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1
Q

Key terms: Atom, Element, Ion, Molecule, Compound, Empirical formula and Molecular formula

A

-atom: is the smallest part of an element that can exist, all substances are made up of atoms
-element: a substance made up of only one type of atom, all with the same number of protons
-ion: is what’s formed when an atom either gains or loses electrons, this gives it an overall charge + a positive if it has lost at east one electron and a negative charge if it has gained at least one electron
-molecule: consists of two or more atoms that have been bonded together chemically
-compound: is a substance that combines two or more different elements through the formation of chemical bonds
-empirical formula: smallest whole number ratio of atoms of each element in a compound
-molecular formula: actual number of atoms of each element in a compound

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2
Q

Mole and the Avogrado Constant:

A

-> the mole is a unit of measurement for substances, it always contains the same number of particles
L = 6.02 x 10^23 particles
-this number is the Avogrado Constant (L) and allows the number of particles present in a sample of a substance with known mass to be found:
Number of particles = n (number of moles) x L (Avogrado Constant)
-the mole is a very important unit of measurement in many calculations:
Moles = mass/Mr = concentration (dm^3) x volume (cm^3) /1000

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3
Q

Balanced Equations:

A

-full or ionic chemical equations must be balanced before they can be used in calculations
-this is because the reaction ratios must be correct
-for a chemical equation to be balanced, it must have the same number and type of each atom present on both sides of the equation
-it can be useful to also include state symbols so it is clear what might be observed during the reaction, for example, bubble of gas, a precipitate forming, or a colour change that may infer a displacement reaction
-there are four state symbols:
-(s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water)
-these balanced equations can then be used to calculate reacting masses, percentage yield and atom economy

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4
Q

Relative Atomic Mass:

A

-> the mean mass of an atom of an element, divided by 1/12 of the mass of an atom of the carbon-12 isotope

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5
Q

Relative Molecular Mass and Relative Formula Mass:

A

-> the mean mass of a molecule of a compound, divided by 1/12 of the mass of an atom of the carbon-12 isotope
-the molecular mass can be calculated by adding together the atomic masses (Ar) of all the atoms in that compound
-example: To calculate the Mr of the compound C2H5OH the Ar’s must be used:
C = 12, O = 16, H = 1
C2H5OH = (2x12) + (6x1) + (16x1) Mr = 46
-relative formula mass refers to compounds that have a giant structure

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6
Q

Empirical and Molecular formula:

A

True Mr = Mr of empirical formula x multiplier
-example: The empirical formula of a molecule containing 5 atoms of oxygen for every 2 atoms of phosphorus has an Mr of 284. What is its molecular formula?
Empirical formula = P2O5
Mr of empirical formula = (31x2) + (16x5) = 142
Multiplier = 284 divided by 142 = 2
Molecular formula = 2(P2O5) = P4O10

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7
Q

Molar Mass:

A

-> the molar mass of a substance is its mass in grams per mole and has the units g mol^-1. It can be calculated using the following equation:
Molas mass = mass/number of moles

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8
Q

Parts Per Million (ppm):

A

-concentration can be given in parts per million (ppm)
-this gives the units of mass of that particular species within 1,000,000 total units of mass
-it is most commonly used to represent the concentrations of gases

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9
Q

Concentration calculations:

A

-the concentration of a solution can be measured in mol dm^-3 and g dm^-3 which can be calculated using the following equations:
-Concentration mol dm^-3 = Number of moles (mol) / volume (dm^3)
-Concentration (g dm^3) = Mass (g) / volume (dm^3)

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10
Q

Experimental Data:

A

-can be used to work out empirical and molecular formulas and reaction stoichiometries
-these calculations require the use of equations given in this section, along with some others, these include:
-Mol = concentration x volume
-True Mr = Mr of empirical formula x multiplier
-Volume of gas (dm^3) = 24 x number of moles (at room temperature and pressure)
-Number of particles = n x L
-Mass = Mr x mol

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11
Q

Volume of gases: Molar volume of gases

A

-one mole of any gas at room temperature and pressure will take up the same volume, regardless of its composition
-this volume is 24,000 cm^3, or 24 dm^3, and is known as the molar volume of gases
-this relationship gives the following equation that can be used to work out the volume of a gas if its amount (number of moles) is known and vice versa
Volume of gas (dm^3) = 24 x number of moles (at room temperature and pressure)

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12
Q

The Ideal Gas Law:

A

-when under standard conditions, gases and volatile liquids follow certain trends:
-pressure is proportional to temperature
-volume is proportional to temperature
-pressure and volume are inversely proportional
-these relationships can be combined to give the ideal gas equation:
pV = nRT = mRT/Mr
-in order to use this equation, the variables must be in the correct standard units:
p = pressure in pascals
V = volume in m^3
T = temperature in Kelvin
n = moles
m = mass in grams
R is the ideal gas constant, equal to 8.31 JK^-1 mol^-1

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13
Q

Percentage Yield:

A

-> indicates how much of the maximum amount of product you obtained during an experiment
-a low percentage yield could indicate an incomplete reaction, or the loss of product during purification
% yield = experimental mass/theoretical mass x 100

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14
Q

Atom Economy:

A

-> is a measure of efficiency since it measures the proportion of reactant atoms which are converted into the desired product
% atom economy = Mr of desired product/total Mr of all products x 100
-in industrial chemical processes, it is desirable to have a high atom economy for a reaction
-this means there is little or no waste product, only the desired product
-therefore it means the process is more economically viable for industrial-scale manufacture

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15
Q

Displacement reactions:

A

Br2 (aq) + 2Kl (aq) -> I2 (aq) + 2KBr (aq)
-in this reaction, the more reactive bromine displaces the less reactive iodide in potassium iodide
-this can also be seen in the ionic equation for the reaction
Br2 (aq) + 2I^- (aq) -> I2 (aq) + 2Br^-(aq)

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16
Q

Neutralisation reactions:

A

-these can be identified by the presence of reactant acids and bases as well as the formation of a neutral salt solution and water (and sometimes other compounds such as carbon dioxide)
HCl (aq) + NaOH3 (aq) -> NaCl (aq) + H2O (l)
Na2CO3(aq) + 2HNO3 (aq) -> 2NaOH3 (aq) + H2O (l) + CO2 (g)
-the ionic equations can more clearly demonstrate the neutralisation of an acid and a base:
H^+ (aq) + OH^- (aq) -> H2O (l)
2H+ (aq) + CO3 ^2- (aq) -> H2O (l) + CO2 (g)

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17
Q

Precipitation Reactions:

A

-there are shown by the reaction of two aqueous solutions to form products which include one solid
BaCl2 (aq) + Na2SO4 (aq) -> BaSO4 (s) + 2NaCl (aq)
-the ionic equation shows the precipitation reactions more clearly as there are no other products considered
Ba^2+ (aq) + SO4^2- (aq) -> BaSO4 (s)

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18
Q

Fundamental particles:

A

-the atomic structure model has evolved over time, as knowledge and scientific understanding changes
-the current, accepted model of the atom consists of a small, dense central nucleus surrounded by orbiting electrons in electron shells
-this was discovered in the Rutherford scattering experiment in 1911
-the nucleus consists of protons and neutrons giving it an overall positive charge
-it contains almost the entire mass of the atom
-in a neutral atom, the number of electrons is equal to the number of protons due to the relative charges

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19
Q

Protons, Neutron and Electrons:

A

-relative charge: proton -> +1, neutron -> 0, electron -> -1
-relative mass: proton -> 1, neutron -> 1, electron -> 1/1840
-the maximum number of orbiting electrons that can be held by any single shell, depends on the number of the shell, this can be calculated using 2n^2 where n is the number of the shell
-example: Electrons in shell 2 = 2(2^2) = 8 electrons
-each electron shell must fill before the next one can hold any electrons

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20
Q

Atomic Number and Mass Number:

A

-mass number is represented using A and can be calculated as the sum of protons and neutrons in an atom
-atomic number is represented using Z and is equal to the number of protons in an atom, hence it can be referred to as proton number

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21
Q

Isotopes:

A

-> isotopes are atoms of the same element with the same atomic number, but with a different number of neutrons, resulting in a different mass number
-neutral atoms of isotopes will react chemically in the same way because their proton number and electron configuration is the same
-the sharing and transfer of electrons is unaffected
-however, the different mass numbers means they have different physical properties
-example: the following are all isotope of hydrogen:
Hydrogen = 1 proton and 1 neutron
Deuterium = 1 proton and 2 neutrons
Tritium = 1 proton and 3 neutrons

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22
Q

Relative Masses:

A

-Relative atomic mass (Ar) is defined as:
-> the mean mass of an atom of an element, relative to 1/12 of the mean mass of an atom of the carbon-12 isotope
-this takes the relative abundances of the different isotopes of an element into account
Ar = means mass of an atom of an element/ 1/12 x mean mass of C-12 isotope

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23
Q

Relative isotopic mass:

A

-> the isotopic mass of an isotope relative to 1/12 of the mean mass of an atom of the carbon-12 isotope

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24
Q

Relative molecular mass (Mr):

A

-> the mean mass of a molecule of a compound, relative to 1/12 of the mean mass of an atom of the carbon-12 isotope
-it can be calculated for a molecule by adding together the separate Ar values of the component elements
-example: Mr of H2O = (2x1.0) + (1x16.0) = 18.0
-calculating the relative molecular mass can be used to identify the molecules present in a sample
-relative formula mass is similar to Mr but is used for compounds with giant structures

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25
Q

Ions and Mass Spectrometry:

A

-ions are formed when an atom loses or gains electrons meaning it is no longer neutral and will have an overall charge
-they are very useful in the analytical technique of mass spectrometry
-it is used to identify different isotopes and find the overall relative atomic mass of an element

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26
Q

Time of Flight (TOF): Mass Spectrometry

A

-this form of mass spectrometry records the time it takes for ions of each isotope to reach a detector, using this spectra can be produces showing each isotope present
1. Ionisation - a sample of an element is vaporised and injected into the mass spectrometer where a high voltage is passed over the chamber
-this causes electrons to be removed from the atoms (they are ionised) leaving +1 charged ions in the chamber
2. Acceleration- these positively charged ions are then accelerated towards a negatively charged detection plate
3. Ion drift - the ions are then deflected by a magnetic field into a curved path
-the radius of the path is dependent on the charge and mass of the ion
4. Detection - when the positive ions hit the negatively charged detection plate, they gain an electron, producing a flow of charge
-the greater the current produced, the greater the abundance of that particular ion
5. Analysis - these current values are then used in combination with the flight times to produce a spectra print-out with the relative abundance of each isotope displayed
-during the ionisation process, a 2+ charged ion may be produces which means it will be more affected more by the magnetic field producing a curved path of smaller radius
-as a result, this mass to charge ratio (m/z) is halved which can be seen on spectra as a trace at half the expected m/z value

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27
Q

Ar and Mass Spectrometry:

A

-using this print-out spectra, the Ar (relative atomic mass) can be calculated my multiplying each m/z value by its abundance and adding each of these together, before dividing by the total abundance of all species present
Ar = (m/z x abundance)/total abundance
-using this calculated value of Ar, the element can be identified by referring to the Periodic Table

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28
Q

Predicting Mass Spectra:

A

-if you know the abundance of an isotope you can generate the mass spectra for its molecules, including relative speak heights
-example: The relative abundance of ^35Cl atoms is 75% and ^37Cl atoms is 25%. In other words, for every 100 atoms of chlorine, 25 would be ^35Cl and 75 would be ^37Cl. Spectra produced by the mass spectrometry of chlorine display a characteristic pattern in a 3:1 ratio for Cl^+ ions and a 3:6:9 ratio for Cl2^+ ions. This is because one isotope is more common than the other and the chlorine molecule can form in different combinations
^70CL2^+ = 35 + 35
^72Cl2^+ = 35 + 37 OR 37 + 35
^74Cl2^+ = 37 + 37

29
Q

Ionisation Energy: first ionisation

A

-> the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state, it is measured in kJmol^-1
Na (g) -> Na^+ (g) + e^-
-successive ionisation energies occur when further electrons are removed
-this usually requires more energy because, as electrons are removed, the electrostatic force of attraction between the positive nucleus and the negative outer electron increases
-more energy is therefore needed to overcome this attraction, causing ionisation energy to increase
-all ionisation energies are endothermic, as the removal of electrons requires an energy input

30
Q

Ionisation energy: second ionisation

A

-> the minimum energy required to remove one mole of electrons from one mole of +1 ions in a gaseous state, it is measured in kJmol^-1
Na^+ (g) -> Na^2+(g) + e^-
-the definition can be used to define successive ionisation energies
-ionisation energy is influenced by 3 factors:
1. The number of protons
2. The electron shielding
3. The sub shell from which the electron is removed

31
Q

First ionisation energy: trends

A

-first ionisation energy follows trends within the Periodic Table as they are influenced by proton-electron forces of attraction, and electron shielding
-along a period: first ionisation energy increases due to a decreasing atomic radius and greater electrostatic forces of attraction, electrostatic forced of attraction increase since there is an increasing number of protons
-down a group: first ionisation energy decreases due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction
-when successive ionisation energies are plotted on a graph, a sudden large increase indicates a change in energy level which is because the electron is being removed from an orbital closer to the nucleus so more energy is required to do so
-this large energy increase provides supporting evidence for the atomic orbital theory
-the first ionisation energy of Aluminium is lower than expected due to a single pair of electrons with opposite spin
-as a result there is a natural repulsion which reduces the amount of energy needed to be put in to remove the outer electron

32
Q

Electron orbitals:

A

-electrons are held in clouds of negative charge called orbitals
-there are different types of orbital: s,p, d and f, each orbital can hold 0,1 or 2 electrons
-when an orbital holds 2 electrons the electrons must have opposite spins which have different shapes

33
Q

Electron Orbitals: spin

A

-within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible
-electrons in the same orbital must have opposite spins, spin is represented by opposite arrows
-overall there are three rules for writing out electron configuration:
1. The lowest energy orbital is filled first
2. Electrons with the same spin fill up an orbital first before pairing begins
3. No single orbital holds more than 2 electrons

34
Q

Electron orbitals: Exceptions to the rules

A

-if electron spins are unpaired and therefore unbalanced, it produces a natural repulsion between the electrons making the atom very unstable
-if this is the case, the electrons may take on a different arrangement to improve stability
-example: the 3p^4 orbital contains a single pair of electrons with opposite spins, making it unstable: therefore the electron configuration changes to become 3p^34s^1 which is a much more stable arrangement

35
Q

Electron Configurations:

A

-scientific ideas on electronic configuration have developed over time as new discoveries are made, the current, accepted model is based on the following evidence:
1. Emission spectra provide evidence for the existence of quantum shells
2. Successive ionisation energies provide evidence for quantum shells within atoms and suggest the group to which the element belongs
3. First ionisation energy of successive elements provides evidence for electron sub shells
-these orbitals correspond with blocks on the Periodic Table, each element in the block has outer electrons in that orbital

36
Q

Electron configuration: different sub-shells

A

-each sub shell has a different number of orbitals and therefore can hold a different number of electrons before the next one is filled: s-sub shell = 2 electrons, p-sub shell = 6 electrons, d-sub shell = 10 electrons
-the energy of the orbitals increases from s to d meaning the orbitals are filled in this order, each orbital is filled before the next one is used to hold electrons
-example: Sodium has 11 electrons, these would be written in the following configuration -> Na = 1s^2, 2s^2, 2p^6, 3s^1
-it has 3 energy levels and 4 orbitals holding the 11 electrons
-there are two main exceptions to electron configuration, a completely full or half-dull d sub level is more stable than a partially filled d sub level, so an electron from the 4s is excited to the 3d orbital
Chromium: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^5, 4s^1
Copper: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^10, 4s^1

37
Q

Periodicity:

A

-> refers to the study of patterns of physical, atomic and chemical properties within the Periodic Table that repeat regularly
-the Periodic Table arranges the known elements according to proton number,
-all the elements along a period have the same number of electron shells
-all the elements down a group have the same number of outer electrons, this number is indicated by the group number
-elements are classified into blocks within the Periodic Table that show electron configuration:
S-block = groups 1 and 2
P-block = groups 3 and 0
D- block = transitions metals
F- block = radioactive elements
-these different electron configurations are often linked to other trends within the Periodic Table, periodicity is the study of these trend

38
Q

Periodicity Data:

A

-the trends of properties like melting point, atomic radius and ionisation energy can be represented on a graph, these trends are known as periodic properties
-ionisation energies may be shown on a logarithmic graph, which has a scale where a one unit increase represents an increase by a factor of 10
-for instance, the numbers 10 and 100 would be the same distance apart as the numbers 100 and 1000, this enable you to plot data that has a large range

39
Q

Atomic radius:

A

-along a period, atomic radius decreases
-this is due to an increased nuclear charge for the same number of electron shells
-the outer electrons are pulled in closer to the nucleus as the increased charge produces a greater attraction, as a result, the atomic radius for that element is reduced
-down a group, atomic radius increases, with each increment down a group, an electron shell is added each time
-this increases the distance between the outer electrons and the nucleus, reducing the power of attraction
-more shells also increases electron shielding where the inner shells create a ‘barrier’ that blocks the attractive forces
-therefore, the nuclear attraction is reduced further and atomic radius increases

40
Q

Ionisation energy: as a trend

A

-along a period, ionisation energy increases, the decreasing atomic radius and increasing nuclear charge means that the outer electrons are held more strongly and therefore more energy is required to remove the outer electron and ionise the atom
-down a group, ionisation energy decreases, the nuclear attraction between the nucleus and outer electrons reduces and increasing amount of shielding means less energy is required to remove the outer electron

41
Q

physical evidence for exisiting ions:

A

-ionic substances are brittle and they are unable to conduct electricity in solid form
-however, when ionic substances are molten or aqueous the ions are free to move and so can conduct electricity in these states
-these properties, in comparison to metals which are malleable and conduct electricity when solid, can be explained by the model of oppositely charged ions in a giant ionic lattice

42
Q

electron density maps: ionic bonding

A

-electron density maps show the region around the nucleus in which electrons are distributed
-a high density corresponds to a high probability of an electron being there
-different types of chemical bonds have different electron density maps
-in covalent bonds, for instance, there is a high electron density between the bonding nuclei, whereas in ionic bonds there is a low electron density between ions
-this is explained by the fact that ionic bonds are formed by the physical transfer of electrions whereas covalent bonds are formed by the sharing of electrons
-electron density maps therefore allow chemists to work out the type of bondingpresent, and also the distance between ions in an ionic lattice

43
Q

ion migration:

A

-in electrolysis, ions are free to move and ions move to the oppositely charged electrode- where they gain or lose electrons to form atoms
-this ion migration can be explained by the electrostatic attraction that results between oppositely charged species

44
Q

ionic bonding:

A

-ionic bonding occurs between a metal and a non-metal
-electrons are transferred from the metal to the non-metal to achieve full outer shells, according to the octet rule
-the octet rule states that atoms will often react to form outer shell containing 8 electrons
-when the electrons are transferred, it creates charged particles called ions
-oppositely charged ions attract through strong electrostatic forces to form a giant ionic lattice
-positive ions are formed when an atom loses at least one electron
-negative ions are formed when an atom gains at least one electron

45
Q

ionic bond strength:

A

-the charge of an ion is related to the strength of the ionic bond that forms
-ions with a greater charge will have a greater attraction to the other ions, resulting in stronger forces of attraction and therefore stronger ionic bonding
.larger ions that have a greater ionic radius will have a weaker attraction to the oppositely charged ion because the ttractive forces have to act over a greater distance

46
Q

dot and cross diagrams:

A

-cations (+ve) and anions (-ve) can be represented using dot and cross diagrams-and so can ionic bonding
-the electrons being transferred from the cation can be displayed on the outer shell of the anion

47
Q

ionic radii: down a group

A

-as you move down the group, the number of electron shells increases, and therefore the ionic radius increases

48
Q

ionic radii: across a period

A

-groups 1 and 2: the atoms in these groups lose electrons to form positive ions
-as you go across the period, the magnitude of the positive charge increases for the same amount of electron shielding
-this means there is a greater electrostatic attraction which pulls the outer electrons to form this configuration
-the ions formed are said to be isoelectronic as they have the same electronic configuration (but lose different numbers of electrons to form this configuration)
-groups 6, 7 and 8: the atoms in these groups gain electrons to form negative ions
-as you go across the period the additional proton makes less of a difference, so the only change is a slight decrease in ionic radius

49
Q

polarisation:

A

-in ionic bonds, polarisation refers to the distortion of the electron cloud
-in the perfect ionic model, ions are spherical with no distortion, in reality, the positive ion will attract some of the negative ion’s electron cloud towards itself

50
Q

polarisation power:

A

-the polarising power of a cation and the polarisability of an anions depends on the radius and charge of ions
-a high charged cation will have a strong electrostatic attraction with an anion which will distort the anion’s electron cloud, if a cation had a smaller radius, there will generally be a greater concentration of charge in a small area, causing a distortion of the electron could therefore, the cation’s polarising power increases with an increase in charge and decrease in radius
-the more negatively charged an anion is, the greater the electron cloud, the larger the electron cloud, the more easily it is distorted. Also a larger radius means the outer electrons are held more loosely, allowing them to be more easily distorted however, the anion’s polarisability increases as the radius increases and charge becomes more negative

51
Q

covalent bonding: physical evidence

A

-giant covalent structures have high melting and boiling points due to strong covalent bonds
-they have no free ions or electrons (except graphite and graphene) so generally cannot conduct electricity
-simple covalent molecules have low melting and boiling points because, although the covalent bonds are strong, the intermolecular forces between molecules are weak
-the properties of these covalent structures can all be explained by the model of covalent bonding

52
Q

electron density maps: covalent bonding

A

-as described before, electron density maps show the regions around nuclei in which electrons are likely to be found
-they can be used to identify the type of bonding present
-covalent bonds involve the sharing of electrons between nuclei, so the electron density map shows a high density of electrons between bonding atoms

53
Q

covalent bonding:

A

-covalent bonds form between two non-metals
-there is a strong electrostatic attraction between the two nuclei and the shared electrons between them
-electrons are shared between the 2 outer shells in order to form a full outer shell, multiple electron pairs can be shared to produce multiple covalent bonds

54
Q

Covalent bonding: dot and cross diagrams

A

-the shared paired of electrons can be represented using dot and cross diagrams
-the overlap includes a covalent bond
-the number of electrons within the overlap tells you the nature of the covalent bond:
-2 electrons (1 from each atom): single (sigma) bond, displaced formula represented as -
-4 electrons (2 from each atom): double (sigma and pi) bond, displayed formula represented as =
-6 electrons (3 from each atom): triple bond (sigma, pi, pi) bond, displayed formula represented as ≡
-double and triple bonds can also be shown on dot and cross diagrams with the multiple electron pairs being displayed in the shared segment between two atoms
-the length of a covalent bond is strongly linked to its strength
-short bonds tend to be strong as the atoms are held close together so the forces of attraction are greater, requiring more energy to be overcome
-double and triple bonds are shorter than single covalent bonds, explaining why they are so much stronger

55
Q

dative bonding:

A

-dative or coordinate bonds form when both of the electrons in the shared pair are supplied from a single atom, it is indicated using an arrow from the lone electron pair
-once a dative bond has formed, it is treated as a standard covalent bond because it reacts in exactly the same way and has the same properties regarding length and strength
-since both electrons come from the same atom in a dative covalent bond, in dot and cross diagrams both electrons in that bond will have the same shape
-in order words, they will both be dots or both be crosses

56
Q

Giant covalent structures:

A

-macromolecular covalent substances are covalently bonded into a giant lattice structure
-each atom has multiple covalent bonds which are very strong, giving the substance a very high melting point

57
Q

Giant covalent structures: diamond

A

-is a macromolecular structure made up of carbon atoms each bonded to 4 other carbon atoms
-this forms a rigid tetrahedral structure, making diamond one of the hardest, strongest materials known-which is often why it is used on the tips of drills

58
Q

Giant covalent structures: graphite

A

-graphite is another macromolecular structure made up of carbon atoms
-however, in graphite, each carbon atom is bonded to 3 others in flat hexagonal sheets
-this means there is 1 delocalised electron per carbon atom
-these electrons can move freely, allowing graphite to conduct electricity
-graphite can therefore be used in an electrode
-the intermolecular forces between layers are weak and can easily slide over each other, meaning graphite can be used as a lubricant

59
Q

Giant covalent structures: graphene

A

-graphene consists of single, 2D sheets of graphite that are just one atom thick
-these sheets formed of hexagonal carbon rings that create a very strong, rigid material that is extremely lightweight
-delocalised electrons move through each layer allowing it to conduct electricity, making graphene a useful material in electronics

60
Q

Bond polarity in covalent bonds:

A

-the negative charge around a covalent bond is not spread evenly around the orbitals of the bonded atoms

61
Q

Electronegativity: definition

A

The power of an atom to attract the electron pair in a covalent bond towards itself

62
Q

Electronegativity: trend

A

-this ‘power’ is different for every atom depending on its size and nuclear charge
-Electronegativity increases along a period, as you move across a period, atoms have a greater nuclear charge and a smaller covalent radius
-this allows the nucleus to attract the bonding electrons more strongly
-electronegativity decreases down a group, going down a group, atoms increase in size due to the extra electron shells, increasing shielding towards the bonding electrons
-ionic and covalent bonding are the extremes in a continuos scale of bonding
-if the electronegativity between two atoms is great enough, an ionic bond will form between them
-the bond polarity can be permanent or induces depending on the molecule and how it interacts with things around it

63
Q

Permanent dipole:

A

-if the two atoms that are bonded have sufficiently different electronegatives, a polar bond forms
-the more electronegative atom draws more of the negative charge towards itself and way from the other atom, producing a a- region and a+ region which produces a permanent dipole
-hydrogen fluoride is a polar molecule as fluorine is a lot more electronegative than hydrogen, this causes electrons to be drawn towards the fluorine atom
-polar molecules with a permanent dipole can align to form a lattice of molecules similar to an ionic lattice

64
Q

Polarisability of ions:

A

-where electronegativity difference is large enough the electrons will not be shared between atoms but will instead spend most of their time localised around one of the atoms only - forming ions
-a cation is able is able to distort the shape of the electron could on a nearby anion
-the extent by which the electron cloud is distorted is known as polarisability
-smaller cations with a greater charge are the most polarising
-anions with a larger ionic radius can have their electron clouds distorted more easily

65
Q

Polar bonds and polar molecules:

A

-a polar bond results from a large difference in electronegativity between two atoms forming a covalent bond
-tables of electronegativity can be used to workout if a bond will be polar or not
-if the difference in electronegativity between bonding atoms is between 0.4 and 1.7, the bond will be a polar covalent bond
-if the difference in electronegativity is greater than that, the bond will be ionic
-since electronegativity is a periodic trend, elements that are close together on the periodic table will not form polar bonds

66
Q

Polar molecules:

A

-polar molecules arise when there is an overall difference in polarity across the molecule, due to the arrangement of polar bonds and the geometry of the molecule
-polar molecules must have polar bonds, however a molecule with polar bonds may not necessarily be a polar molecule
-CO2 - the C=O bonds in CO2 are polar, however the molecule is linear so the dipoles created by each polar bond cancel out
-H2O - the O-H bonds in water are polar and the geometry of the molecule is bent, so overall there is a polarity and the molecule is polar

67
Q

Metallic bonding:

A

-metallic bonding consists of a giant lattice of positively charged ions surround by a ‘sea’ of delocalised electrons
-there are very strong electrostatic forces of attraction between the oppositely charged particles
-the greater the charge on the positive ion, the stronger the attractive force since more electrons are released into the ‘sea’
-ions that are larger in size, such as barium, produce a weaker attraction due to their greater atomic radius

68
Q

Metallic structure (e.g. aluminium):

A

-substances with metallic structure are often good conductors
-the ‘sea’ of delocalised electrons is able to move and carry a flow of charge
-metals are also malleable because the uniform layers of positive ions are able to slide over one another
-the delocalised electrons prevent fragmentation as they can move around the lattice
-the electrostatic forces of attraction between the positive ions and delocalised electrons are very strong and therefore require a lot of energy to overcome
-this means metallic substances have high melting points and are nearly always solid a room temperature
-Mercury is the only metal which is liquid at room temperature

chem year 1 (5 decks)

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