Elements of Life Flashcards
1
Q
Hydrochloric Acid
A
HCl
2
Q
Nitric Acid
A
HNO3
3
Q
Sulphuric Acid
A
H2SO4
4
Q
Phosphoric Acid
A
H3PO4
5
Q
Methanoic Acid
A
HCOOH
6
Q
Ammonia
A
NH3
7
Q
Isotopes
A
Different atomic number same mass number
Different number of neutrons same number of protons
8
Q
Relative Atomic Mass
A
A measure of the average mass of an element compared to a standard unit of mass
Calculated by adding all masses of an element then averaging them based on their abundance
9
Q
Atomic Mass
A
The actual mass of an individual isotopes - a fixed value
10
Q
How to calculate RAM
A
Need abundance of each isotope
Need the mass number
(mass * abundance) + (mass * abundance)/(abundance + abundance)
11
Q
Example of RAM with Chlorine
A
Cl-35 75%
Cl-37 25%
(75 * 35) + (37 * 25)/ 75 + 25 = 35.5
11
Q
Atom
A
The smallest piece of an element, made up of protons, neutrons and electrons and has a neutral charge
12
Q
Molecule
A
A group of atoms held together by covalent bonds
13
Q
Ion
A
An atom or molecule which has lost or gained electrons so is positively or negatively charged
14
Q
Element
A
A substance made from one type of atom only
15
Q
Compound
A
A substance formed from two or more chemically bonded elements in a fixed ratio, shown by a chemical formula
16
Q
Simple Structure
A
Made up of small molecules held together by weak intermolecular forces
17
Q
Giant lattice structure
A
A 3-dimensional structure of particles held together by strong bonds (covalent or metallic)
18
Q
Group
A
A vertical column in the periodic table. The element has the same number of outer electrons hence similar chemical properties
19
Q
Intermolecular Forces
A
Attractive forces between neighbouring molecules
20
Q
Neutralisation Reactions
A
Acid + Base –> Salt + Water
21
Q
Combustion
A
Complete - +O2 —> CO2 + H2O
Incomplete - +O2 —> CO + H2O
22
Q
Acid + Metal Carbonate
A
—> Salt + Water + Carbon Dioxide
23
Q
Acid + Metal Oxide
A
—> Salt + Water
23
Acid + Metal Hydroxide
---> Salt + Hydrogen + Water
24
Common Acids
H+ Donors
Hydrochloric - HCl
Nitric - HNO3
Sulfuric - H2SO4
Phosphoric - H3PO4
24
Acid + Metal
---> Salt + Hydrogen
24
Common Alkalis/Bases
H+ Acceptors
Group 1 hydroxides + Oxides = Na2O
Group 2 hydroxides + Oxides = MgO
25
Polyatomic Ions
Hydrogen Sulfate - HSO4-
Carbonate - CO3 2-
Ethanoate - CH3COO-
Hydroxide - OH-
Sulphate - SO4 2-
Nitrate - NO3-
Hydrogen Carbonate - HCO3-
Ammonium - NH4+
26
Ar
Relative Atomic mass (average mass of an atom)
27
A
Atomic Mass (number of P+ and e-)
28
Z
Atomic Number (number of p+)
29
Nucleons
Protons and neutrons as they're found in the nucleus
30
Avogadro's constant
6.022 * 10^23
31
Molecular Formula (Mr)
Shows the actual number of atoms of each element present in a compound or molecule
Glucose's Mr = C6H12O6
32
Empirical Formula
The simplest whole ratio of atoms of each element
Glucose = CH2O
33
Water of Crystallisation
Water that's trapped between ions when an ionic solid forms
Ionic solids with woc = hydrated
Ionic solids without woc = anhydrous
34
woc shown in a formula
.xH2O where x = a number
35
Simple experiment for woc
1) Weigh crucible before and after hydrated solid is added
2) Place above bunsen burner
3) Stop heating, cool it, weigh it
4) Repeat until constant mass
36
Nuclear Fusion
2 small nuclei join to make a larger nucleus
Only occurs at a high temp and high pressure - lots of energy needed to overcome repulsion between 2 positive nuclei
37
Nuclear Fusion equations
add top numbers together and bottom numbers
12C + 4 He ---> 16 BE
6 2 8
38
Solute
The solid in the solution
39
Solvent
The liquid that the solute is getting dissolved in
40
1dm^3 in litres
1 litre
41
Concentration equation
Conc=mol/volume (volume = dm3 (cm3/1000))
42
Titration use
Used to find the concentration of one solution
43
Volumetric Pipette
Rinse with what it's measuring
Touch tip to the surface when emptying
Normally 25cm3
44
Burette
Rinse with what it's measuring
Read to 2d.p - last digit = 5 or 0
Bottom of meniscus
45
Method of titration calculations
1) Write a balanced equation
2) Write information underneath (average titre or 25cm3)
3) Calculate the moles = conc * dm3
4) Use Stoichiometric ratio to determine moles of other reactant
5) Calculate conc (mol/dm3) give to 3s.f
46
Acid
H+ donor
47
Base
H+ receiver
48
gcm-3
Convert to moles then times by 1000 - mass/mr * 1000
49
gmdm-3
Convert to moles - mass/mr
50
v/v
Find percentage then times by 1%
51
PPM
Parts per million
ppm/1000 = dm3
Then moles = mass/mr
52
Ionic Compounds
Giant lattices
Regular structure
Ions of opposite charge in a fixed ratio
53
Ionic Compounds definition
The electrostatic attraction between oppositely charged ions
No e- transfer
Not always metal and non-metal
54
Proof of ionic compound
Presence of a polyatomic ion
One from LHS and one from RHS
55
Ionic compound conductivity
Conduct when molten
Conduct if dissolved
No conduction if solid
-Ions are free to move around when not solid
56
Ionic Compound Mpt/Bpt
High due to electrostatic forces requiring energy to overcome
Also has a giant lattice that requires lots of energy to overcome
57
Ionic Compound Solubility
Many dissolve in water
Water molecules interact with salts surface ions.
If soluble, the H+ attracts the anion and the O2- attracts the cation.
Ions that are pulled away are called hydrated ions
Held in place by ion-dipole forces
58
Ionic dot and cross
Coefficient [M] charge Coefficient[NM]charge
59
Standard Solution
We know the concentration very accurately
60
s orbitals
Spherical, one on each energy level
61
p oribtals
Dumbell shaped, 3 on each energy level except n = 1
62
d orbitals
Crazy shapes, 5 on each energy level starting at n = 3
63
f oribtals
Crazy shapes, 7 on each energy level starting at n = 5
64
D block when writing
Always lags one behind
65
Rule of noble gases in configuration
Use noble gas before and add remaining letters
66
Ionisation energy
The energy required to remove the outermost electron from 1 mole of gaseous atoms
67
Ionisation energy equations
Mg(g) --> Mg+(g) + e-
Cl (g) --> Cl+(g) + e-
68
Successive ionisation energy
Electrons are removed one after the other from the same atom
e.g
1) Cl(g) --> Cl+(g) + e-
2) Cl+(g) --> Cl2+ (g) + 2e-
69
Ionisation energy trend
Each time, more is required as same number of protons is attracting fewer electrons
70
Absorption Spectra
Black lines
Rainbow background
When electrons move from a low EL to a high EL
71
Emission Spectra
Coloured lines
Black Background
When electrons move from a high El to a low EL
72
Spectra line spacing
Lines are always closer on the high frequency end of the scale (purple)
73
Absorption and Emission lines same element
Same elements will have the spectra lines in the same place as the ELs are the same so the frequency released matches the frequency absorbed
74
Absorption and Emission lines different element
Different elements have different spectra lines as the gaps between the ELs are different
75
Model Answer for for spectrum
Only some frequencies of light are emitted or absorbed for a particular element as the energy levels are fixed (quartised) and the electrons can only move between energy levels
76
Frequency equation
change in energy = plancks constant * frequency
77
Wavelength calculation
speed of light = wavelength * frequency
78
wavelength value
nanometres (*10^-9)
79
Metallic bonding definition
The electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons
80
Metallic structure
-Giant
-Delocalised electrons (constantly moving)
-Charge written inside the circles (+, 2+ etc..)
-As positive charge increases, size of ions decrease
81
Metallic Bonding Mpt/Bpt
Very high due to giant structure
Very high due to strong attraction between ions and delocalised electrons
82
Metallic Bonding Conductivity
Very conductive due to delocalised electrons being free to move
83
Metallic Bonding Solubility
Insoluble in H2O, forces of attraction between ions and e- are very strong
84
Covalent Bonding Definition
The electrostatic Attraction between shared pairs of electrons and two positive nuclei
85
Covalent Structures
Can be simple or giant. Simple is much more common than giant
86
Examples of giant covalent structures
Diamond
Graphite
Silicon dioxide
87
Covalent bonding trends
No of pairs increases
Type of bond increases (single, double and triple)
Relative bond length decreases
Relative bond strength increases
Two double bonds next to each other are not possible
88
Explanation of dative bond
Electrons are shared so in most cases the atoms have a complete outer energy level. Most bonds have an electron from each atom but in some covalent bonds, both electrons come from one atom (a dative bond)
89
Covalent bonding rules
No of covalent bonds normally = 8 - group number
e.g C = 8 - 4 = 4
H only makes 1 bond
Be makes 2 bonds due to 2 outer electrons
Al makes 3 bonds due to 3 outer electrons
P makes 8 - 5 =3 bonds or 5 bonds due to outer electrons
S makes 8 - 6 = 2 bonds of 6 bonds due to outer electrons
90
Dot and Cross diagram
1) Check it's covalent (RHS + Beryllium)
2) Draw the molecule with lines and check all atoms have the correct number of bonds
3) Replace each lines with a dot or a cross
4) Add lone pairs (group number - e- used for bonding)
91
Dative bonds
A covalent bond where one atom donates both e-
Shown by an arrow
