Foundations in Chemistry Flashcards
(113 cards)
1
Q
what is atomic number
A
number of protons and neutrons in the nucleus of an atom
2
Q
what is mass number
A
total number of neutrons and protons in the nucleus of an atom
3
Q
what is relative abundance
A
the amount of one substance compared with another
4
Q
what is relative atomic mass
A
weighted mean mass of an atom compared with 1/12 of the mass of ana tom of Carbon-12
5
Q
what is 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
6
Q
what is relative formula mass
A
the mass of the formula unit of a compound with a giant structure
7
Q
what is relative molecular mass
A
mass of a simple molecule
8
Q
what is amount of substance
A
a quantity that uses mols as units
9
Q
what does anhydrous mean
A
crystalline compound containing no water
10
Q
what is Avogadro’s constant
A
number of atoms per mole of a substance
6.02 x 10^23
11
Q
what is empirical formula and how do you calculate it
A
simplest whole number ratio of atoms of each element in a compound
1. given mass
2. divide by molar mass of each element
3. divide by smallest number
4. may need to multiply until all are whole numbers
12
Q
what does hydrated mean
A
a crystalline that contains water
13
Q
what is ideal gas and its equation
A
a gas which has molecules that occupy negligible space with no interactions
between them
pV = nRT
14
Q
what is relative molecular formula
A
the average mass of one molecule of an element or compound compared to 1/12th the mass of an atom of carbon-12.
15
Q
what is the meaning of standard solution
A
a solution of known concentration
16
Q
rearrange pV = nRT to give the value of volume
A
v = nRT/P
17
Q
rearrange pV = nRT to give the value of pressure
A
p - nRT/V
18
Q
rearrange pV = nRT to give the value of moles
A
n = pV/RT
19
Q
rearrange pV = nRT to give the value of temperature
A
T = PV/nR
20
Q
what does each symbol in pV = nRT stand for
A
p = pressure (Pa/Pascals)
v = volume (m^3)
n = mols
R = ideal gas (8.314 J K-1 mol-1)
T = Temperature (K / Kelvin)
21
Q
what is atom economy
A
conversion efficiency of a chemical process in terms of all atoms involved and the desired products produced.
22
Q
what is the equation for atom economy
A
mr of desired product/ mr of all products/reactants
23
Q
what is % yield
A
the ratio between what is experimentally obtained and what is theoretically calculated, multiplied by 100%
24
Q
what is the equation for % yield
A
actual mass/theoretical mass x 100
25
why will actual mass always be lower than theoretical mass
- some product is left in apparatus
- some product/reactant is lost to the surroundings
- incomplete reaction
- reactants are not pure
26
What is a dative covalent bond
when one of the bonding atoms donates both of the hared pair of electrons
27
what is an ionic bond
the electrostatic forces of attraction between oppositely charged ions
28
what are the properties of ionic bonding
high melting and boiling points - hlots of energy required to overcome the strong electrostatic forces f attraction
most are soluble in polar solvents as the ions are attracted to polar water molecules and break apart the lattice allowing water molecules to surround the ions
doesn't conduct electricity as solid as ions are in fixed positions with no mobile charge carrier
conducts when molten/aqueous as the charge can flow
29
what is covalent bonding
the strong electrostatic forces of attraction between a shared pair of electrons and the nuclei of bonding atoms
the overlap of atomic orbitals
30
is attraction localised in a covalent bond
yes, it acts solely between the shared pair of electrons and the nuclei of bonding atoms
31
shape of molecule with
2 bonding pairs 0 lone pairs
2 bonding pairs and 2 lone pairs
linear
bent/non-linear
32
shape of a molecule with
3 bonding pairs and 0 lone pairs
3 bonding pairs and 1 lone pair
trigonal planar
triangular pyramid
33
what is the shape of a molecule with
4 bonding pairs and 0 lone pairs
5 bonding pairs and 0 lone pairs
tetrahedral
trigonal byprimid
34
what is the shape of a molecule with
6 bonding pairs and 0 lone pairs
octahedral
35
what does a solid line mean
a bond in the plane of the paper
36
what does a solid wedge mean
a bond that comes out of the plane of the paper
37
what does a dotted wedge mean
a bond that is going into the plane of the paper
38
what is electron-pair repulsion theory
the electron pairs that surround the central atom determine the shape of the molecule or ion
bc the electron pairs repel one another so that they are arranged as far apart as possible
the arrangement of electron pairs and minimises repulsion and thus holds the bonded atoms in a definite shape
39
how much is a bond angle reduced by for each lone pair
2.5 degrees
40
what is the bond angle for
tetrahedral
triangular pyramidal
non-linear/bent
109.5
107
104.5
41
what is the bind angle for
linear
trigonal planar
trigonal bipyramid
octahedral
180
120
109.5
90
42
what is electronegativity
the ability of an atom to attract a pair of electrons in a covalent bond towards itself
43
what is a non polar bond
when the bonded electron pair is shared equally and occurs when the atoms are either the same or have the same/similar electronegativity
44
what is a polar bond
when the electron pair is shared unequally between the bonded atoms and occurs when the bonded atoms are different or have different electronegativity values
45
does a dipole in a covalent bond change
no, it is a permanent dipole
46
what is a polar molecule and when does it occur
when a molecule has a net dipole as the dipoles present may reinforce one another to form a larger dipole over the whole molecule
47
when do you have polar bonds and a non-polar molecule
when a molecule's net dipole is cancelled out as its dipoles are acting in opposite directions
48
what are intermolecular forces and what are their three categories
weak interactions between dipoles if different molecules
London forces
permanent dipole-dipole interactions
Hydrogen bonding
49
what are london forces
weak intermolecular forces that exist between all molecules
act between induced dipoles of different molecules
50
describe the origin of an induced dipole
1. movement of electrons produces a changing dipole in a molecule
2. at any instant, an instantaneous dipole will exist but its position is constantly shifting
3. the instantaneous dipole induces a dipole on a neighbouring molecule
4. the induced dipoles induces further dipoles in neighbouring molecules which then attract one another
51
london forces result from interactions of electrons between molecules
the more electrons…
- the larger the instantaneous and induced dipoles
- the greater the induced dipole-dipole interactions
- the stronger the attractive forces between molecules
52
compare fluorine molecules and HCL molecules
fluorine non polar and only have london forces
HCl polar and have london forces and have permanent dipole dipole interactions between molecules
so bp of HCl is higher than F
53
describe a simple molecular lattice
molecules are held in place by weak intermolecular forces
the atoms within each molecule are bonded together strongly by covalent bonds
54
why do simple molecular substances have low melting and boiling points
weak intermolecular forces can be broken down by the energy present at low temperatures
only the weak intermolecular forces break
the covalent bonds are strong and do not break
55
describe the solubility of non-polar simple molecular substances in non polar solvents
non polar simple molecular substances tend to be soluble in non polar solvents
- intermolecular forces form between molecules and solved and weaken intermolecular forces in the simple lattice
56
describe solubility of non polar substances in polar solvents
little interaction between molecules in lattice and solvent molecules
intermolecular binding within polar solvent is too strong to be broken
57
what does solubility of polar substances depend kn
strength of the dipole
58
what is the electrical conductivity of simple molecular substances
no mobile charged particles in structure
nothing to complete an electrical circuit
59
where do you find hydrogen bonds
between molecules containing an electronegative atom with a lone pair of electrons
O N F
or
in molecules containing a hydrogen atom attached to an electronegative atom
H-O H-N H-F
60
why is ice less dense than water
hydrogen bonds hold water molecules apart in an open lattice structure
the water molecules are further apart than in water
solid ice is less dense than liquid water and flits
61
why does water have a high mp
and bp
hydrogen bonds are strongmg forces - stronger than London forces
H bonds require a lot of energy to break so water has a higher than anticipated m/bp than from just London forces
when ice lattice breaks, the arrangement of H bonds is broke, when water boils, the H bonds break completely
62
what are the other odd properties of water
high surface tensions and viscosity
63
what is Dalton's atomic theory
atoms are tiny particles made up of elements
atoms cannot be divided
all atoms in an element are the same
atoms of one element are different to atoms of another element
64
why do different isotopes of the same element react in the same way
neutrons have no impact on reactivity
reactions involve electrons and isotopes have the same number of the electrons
65
what are the uses of mass spectrometry
identify unknown compounds
find relative abundance of each isotope in an element
determine structural information
66
how does a mass spectrometer work
sample is made into positive ions
they pass through the apparatus and are separated according to mass: charge ratio
computer analyses the data and produces a mass spectrum
67
do metals normally lose or gain electrons
lose electrons
68
state avogadro's law
under the same temperature and pressure, one mole of any gas would occupy the sane volume
69
why do different gas particles occupy the same volume
gas particles are very spread out, individual differences have no effects
70
what are the ideal ways gases behave
are in continuous motion
no intermolecular forces
exert pressure when colliding with each other/container
no kinetic energy lost in the collisions
when temperatures increase, kinetic energy of particles increase
71
what is a standard solution
a solution of a known concentration
72
how do you make up a standard solution
1. weigh solute using weigh-by-difference method
2. in a beaker, dissolve solute using solvent
3. pour solution into a volumetric flask
4. rinse beaker using solution and add to flask
5. add solvent to flask carefully until it reaches graduation line
6. mix solution thoroughly to ensure complete mixing
73
does 100% yield mean 100% economy
no, even if all reactants are converted into products, not all products of the reactions will be the require products
74
what type of reaction has 100% atom economy
addition reactions
75
define acid
proton donor
76
define base
proton acceptor
77
what are alkalis
a base that dissolves in water releasing OH- ions
NaOH(s) + aq --> Na+ (aq) + OH- (aq)
78
what is a salt
a compound that is formed when H+ of an ion is replaced by a metal ion/ cation
79
what are hydrated crystals
a crystalline structure containing water
80
write the method to carry out a titration
1. using a pipette, measure the volume of a solution
2. add the solution into a conical flask and add an indicator to it
3. add the other solution into a burette and record the volume
4. slowly add the solution from the burette into the conical flask
5. swirl the mixture continuously until the end point is reached
6. repeat until concordant results are reached
81
what is the colour of methyl orange in
a. an acid
b. a base
c. end point
red
yellow
orange
82
what is an oxidation number
number of electrons an atom uses to bond with any other atom
83
what is the oxidation number of
a. O in H2O
b. O in peroxides
c. hydrogen in NH3 and H2S
a. -2
b. -1
c. +1
84
how is it indicated 2 when an ion has more than one stable oxidation number
written in roman numerals
85
what is the oxidation number of
a. H in metal hydrides
b. O bonded to F
c. Cl-, Br-, I-
a. -1
b. +2
c. -1
86
What is the oxidation number of Fe in iron(III) chloride
+3
87
define oxidation in terms of electron transfer and oxidation number
- loss of electrons
- increase in oxidation number
88
define reduction in terms of electron transfer and oxidation number
- gain of electrons
-decrease of oxidation number
89
what is a redox reaction/disproportionation reaction
a reaction in which both oxidation and reduction takes place
90
why is the oxidation number of metals 0
they are uncombined elements
91
how many electrons can the first shell hold
2
92
how many electrons can the second shell hold
8
93
how many electrons can the 3rd shell hold
18
94
what is an orbital
a region around the nucleus that can hold up 2 electrons with opposite spins
95
how many electrons can an orbital hold
2
96
what are the types of orbitals
s,p,d,f
97
what is the shape of
a. the s orbital
b. the p orbital
a. spherical
b. dumbbell
98
what are the rules by which electrons are arranged in a shell
electrons are added one at a time
lowest available energy must be filled first
each energy level must be filled before the next starts to fill
each orbital is filled singly before pairing
4s is filled before 3d
99
why does the 4s orbital fill before the 3d
because it has a lower energy level than 3d before it is filled
100
which electrons are lost when an atom becomes a positive ion
electrons in the highest energy level
101
what is metallic bonding
electrostatic attraction between positive metal ions and sea of delocalised electrons
102
in what type of solvents do ionic lattices dissolve
polar
103
why are ionic compounds soluble in water
water has polar bonds
hydrogen atoms have δ+ and oxygen atoms have δ-
charges are able to attract charged ions
104
how many covalent bonds does carbon form
4
105
how does graphite conduct electricity
has delocalised electrons that are able to move freely and carry the charge
106
what is electronegativity
relative tendency of an atom in a covalent bond in a molecule to attract electrons to itself
107
what factors affect electronegativity
increases across a period because: increasing proton number - atomic radius decreases and electrons in the same shell are pulled in more
Decreases downs group because: the distance between the nucleus and the other outer electrons increases and shielding of inner electrons increases
108
how do you explain the shape of a molecule
1. state the shape and number of bonding pairs
2. state that electrons repel and try to get as far away from each other as possible
3. if no lone pairs, state that electrons repel equally
4. if there are lone pairs, state that they repel more than bonding pairs
5. state actual shape and bond angle
109
when does a polar bond form
when the electronegativities of each element in the bond is different
110
when is a bond purely covalent
when each element has similar electronegativities
111
where do induced dipole dipole interactions form
between every molecule
112
what factors affect dipole-dipole interactions
aka london forces
the more electrons there are in a molecule
the higher the chance that temporary dipoles will form
makes the induced dipoles stronger between molecules and so boiling points will be greater
113
can permanent dipole interactions be in addition to induced dipoles-dipole interactions
yes
