Sem 1 Chem Flashcards
(126 cards)
1
Q
Avagadro’s Number
A
6.022 * 10^23
2
Q
Particle - Mole Relationship
A
n (mol) = n of particles / Avagadro’s number
3
Q
Mole - Mass Relationship
A
n(mol) = mass/molar mass
4
Q
Atom
A
Smallest unit of matter that can be divided without the release of electrically charged particles
5
Q
Compound
A
A substance made up of different atoms chemically bonded together
6
Q
Molecule
A
Smallest particle of a substance that has all of the physical and chemical properties of that substance
7
Q
Molecular Element
A
Elements that exist in molecules or more than one atom such as diatomic gases, sulphur. They are still an element, just in a molecular form
8
Q
Pure substance
A
A pure substance is a substance where there is one type of chemical present (only molecule, only compound, such as only H2O or only NaCl). Cannot be separated by physical means
9
Q
Mixture
A
Substance containing two or more different substances in varying proportions
10
Q
Homogenous Mixtures
A
Have uniform composition throughout, include all pure substances and some mixtures (like salt dissolved in water or sugar dissolved in water)
11
Q
Heterogenous Mixtures
A
Have non-uniform composition and you can see differences from one section to another. The constituents are physically separated in a different phase (e.g granite, milk, toothpaste as toothpaste solid particles suspended in liquid)
12
Q
Ammonium
A
NH4 +
13
Q
Dihydrogenphosphate
A
H2PO4 -
14
Q
Ethanoate/Acetate
A
CH3COO -
15
Q
Hydrogencarbonate
A
HCO3 -
16
Q
Hydrogensulfate
A
HSO4 -
17
Q
Nitrite
A
NO2 -
18
Q
Nitrate
A
NO3 -
19
Q
Permanganate
A
MgO4 -
20
Q
Hydrogenphosphate
A
HPO4 (2-)
21
Q
Oxalate
A
C2O4 (2-)
22
Q
Oxide
A
O (2-)
23
Q
Sulfate
A
SO4 (2-)
24
Q
Sulfide
A
S (2-)
25
Sulfite
SO3 (2-)
26
Nitride
N (3-)
27
Phosphate
PO4 (3-)
28
Ammonia
NH3
29
Hydrogen Peroxide
H2O2
30
Ethanoic/Acetic Acid
CH3COOH
31
Chromate
CrO4 (2-)
32
Dichromate
Cr2O7 (2-)
33
Nitric Acid
HNO3
34
Carbonic Acid
H2CO3
35
Sulfuric Acid
H2SO4
36
Sulphurous Acid
H2SO3
37
Phosphoric Acid
H3PO4
38
What are nano-materials?
Materials with at least one dimension between 1 and 100 nanometres
39
How do properties of nanomaterials differ from bulk materials?
Due to the small particle size of nano-materials, it causes quantum effects to arise. E.g: Zinc Oxide and Titanium Oxide in macro form cause an opaque layer on sunscreen. In nano form, they don't.
40
Advantages and disadvantages of nano-materials
Advantages: Advances in computing (faster computing), medicine (cell targeting) and more efficient power sources.
Disadvantages: Not all quantum effects are known and they can be detrimental, bad for environment and nanoparticles can get into body
41
What is Dalton's Atomic Theory?
Proposed the Atomic Theory, stating:
- Elements are composed of extremely small particles called atoms
- All atom of a given element are identical (have the same size, mass and chemical properties)
- Atoms are created nor destroyed nor changed into other atoms during a chemical reaction
- A chemical reaction involves only separation, combination or rearrangement of atoms
- Compounds are formed when atoms of more than one element combine in a specific ratio
42
Date of Dalton's Discovery
1803-1808
43
What did Dalton's theory not account for or could not explain?
- Could not account for structures of atoms (existence of subatomic particles)
- No account of isotopes
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Dalton's Atomic Model Name
Billiard Ball model
45
J.J Thomson's discovery
- The electron by observing cathode rays and determining that rather than being radiation, they were streams of charged rays with mass, which were called electrons
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J.J Thomson's Atomic Model name
Plum pudding model
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Date of J.J Thomson's Discovery
1897
48
Ernest Rutherford Discovery Date
1909-1913
49
Ernest Rutherford Discovery
Discovery of nucleus/protons by firing alpha particles at gold foil. Most passed through but some deflected, indicating most of an atom is empty space with charge concentrated in area called nucleus
50
Ernest Rutherford Atomic Model Name
Planetary Model
51
Date of Sir James Chadwick's Discovery
1932
52
James Chadwick's Discovery
Discovered neutron by bombarding beryllium with alpha particles
53
Niels Bohr Discovery
Bohr proposed that electrons orbited in specific orbits, which would explain the emission and absorption spectra of atoms.
54
Niels Bohr Discovery Date
1913
55
What is relative atomic mass?
The atomic mass of an element, relative to it's abundance. Calculated using the mass number multiplied by the percentage abundance and then adding that for all isotopes of that element.
56
Which instrument can be used to separate and detect different isotopes in a sample?
Mass spectrometer
57
Name and describe the steps in mass spectrometry
1. Vaporisation: Sample is vaporised (becomes gas)
2. Ionisation: Sample is ionised via the firing of a beam of electron, which converts the atoms into 1+ or 2+ ions.
3. Acceleration/Separation: The sample is accelerated via a field, and then through the use of of an electric field. The sample is then deflected using a magnetic field, with different isotopes in the sample deflecting differently due to their mass-charge ratio. (Ions with a lower mass number experience more deflection)
4. Detection: A detector at the end of the spectrometer measure the intensity and radius of deflection of each of the ion beams. The count/data is of the isotope abundance in the sample
58
What determines an elements chemical properties?
Number of electrons in valence shell
59
Where do electrons exist?
Shells or energy levels around the nucleus, with increasing levels being further away from the atom. The letter n is used to denote energy levels.
60
What is the octet rule?
The tendency of atoms to achieve a full outer shell by having 8 electrons in their outer shells (simplified). They do this by gaining, losing or sharing electrons.
61
How is the absorption spectrum produced?
Every element has a different number of electrons. By exposing these atoms and by extension the electrons to light, it is possible to excite the electron and make it 'jump' to the next energy level, say from n=1 to n=2. This will only occur if the light has the same amount of energy as the difference in energy between the energy levels. The specific energy can be visualised as the different colours of light. This is represented as black lines on the visible spectrum. The black lines are the specific wavelengths/energies of light that will excite the atoms. The absorption spectrum of elements is identical.
(Spectrum goes from right to left, lines get closer together as less energy is needed to jump due to less attraction from nucleus)
E.g: Take hydrogen's n=1 level. We can make it jump up to the 4th shell, provided the wavelength of the light we expose it to is dark blue.
62
How is the emission spectrum produced?
After an atom is excited, it cannot remain at the energy level it has jumped to. The exact amount of energy it used to jump is the exact energy is emits. The electrons returns to the ground state and the emission of light is exact as an exact amount of energy was used to jump, so we can match the black lines of the absorption spectrum to the emission spectrum. Again, the lines get closer as we go towards the left.
Note: An electron may not jump back to the ground state immediately, it may jump to n=2 before n=1
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What do the lines in the absorption spectrum represent?
Electrons gaining energy as they move to higher energy levels
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What do the lines in the emission spectrum represent?
Electrons losing energy as they move back down to lower energy levels in the atom.
65
How can flame tests be used to identify elements?
Flame tests work by burning a substance in a blue flame of a bunsen burner.
Different elements emit different colours when burned so we can identify which element is in our substance based on the colour that we can see.
This works because different elements have different emission and absorption spectrum so different elements emit different colours for the same energy because of their individual spectrums.
66
How can AAS be used to identify element? (what is it, how it identifies elements/procedure and why it works.)
AAS or Atomic Absorption Spectroscopy is when a substance is analysed using light as there is a relationship between the concentration of the sample being analysed and the intensity of the light it absorbs.
In AAS, the sample is first vaporised in a flame to produce atoms in a ground state. This cloud of atoms then passes through a beam of light, which causes the substance to absorb the light and emit it. This produces the emission spectrum of the chemical or substance being analysed. The intensity of the light absorbed and emitted by the substance is detected by the machine. This allows us to determine the concentration of a particular element or substance within a chemical as the intensity of the light emitted/absorbed will indicate the concentration of the element in the sample.
E.g: If we are trying to detect whether or not water has been contaminated with bromine, we can vaporise the water, pass it through the AAS and if we detected bromine's emission spectrum, then there is some bromine present and the intensity of the light emitted will tell us the concentration.
67
How do the valence electrons affect the reactivity of an element?
The closer an element is to having a full outer shell, the more reactive it is.
68
What is the relationship between an element's group and it's valence electrons?
Elements in the same group have the same number of valence electrons, thus they have similar chemical properties.
69
What are all the trends in the periodic table? (list them)
- First ionisation energy
- Electronegativity
- Atomic Radius
- Valency
70
Define first ionisation energy
- The amount of energy required to remove one mole of electrons from one mole of atoms in the GASEOUS state.
71
What is the trend in ionisation energy and why?
- Ionisation energy decreases as we go down a group because there are more electrons shells so the valence shell and electron is further away from the nucleus, so it is easier to remove the electron
- Ionisation energy increases as we go to the right of a period because there is the same number of shells, but different number of electrons so there is more of an electrostatic attraction as the nuclear charge increases so it becomes harder to remove an electron
72
Define electronegativity
The measure of an atom's attraction for a shared pair of electrons in a bond
73
What is trend in electronegativity and why?
- The electronegativity increases as you go towards the right of a period
- This is because the nuclear charge increases across a period so they attract the electrons greater
- As you go down a group, the electronegativity decreases because the atoms have a bigger atomic size/radius so the nucleus is further away from the bonding electrons and is shielded by the extra electron shells
74
What is the trend in atomic radius and why?
- Atomic Size increases the further down a group you go and the further right you go, the atomic size decreases
- This is is because, the further down a group you go, more electron shells are added so the outermost electron becomes further out from the nucleus which increases the atomic radius
- The atomic size decreases the further right you go because as you go righter, there is an increase in nuclear charge (protons increase) and more electrons are also added to the same shell (periods have similar shells) which increases electrostatic attraction so the electrons are pulled in closer which decreases the atomic size
75
What is valency?
The number of valence electrons
76
What is the trend in valence electrons?
(Note metals are exceptions), Elements in the same group have the same valency and valency increases as we go to the right of a period as atomic number increases.
77
What is ionic bonding? (which elements, mechanism and structure)
Ionic bonding is when a metal bonds with a non-metal. It occurs when the metal loses electrons to the non-metal forming cations and anions respectively. This creates an electrostatic attraction which causes the elements to bond together. The ions are held together in a lattice with the metal ions surrounded by non-metal ions. It is non-directional
78
Properties of Ionic Compounds (just list them)
- High melting point
- Conductors of electricity in the molten and aqueous state
- Hard but brittle
- Most are soluble in water
79
Why do ionic compounds have high melting points?
Lot of energy is required to break the electrostatic energy between the ions as there is a strong electrostatic attraction.
80
Why do ionic compounds only conduct electricity in very specific circumstances?
In the solid state, the ionic compound has no free charged particles to conduct electricity. In the molten and aqueous state the particles are free to be mobile and thus can conduct electricity.
81
Why are ionic compounds hard yet brittle?
The compound is hard as it resists being scratched due to the incredibly high electrostatic attraction in the lattice but once sufficient force is applied, the lattice shifts position and the cations and anions come close together, repelling each other and shattering the lattice structure.
82
What is a hydrous salt?
A hydrous salt is a salt which contains water in its structure. Usually, the salt crystallises with water as a part of its structure.
83
What is an anhydrous salt?
A salt with no water in its structure.
84
What is the role of crystallisation?
To speed up the process of crystallising
85
What is metallic bonding?
Metallic bonding occurs between two or more metals. It occurs when metals come together and release their electrons, becoming cations. These ions come together in a 3-dimensional lattice and the attraction between the free electrons and the ions holds the atoms in place.
86
Properties of Metallic Bonding (list them)
- Good conductors of heat and electricity
- Malleable and ductile
- High melting and boiling points
87
Why are metallic substances good conductors of heat and electricity?
The valence electrons which have been released are free to move so they can conduct a current once voltage has been applied.
The electrons also can carry heat energy, in the form of kinetic energy, because they are free to move around. The vibration of the metallic lattice also contributes to the conduction of heat.
88
Why are metallic substances malleable and ductile?
They are malleable and ductile because metallic bonding is non-directional meaning that the metal ions can move around without breaking the bond between itself and the electrons. This allows the metal to bend and be able to be drawn out into a wire.
89
Why do metallic substances have a high melting point?
Metallic substances have a strong attractive force between the cations and the sea of delocalised electrons which requires a lot of energy/heat to break.
90
What is covalent molecular bonding?
Covalent molecular bonding occurs when two non-metals share electrons. The bond is directional. The bonding occurs due to the attraction between the positive nuclei of the atoms and the shared electrons.
91
What are the properties of covalent molecular substances? (list them + exceptions if applicable)
- Do not conduct electricity in either solid, liquid or aqueous phase (except for substances which are acidic or basic like sulphuric acid)
- Soft and weak
- Low to moderate melting points
92
Why are covalent molecular substances non-conductive? (save for a few acids or bases)
The electrons in the substance are localised between the atoms as they share the electrons. None of these electrons are free to move so they cannot conduct electricity.
93
Why are covalent molecular substances soft and weak?
The intermolecular forces between the molecules in the substance are weak as the intermolecular forces are weak compared to metals and ionic compounds.
94
Why is the melting point of covalent molecular substances low or moderate?
There are weak intermolecular forces that are holding a covalent molecular substances together and they don't require much energy to separate, thus the melting point is low.
95
What is covalent network bonding?
Covalent network bonding is when atoms, usually non-metals, bond chemically to neighbouring atoms, resulting in a massive array forming a huge network that fulfils its individual atom's shells covalently.
96
What are the properties of a covalent network substance? (list them + exceptions, if applicable)
- Non-conductors of electricity (graphite is an exception)
- Very hard and brittle (graphite is an exception)
- Very high melting and boiling points
97
Why are covalent network substances not conductors of heat and electricity? (if there's an exception, explain what and why)
As the atoms are bonded covalently, the electrons are held/localised and cannot freely move due to the attraction between the electrons and the positive nuclei of the atoms, which means that electricity cannot be conducted.
Graphite is an exception as it is made up of sheets of graphene, the graphene has some electrons which are delocalised and not being shared. The sheets are held together weakly which allows the electrons to move across the sheets and conduct electricity.
98
Why are covalent network substances very hard but brittle? (if there's an exception explain what it is and why)
CT Network substances are hard due to the incredibly strong covalent bonds between the atoms. As these bonds are everywhere, it creates a strong array so it is difficult to scratch the substance.
It is brittle because when sufficient force is applied, due to the bonding being directional, the atoms are unable to bend and the bonds break, shattering the array and the substance.
An exception to all this is graphite, which is made up of 2-d sheets. These sheets have weak intermolecular forces and thus can slip over each other pretty easily. This causes graphite to be slippery.
99
What is the difference between intramolecular bonding and intermolecular forces?
IM Bonding is the type of bonding within the molecule, whether it is ionic, metallic or covalent. Intermolecular forces are the interactions between molecules.
100
Name 3 covalent network substances
Diamond, graphite and silicon dioxide
101
What are the 3 intermolecular forces
- Dipole - dipole forces
- Hydrogen bonding
- Dispersion forces
102
List the strength of intermolecular forces
Hydrogen bonds > Dipole- dipole > Dispersion forces
103
What happens to boiling point as IMFs increase?
Boiling point increases
104
What happens to vapour pressure as IMFs increase?
Vapour pressure decreases
105
How does the length of a molecule affect the dispersion forces?
Larger molecules will have stronger dispersion forces as they have more electrons which increases the chances for a temporary dipole to occur.
106
What are electrolytes?
Solutes that produce ions when dissolved such as NaCl
107
List the unique properties of water
- It has a surprisingly high boiling point despite a small molecular size
- Water is denser in liquid form than solid form
- Water has the highest surface tension known
108
Why is water denser in solid form than liquid form?
Water in liquid form has hydrogen bonding, so they molecules are close together due to the strong attraction between the hydrogens and oxygens. However, in the solid form, each water molecule bonds with 4 water molecules creating a lattice.
This leads to the hydrogen bonding in the solid phase being maximised but more space is taken up, decreasing the mass to volume ratio and thus being less dense overall.
109
Why does water have high surface tension?
Water has strong IMFs from the hydrogen bonding combined with the dispersion forces. At the surface of water there is an imbalance of forces, causing the surface molecules to be pulled inwards towards the majority of the liquid
This leads to the surface trying to contract and achieve a minimum area
110
Why are water droplets spherical?
As water has a high surface tension, the molecules near the outside try to come inwards, which causes the droplet to contract. This causes the water to want to have the minimum amount of surface area so by becoming a sphere, the surface area to volume ratio is diminished.
111
What is a solution?
A homogenous mixture of a solute dissolved in a solvent
112
What is meant by the solubility of a solute?
A mass of the solute that can dissolve in 100 g of water
113
What happens to solubility as temperature increases?
Generally, solubility increases but if the solute is a gas, the solubility decreases
114
What is meant when a solution is unsaturated?
More solute can be dissolved in the solvent
115
What is meant when we say a solution is saturated?
A solute can no longer dissolve in a solvent
116
How can a solution become supersaturated?
A solution becomes supersaturated when more solute dissolves in the solvent than is expected at a certain temperature
117
Acid - Metal Reaction
Acid + Metal --> Salt + Hydrogen gas
118
Acid - Metal Hydroxide Reaction
Acid + Metal Hydroxide --> Salt + Water
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Acid - Metal Oxide Reaction
Acid + Metal Oxide --> Salt + Water
120
Acid - Carbonate Reaction
Acid + Carbonate --> Salt + Water + Carbon dioxide gas
121
Acid - Hydrogen carbonate reaction
Acid + Hydrogencarbonate --> Salt + Water + Carbon dioxide gas
122
Acid - Metal sulfite reaction
Acid + Metal sulfite --> Salt + water + sulphur dioxide gas
123
Base - Ammonium Salt reaction
Base + Ammonium Salt --> Salt + Water + Ammonia (NH3) gas
124
Base - Non-metal oxide
Base + non-metal oxide --> Salt + Water
125
Acid - Base reaction
Acid + Base --> Salt + Water
126
Which metals don't react visibly with acids?
Copper, silver, gold and platinum
