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Chemistry 🧬 > Key Concepts In Chemistry > Flashcards

Key Concepts In Chemistry Flashcards

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1
Q

What are ionic bonds?

A
  • Strong electrostatic forces of attraction
  • Between oppositely charged ions
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2
Q

Do ionic compounds usually have high or low melting points and why?

A
  • High
  • Many strong ionic bonds so large amounts of energy must be transferred to lattice structure to break these bonds
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3
Q

Are ionic compounds usually soluble in water?

A

Yes they dissolve to form aqueous solutions

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4
Q

Do you giant molecular substances usually have high or low melting points?

A

High, a lot of energy must be transferred to break string covalent bonds

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5
Q

In a non-enclosed system, what will happen to the mass of a reactive metal it is heated in air?

A

Increase - because oxygen atoms in the air will combine with metal atoms to form a metal oxide

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6
Q

In a non-enclosed system what will happen to the mass of a reactive non-metal or a fuel if it is heated in air?

A

Decrease - because products in the gas state will escape from the container

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7
Q

In a non-enclosed system what will happen to the mass of a metal carbonate if it is heated?

A

Decrease - because carbon dioxide gas is produced and this escapes from the container

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8
Q

What is Avogadro’s constant?

A

6.02 x 10^23

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9
Q

Which elements are written with a subscript number 2?

A
  • Iodine
  • Bromine
  • Chlorine
  • Fluorine
  • Oxygen
  • Nitrogen
  • Hydrogen
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10
Q

What did Dalton’s ideas state about the atom?

A
  • Solid atom model
  • Atoms cannot be broken down into anything simpler
  • The atoms of an element are identical to each other
  • The atoms of different elements are different from each other
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11
Q

How did the model of the atom change?

A
  • Thompson: discovered the electron
  • Thompson: plum pudding model, spheres of positive charge with negative electrons dotted inside
  • Rutherford: solar system model, positive nucleus surrounded by negative electrons in orbits
  • Bohr: electron shell model, electrons occupy shells or energy levels around the nucleus
  • Rutherford: discovered the proton
  • Chadwick: discovered the neutron
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12
Q

What is the relative mass of an electron?

A

1 / 1836

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13
Q

How do you calculate the relative atomic mass of an element?

A

(Mass number of first isotope x abundance)
+
(mass number of second isotope x abundance)
/ 100

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14
Q

How did Mendeleev arrange the elements in his periodic table?

A
  • In order of increasing relative atomic masses
  • He noted that the chemical properties of the elements and showed a trend
  • He swapped the places of some elements so that elements with similar properties lined up
  • He left gaps we thought there were other elements and predicted their chemical and physical properties (they were very accurate)
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15
Q

Why did Mendeleev put tellurium in group 6 and iodine in group 7?

A
  • He did not know about isotopes
  • Iodine has one naturally occurring isotope, iodine–127
  • The most abundant tellurium isotopes are tellurium-128 and tellurium-130
  • This means the relative atomic mass of tellurium is greater than iodine
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16
Q

How does the electronic configuration of an element determine which period it is in?

A

The number of occupied shells

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17
Q

How does the electronic configuration of an element determine what group it is in?

A

The number of electrons in the outer shell (except for group 0 which have full outer shells)

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18
Q

Are cations positive or negative?

A

Positive

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19
Q

Are anions positive or negative?

A

Negative

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20
Q

Do ionic compounds conduct electricity and why?

A
  • Yes, but only when molten or dissolved in water
  • These processes make the ions free to move
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21
Q

How do you name positively charged ions formed from hydrogen or metal atoms?

A

They take the name of the element

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22
Q

How do you name negatively charged ions formed from single non-metal atoms?

A

They take the name of the element the end in –ide

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23
Q

How do you name negatively charged ions in compounds containing three or more elements, one of which is oxygen?

A

They end in -ate

OH- does not follow this rule (hydroxide)

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24
Q

Do simple molecular substances usually have low or high boiling points and why?

A
  • Low
  • Although they have strong covalent bonds between atoms
  • There are weak intermolecular forces between molecules which can easily be overcome
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25
Q

Do simple molecular substances conduct electricity and why?

A
  • No
  • They have no overall charge, or any charged particles that can separate
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26
Q

What is the typical size of an atom?

A

1 x 10^-10

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27
Q

What is the typical size of a simple molecule?

A

1 x 10^-9

28
Q

Are simple molecular substances soluble and why?

A
  • No many are not
  • Because the intermolecular forces between water and the substances are weaker than those between water molecules and the molecules of the substance itself
  • They can if they form strong enough attractions with water molecules
29
Q

Do giant molecular substances usually have low or high boiling point and why?

A
  • High
  • A lot of energy must be transferred to break the many strong covalent bonds
30
Q

Do giant covalent substances conduct electricity and why?

A
  • No
  • They have no overall charge
  • Graphite is an exception
31
Q

What is the structure of diamond?

A
  • Each carbon atom is bonded to 4 others by covalent bonds
  • The atoms form a regular tetrahedral network structure
  • There are no free electrons
32
Q

What is the structure of graphite?

A
  • Each carbon atom is bonded to 3 others by covalent bonds
  • The atoms form a hexagonal layered network structure
  • The layers have weak intermolecular forces between them so they can slide
  • The unbonded electrons are delocalised so are free to move
33
Q

Why is graphite used to make electrodes and lubricants?

A
  • Electrodes because it has delocalised electrons to conduct electricity
  • Lubricant is because the layers can slide over each other so it is slippery
34
Q

Why is diamond used to make cutting tools?

A

It has a rigid lattice structure and a strong covalent bonds so it is very hard

35
Q

What useful properties does graphene have?

A
  • Conduct electricity because it has delocalised electrons that can move through the structure
  • Strong and flexible because it contains many strong covalent bonds
  • Almost transparent because its layers are just one atom thick
36
Q

What are fullerenes?

A
  • They resemble a sheet of Graphene rolled to form buckyballs or nanotubes
37
Q

What properties do materials made from buckyballs have?

A
  • Conduct electricity because they have delocalised electrons
  • Are soft when solid because they have weak intermolecular forces
38
Q

What properties do nanotubes have?

A
  • Conduct electricity because they have delocalised electrons
  • Have high tensile strength because the structure has many strong covalent bonds
39
Q

How many carbon atoms does buckminsterfullerene have?

A

60

40
Q

What is the structure of polyethene?

A

Consists of large molecules containing chains of carbon atoms which are joined together and to hydrogen atoms by covalent bonds

41
Q

What are the typical properties of metals?

A
  • Shiny
  • Good electrical conductors
  • High melting points (not mercury)
  • High density
  • Good heat conductors
  • Malleable 
42
Q

What are the typical properties of nonmetals?

A
  • Dull
  • Poor electrical conductors (not graphite)
  • Low density
  • Low melting points (not graphite)
  • Poor heat conductors
  • Brittle
43
Q

What is the structure of metallic bonding?

A
  • Giant lattice of positively charged metal ions
  • ‘Sea’ of delocalised electrons
  • Strong electrostatic forces of attraction between them
44
Q

Why are metals malleable?

A
  • Layers of positive ions can slide over each other
  • This allows the metal to change shape without shattering
45
Q

Why do metals conduct electricity?

A

They have delocalised electrons which can move through the structure of the metal

46
Q

What are the limitations of the space filling models?

A
  • Does not show how covalent bonds are formed
  • Does not show which elements are present (unless given at colour key)
  • You may not be able to see all the atoms if the model is complex
47
Q

What are the limitations of ball and stick models?

A
  • Does not show how covalent bonds are formed (the bonding and non-bonding electrons)
  • Bonds are shown as sticks rather than forces
48
Q

What are the limitations of structure models for simple molecular substances?

A
  • Does not show the relative sizes of the atoms
  • Does not show how covalent bonds are formed (the bonding and non-bonding electrons)
  • Does not show the relative sizes of the atoms
49
Q

What are the limitations of dot and cross diagrams for simple molecular substances?

A
  • Does not show the intermolecular forces
  • Does not show the relative sizes of the atoms
50
Q

What are the limitations of chemical formula for simple molecular substances?

A
  • Does not show the type of bonding
  • Does not show the shape of the molecule
51
Q

What are the limitations of empirical formula?

A
  • Does not show how the atoms are arranged
  • Does not show the actual number of atoms usually
52
Q

What are the limitations of 2D diagrams for ionic bonding?

A
  • Does not show how ions were formed
  • Does not show more than one layer
53
Q

What are the limitations of 3D diagrams for ionic bonding?

A
  • Does not show the charges
  • Does not show that there are no spaces between ions
54
Q

What are the limitations of dot and cross diagrams for ionic bonding?

A
  • Does not show the lattice structure
  • Does not show the ionic bonds
55
Q

What is the empirical formula of a substance?

A

The simplest whole number ratio of the atoms of each element in a compound

56
Q

How do you calculate an empirical formula?

A
  • Write down the symbol of each element
  • For each element calculate mass divided by relative atomic mass
  • Divide each answer by the smallest answer
  • Multiply (if needed) to get whole numbers
57
Q

How do you convert empirical formula to molecular formula?

A
  • Calculate the relative formula mass of the empirical formula
  • Divide the relative formula mass given by the relative formula mass of the empirical formula
  • Multiply each number in the empirical formula by the answer
58
Q

Experiment: how do you determine the empirical formula mass of magnesium oxide?

A
  • Measure and record the mass of an empty crucible with its lid
  • Put a length of magnesium ribbon into it
  • Measure and record the total mass of the crucible, it’s a lid and contents
  • Place to crucible on a tripod with a pipeclay triangle, strongly heat for several minutes using a Bunsen burner
  • When the magnesium has stopped glowing, turn off the Bunsen burner and allow the Crucible to cool
  • Measure the total mass again
59
Q

What are some examples of reactions in closed systems?

A
  • Reactions in a sealed container
  • Precipitation reactions, where two soluble reactants from the insoluble precipitate
60
Q

What is the stoichiometry of a reaction?

A

The ratio of the amount of each substance in a balanced chemical equation.

61
Q

How do you calculate concentration?

A

Mass of solute (g) / Volume of solution (dm³)

62
Q

What is a solution?

A

A mixture of a solute in a solvent

63
Q

What is a solute?

A

The dissolved substance in a solution

64
Q

What is a solvent?

A

The liquid which the solute dissolves in

65
Q

How do you convert cm³ into dm³ ?

A

Divide by 1000

66
Q

How do you calculate the number of moles in a substance?

A

Mass (g) / Relative atomic mass

Chemistry 🧬 (8 decks)

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