2-Foundations in Chemistry Flashcards
1
Q
Proton
A
M=1
Q=1+
2
Q
Neutron
A
M=1
Q=0
3
Q
Electron
A
M=1/1836
Q=1-
4
Q
Isotopes
A
Atoms of the same element with different numbers of neutrons and different masses
5
Q
Reactions with isotopes
A
Different isotopes of the same element have the same amount of electrons and therefore have the same chemical properties and so react in the same way
6
Q
Positive ions
A
Cations - fewer electrons than protons
7
Q
Negative ions
A
Anions - more electrons than protons
8
Q
Relative isotopic mass
A
The mass of an isotope relative to 1/12th of the mass of an atom of carbon 12
9
Q
Relative atomic mass
A
The weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12 (Ar)
10
Q
Mass Spectrometry
A
- sample is vaporised and then ionised to form cations
- the cations are accelerated, heavier ions move more slowly and are more difficult to deflect than lighter ions so ions of each isotope are separated
- the ions are detected on a mass spectrum as a mass-to-charge ratio. The greater the abundance the larger the signal
11
Q
Ammonium Hydroxide Nitrate Carbonate Sulphate
A
NH+ OH- NO3- CO3(2-) SO4(2-)
12
Q
1 mole
A
The amount of any substance containing as many elementary particles as there are carbon atoms in exactly 12g of carbon-12, which is 6.02 x 10^23
13
Q
Avogadro constant
A
6.02 x10^23 mol^-1, the number of particles in each mole of carbon-12
14
Q
Molar mass
A
The mass per mole of a substance, in units of g mol^-1
15
Q
Equation link molar mass, moles and mass
A
N=mass/molar mass
16
Q
Molecular formula
A
Formula that shows the number and type of atoms of each element present in a molecule
17
Q
Empirical formula
A
The formula that shows the simplest whole-number ratio of atoms of each element a compound
18
Q
Relative molecular mass
A
The weighed mean mass of a molecule of a compound compared with 1/12th of the mass of an atom of carbon-12
19
Q
Relative formula mass
A
The weighted mean mass of the formula of a compound compare to 1/12th of the mass of an atom of carbon-12
20
Q
Hydrated
A
Water molecules are part of the crystalline structure
21
Q
Anhydrous
A
When hydrated crystals are heated and water is driven off
22
Q
Negatives about carrying out experiment to find the formula of a hydrated salt
A
- not all water may be lost
- many salts decompose further when heated
23
Q
Concentration
A
The amount of solute, in moles, dissolved in each 1dm^3 of solution
24
Q
Equation linking moles, volume (in cm^3), and concentration
A
n=VxC/1000
25
Steps to work out the mass needed to make a standard solution
- work out the amount in moles (of the compound) required, using n=VxC/1000
- then workout the molar mass of compound
- use n=M/m to find mass of substance
26
RTP conditions
20°C and 101kPa (1 atm)
27
How to work out moles of gas at RTP
n=V(dm^3)/24
28
Ideal gas equation
pV=nRT
p(kPa)
V(m^3)
R(8.314mol^-1 K^-1)
T(°C+273 = K)
29
Percentage yield
(Actual yield/theoretical yield)x100
30
Limiting reagent
The reactant that is not in excess that will be completely used up first and stop the reaction
31
Atom economy
(Sum of molar masses of desired products/sum of molar masses of all products) x 100
32
Other factors for sustainability other than atom economy
- processes sometimes uses reactants that are readily available, e.g. Carbon from coal and steam from water therefore costs for obtaining materials are low
- some reactions may have much larger atom economy but poor %age yields; efficiency depends on both factors
33
Acids
A species which releases H+ ions in an aqueous solution
34
Base
A compound which neutralises an acid to form a salt
35
Alkali
A type of base that dissolves in water forming hydroxide ions, OH- ions
36
Strong acids
Release all H+ ions into a solution; completely dissociate in aqueous solution
37
Weak acids
Only releases a proportion of its available H+ ions; partially dissociates in aqueous solution
38
Neutralisation of an acid
H+ ions react with a base to form a salt and neutral water;
| This can either be a metal oxide or a metal hydroxide, or a carbonate (also produces CO2)
39
Titration
Technique used to accurately measure the volume of one solution that reacts exactly ugh another solution. They can be used to force ne the concentration of something, identify an unknown chemical, and *find the purity of a substance.
Finding purity is very important in the world, e.g. For compounds manufactured for human use (medicines)
40
Preparing a standard solution
- Weigh solid accurately
- dissolve solid in a beaker using less distilled water than will be needed to fill volumetric flask the mark
- transfer to volumetric flask rinsing the last traces into the flask with distilled water
- carefully fill the flask with distilled water up to the mark exactly so that the bottom of the meniscus is just touching the mark
- slowly invert the flask several times
41
Oxidation number rules
Atoms of the same element = 0
Compound and ions = oxidation number
O= -2. O in peroxides= -1 (H2O2) O bonded to F= +2.
H= +1. H in metal hydrides= -1 (NaH/CaH2)
F=-1
Groups 1&2 metals = +1 & +2 respectively
Halides=-1
42
Reduction
Gain of electrons
| Decrease in oxidation number
43
Oxidation
Loss of electrons
| Increase in oxidation number
44
Shell number or energy level is known as
Principle quantum number (n)
45
Atomic orbital
Region around the nucleus that can hold up to 2 electrons, with opposite spins
Models visualise them as a region in space where there is a high probability of finding an electron
46
Electron
A negative-charge cloud with the shape of the orbital, referred to as an electron cloud
47
s-orbitals
Electron cloud is within the shape of a sphere
Each s-sub shell contains 1 s-orbital (2 electrons)
The greater the shell number (n) the greater the radius of its s-orbital
48
p-orbitals
Electron cloud within the shape of a dumbbell
Each p-sub shell contains 3 p-orbitals (6 electrons)
The greater the shell number (n) the further away the p-orbital is from the nucleus
49
d-orbitals
Each d-sub shell contains 5 d-orbitals (10 electrons)
50
f-orbital
Each f-sub shell contains 7 f-orbitals (14 electrons)
51
Formula for amount of electrons in each principle quantum number
2n^2. (Square n first, then x2)
52
Opposite spin of electrons in orbitals reasoning
The opposite spin helps counteract the repulsion between the negative charges of the two electrons
53
s-block
Highest energy electrons are in the s-sub-shell (left block of periodic table; group 1&2)
54
p-block
Highest energy electrons are in the p-sub-shell (right block of periodic table; group 3-8)
55
d-block
Highest energy electrons in the d-sub-shell (centre block of periodic table; between group 2&3 (10 blocks))
56
Ions of d-block elements
- The 4s sub shell is at a lower energy than the 3D sub-shell, so is filled first
- The energies of the 4s and 3d sub shells are very close together and, once filled, the 3D energy level falls below the 4s energy level
- This means the 4s sub shell fills before the 3D sub shell but also empties before the 3D sub shell
57
Ionic bonding
The electrostatic attraction between positive and negative ions. It holds cations and anions together in ionic compounds (metals and non metals)
58
Structure of ionic compounds
Each ion attracts oppositely charge ions in ALL directions
| Therefore a giant ionic lattice is formed
59
Properties of ionic compounds
- High melting points and boiling points; high ion charges mean higher melting points
- Tend to dissolve in polar solvents such as water; some do not due to higher ion charges
- Conduct electricity only as liquid or gas in solution; due to mobile charge carriers being available (ions are free to move).
60
Required processes for solubility
- Ionic lattice must be broken down
| - H2O must attract and surround the ions
61
Covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
Between: elements of non-metals (O2)
compounds of non metallic elements (H2O)
Polyatomic ions (NH4^+)
62
The covalent bond (how it works)
- Atomic orbitals overlap, each containing one electron, to give a shared pair of electrons
- The shared pair of electrons is attracted to the nuclei of both the bonding atoms
- A covalent bond is localised, acting solely between the shared pair of electrons and the nuclei of the two bonding atoms; can result in a small unit called a molecule
63
Boron's covalent bond properties
Boron's electron configuration is 1s2, 2s2, 2p1, so only 3 outer shell electrons can be pared, leaving only 6 electrons in its outer shell(e.g. BF3); not all elements follow the octec rule.
64
Phosphorus, sulphur and chlorine
For phosphorus, sulphur and fluorine, the n=3 outer shell can hold 18 electrons, so more electrons are available for bonding, this means that the electrons in the outer shell can arrange themselves how they please and make as many covalent bonds as they want; (e.g. SF2 or SF4 or SF6) this is an expansion of the octet and is only available in the n=3 shell, when a d-sub-shell becomes available for the expansion
65
Double covalent bond
The electrostatic force of attraction between 2 shared pairs of electrons and the nuclei of the bonding atoms
66
Triple covalent bond
Electrostatic attraction is between three shared pairs of electrons and the nuclei of the bonding atoms
67
Dative covalent bond
A covalent bond in which the shared pair of electrons have been supplied by one of the bonding atoms only.
It is shown by an arrow going away from the atom that supplies the electron pair
68
Electron pair repulsion theory
Electron pairs repel one another and so they arrange themselves as far apart as possible from each other to minimise repulsion and thus hold the bonded atoms in a definite shape; this determines the shape of molecule
69
Shape and bond angle of CH4
Tetrahedral
109.5
4 bonding pairs
70
Shape and bond angle of CO2
Linear
180°
2 bonding pairs
71
Shape and bond angle of BF3
Trigonal planar
120°
3 bonding pairs
72
Shape and bond angle of SF6
Octahedral
90°
6 bonding pairs
73
Shape and bond angle of NH3
Pyramidal
107°
3 bonding pairs
1 lone pair (reduce bond angle by 2.5°)
74
Shape and bond angle of H2O
Non-linear
104.5°
2 bonding pairs
2 lone pairs (reduce bond angle by 2.5°x2)
75
Electronegativity
A measure of the attraction of a bonded atom to a pair of electrons in a covalent bond
76
Electronegativity explanation
In molecules of the same element the electron pair is shared evenly. This changes when the bonded atoms are different elements:
-the nuclear charges are different
-the atoms may be different sizes
- the shared pair of electrons may be closer to one nucleus than the other
Therefore the shared pair of electrons may experience more attraction from one of the bonded atoms than the other
77
Non-polar bonds
the electron pair is shared equally between bonded atoms. This happens when:
- the bonded atoms are the same or
- the bonded atoms have the same or similar electronegativity
* When this happens this is a pure covalent bond
78
Polar bond
When the electron pair is shared unequally between the bonded atoms. This happens when:
- the bonded atoms are different and have different electronegativity values
* this is called a polar covalent bond
79
Permanent dipoles
This is in polar covalent bonds and is the permanent separation of opposite charges; this is called a permanent dipole
80
Polar solvents and ionic compounds
NaCl dissolves in water because H2O molecules are polar (because the shape of the molecule is non-linear). This means the hydrogen end is positive, oxygen end is negative. This means that the negative Oxygen end of the H2O will attract the Na+ ions and the positive hydrogen end of the H2O will attract the Cl- ions. This means that the ionic lattice breaks down and the H2O molecules will surround the Na+ ions and Cl- ions
81
Intermolecular forces
weak interactions between dipoles of different molecules
82
Types of intermolecular forces
- induced dipole-dipole interactions (London forces
- permanent dipoles-dipole interactions
- Hydrogen bonding
*these forces are largely responsible for physical properties such as melting and boiling points, whereas covalent bonds determine the identity and chemical reactions of the molecules
83
Induced dipole-dipole interactions (London forces)
Weak intermolecular forces that act between ALL molecules, whether polar or non-polar. They act between induced dipoles in different molecules
84
The origin of induced dipoles
- movement of electrons produces a changing dipole in a molecule
- at any moment, an instantaneous dipole will exist, but its position is constantly shifting
- the instantaneous dipole induces a dipole on a neighbouring molecule
- the induced dipoles induces further dipole on neighbouring molecules which then attract one another
85
The strength of London forces
The more electrons in each molecule:
- the larger the instantaneous dipole-dipole interactions
- the greater the induced dipole-dipole interactions
- the stronger the attractive forces between molecules
86
Permanent dipole-dipole interactions
The permanent dipoles in polar molecules attract the other permanent dipoles in other molecules in the atoms of opposite partial charge
87
Solubility of simple non-polar molecular substances
When added to a non-polar solvent, intermolecular forces form between the molecules and the solvent. The interactions weaken the intermolecular forces in the simple molecular lattice. The intermolecular force break and the compound dissolves.
When added to polar solvents, however, there is little interactions between the molecules in the lattice and the solvent molecules. The intermolecular bonding between the polar solvent is too strong to be broken and therefore non-polar molecular substances tend to be insoluble in polar solvents
88
Solubility of polar simple molecular substances
When added to polar solvents the may dissolve because he solute molecules and the polar solvent molecules can attract each other. The process s very similar to that of ionic compounds in polar solvents.
The solubility depends on the strength of the dipole and can be hard to predict. Some compounds (like C2H5OH) have polar and non-polar parts and so can dissolve m bother polar and non-polar solvents
90
Hydrogen bonds
A strong dipole-dipole attraction between an electron-deficient hydrogen atom of -NH, -OH or HF molecule and a lone pair of electrons on a highly electronegative atom containing N, O, or F on a different molecule
92
Anomalous properties of water (due to Hydrogen bonds)
- Solid (Ice) is less dense than liquid (water)
- relatively high melting and boiling point
- relatively high surface tension
- high viscosity
92
Explanation for Ice being less dense than water
- Hydrogen bonds hold H2O molecules apart in an open lattice structure
- the water molecules in ice are further apart than in water
- solid ice is less dense than liquid water and therefore floats
92
Explanation for high melting and boiling points
- Hydrogen bonds are extra forces, over and above landing forces
- therefore more energy is required to over come the strong hydrogen bonds acting between the H2O molecules
- so it has a higher melting and boiling point
{without Hydrogen bonds water would exist as a gas at room temp, having a boiling point of about -75°C, meaning almost no liquid water on earth and there would be no life as we know it}
