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Chemistry A LEVEL > Both > Flashcards

Both Flashcards

1
Q

Ammonium

A

NH4 +

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2
Q

Silver

A

Ag +

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3
Q

Zinc

A

Zn 2+

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4
Q

Scandium

A

Sc 3+

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5
Q

Hydroxide

A

OH -

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6
Q

Nitrate

A

NO3 -

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7
Q

Manganate

A

MnO4 -

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8
Q

Hydrogencarbonate

A

HCO3 -

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9
Q

Hydrogensulfate

A

HSO4-

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10
Q

carbonate

A

CO3 2-

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11
Q

Sulfate VI

A

SO4 2-

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12
Q

Sulfate IV

A

SO3 2-

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13
Q

Dichromate

A

Cr2O7 2-

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14
Q

Phosphate

A

PO4 3-

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15
Q

An acid is

A

A proton doner

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16
Q

Strong acid

A

Completely dissociates into its ions in aqueous solution

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17
Q

Weak acid

A

Partially dissociates into it ions in aqueous solution

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18
Q

Base

A

Proton acceptor

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19
Q

Alkali is

A

A base that dissolves in water and releases 0H-ions in aqueous solution

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20
Q

Salt

A

Produced when the H+ ion of an acid is replaced by a metal ion or NH4+

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21
Q

Acid + ammonia

A

Ammonium salt

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22
Q

Acid + metal oxide

A

Salt + water

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23
Q

Acid + metal hydroxide

A

Salt + water

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24
Q

Acid + metal

A

Salt + hydrogen

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25
Q

Acid + carbonate

A

Salt + water + carbon dioxide

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26
Q

Atom consists of

A

Protons and neutrons in the nucleus

Electron which orbit the nucleus

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27
Q

Proton, neutron, electron

Charge and mass

A

Proton +1 mass 1
Neutron 0 mass 1
Electron -1 mass negligible

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28
Q

Atomic number

A

Equal to the number of protons in the nucleus of an atom

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29
Q

Electrons are held in place by t

A

Attractive forces from the nucleus

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30
Q

Atoms are neutral because

A

Same number of protons and electrons

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31
Q

Isotopes

A

Atoms of the same element with different numbers of neutrons

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32
Q

Mass number

A

Equal to the number of protons and neutrons in the nucleus

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33
Q

Ions are formed by

Ions are charged because

A

Losing or gaining electrons

Because they have different numbers of Electrons from protons

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34
Q

Relative isotopic mass

A

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

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35
Q

Weighted mean mass takes into account

A

The percentage abundance of each isotope

The relative mass of each isotope

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36
Q

Relative atomic mass

A

The weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon 12

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37
Q

Mole

A

The amount of substance containing as many particles as there are carbon atoms in exactly 12 g of carbon 12

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38
Q

Molar mass

A

The mass in grams per mole of a substance

Units gmol^-1

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39
Q

Mile equation solid

A

Mole=mass/mr

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40
Q

Avagadros constant

A

6.02x10^23

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41
Q

Mol équation gas

A

Mole= volume/molar gas volume

24(dm3)
24000(cm3)

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42
Q

Mol équation solution

A

Dm3= c x v

Cm3= c x v/1000

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43
Q

Mol dm3 —> gdm3

A

X mr

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44
Q

Titration reading

4

A

To 2dp
End in 0 or 5
Two concordant results within 0.1
Don’t round mean titres

45
Q

Molecular formula

A

The actual number of atoms of each element in a molecule

46
Q

Empirical formula

A

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

47
Q

Ideal gas equation

A

PV=nRT

48
Q

Atom economy

A

Molecular mass of desired product/ molecular mass of all products

X100

49
Q

High atom economy

A

Less waste product

50
Q

Yield

A

The mass of a product obtained from a reaction

51
Q

% yield

A

Actual mass or mol of product/ theoretical mass or mol of product

X100

52
Q

Oxidation number

A

A measure of the number of electrons that an atom uses to bond with atoms of another element

53
Q

Oxyanions

A

Negative ions containing oxygen and one or more element

-ate

54
Q

Redox

A

A reaction where both oxidation and reduction takes place

55
Q

Oxidation

A

Loss of electrons

Increase in oxidation number

56
Q

Reduction

A

Gain of electrons

Decrease in oxidation number

57
Q

OIL RIG

A

Oxidation is loss of electrons

Reduction is gain of electrons

58
Q

Oxidising agent

A

A reagent that oxidises, takes electrons from another species

59
Q

Reducing agent

A

Reagent that reduces, adds electrons to another species

60
Q

First ionisation energy

A

The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions

61
Q

3 factors affecting ionisation energy

A

Nuclear charge

Distance of the outermost electron from the nucleus

Electron shielding

62
Q

Nuclear charge

Effect on ionisation energy

A

The more protons the greater the nuclear charge

The greater the nuclear charge of stronger the nuclear attraction therefore a greater nuclear charge means more energy would be needed to overcome the attraction between the nucleus and the outer most electron

63
Q

Distance of the outermost electron from the nucleus (atomic radius)

Effect on ionisation energy

A

As the distance between the nucleus and the outer electron increases the attraction between them decreases the weaker the nuclear attraction the less energy needed to remove the outer electron

64
Q

Electron shielding

Effect on ionisation energy

A

Electron shielding is the repulsion between electrons from the inner shell and outer shell

This shielding effect reduces the overall nuclear attraction

The more inner shells there are the greater the shielding effect and the weaker the nuclear attraction meaning low ionisation energy

65
Q

Why boron ( group 3) has a lower first ionisation energy than beryllium (group 2)

A

The 2p sub shell in B has higher energy than the 2s subshell
The 2p1 electron in B needs less energy to be removed, giving Boron a lower first ionisations energy than beryllium

66
Q

Atomic orbital

A

A region within an atom, around the nucleus, that can hold up to two electrons, with opposite spins

67
Q

S orbital shape

A

Spherical

68
Q

P orbital shape

A

Dumb-bell

69
Q

Order of filling orbitals

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

70
Q

A transition element

A

A d block element that forms an ion with an incomplete d sub shell

71
Q

Transition element

A

Atoms lose their 4s electrons first when forming positive ions

72
Q

Cr electrons configuration

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5

73
Q

Cu electron configuration

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10

74
Q

Ionic bonding

A

The electrostatic attraction between oppositely charged ions

75
Q

Covalent bonding

A

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

76
Q

Dative covalent bonding

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms, where only one of he atoms supplies both of the electrons shared.

77
Q

Average bond enthalpy can be used as

A

A measurement of covalent bond strength

78
Q

Shapes of molecules (3)

A

Electron pairs repel each other to get as far apart as possible

Lone pairs repel more strongly than bonded pairs

Shape is determined by the number and type of electron pairs around the central atom

79
Q

2 bonding pairs

A

Linear

180

80
Q

3 bonding pairs

A

Trigonal planar

120

81
Q

4 bonding pairs

A

Tetrahedral

109.5

82
Q

6 bonding pairs

A

Octahedral

90

83
Q

3 bonding pairs

1 lone pair

A

Pyramidal

107

84
Q

2 bonding pairs

2 lone pairs

A

Non linear

104.5

85
Q

2 bonding pairs 1 lone pair

A

Non linear

117.5

86
Q

Electronegativity

A

The ability of an atom to attract the bonding electrons towards itself in a covalent bond

87
Q

Symmetrical molecules

A

Non polar because the dipoles cancel out

88
Q

Unsymmetrical molecules

A

Contain one or more polar bonds, polar because the dipoles do not cancel out

89
Q

Intermolecular forces

3+3

A

Between ,molecules

Weaker than ionic or covalent bonds

Only found in covalent structures

Hydrogen bonds

Permanent dipole-dipole interactions

Induced dipole-dipole interactions

90
Q

Induced dipole-dipole interactions

A

Very weak

Between all molecules, polar and non-polar

An uneven distribution of electrons in a molecule causes a temporary(instantaneous) dipole
The temporary dipole causes an induced dipole in a neighbouring molecule
The delta+ of a dipole in one molecule attracts the delta- of a dipole in a neighbouring molecule to produce a London force

91
Q

Permanent dipole dipole interactions

A

Weak

Between polar molecules

Eg; the delta+ H of one HCl molecule attracts the delta- Cl of a neighbouring HCl molecule to produce a permanent dipole-dipole force of attraction between the molecules

92
Q

Which are stronger permanent or induced dipole-dipole

A

Permanent

93
Q

Hydrogen bonds

A

Strong dipole-dipole attraction between molecules containing O-H, N-H or F-H bonds

A hydrogen bond exists betweeen
Delta+ H atom in one molecule
A lone pair on a highly electronegative atom (delta- O or N or F) on another molecule

Strongest type of inter molecule force

94
Q

Anomalous properties of water

2

A

Ice is less dense than water;
In soli H2O molecules are held further apart by the hydrogen bonds , gives the ice an open lattice structure

Water has relatively high melting and boiling points;
Water has a higher than expected melting and boiling point
Hydrogen bonds are relatively strong and therefore stronger than other intermolecular forces, so more energy is needed to break the hydrogen bonds

95
Q

Metallic bonding

A

The strong electrostatic attraction of a lattice of cations to a ‘sea’ of delocalised electrons

96
Q

Giant

A

Strong forces

97
Q

2 physical properties of metals h

A

High melting and boiling point= strong electrostatic attraction between the lattice of cations and delocalised electrons which require lots of energy to break

Good electrical conductors= in both solid and liquid states as delocalised electrons can move and carry charge

98
Q

Ionic compound

A

The solid structure of a giant ionic lattice results from the electrostatic attraction between oppositely charged ions

99
Q

4 properties of ionic compounds

A

High melting an boiling point= strong electrostatic attraction between oppositely charged ions splits of energy needed to break ionic bonds, all solid at room temperature

Non electrical conductor when solid= ions are fixed in their position in the lattice and cannot move

Good electrical conductors when molten or in aqueous solution= ions are able to move and carry charge

Soluble in water-= polar water molecules are attracted to the positive and negative ions and the giant ionic lattice breaks down. Water molecules completely surround the ions

100
Q

Simple means

A

Weak forces

101
Q

Simple covalent compounds

A

Simple covalent lattices are solid structures made up of small, simple molecules with weak forces of attractions betweeen the molecules (INTERmolecular forces)

The atoms within the molecules are bonding together by strong covalent bonds (INTRAmolecular covalent bonds)

102
Q

3 physical properties of simple molecular lattices

A

Low melting and boiling point- little energy is needed to overcome the weak forces of attraction between the molecules(weak intermolecular forces)

Non conductors of electricity- no mobile ions or electrons (charge carriers) all electrons are fixed in covalent bonds and cant move

Soluble in non polar solvents- the simple molecules can form induced dipole-dipole forces with the non polar solvent molecules

103
Q

Giant covalent lattice structures

A

Solids which are networks of atoms bonded together by strong covalent bonds

104
Q

Diamond

3

A

C atom covalently bonds to 4 Others in tetrahedral shape

Very high melting point= all strong covalent bonds between C atoms so lots of heat energy needed to break the covalent bonds

Doesnt conduct electricity= no ions or mobile electrons, all electrons fixed in covalent bonds

105
Q

Graphite

3

A

Carbon, three of the four outer electrons are used n covalent binding form planar hexagonal layers

One electron from each atom is delocalised and can move between the layers of hexagons, allowing graphite to conduct electricity

High melting point as there are many strong covalent bonds between C atoms so lots of heat energy needed to break the covalent bonds

106
Q

Graphene

2

A

Form of carbon, consisting of a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds

Can be rolled into nanotubules

107
Q

Graphene properties

A

One of the thinner and strongest materials as only one atom thick

Extremely light

Transparent

V high melting point

Excellent conductor of heat and electricity

One of the toughest almost flexible materials

108
Q

Silicon

A

Same structure as diamond, four electrons to form covalent bonds to other silicon atoms

Tetrahedral structure

Silicon dioxide also shows a tetrahedral structure where silicon and oxygen atoms are. Bonded covalently

Chemistry A LEVEL (10 decks)

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