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module 2: foundations in chemistry Flashcards

1
Q

which subatomic particles are found in the nucleus?

A

protons and neutrons

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2
Q

where are electrons found in an atom?

A

arranged in shells surrounding the nucleus

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3
Q

what is the relative mass of an electron?

A

1/1836

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4
Q

what are isotopes?

A

atoms with the same number of protons and electrons, but with different numbers of neutrons

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5
Q

what is the mass number of an atom?

A

the number of protons and neutrons

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6
Q

what is the atomic number of an atom?

A

the number of protons

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7
Q

what is an ion?

A

a charged atom

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8
Q

what are cations?

A

positive ions

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9
Q

what are anions?

A

negative ions

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10
Q

how are ions formed?

A

when an atom gains or loses electrons

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11
Q

what is meant by the relative atomic mass?

A

the weighted mean mass of an atom of an element, relative to 1/12th of the mean mass of an atom of carbon-12

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12
Q

what is meant by the relative isotopic mass?

A

the mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12

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13
Q

what is meant by the relative molecular mass?

A

the weighted mean mass of a molecule of a compound compared with 1/12th of the mass of an atom of carbon-12

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14
Q

state the similarities and differences between isotopes of the same element

A

similarities:
- same number of protons
- same number of electrons

differences:
- different number of neutrons
- different mass numbers

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15
Q

why do isotopes have the same chemical properties?

A

they have the same number of electrons in the outer shell

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16
Q

when does ionic bonding occur?

A

between a metal and a non-metal

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17
Q

what is the charge on a carbonate ion?

A

2-

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18
Q

what is the charge on a sulfate ion?

A

2-

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19
Q

what is the empirical fomula?

A

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

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20
Q

what is the molecular formula?

A

shows the number and types of atoms of each element present in a molecule

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21
Q

what is a standard solution?

A

a solution of known concentration

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22
Q

what is the limiting reagent?

A

the reactant that is not in excess, which will be used up first and stop the reaction

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23
Q

what is the ideal gas equation?

A

pV = nRT

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24
Q

what is an acid?

A

a species that releases H+ ions in aqueous solutions

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25
Q

what is an alkali?

A

a type of base that dissolves in water forming hydroxide ions

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26
Q

what is a strong acid?

A

an acid which fully dissociates in aqueous solutions

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27
Q

what is a weak acid?

A

an acid which partially dissociates in aqueous solutions

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28
Q

what is oxidation?

A

the loss of electrons

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29
Q

what is reduction?

A

the gaining of electrons

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30
Q

what is the effect of oxidation and reduction on the oxidation number?

A

oxidation - increase in oxidation number
reduction - decrease in oxidation number

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31
Q

what is an oxidising agent?

A

a reagent that accepts electrons from another species

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32
Q

what is disproportionation?

A

a redox reaction in which the same element is both oxidised and reduced

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33
Q

what is the formula for the maximum number of electrons a shell number n?

A

maximum number of electrons = 2n²

34
Q

what is an atomic orbital?

A

a region around the nucleus that can hold up to two electrons, with opposite spins

35
Q

what is a shell?

A

a group of atomic orbitals with the same principal quantum number

36
Q

what is the principal quantum number?

A

a number representing the relative overall energy of each orbital

37
Q

what are the four different types of orbitals?

A
  • s-orbitals
  • p-orbitals
  • d-orbitals
  • f-orbitals
38
Q

what is the shape of a s-orbital?

39
Q

what is the shape of a p-orbital?

40
Q

what is the maximum number of electrons that can fill a s-subshell?

A

2 electrons

41
Q

what is the maximum number of electrons that can fill a p-subshell?

A

6 electrons

42
Q

what is the maximum number of electrons that can fill a d-subshell?

A

10 electrons

43
Q

how many orbitals does a s-subshell have?

44
Q

how many orbitals does a p-subshell have?

45
Q

how many orbitals does a d-subshell have?

46
Q

what are the rules by which electrons are arranged in shells?

A
  • electrons are added one at a time
  • lowest available energy level is filled first
  • each energy level must be filled before the next one can fill
  • each orbital is filled singly before pairing
  • 4s is filled before 3d
47
Q

why does the 4s orbital fill before the 3d orbital?

A

the 4s orbital has lower energy than the 3d orbital, before it is filled

48
Q

why does each orbital fill singly before pairing?

A
  • electrons with the same spin repel each other (spin-pair repulsion)
  • therefore, electrons will occupy separate orbitals in the same subshell first to minimise this repulsion and have their spin in the same direction
49
Q

what is ionic bonding?

A

the electrostatic attraction between positive and negative ions

50
Q

what are sub-shells?

A

a group of orbitals of the same type within a shell

51
Q

explain why ionic compounds have high melting and boiling points

52
Q

describe the relative energies of the 2s orbital and each of the three 2p orbitals in a nitrogen atom

A
  • p-orbitals have greater energy than s-orbitals
  • p-orbitals have equal energy
53
Q

what is covalent bonding?

A

the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

54
Q

what is a covalent bond?

A

a shared pair of electrons

55
Q

where does covalent bonding occur?

A

between non-metals

56
Q

what is the shape and bond angle of a molecule with 4 bonded pairs and 0 lone pairs?

A

tetrahedral, 109.5°

57
Q

what is the shape and bond angle of a molecule with 3 bonded pairs and 1 lone pair?

A

pyramidal, 107°

58
Q

what is the shape and bond angle of a molecule with 2 bonded pairs and 2 lone pairs?

A

non-linear, 104.5°

59
Q

what is the shape and bond angle of a molecule with 3 bonding regions and 0 lone pairs?

A

trigonal planar, 120°

60
Q

what is the shape and bond angle of a molecule with 2 bonding regions and 0 lone pairs?

A

linear, 180°

61
Q

what is electronegativity?

A

a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

62
Q

what is a dative covalent bond?

A

a shared pair of electrons in which the bonded pair has been provided by one of the bonded atoms only, also called a coordinate bond

63
Q

what is the most electronegative atom?

64
Q

what is the trend of electronegativity across a period of the periodic table?

A

electronegativity increases across a period of the periodic table

65
Q

what is the trend of electronegativity up a group of the periodic table?

A

electronegativity increases going up a group of the periodic table

66
Q

what is the trend of nuclear charge and atomic radius going across the periodic table?

A
  • nuclear charge increases
  • atomic radius decreases
67
Q

what is the electronegativity difference between atoms in a covalent bond?

68
Q

what is a non-polar bond?

A

a type of covalent bond where the bonded electron pair is shared equally between the bonded atoms

69
Q

when will a bond be non-polar?

A
  • when the bonded atoms are the same, or
  • the bonded atoms have the same or similar electronegativity
70
Q

why are hydrocarbons non-polar?

A

carbon and hydrogen atoms have very similar electronegativities, so form non-polar bonds

71
Q

what is a polar bond?

A

a type of covalent bond where the bonded electron pair is shared unequally between the bonded atoms

72
Q

when will a bond be polar?

A
  • the bonded atoms are different, and
  • have different electronegativity values
73
Q

what is a dipole?

A

a separation in electrical charge so that one atom of a polar covalent bond, or one end of a polar molecule, has a small positive charge δ+ and the other has a small negative charge δ−

74
Q

what is a permanent dipole?

A

a small charge difference that does not change across a bond, with δ+ and δ− partial charges on the bonded atoms

75
Q

what are the three main categories of intermolecular forces?

A
  • induced dipole-dipole interactions (london forces)
  • permanent dipole-dipole interactions
  • hydrogen bonding
76
Q

what are intermolecular forces?

A

weak interactions between dipoles of different molecules

77
Q

in which type of molecules do you find london forces?

A

all molecules, whether polar or non-polar

78
Q

explain how induced dipoles are created

A
  • movement of electrons produces a changing dipole in a molecule
  • at any instant, an instantaneous dipole will exist, but its position is constantly changing
  • the instantaneous dipole induces a dipole on a neighbouring molecule
  • the induced dipole induces further dipoles on neighbouring molecules, which then attract one another
79
Q

how does the number of electrons affect the strength of induced dipole-dipole interactions?

A

the more electrons in each molecule:
- the larger the instantaneous and induced dipoles
- the greater the induced dipole-dipole interactions
- the stronger the attractive forces between molecules

80
Q

what is an ionic lattice?

A

a repeating pattern of oppositely charged ions

81
Q

state and explain two anomalous properties of ice caused by hydrogen bonding

A
  • ice is less dense than water because the molecules in ice are held apart in an open lattice structure by hydrogen bonds
  • ice has a relatively high melting point because hydrogen bonds are relatively strong

chemistry (9 decks)

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