Module 2 Flashcards
Basics in Chemistry (155 cards)
1
Q
what is atomic number and atomic mass?
A
number= number of protons
mass= number of protons and neutrons
2
Q
charge, mass and position of subatomic particles?
A
proton= nucleus, +1 charge, 1 mass
neutron= nucleus, 0 charge, 1 mass
electron= shells/ energy levels, -1 charge, 1/1846 mass
3
Q
state the forces present within an atom
A
electrostatic forces of attraction between positive nucleus and negative electrons in shells
4
Q
what are isotopes?
A
atoms of the same element with different number of neutrons and different atomic masses
-can be seen as O^16 and O^17, where atomic number is ommited as is same in both.
5
Q
what is relative atomic mass?
A
the weighted mean mass of an atom of an element compared to 1/12th the mass of an atom of Carbon-12
6
Q
equation for calculating relative atomic mass?
A
sum of (relative mass x percentage abundance) / total abundance
7
Q
how can you determine the percentage abundance of an isotope in a sample
A
- using a mass spectrometer:
1) place sample in mass spectrometer
2) sample is vaporised and ionised to form positive ions
3) ions or accelerated, where heavier ions move more slowly and are more difficult to deflect than lighter ions, so ions of each isotope are separated
4) ions are detected on a mass spectrum as a mass to charge ratio, and add to signal
5) greater the abundance, greater the signal - peak height/total height of all peaks x 100
8
Q
what is relative isotopic mass
A
mass of an atom of an isotope relative to 1/12th the mass of an atom of carbon-12
9
Q
why do isotopes still have the same chemical properties?
A
- the number of electrons don’t change
- it is electrons involved in chemical reactions (neutrons have no impact)
- so different isotopes of the same element react in the same way
- (may be small difference in physical properties, e.g. higher melting/boiling point and density for higher relative mass isotopes)
10
Q
what are cations
A
- positive ions
- have fewer electrons than protons
- e.g. 2+
11
Q
what are anions
A
- negative ions
- have more electrons than protons
- e.g. 2-
12
Q
rules for naming compounds
A
- if two elements bonded only (BINARY COMPOUNDS), metal goes first and ends with -ide
- if element contains oxygen, must end in -ate
13
Q
which elements are diatomic
A
- gases (O2, H2, N2)
- group 7 halogens (F2, Cl2, Br2, I2)
14
Q
common formulas you should know
A
- CO2 = carbon dioxide
- CO = carbon monoxide
- H2O = water
- CH4 = methane
- NH3= ammonia
15
Q
what are polyatomic ions
A
ions made up of more than 1 element bonded together
16
Q
why are showing ions of transition metals more difficult, and how are they displayed
A
- transition metals can form several ions with different charges
- show the ionic charge using roman numerals in brackets
- e.g. Fe (III) = Fe3+
17
Q
sulfate ion
A
SO4 2-
18
Q
nitrate ion
A
NO3 -
19
Q
hydroxide ion
A
OH-
20
Q
carbonate ion
A
CO3 2-
21
Q
ammonium ion
A
NH4 +
22
Q
phosphate ion
A
PO4 3-
23
Q
bicarbonate ion
A
HCO3 -
24
Q
silver ion
A
Ag+
25
zinc ion
Zn 2+
26
phosphor and sulfur ions
P4, S8
27
how to balance equations with state symbols
- must have same number of atoms of each element of either side
- (s) = solid
- (l) = liquid
- (g) = gas
- (aq) = aqueous, dissolved in water
28
what are ionic equations
- show the reacting ions only
29
how do you make ionic equations
1) anything aqueous can be split into ions
2) ions that don't change can be cancelled out
3) left with just reaction ions
4) REMEMBER STATE SYMBOLS
30
what is a shell
group of atomic orbitals having the same principal quantum number n
- regarded as energy levels, and energy increases as the shell number does
31
what is an orbital
region around the nucleus that can hold up to 2 electrons with opposite spins
32
what is a sub-shell
group of the same type of atomic orbitals within a shell (e.g p subshell contains 3 p orbitals)
33
characteristics of each orbital
s - orbital: sphere shape, 1 orbital (2 electrons)
p - orbital: dumb-bell shape, 3 orbitals (px, py, pz) in sub-shell (6 electrons)
d - orbital: 5 orbitals in sub-shell, (10 electrons)
f - orbital: 7 orbitals in sub-shell, (14 electrons)
34
how does increasing shell number effect the location of the orbitals
the greater the shell number, n, :
- the greater the radius of its s-orbital
- the further its p-orbital is from the nucleus
35
what are the max number of electrons in each shell 1-4
2 (s orbital only)
8 (s and p orbital only)
18 (s,p and d orbitals only)
32 (2,p,d and f orbitals)
36
how do electrons fill shells
- fill sub-shells in order of increasing energy
- 4s is filled before 3d:
- 4d subshell is at a lower energy level than 3d subshell
37
how to draw electrons in orbitals as boxes
- 1 orbital = shown as box
- electrons = shown as arrows (2 in same one are in opposite spins, as this helps to counteract the repulsion between the 2 negatively charged electrons)
- electrons only pair when no empty orbitals left in subshell (prevents repulsion between paired electrons until there is no further empty orbital available at same energy level)
38
what are all the blocks of the periodic table
- determined by the highest energy sub-shell where outermost electrons are
- s block = group 1, 2 + hydrogen and helium
- p block = group 3, 4, 5, 6, 7, 8
- d block = middle block ( 4s BEFORE 3d)
39
what order are shells filled
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6
40
what are the two anomalies to electron structure
Chromium= 4s1, 3d5 (instead of 4s2, 3d4)
- 4s subshell gives electron to 3d subshell to give a fully 1/2 filled subshell, giving extra stability
Copper = 4s1, 3d10 (instead of 4s2, 3d9)
- 4s gives one electron to 3d subshell to give a full outer shell, providing extra stability
41
how do you write a shorthand electronic configuration
- use noble gas before in brackets and add on additional subshell electrons
- can also be done with noble gases themselves, using the one before
- e.g. 11-Na : [Ne] 3s1
42
how do you write electron configuration for ions
- consider the amount of electrons and do accordingly
- for ions which lose electrons, remember to remove 4s subshell electrons first
-e.g. chloride ion = [Ne] 3s2, 3p6 = [Ar]
- zinc ion = [Ar] 3d10 (you take away the 2 on 4s first)
- remove 4s first because the energies of 4s and 3d are close together, and once filled, 3d falls below 4s energy level
43
what is the meaning of isoelectronic
- containing the same amount of electrons, like between an ion and atom
44
what is a mole
amount of any substance containing as many particles as there are carbon atoms in exactly 12g of Carbon-12
45
what is avogadro's constant
the number of atoms per mole o the Carbon-12 isotope = 6.02 x 10^23
46
why is the term "particles" important when talking about moles
1 mole of H : 1 mole of H atoms
1 mole of H2 : 1 mole of H molecules (so x2 for number of atoms)
47
what is molar mass
mass per mole of a substance (gmol^-1)
48
equation linking mass, molar mass, moles
mass (g) = moles x molar mass (gmol^-1)
49
what is avogadro's law
at room temperature and pressure, 1 mol of gas occupies 24dm^3
50
equation linking volume and moles
volume (dm^3) = 24 (mol dm^-3) x moles
24 = the molar gas volume (volume per mole of gas molecules)
51
what is concentration
quantity of a substance, in moles, in a given volume
52
what is an ideal gas
a gas with:
- no intermolecular forces
- particles of random motion
- particles with negligible size
- elastic collisions
53
what is the ideal gas equation
pV= nRt
p= pressure (Pa/ pascals)
V= volume (m^3 (x10^-6 to dm))
n= moles
R= ideal gas constant, (8.314 Jmol-2K-2)
t= temperature (K, +273 to C)
54
what is molecular formula
the number of atoms of each element in a molecule
55
what is empirical formula
the simplest whole number ratio of each element present in a compound
- useful for giant crystalline structures such as ionic structures, metals, giant covalent structures, where actual number of atoms would be VERY large
56
what is relative molecular/formula mass
compares the mass of a molecule/formula unit with the mass of Carbon-12
(add up all RAMs)
- molecular = for covalent structures
- formula = for (giant) ionic compounds, where you use the empirical formula
57
how do you calculate empirical formula
1) find the moles of each element present
2) divide by the smallest value of moles present
3) use to calculate the ratio of elements, therefore formula
58
what is hydrated salt
crystalline compound containing water molecules
59
anhydrous salts
salt containing no water molecules
60
water of crystallisation
-H2O: water molecules in hydrated salts
61
explain PAG 1: mass change in crucible to find water of crystallisation
1) weigh an empty crucible
2) add hydrated salt and reweigh
3) place crucible onto pipe-clay triangle on top of tripod
4) heat gently for 1 minute and strongly for 5 minutes
5) leave crucible to cool and reweigh
6) heat crucible again for 1 minute, leave to cool and reweigh
7) repeat until no mass change- CAREFUL TO NOT FOR TOO LONG, MAY DECOMPOSE SALT
62
what are the assumptions made in the water of crystallisation practical, and how does this effect the accuracy
1) you assume all the water has been lost:
- can use colour change in compound, but you can still only see the surface, and there may be extra water inside
- if similar colour compounds, very difficult
- SO, heat to constant mass
2) the salt may decompose further when heaters:
- could cause a colour change (copper (II) sulfate turns into black copper (III)
- very difficult to judge if no colour change occurs when decomposed
63
how are molar ratios used
- used to work out limiting reagents (reactant not in excess which will be used up and stop the reaction)
- always base calculation off limiting reagent
64
equation for percentage yield
actual yield / theoretical yield
65
what is theoretical yield, and what can it be prevented by
maximum possible amount of product:
- reaction not completed
- side reactions taken place instead
- purification causing loss of some product
66
what is the limiting reagent
the reactant that is used up first and finishes the reaction (not in excess)
- find out by finding moles of all reactants, and comparing to how many moles there SHOULD be present using equation
67
what is atom economy and explain how it is used
a measure of how well atoms have been utilised
- makes up part of argument when discussing which reactions to use, BUT, may still use low AE if
- reactants are readily available and won't require much energy to get
- the percentage yield is high
68
equation for atom economy
- Mr of desired product/ Mr of all products
69
what are the pros of high atom economy experiments
- sustainable, as preserves raw materials by making best use of natural resources
- more DP, and less WP, reducing waste and not needing to dispose of it if harmful
70
equation for percentage uncertainty
error of equipment x number of readings/ amount measures x 100
71
how can you reduce percentage uncertainty
- take fewer readings
- increase amount measured
- use equipment with lower error
72
what is an acid
-proton (H+) donor
-pH < 7
- contain hydrogen in formula, and release H+ ions in aqueous solutions
- e.g. HCl(g) + aq ===> H+(aq) + Cl-(aq)
- where + aq shows that there is an excess of water
73
what is a strong acid
- acid that fully dissociates in water, releasing all H+ ions
- e.g. HCl(aq) ====> H+(aq) + CL-(aq)
74
what is a weak acid
- acid that only partially dissociates in water, only releasing a small amount of H+ ions
- e.g. CH3COOH(aq) ===> H+(aq) + CH3OOH-(aq)
75
what is a base
proton (H+) acceptor, which neutralises acid to form a salt
76
what is an alkali
a base that dissolves in water releasing hydroxide ions (OH-) into the solution
-e.g. NaOH(s) + aq ===> Na+(aq) + OH-(aq)
77
what is a salt
a product from a reaction in which H+ ions from the acid are replaced with metal or ammonium ions
78
what are example of bases
metal oxide, metal hydroxide, metal carbonates, alkali
79
types of neutralisation reactions
acid + base ===> salt + water
acid + carbonate ===> salt + water + CO2
acid + metal ===> salt + hydrogen
acid + ammonia ===> ammonium salt
80
what is the neutralisation ionic equation
H+(aq) + OH-(aq) ===> H2O(l)
81
what are common acids
-HCl = hydrochloric acid
-H2SO4 = sulfuric acid
-HNO3 = nitric acid
-CH3COOH = ethanoic acid
82
what are common alkalis
-NaOH (sodium hydroxide)
-KOH (potassium hydroxide)
-NH3 (ammonia)
83
what do titrations tell us and what is their use
- used to measure the volume of one solution that reacts exactly with another solution
1) finding the concentration of a solution
2)identification of unknown chemicals
3) finding the purity of a substance
84
what is a standard solution
a solution of a known concentration
85
PAG 2: making a standard solution
1) weigh solid accurately (to find amount, use equations)
2) dissolve solid into a beaker with less distilled water than needed to fill the flask fully
3) move to VOLUMETRIC FLASK and rinse last traces of solution from beaker into flask also
4) fill flask carefully to graduation line so BOTTOM of meniscus touchers mark (if goes over line, start again)
5)invert flask a few times to mix solution thoroughly
86
PAG 2: carrying out a titration
1) add one solution to conical flask using pipette
2) add second solution to burette and record initial reading
3) add few drops of indicator into flask
4) run solution from burette into flask, whilst swirling flask
5) stop when indicator colour changes, so reaches end point
6) record final reading and calculate titre (volume added from burette)
7) repeat and calculate mean with concordant results
87
how do you take measurements from a burette
- to nearest 0.05cm^3
- if bottom of meniscus is in between two lines, use 0.05, and if on line use 0.00
- read from top down of burette
88
what are concordant titres
- titres within 0.10cm^3 of eachother
89
how do you calculate mean titre
- only take concordant results and calculate
90
whats the difference between oxidation number and charge number
- oxidation number: can be thought of as the number of e- involved in bonding to another element
- in oxidation, sign always goes first (+2 vs 2+)
91
what are the rules of assigning oxidation numbers
- all elements in their natural state have a number 0
- sum of all oxidation numbers = charge of particle
92
what is the order for assigning oxidation numbers
1) group 1,2,3 metals = charge number
2) fluorine is usually -1 (most electronegative)
3) hydrogen in usually +1 (except metal hydrides)
4) oxygen is usually -2 (except peroxides and flouride)
5) chlorine is usually -1
93
how do you name ions which have different charges using oxidation numbers
- use roman numerals to indicate oxidation state (number)
94
how do you name polyatomic ions using oxidation number
- always end in -ate, and bracket with roman numerals indicates oxidation number of the non-oxygen element present
- used to be use of -ite as well, but erased now with just -ate
- the bracket is usually emitted when using the common form of -ate
95
what are redox reactions
- reactions containing both reduction and oxidation
96
what is oxidation
- loss of electrons and increase in oxidation number
97
what is reduction
- gain of electrons and decrease in oxidation number
98
how can you tell what has been oxidised and what has been reduced, and what is the pattern in numbers
- work out the oxidation number of each ion present, and see the overall change between it as a product and as a reactant
- the TOTAL reduction change will always equal the TOTAL oxidation change (total changes is oxidation number balance)
99
how would you word if something has been oxidised or reduced
_____ has been oxidised/reduces because its oxidation number has increased/decreased from _____ to _______
100
what is disproportionation
when an element is both oxidised and reduced in same reaction
101
what is ionic bonding
electrostatic attraction between positive and negative ions (transfer of electrons)
- ions formed often have same electron configuration as nearest noble gas
102
what are cations
positive ions
103
what are anions
negative ions
104
how are ionic compounds structured?
giant ionic lattice: repeating pattern of alternatively charged ions, held by electrostatic forces of attraction between ions
105
explain boiling point of ionic compounds
-lots of energy required to overcome the strong electrostatic attraction between +/- ions
-dependent on charge density of an ion (greater charge density increases boiling point):
1) smaller ionic radius (across period)
2)containing ions of greater charge
106
solubility of ionic compounds
-dissolve in polar solvents, like water
-polar solvent molecules break down the lattice and surround each ion in solution
107
trend of solubility in ionic compounds, and why it’s hard to predict
depends on: attraction of solvent itself and attraction with water molecules:
1) compound with high charge ions may have attraction too large for water to break down, so not very soluble
2) BUT if higher charge, will be able to attract water molecules more too
108
explain electrical conductivity in ionic compounds
- changes dependant on state
-solid: ions held in fixed position in giant ionic lattice, no mobile charge carriers so does NOT conduct
-liquid/molten: ionic lattice breaks and ions are free to move as mobile charge carriers, so DO conduct
109
what is covalent bonding
strong electrostatic forces of attraction between nuclei of bonded atoms and shared pair of electrons
110
explain covalent bonding in terms of orbitals, and the type of attraction
- the overlap of atomic orbitals, each one containing one electron and therefore giving the shared pair
- the attraction is LOCALISED in covalent bonding, and solely between the shared pair of e- and the positive nuclei of bonded atoms - called a molecule
111
what does covalent bonding form
- simple molecular compounds
- giant molecular structures
112
what is the usual structure of the atoms once in a covalent bond
same electron structure as the nearest noble gas
-USUALLY, but cannot always trust the model, e.g. with BF3
113
how do you draw dot and cross for covalent bonding
- electrons in middle of shared shells
- lone pairs shown as dots in displayed formula
- atoms beyond fluorine (which have outer shell n=3 (can hold 18 electrons due to the d-subshell allowing expansion of the octect)) can hold up to 18 electrons in outer shell, all in different arrangements
114
how do you show the displayed formula of covalent bonds
- use lines to show bonds and show the relative positioning of all the atoms
- use : for a lone pair of electrons (paired electrons that are not shared)
115
what is multiple covalent bonds
electrostatic attraction between 2/3 shared pairs of electrons and the nuclei of bonding atoms
-e.g. O=C=O and O=O
-e.g. N≡N and H-C≡N
116
what is a dative covalent bond
(may be called coordinate too)
-covalent bond in which the shared pair of electrons has been provided by one of the bonded atoms only
- the shared pair of electrons comes from an original lone pair which has been donated
- if an ion, put brackets and charge outside like ionic diagram, e.g. NH4+ ion
- if displayed, put arrow showing where electrons are going to
117
what is average bond enthalpy
- measure of strength of covalent bond
118
what is electron pair repulsion theory
1) bonding pair regions repel each other as far as possible
2) lone pairs of electrons repel more than bonding pairs, reducing the angle by 2.5 degrees each time ((as slightly closer to nuclei))
- means that repulsion is minimised, so held in definite shape
- bonding pair REGION = accounts for multiple bonds too, as they are treated same as a single pair
119
how do you show the 3D shape of a molecule on paper
1) solid line= bond in plane of paper
2) solid wedge= bond coming out of plane of paper
3) dotted wedge= bond going into plane of paper
120
2 bonding pairs - shape of molecule
-linear
-180 degrees bond angle
121
3 bonding pairs- shape of molecule
-trigonal planar
-120 degree bond angle
122
4 bonding pair - shape of molecule
-tetrahedral
-109.5 degrees bond angle
123
5 bonding pairs- shape of molecule
-octahedral
-90 degree bond angle
124
3 bonding pairs
1 lone pair
shape of molecule?
-pyramidal
-107 degree bond angle
-lone pair = top line = 2 dots
125
2 bonding pairs
2 lone pairs
shape of molecule
- non-linear
- 104.5 degree bond angle
- top and left lone pairs = 4 dots
126
what is electronegativity
the relative ability of an atom to attract the bonded pair of electrons in a covalent bond
127
what scale determines electronegativity
Pauling electronegativity values
128
what does electronegativity depend on, and where is it highest?
1) nuclear charge (increases across periodic table)
2) atomic radius (decreases across periodic table)
- highest at top left of periodic table
129
most electronegative and least electronegative ions
- Fluorine = 4.0 (most)
- nitrogen, oxygen, chlorine
-group 1 metals (least)
130
how do you calculate with electronegativities the type of bond present
- covalent = EN difference of 0
- polar covalent = EN difference of 0-1.8
- ionic = EN difference of over 1.8
131
what are non polar bonds
- where the bonded pair of electrons is shared equally between bonded atoms
-e.g. diatomic molecules or hydrocarbons
132
what are polar bonds
where the bonded pair of electrons are unevenly distributed between the atoms
133
what is a dipole
- the separation of opposite charges in a polar bond, where one side is slightly negative and one side is slightly positive
134
how do you show a dipole
- arrow with line facing least to most electronegative (from δ+ to δ-)
135
explain why a molecule may have polar permanent dipoles, but not be polar overall
- if the dipoles cancel out
- by acting in opposite directions
- so overall, no dipole
- e.g. CO2
- RATHER THAN H2O, which would be polar as the dipoles reinforce each other (don't oppose, contributing to a larger overall dipole)
136
what are intermolecular forces
weak interactions between dipoles of different molecules
137
difference in covalent bonds and intermolecular forces
covalent - identity and chemical reactions of a molecule
intermolecular - physical properties such as melting and boiling points
138
what are induced dipole dipole interactions
- london forces
- intermolecular forces which occur between ANY and ALL molecules, whether polar or not, acting in between induced dipoles
139
how do london dispersion forces arise
- electrons are constantly moving
- causes an uneven distribution of charge
- resulting in an instantaneous dipole
- dipole induces another dipole in neighbouring particle and so on
- dipoles attract each other
140
what is the nature of instantaneous dipoles
- it is constantly changing
- at any instant, will exist, but position is constantly shifting
141
how can you increase the strength of london dispersion forces
- increase the number of electrons
- larger instantaneous and induced dipoles
- stronger london dispersion forces (so stronger attractive forces between molecules)
- more energy needed to overcome the IMF
142
what are permanent dipole-dipole interactions
- difference in electronegativity of bonded atoms causes a small charge difference across the bond, resulting in a permanent dipole
- permanent dipole-dipole interactions act between the permanent dipoles in polar molecules (dotted line in between δ+ and δ-)
- REMEMBER, they still have london forces too
143
3 intermolecular forces in order of strength, most to least
- hydrogen bonding
- permanent dipole dipole interactions
- induced dipole dipole interactions
144
what are simple molecular substances
- made of simple molecules
- small units containing a definitive number of atoms with a definitive molecular formula
145
explain forces present in simple molecular substances
- molecules held inlattice with WEAK intermolecular forces
- atoms in each molecule bonded with STRONG covalent bonds
146
melting point of simple molecular substances
- low melting or boiling point
- as only weak intermolecular forces holding them together in lattice break
- not the strong covalent bonds which stay
147
electrical conductivity of simple molecular substances
- do not conduct electricity
- as no mobile charge carriers present to carry charge
148
solubility of polar + polar (vice versa) molecular substances
1) new interactions form between the solute and solvent
-for IONIC, between the δ+ and δ- ends of molecule and the positive and negative ions present, surrounding them
- for POLAR, the molecule and solvent can attract each other
-for NON-POLAR, e.g. simple molecular substances, IMF form between the molecules and solvent
2) the existing IMF weaken and break if new interactions strong enough
3) compound dissolves
149
solubility of polar + non polar substances
- new interactions which form between solute and solvent are weak and insignificant
- insufficient to break existing IMF (e.g. within the polar solvent)
- solute does not dissolve
150
what is hydrogen bonding
- type of permanent dipole-dipole
- between lone pair of electrons on electronegative atom
- and hydrogen atom attached to electronegative atom
151
what are the atoms hydrogen bonding can occur with
N
O
F
(or OH, NH2 groups)
152
how do you draw hydrogen bonding
- with dashed lines (straight, as shape around H is linear)
- between lone pair of electrons on the EN atom in one molecule
- and hydrogen on another molecule
- always include slightly positive and negative signs
153
how can you increase the hydrogen bonding
- use atom with more electron lone pairs
- so more hydrogen bonds can be made
154
water's density
- ice (solid) is less solid than water and floats
- hydrogen bonding forms 4 bonds (as 2 lone pairs on O2 and 2 Hs attached)
- holding water molecules apart in an open lattice structure
- so is spread out as solid as molecules are further apart (than as a liquid)
155
water's melting and boiling point
- has strong hydrogen bonding as well as given London forces
- strongest IMF, so higher amounts of energy needed to break it down
- so high melting and boiling points
