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Chemistry > Module 1 Yr 11 > Flashcards

Module 1 Yr 11 Flashcards

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1
Q

Define colloid

A

Colloid: homogenous mixture where fine particles of insoluble substance are dispersed and suspended in another substance such as milk (dispersed fat molecules in water)

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2
Q

Define suspension

A

Suspension; heterogenous mixture with large solute particles that don’t dissolve in the solvent. Particles are dispersed temporarily then settle to the bottom (sand in water as sand is suspended in water).

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3
Q

Define physical properties

A

Physical properties: can be observed without changing chemical identities or the composition of the matter

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4
Q

Define chemical properties

A

Chemical properties: When a substance undertakes a chemical reaction (chemical bonds broken or formed) to form new stable compounds

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5
Q

Define reactivity

A

Reactivity: Ease of element to achieve noble gas configuration

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6
Q

Define isotope

A

Isotope: Variants of a nuclide consisting of different no of neutrons but with same no of protons

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7
Q

Define mixture

A

Substance of variable composition, containing two or more elements or compounds that are physically combined. Are impure substances.

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8
Q

How to test purity of mixture

A

Purity can be determined from m.p with pure substances melting at a single temperature (or small temp range) and impure substances melting over a much larger range.

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9
Q

Define homogenous and heterogenous mixture
Provide examples

A

Consistent proportions of constituent elements or compounds throughout the mixture
Inconsistencies in the proportions of constituent elements or compounds throughout the mixture
Homogenous: Homogenised milk, saltwater, air, steel
Heterogenous: Sand in water, blood, cereal in milk, mixed nuts

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10
Q

Name separation techniques

A

Evaporation to dryness, sieving, ferromagnetism, froth flotation, distillation, fractional distillation, filtration, sedimentation and decantation, centrifugation, separating funnel, liquefaction and fractional distillation

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11
Q

Describe evaporation to dryness

A

Process occurs at temperatures that facilitate evaporation of solvent to ensure solid is left to crystallise. Often used for separating solutions where solid solute is dissolved in liquid.

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12
Q

What is sedimentation and decantation

A

Large solid particles with high density often settle to the bottom when mixed with liquid. Process occurs over periods of time (sedimentation). Above sediment, liquid is poured off the top without disturbing sediment (decantation)

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13
Q

What is centrifugation

A

Mixture is spun at high speeds causing denser particles to settle at base of container. Often done with colloids in which insoluble particles are too usually fine to settle naturally.

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14
Q

Define distillation

A

The liquid is simply passed from a container under high atmospheric pressure to one under lower pressure. The reduced pressure causes the liquid to vaporize rapidly; the resulting vapour is then condensed into distillate.

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15
Q

Define fractional distillation

A

Hence, a fractionating column is attached to the reaction vessel and the process becomes fractional distillation. The column consists of glass beads, whose purpose is to disrupt the convection currents formed by the heat vapours, thus creating a smooth temperature gradient up the column where the top is cooler than the bottom. The beads also serve as a surface for the vapour to condense and then evaporate, rendering the substance with the higher boiling point unable to pass through the column as it requires greater energy to do so. Accordingly, the substance with the lower boiling point passes through the apparatus as distillate while the other substance remains as residue.

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16
Q

Describe separating funnel process

A

Less dense liquid will settle on top of denser liquid. After removing the lid, stopcock is turned gradually to allow for the bottom liquid to drain off.

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17
Q

What is Liquefaction and Fractional Distillation

A

Gases are initially liquefied by cooling them below their boiling point. Liquid is separated by fractional distillation. In case of air, nitrogen vaporises first followed by argon and oxygen respectively.

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18
Q

What is froth flotation

A

Used in mining industry. When extracting copper ore, the ore is initially combined with gangue. Detergent is added to the mixture to reduce surface tension with air then bubbled through. This causes light copper ore grains to rise to the top to be removed

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19
Q

Difference between miscible and immiscible liquids

A

Liquids which mix together in all proportions and form a single layer are called miscible liquids. Liquids which do not mix with each other and form separate layers are called immiscible liquids.

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20
Q

How are ionic compounds formed?

A

Formed due to electrostatic attraction between cation and anion. This electrostatic attraction is known as ionic bond. Cation is usually metal ion which loses one or more electrons. Anion is usually non-metal which gains electrons. Electrons lost by cation are transferred to anion forming an ionic compound.

Ionic compounds are named by first using the name of the cation and then anion.

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21
Q

How are covalent compounds formed?

A

Covalent compounds are formed when two or more elements, usually non-metals, share electron pairs in one or more covalent bonds

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22
Q

State prefixes from 1-10

A

mono: 1
di: 2
tri: 3
tetra: 4
penta: 5
hexa: 6
hepta: 7
octa: 8
nona: 9
deca: 10

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23
Q

Define polyatomic ions and balmer series

A

Ions that consist of more than one atom covalently bonded together.
Balmer series: visible lines of hydrogen spectrum

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24
Q

Name of chemical formula and charge of polyatomic ions
Nitrate
Hydroxide
Hydrogen carbonate/bicarbonate
Carbonate
Sulfate
Phosphate
Ammonium

A

Chemical Formula Charge
NO3^- -1
OH^- -1
HCO3^- -1
CO3^2- -2
SO4^2- -2
PO4^3- -3
NH4^+ 1

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25
Q

Trends in periodic table

A

Groups I-III, reactivity increases downwards
Groups V-VII, reactivity decreases downwards
Most reactive elements: (Bottom Left (Caesium & Francium)) & (Top right (Fluorine and Chlorine))
Bigger the atom, more reactive due to larger atomic radius
Rules don’t apply to transition metals due to them being semimetals

Across period, electronic configuration of atom contains one more valence electron than previous. Down group, one more principle quantum shell is added to electronic configuration.

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26
Q

Factors of isotope stability

A

Atomic number > 83

Neutron to proton ratio is not within the zone of stability

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27
Q

Analysis of zone of stability graph

A

Stable neutron to proton ratio is approximately
1:1 for elements 1-20
1.3:1 for elements 21-50
1.5:1 for elements 51-83
B- decay is above band of stable nuclei
B+ decay is below band of stable nuclei
Alpha decay is above zone of stability

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28
Q

Define alpha decay

A

Alpha decay: the release of alpha particles (helium nuclei) from nucleus

Helium nuclei has 4 mass number and atomic number of 2

Likely to be emitted from nucleus when too large (atomic number > 83)

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29
Q

Difference between beta+ and beta-

A

Occurs when neutron to proton ratio is too low
Proton changes to neutron and positron where latter is emitted
nucleus has 0 mass no and -1 atomic no
Whilst:
Neutron in the nucleus changes into a proton and an electron where latter is eliminated
Process occurs when neutron to proton ratio is too high
nucleus has 0 mass no and 1 atomic no

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30
Q

Define gamma decay

A

The emission of gamma rays.
Occurs when nucleus is in an excited state and gamma ray release brings it back to stability

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31
Q

Define ionising power

A

Ionising power: the ability to remove electrons from an atom producing an ion. The ability of radiation to harm humans is a result of ionising power.

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32
Q

Penetration power vs ionising power and charge of types of radiation

A

Alpha (+2): Low pp, high ip
Beta- (-1): Medium pp and ip
Gamma (0): High pp and low ip

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33
Q

Difference between nuclear fission and fusion

A

Fusion: process when two small nuclei combine forming a larger nucleus (more expensive due to large energy need to power it)

Fission: process when one large nucleus splits to form two or smaller nuclei

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34
Q

What characterises an orbital

A

Principle quantum number: Are orbitals in every principle electron shell. Orbitals are described by which principle level it occupies

Shape: Orbitals can take the shape of a sphere (s orbital), hourglass (p orbital), cloverleaf (d orbital) and even more complex (f orbital)

Orientation: point in different directions

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35
Q

What is heisenburg’s principle?

A

Heisenburg’s uncertainty principle states one cannot determine with certainty the exact position and velocity of an electron

36
Q

What is Pauli’s exclusion principle

A

States that no two bound electrons within one atom can have the same set of four quantum numbers.

37
Q

What is Aufbau’s principle?

A

electrons fill lower-energy atomic orbitals before filling higher-energy ones

38
Q

What is hund’s rule

A

Every orbital in a sublevel is singly occupied before any orbital is doubly occupied.

39
Q

Characteristic of excited species

A

If a species is excited by quanta (packets of energy of specific size), electrons can jump from lower energy level to higher energy level. An excited species can violate the aufbau principle and hund’s rule but must obey Pauli’s exclusion principle

40
Q

Purpose of flame tests

A

Used to identify presence of certain cations since a distinctive colour is produced when an ionic solution containing that cation is placed into a flame. When solution is placed in flame, metal cations are atomised.

Visible light created when electrons transition to higher energy level due to heating, electrons lose energy and relaxes to original energy level, em radiation in form of visible light is released

Not all metals have a flame test colour in visible range. Colour produced depends on the difference in energy between energy levels

41
Q

Colours produced by metals

A

Lithium: Red
Sodium: Yellow-orange
Potassium: Lilac
Calcium: Brick-red
Barium: Apple green
Strontium: Red
Copper: Blue-green

42
Q

Differences between type of atomic spectra

A

Absorption spectrum:
Produced when white light passes through cool gaseous sample of an element with element absorbing specific wavelengths of that light

Emission spectrum:
Produced when a gaseous sample of an element is heated. Causes excitation & de-excitation of electrons, with de-excitation releasing energy in the form of photons. Correspond to specific wavelengths of light (spectral lines).

43
Q

Characteristics of Bohr’s model

A

Electrons are in stationary states, don’t emit em radiation. Electrons exist at discrete energy levels with transitions between energy levels caused by absorption or release in energy. Absorption spectrum of hydrogen can be explained in terms of electron between specific energy levels.

44
Q

Characteristics of Schrodinger’s model (spdf model)

A

Proposed electrons behave like waves and particles. Electrons don’t orbit, energy and mass is spread out over region (orbital)

45
Q

Define Zeff

A

Effective Nuclear Charge: The net positive charge experienced by electrons.
=to the charge of the nucleus subtracted by no of shielding electrons.

46
Q

What influences ENC

A

Atomic radius, IOE1 and electronegativity

Zeff increases across period due to charge of nucleus increasing due to increased no of protons whilst the no of shielding inner electrons remain constant due to the no of shells remaining the same

Zeff increases down the group as even though the no of shielding inner electrons increases, the effect of shielding diminishes with increased electron shells. Whilst this occurs, the charge increases due to the increased no of protons in the nucleus.

47
Q

Define atomic radius

A

Total distance from atomic nucleus to the outermost occupied electron orbital.
Down group, Ar increases due to increased atomic number meaning there are more protons in the nucleus and the zeff on all electrons including valence electrons have increased. Electrons are drawn together to the nucleus due to a stronger electrostatic force causing the atom size to shrink.

Across period, Ar decreases
The zeff on valence electrons increases as down the group, the valence electrons occupy higher energy energy shells due to an increased principal quantum number consequently causing an increased distance away from nucleus

47
Q

Define IE1

A

First ionisation energy: energy required to remove an outer valence electron from an element in its gaseous state. Prior to measurement, element must be vaporised.
It increases across period due to the number of protons increasing causing a greater zeff making it more strongly bound. Thus more energy is required to remove the electron.
Decreases down the group due to increased atomic radius as the highest energy valence electron is further away from nucleus causing less energy to be put to remove the electron.

48
Q

Define electronegativity

Describe when it is high or low

A

A measure of tendency of an atom in a compound to attract shared electrons towards itself.
Elements on left side have low electronegativity as they want to lose electrons to attain noble gas configuration.
Elements on right side have high electronegativity as they want to gain electrons to achieve stability
Across period, it increases as increased no of protons in nucleus increases Zeff. A stronger zeff results in a greater electrostatic attraction with electrons increasing element’s electronegativity.
Decreases down group as valence electrons are further away from nucleus. Effect of increased radius is much greater than the increase in ENC decreasing element electronegativity.

Electronegativity increases to top right of periodic table
Fluorine most electronegative whilst francium is least

49
Q

Zeff vs principle quantum number

A

For trends concerning behaviour of valence electrons
Across period, effect of zeff dominates has principle quantum number is unchanged
Down group, effect of principle quantum no dominates over zeff

50
Q

What increases bond polarity

A

Greater difference in electronegativity of two bonding atoms, the more polar the bond is

51
Q

Allocation of each type of bond in terms of electronegativity

A

> 1.8: Large difference in electronegativity (Ionic Bond)
0.4-1.8: Moderate difference in electronegativity (Polar Covalent)
<0.4 Small or no difference in electronegativity (Non-polar)

52
Q

Formation of ionic, polar covalent, non-polar explanation in terms of electronegativity

A

When atoms with large difference in electronegativity participate in bonding, electron transfer occurs leading to one atom becoming positive and the other negative

When atoms have a moderate difference, electrons are shared unequally drawing electrons closer to the atom with higher electronegativity. Leads to formation of electric dipole where more electronegative element is slightly negative and the other positive.

When atoms have a small difference, electrons are shared equally and no polarity develops

53
Q

Description of valence electrons in ionic compounds

A

In ionic compounds, all atoms have full valence electrons achieved through transfer of valence electrons from one species to another

54
Q

Description of valence electrons in covalent compounds

A

In covalent compounds, all atoms similarly have a full valence shel of electrons which is attained through sharing valence shell electrons between two atoms.

55
Q

Difference between empirical and molecular formula

A

Ionic compounds are defined by empirical formulas whilst covalent compounds are defined by their molecular formula. This specifies the actual no of atoms present in a molecule in a molecule of the compound.

56
Q

What do lewis structures show in ionic vs covalent?

A

The electrostatic attraction between positively and negatively charged components which arises due to transfer of electrons

The bonds between atoms as pairs of electrons shared between atoms.

57
Q

Shape of ionic and covalent compounds

A

Exist as lattice structures
Covalent compounds exist in a variety of different shapes

58
Q

What does the VSEPR Model state

A

Only valence electrons (typically s and p orbital electrons) are involved in determining molecular shape.
Main idea is that the structure around the given atom is determined by minimising electron pair repulsion. Bonding groups and lone pairs around given atom should be as far away as possible.

59
Q

Assumptions of VSEPR model

A

Atoms in species are held together by pairs of electrons (bonding pairs)
Some atoms within species may have pairs of electrons that are not involved in bonding (lone pairs)
As electron pairs are negatively charged, they repel each other. On each atom, the electron pairs adopt positions as far away from one another as possible.

60
Q

What is the electron group geometry (linear, trigonal planar) and molecular geometry determined by?

A

The electron group geometry is determined by all the electron groups in the molecule
The molecular geometry is the shape governed by atoms in the molecule and is determined by the bonding groups

61
Q

Names of molecular geometries with:
2 BG + 0 LP
3 BG + 0 LP
2 BG + 1 LP
4 BG + 0 LP
3 BG + 1 LP
2 BG + 2 LP
5 BG + 0 LP
6 BG + 0 LP

A

Linear
Trigonal planar
Bent
Tetrahedral
Trigonal pyramid
Bent
Trigonal Bipyramidal
Octahedral

62
Q

Define ionic networks

A

Ionic networks consist of cations and anions bonded by electrostatic attraction (positives and negatives attracted to each other)

63
Q

Ionic networks characteristics

A

Hard and brittle (when crystal is hammered, layers of ions slide over each other causing like ions to come into close proximity, repelling each other causing it to shatter)
High mp and bp (strong ionic bonds must be broken to melt ionic network requiring lots of heat energy)
Cannot conduct electricity in solid state
Can conduct in molten or aqueous state (as the ions are liberated from lattice and are free to conduct electricity)

64
Q

What are covalent networks create from?

A

Covalent network are formed by extensive networks or chains of atoms bonded together by strong covalent bonds

65
Q

Covalent network characteristics

A

No molecules as it is a continuous crystal such as diamon, sand (silicon dioxide) and graphite
Are hard (presence of very strong covalent bonds throughout lattice)
Are brittle (covalent bonds are directional and will break)
High mp and bp (Strong covalent bonds must be broken to melt covalent network which requires lots of heat energy)
Most don’t conduct electricity
Don’t dissolve in water and don’t melt rather decomposing at high temperatures
Hence, most covalent networks don’t conduct electricity
Graphite is exception

66
Q

Define metals

A

Metals are composed of 3D lattice of metal cations embedded in a sea of delocalised valence electrons

67
Q

Characteristics of metallic structure

A

Cations are arranged in a regular 3D lattice whilst valence electrons freely occupy space around them, offering multi-directional metallic bonding. This provides a strong cohesive force that holds the metals cations together.
Metals are hard but softer than covalent networks or ionic networks
Metals are malleable (can be hammered into sheets)
Metals are ductile (can be drawn into wires)
When hammered, metal cation lattice structure is disturbed but delocalised valence electrons cushion this deformation since they provide multi-directional bonding
High mp and bp (electrostatic interactions between valence electrons and metal cations must be broken which requires a lot of energy)
Metals conduct electricity in solid or molten state
Metals are lustrous (delocalised electrons reflect light)

68
Q

Define covalent molecular structure

A

Covalent molecular structures are composed of discrete covalent molecules with no chemical bonds between them. An example of this is ice.

69
Q

Covalent molecular structure characteristics

A

Are held together not by chemical bonds but intermolecular forces/bonds
Intermolecular forces/bonds are weaker than chemical bonds and act between molecules.
Are distinct from intramolecular bonds which are covalent bonds
Covalent molecular structures are soft (intermolecular bonds are weaker than ionic or covalent bonds)
Have low mp and bp (weaker intermolecular bonds between molecules need to be broken to melt constant molecular structure which doesn’t need a lot of energy.
Don’t conduct electricity

70
Q

Define allotrope

A

Allotropes are different structural forms of the same element in the same state. They exhibit different physical properties

71
Q

what are allotropes of carbon

A

Diamond, buckminsterfullerene, graphite

72
Q

Characteristics of diamond

A

Has covalent network structure
Hard
Every carbon atom is bonded to its neighbours by four strong covalent bonds arranged in tetrahedral geometry.
Doesn’t conduct electricity
High mp and bp (in order to melt or boil diamond, strong covalent bonds need to be broken which requires immense energy)
Colourless, transparent and lustrous

73
Q

Characteristics of Graphite

A

Composed of layers of covalent networks
Within each layer, each carbon atom is bonded to only three other carbons forming trigonal planar geometry
This allows one valence electron per carbon atom to be delocalised between layers, creating weak interactions that hold the layers together
Soft and slippery (layers are weakly held together consequently they can easily slip over one another
Reasonable electrical conductor
Very high mp and bp (strong covalent bonds need to be broken)

74
Q

Characteristics of Buckminsterfullerene

A

A spherical cage-like structure of carbon atoms with chemical formula of C60
Has covalent molecular structure
Poor electrical conductor
Low mp and bp
Weak intermolecular forces need to be broken which don’t require as much energy to break.

75
Q

How is polarity of covalent molecules dictated

A

A covalent molecule is polar if it possesses a net dipole moment
Is dependent on polarity of bonds and molecular geometry
If there exists no cancellations, then the molecular is polar
If none of the bonds or all the dipoles, cancel out, then it is non-polar

76
Q

Describe Dipole-dipole interactions

A

Occur between two polar molecules. Polar molecules have a net dipole. Molecules arrange themselves such as opposite charges attract each other. These electrostatic interaction are named dipole-dipole interactions
Are considerably weaker than ionic or covalent bonds
Strength depends on size of overall dipole moment.
Elements with large electronegativity differences tend to have larger dipole moments
Greater dipole moment, stronger the dipole-dipole interactions

77
Q

Hydrogen bonding characteristics

A

Special type of strong dipole-dipole interactions
Only occurs when hydrogen is directly bonded to highly electronegative elements: F, O or N (Large difference in electronegativity results in a greater bond polarity. Small size of H atom allows the close approach of these dipoles)
H-bonding is much stronger than dipole-dipole interactions but weaker than ionic or covalent bonds

78
Q

Dispersion forces characteristics

A

They exist between any molecules regardless of polarity
This is the primary intermolecular force that exists within non-polar molecules
Strength of dispersion forces increases with molecular mass that is larger molecules have larger dispersion forces than smaller molecules
Larger molecules have greater no of electrons and have increased chance of forming momentary dipoles
Are weaker than dipole-dipole and h-bonding.

79
Q

Cause of dispersion force

A

Since electrons move around randomly, a momentary non-uniform distribution of charge causes the formation of a temporary dipole
The instantaneous dipole that forms in one atom/molecule can then affect electron distribution in neighbouring atom/molecule and induce similar dipole
Leads to an intermolecular attraction that is relatively weak and short lived aka dispersion force

80
Q

What is broken when melting or boiling covalent molecules?

A

Intermolecular forces are broken, the stronger the intermolecular forces, the more energy required to break them thus having a higher mp and bp

81
Q

Define intramolecular bonds
Strength
What is boiled

A

Bonds between atoms within molecule or compound
Examples include covalent, ionic or metallic.
Are strong
When boiled, covalent networks, ionic networks and metals are overcome

82
Q

Define intermolecular bonds
Strength
What is overcome once boiled

A

Bonds/forces between molecules
Examples include H-bonding>dipole-dipole>dispersion forces
Are weak
Covalent molecules are overcome when boiled

83
Q

Define reliability, accuracy and validity

A

The extent to which the findings of repeated experiments, conducted under identical or similar conditions, are consistent with each other.
The extent to which a measured value agrees with its true value
Validity refers to whether all the variables within an experiment, apart from the independent and dependent, were controlled.

84
Q

Where does the independent and dependent variable go

A

On graph, independent lies on x axis and dependent lies on y axis
Independent variable must be either first column or row and the
dependent variable can go in the rest of the column or rows.

85
Q

How to calculate error/percent error

A

Error: Measured value - true value
Error percent: |(Measured value - true value)|/true value x 100

Chemistry (8 decks)

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