How did Mendeleev use his table to leave gaps for undiscovered elements?

A

Mendeleev left gaps in his table to place elements not known at the time. By looking at the chemical properties and physical properties of the elements next to a gap, he could also predict the properties of these undiscovered elements.

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16
Q

Why did Mendeleev reverse some pairs in his table instead of following the trend of increasing atomic weight?

A

When he ordered elements to increasing atomic weight, he noted that the chemical properties of the elements and their compounds showed a periodic trend. He then arranged the elements by putting those with similar properties below each other into groups so he had to switch the order of a few elements to keep the groups consistent.

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17
Q

What are groups and periods?

A

The rows of elements in the periodic table, called periods, are in order of increasing atomic number.

The vertical columns of element in the periodic table, called groups, are where the elements have similar properties.

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18
Q

Where are metals and non-metals?

A

The metal elements are found on the left hand side of the periodic table, and the non-metal elements are found on the right. Imagine a zig-zag line, starting at B-Al-Si, separating metals from non-metals.

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19
Q

What is the limit for the number of electrons in the first three shells?

A

1st- 2 electrons
2nd- 8 electrons
3rd - 8 electrons

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20
Q

How can you work out the electronic configuration of an element using its period and group?

A

The number of shells which contain electrons is the same number as the period of the element. The group number tells you how many electrons ocupy the outer shell of an element.

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21
Q

What is a covalent bond?

A

A covalent bond is a bond that forms when a pair of electrons is shared between non-metal atoms.

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22
Q

What are simple covalent structures?

A

These are made of molecules containing a few atoms joined together by covalent bonds.

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23
Q

Explain the structure of hydrogen (scs) in dot cross diagrams

A

Hydrogen - H2. Hydrogen atoms have one electron, so they need one more to complete the first shell. They can form a single covalent bond with another hydrogen atom to achieve this.

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24
Q

Explain the structure of water (scs) in dot cross diagrams

A

Water - H20. In water molecules, an oxygen atom shares a pair of electrons with two H atoms to form each two single covalent bonds

25
Q

Explain the structure of hydrogen chloride (scs) in dot cross diagrams

A

Hydrogen Chloride - HCl. This is very similar to H2 - both atoms only need one more electron to complete their outer shells.

26
Q

Explain the structure of methane (scs) in dot cross diagrams

A

Methane - CH4. Carbon has four outer electrons, which is half a full shell. It can form four covalent bonds with hydrogen atoms to fill up its outer shell.

27
Q

Explain the structure of oxygen (scs) in dot cross diagrams

A

Oxygen - O2. An oxygen atom needs two more electrons to complete its outer shell. In oxygen gas, each oxygen atom forms a double covalent bond (a bond made of two shared electron pairs) with another oxygen atom.

28
Q

Explain the structure of carbon dioxide (scs) in dot cross diagrams

A

In carbon dioxide molecules, a carbon atom shares two pairs of electrons with two oxygen atoms to form two double covalent bonds

29
Q

Give properties of SCS and explain.

A
  • Strong covalent bonds in atoms with molecules but intermolecular forces of attraction are weak
  • Low melting and boiling points as only intermolecular forces need to be broken so low temperature is needed for these weak forces.
  • Don’t conduct electricity as they don’t contain any free ions or electrons
30
Q

Give properties of GCS and explain. (4)

A
  • All atoms are bonded to each other by strong covalent bonds.
  • High melting and boiling points as a lot of energy is needed to break the bonds.
  • They generally don’t contain charged particles, so they don’t conduct electricity.
  • They aren’t soluble in water
31
Q

Give examples of carbon based GCS. (4)

A
  • Diamond
  • Graphite
  • Graphene
  • Nanotubes
32
Q

What is diamond and what are its properties? (4)

A

> Network of carbon atoms that each form four covalent bonds. >The strong covalent bonds take lots of energy to break, so diamond has a high melting point.
The strong covalent bonds also hold the atoms in a rigid lattice structure, making diamond really hard - it’s used to strengthen cutting tools (e.g. saw teeth and drill bits).
It doesn’t conduct electricity because it has no free electrons or ions.

33
Q

What is Graphite and what are its properties? (6)

A

> Sheets of carbon atoms arranged in hexagons. The layers are only held together weakly so they’re free to move over each other. This makes graphite soft and slippery, ideal as a lubricating material.
High melting point - the covalent bonds in the layers need loads of energy to break.
Only three out of each carbon’s four outer electrons are used in bonds, so each carbon atom has one electron that’s delocalised (free) and can move. So graphite conduct electricity and is often used to make electrodes.
Graphite is also used in pencil it slips layer by layer on to the paper leaving a black mark.

34
Q

What is Graphene?

A

Graphene is a sheet of carbon atoms joined together in hexagons. This is one atom thick meaning it is a 2d compound.

35
Q

What are Nanotubes?(3)

A

> Nanotubes are also fullerenes.
They are like tiny cylinders of graphene - so they conduct electricity.
They also have a high tensile strength (they don’t break when stretched) so can be used to strengthen materials without adding much weight. For example, they can be used to strengthen sports equipment that needs to be strong but also lightweight (e.g. tennis racket)

36
Q

What is buckministerfullerene?

A

Buckministerfullerene has the molecular formula C60 and forms a hollow sphere made up of 20 hexagons and 12 pentagons. It’s a stable molecule that forms soft brownish-black crystals

37
Q

What is ionic bonding?

A

Ionic Bonding is when a metal and a non-metal react together, the metal atom loses electrons to form a cation and the non-metal gains these electrons to form an anion. These oppositely charged ions are strongly attracted to one another by electrostatic forces, called an ionic bond.

38
Q

Give six examples of Simple Covalent Structures.

A
  • Hydrogen
  • Water
  • Hydrogen Chloride
  • Oxygen
  • Methane
  • Carbon dioxide
39
Q

Give properties of ionic compounds. (4)

A

> Ionic compounds always have giant ionic lattice structures. The ions form a closely packed regular lattice.
There are very strong electrostatic forces of attraction between oppositely charged ions, in all directions.
Ionic compounds have high melting and boiling points due to the strong attraction between the ions, a large amount of energy needed to overcome this attraction.
Solid ionic compound don’t conduct electricity as the ions are in a fixed place and can’t move. But molten ionic compounds, the electrons are free to move and will carry an electric current. Many ionic compounds dissolve easily in water. The ions separate and are all free to move in the solution, so they’ll carry an electric current.

40
Q

How are cations formed?

A

Cations (positve ions) form when atoms lose electrons (more protons than electrons).

41
Q

What are ions?

A

Ions are charged particles. When atoms lose or gain electrons, all they are trying to do is get a full outer shell. The number of electrons lost or gained is the same as the charge on the ion.

42
Q

How are anions formed?

A

Anions (negative ions) form when atoms gain electrons (more electrons than protons).

43
Q

What does the ending -ide mean in the names of ionic compunds?

A

This means that there are 2 elements in a compound which have only non-metal atoms. E.G. Copper sulfide - CuS

44
Q

What does the ending -ate mean in the names of compunds?

A

This means that there are 3 elements in a compound and the third is always oxygen. These have non-metal atoms and an oxygen atoms. E.G. Copper sulfate - CuSO4

45
Q

Describe metallic bonding

A

In metals, the electrons leave the outer shells of metal atoms, forming positive metal ions and a ‘sea’ of delocalised electrons.

46
Q

Describe the physical properties of metals. (5)

A
  • Electrostatic forces between metal ions and delocalised electrons are very strong, so lots of energy needed to break them. Therefore high boiling and melting points so they are shy solids at room temp.
  • Not soluble in water.
  • More dense than non-metals as the ions are packed closer together.
  • Layers of atoms in a pure metal can slide over each other making metals malleable.
  • Delocalised electrons can carry electrical current and thermal energy, so metals are good conductors of heat and electricity.
47
Q

What is the relative formula mass?

A

The RFM of a compound is the relative atomic masses of all the atom in its formula added together.

48
Q

Find the empirical formula of glucose, C6H12O6

A

Divide by largest factor, 6

6/6=1
12/6=2
6/6=1

So CH2O

49
Q

How can you use empirical formula of a compound with its RFM to find its molecular formula.

A 
Find RFM (Mr) of the empirical formula. C2H6N = 44 
Divide Mr of the compound by the Mr of the empirical formula. 88/44=2
Multiply the answer with the empirical formula. C2H6N x 2 = C4H12N2
50
Q

Describe the law of conservation of mass.

A

During chemical reactions, no atoms are destroyed or created. This means there are the same number and types of atoms on each side of the reaction equation. You can see this is a close system. The total mass before and after doesn’t change. (precipitation reaction)

51
Q

How can the concentration of a solution be calculated?

A

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

52
Q

What is Avagadro’s constant?

A

As the mole quantity is based on the relative masses of atoms, there is a relationship between a mole of a substance and the number of particles it contains. As atoms and molecules are so small the number in a mole is very large. This number is called the Avogadro constant - 6.02 x 10^23.

53
Q

How can the number of particles be calculated using Avagadro’s constant?

A

Avogadro’s constant x amount of substance (mol)

54
Q

How do you calculate mass of a number of moles?

A

The mass of a given number of moles of substance is calculated using: mass in g = Mr x number of moles or Mass = Mr x mol

MASS EQUALS MR MOLE.

55
Q

What are limiting reactants and excess reactants?

A

A reaction finishes when one of the reactants is all used up. The other reactant has nothing left to react with, so some of it is left over:

  • The reactant that is all used up is called the limiting reactant
  • The reactant that is left over is described as being in excess
56
Q

How can you calculate the amount of product from the Limiting Reactant? (5)

A

1) Write out the balanced equation.
2) Work out relative formula masses (M,) of the reactant and product you are interested in.
3) Find out how many moles of the substance you know the mass of.
4) Use the balanced equation to work out how many moles of product will be made by this many moles of reactant.
5) Use the number of moles to calulate the mass

57
Q

Give properties of metals. (5)

A
  • Strong and hard
  • High densities
  • Malleable because layers of atoms in a pure metal can slide over each other.
  • High melting points
  • Good conductors of heat and electricity because the delocalised electrons in metals can carry an electrical current through the materials.
58
Q

Examples of metallic bonding (4)

A

Copper, Lead, Iron and Titanium

GCSE Chemistry P2 (5 decks)

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