Both Flashcards
1
Q
Ammonium
A
NH4 +
2
Q
Silver
A
Ag +
3
Q
Zinc
A
Zn 2+
4
Q
Scandium
A
Sc 3+
5
Q
Hydroxide
A
OH -
6
Q
Nitrate
A
NO3 -
7
Q
Manganate
A
MnO4 -
8
Q
Hydrogencarbonate
A
HCO3 -
9
Q
Hydrogensulfate
A
HSO4-
10
Q
carbonate
A
CO3 2-
11
Q
Sulfate VI
A
SO4 2-
12
Q
Sulfate IV
A
SO3 2-
13
Q
Dichromate
A
Cr2O7 2-
14
Q
Phosphate
A
PO4 3-
15
Q
An acid is
A
A proton doner
16
Q
Strong acid
A
Completely dissociates into its ions in aqueous solution
17
Q
Weak acid
A
Partially dissociates into it ions in aqueous solution
18
Q
Base
A
Proton acceptor
19
Q
Alkali is
A
A base that dissolves in water and releases 0H-ions in aqueous solution
20
Q
Salt
A
Produced when the H+ ion of an acid is replaced by a metal ion or NH4+
21
Q
Acid + ammonia
A
Ammonium salt
22
Q
Acid + metal oxide
A
Salt + water
23
Q
Acid + metal hydroxide
A
Salt + water
24
Q
Acid + metal
A
Salt + hydrogen
25
Acid + carbonate
Salt + water + carbon dioxide
26
Atom consists of
Protons and neutrons in the nucleus
| Electron which orbit the nucleus
27
Proton, neutron, electron
Charge and mass
Proton +1 mass 1
Neutron 0 mass 1
Electron -1 mass negligible
28
Atomic number
Equal to the number of protons in the nucleus of an atom
29
Electrons are held in place by t
Attractive forces from the nucleus
30
Atoms are neutral because
Same number of protons and electrons
31
Isotopes
Atoms of the same element with different numbers of neutrons
32
Mass number
Equal to the number of protons and neutrons in the nucleus
33
Ions are formed by
Ions are charged because
Losing or gaining electrons
Because they have different numbers of Electrons from protons
34
Relative isotopic mass
The mass of an atom of an isotope compared with 1/12 of the mass of an atom of carbon 12
35
Weighted mean mass takes into account
The percentage abundance of each isotope
The relative mass of each isotope
36
Relative atomic mass
The weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon 12
37
Mole
The amount of substance containing as many particles as there are carbon atoms in exactly 12 g of carbon 12
38
Molar mass
The mass in grams per mole of a substance
| Units gmol^-1
39
Mile equation solid
Mole=mass/mr
40
Avagadros constant
6.02x10^23
41
Mol équation gas
Mole= volume/molar gas volume
24(dm3)
24000(cm3)
42
Mol équation solution
Dm3= c x v
Cm3= c x v/1000
43
Mol dm3 —> gdm3
X mr
44
Titration reading
| 4
To 2dp
End in 0 or 5
Two concordant results within 0.1
Don’t round mean titres
45
Molecular formula
The actual number of atoms of each element in a molecule
46
Empirical formula
The simplest whole number ratio of atoms of each element present in a compound
47
Ideal gas equation
PV=nRT
48
Atom economy
Molecular mass of desired product/ molecular mass of all products
X100
49
High atom economy
Less waste product
50
Yield
The mass of a product obtained from a reaction
51
% yield
Actual mass or mol of product/ theoretical mass or mol of product
X100
52
Oxidation number
A measure of the number of electrons that an atom uses to bond with atoms of another element
53
Oxyanions
Negative ions containing oxygen and one or more element
-ate
54
Redox
A reaction where both oxidation and reduction takes place
55
Oxidation
Loss of electrons
Increase in oxidation number
56
Reduction
Gain of electrons
Decrease in oxidation number
57
OIL RIG
Oxidation is loss of electrons
Reduction is gain of electrons
58
Oxidising agent
A reagent that oxidises, takes electrons from another species
59
Reducing agent
Reagent that reduces, adds electrons to another species
60
First ionisation energy
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
3 factors affecting ionisation energy
Nuclear charge
Distance of the outermost electron from the nucleus
Electron shielding
62
Nuclear charge
Effect on ionisation energy
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
Distance of the outermost electron from the nucleus (atomic radius)
Effect on ionisation energy
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
Electron shielding
Effect on ionisation energy
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
Why boron ( group 3) has a lower first ionisation energy than beryllium (group 2)
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
Atomic orbital
A region within an atom, around the nucleus, that can hold up to two electrons, with opposite spins
67
S orbital shape
Spherical
68
P orbital shape
Dumb-bell
69
Order of filling orbitals
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
70
A transition element
A d block element that forms an ion with an incomplete d sub shell
71
Transition element
Atoms lose their 4s electrons first when forming positive ions
72
Cr electrons configuration
1s2 2s2 2p6 3s2 3p6 4s1 3d5
73
Cu electron configuration
1s2 2s2 2p6 3s2 3p6 4s1 3d10
74
Ionic bonding
The electrostatic attraction between oppositely charged ions
75
Covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nucleus of the bonded atoms
76
Dative covalent bonding
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
Average bond enthalpy can be used as
A measurement of covalent bond strength
78
Shapes of molecules (3)
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
2 bonding pairs
Linear
180
80
3 bonding pairs
Trigonal planar
120
81
4 bonding pairs
Tetrahedral
109.5
82
6 bonding pairs
Octahedral
90
83
3 bonding pairs
| 1 lone pair
Pyramidal
107
84
2 bonding pairs
| 2 lone pairs
Non linear
104.5
85
2 bonding pairs 1 lone pair
Non linear
117.5
86
Electronegativity
The ability of an atom to attract the bonding electrons towards itself in a covalent bond
87
Symmetrical molecules
Non polar because the dipoles cancel out
88
Unsymmetrical molecules
Contain one or more polar bonds, polar because the dipoles do not cancel out
89
Intermolecular forces
3+3
Between ,molecules
Weaker than ionic or covalent bonds
Only found in covalent structures
Hydrogen bonds
Permanent dipole-dipole interactions
Induced dipole-dipole interactions
90
Induced dipole-dipole interactions
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
Permanent dipole dipole interactions
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
Which are stronger permanent or induced dipole-dipole
Permanent
93
Hydrogen bonds
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
Anomalous properties of water
2
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
Metallic bonding
The strong electrostatic attraction of a lattice of cations to a ‘sea’ of delocalised electrons
96
Giant
Strong forces
97
2 physical properties of metals h
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
Ionic compound
The solid structure of a giant ionic lattice results from the electrostatic attraction between oppositely charged ions
99
4 properties of ionic compounds
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
Simple means
Weak forces
101
Simple covalent compounds
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
3 physical properties of simple molecular lattices
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
Giant covalent lattice structures
Solids which are networks of atoms bonded together by strong covalent bonds
104
Diamond
3
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
Graphite
| 3
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
Graphene
| 2
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
Graphene properties
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
Silicon
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
