RMK- atomic structure

Question 1

Pure substances

Answer 1

elements and compounds are pure substances as they are made up of one type of substance and have a fixed composition

Question 2

Impure substances

Answer 2

Mixtures are made by combining two or more pure substances together. they do not have a fixed composition

Question 3

Elements

Answer 3

a chemical element is a neutral substance that cannot be broken down into simpler substances using chemical methods. it consists of only atoms that contain the same number of protons, although the atoms can contain a different number of neutrons. there are 118 known chemical elements

Question 3

Compounds

Answer 3

a compound is a pure substance formed when two or more chemical elements are chemically bonded together in a fixed ratio. a compound has different properties from its component elements.

Question 4

Mixtures

Answer 4

contain pure substances that aren’t chemically bonded so can be separated by physical means. Mixtures aren’t arranged in a fixed ratio. mixtures are categorised into two sections.

Homogenous: all elements are in the same phase e.g. salt and water

Heterogenous: all elements are in different phases e.g. oil and water

Question 5

States of matter

Answer 5

solids: occupy a fixed space and fixed volume. particles are held closely together in a lattice. In the solid state particles are held closely together in a fixed position in a lattice. the particles can vibrate about a fixed point but possess no individual translational velocity

liquid: fixed volume that takes up the shape of its container. particles are close together but moving in random motion.

Gas: completely fills its container. widely spaced particles moving with rapid, random motion.

Question 6

what is heat?

Answer 6

heat is a measure of the total amount of energy in a given amount of substance and therefore depends upon the amount of substance present. it is measured in joules

Question 7

What is temperature?

Answer 7

measure of the hotness of a substance and therefore is independent on the amount of substance present. it is a measure of the average kinetic energy of the substance. absolute temperature is measured in Kelvin. 0 kelvin is when all motion has stopped.

Question 8

Kelvin and Degrees

Answer 8

0 kelvin= -273 degrees celsius

273 kelvin= 0 degrees celsius

Question 9

Explain the changes of state from solid to liquid to gas

Answer 9

In the solid state particles are held closely together in fixed position in a lattice. The particles can vibrate about a fixed point but possess no individual translational velocity.
As heat is supplied the vibration intensifies and eventually the particles will have sufficient energy for the lattice to break and the particles are free to move. This is the liquid state. There are still attractive forces between the particles.
As more heat is added the particles move faster, i.e. gain in kinetic energy.
Eventually they will have sufficient energy to overcome the attractive forces and escape as a vapour. When the vapour pressure is equal to the external pressure the liquid boils. Hence the boiling point depends upon the external pressure. At 100 kPa pressure water boils at 373 K (100 °C).

Question 10

Changes of State (names)

Answer 10

Solid —> Gas = sublimation

Gas—>Solid = deposition

Solid—>Liquid =melting

Liquid—>Solid = freezing

Liquid—>Gas = vaporisation (evaporation)

Gas—>Liquid = condensation

Question 11

Describe cooling and heating curves in graphs

Answer 11

in a graph that shows a change of state during the change of state there will be flat sections in the graph, this is because, when changing from a liquid to a solid, the temperature is dropping, but bonds are being made which is exothermic; the heat energy given out balances the decreasing temp leading to the flat line. When changing from a solid to a liquid the opposite is true, bonds are being broken which is endothermic, this results in heat being taken in; and so the energy taken in balances the increasing temp leading to a flat line.

Question 12

Separation techniques (filtration)

Answer 12

to separate solid particles from a liquid. it may be carried out under gravity at atmospheric pressure or for finer particles under reduced pressure.

Question 13

Separation techniques (solvation and evaporation)

Answer 13

where one component of a solid mixture is soluble the mixture can be warmed with the solvent, and filtered to remove insoluble particles then the solvent is removed by evaporation

Question 14

Separation techniques (recrystallisation)

Answer 14

an impure solid can be dissolved in a suitable solvent by heating then allowed to cool slowly so that the solid separates out as crystals which can be washed to remove any lingering solvent and dried

Question 15

Separation techniques (distillation)

Answer 15

used to separate a volatile liquid from dissolved components or less volatile liquids

Question 16

Separation techniques (chromatography)

Answer 16

used to separate different components in a mixture. in paper chromatography a spot of the mixture is placed on the paper and the solvent allowed to rise up the paper. different components partition between the solvent and the paper by different amounts so rise at different rates. A developer can be used for colourless components. the retardation factor Rf= distance travelled by component divided by the distance travelled by the solvent from the original spot.

Question 17

Key factors of a nucleus

Answer 17

1. it is very small in comparison to the atom
2. it has a highly dense structure containing virtually all the mass of the atom
3. it has a positive charge

Question 18

the Nucleus- relative charges, relative masses and location

Answer 18

proton: 1, +1, nucleus
neutron: 1, 0, nucleus
electron: negligible, -1, outside nucleus

Question 19

Isotopes

Answer 19

isotopes are different atoms of the same element with a different number of neutrons. as a result, they have different mass numbers, A, but the same atomic number, Z.

Chlorine for example has two isotopes: one with mass number 35, and one with mass number 37. they have similar chemical properties, as they are both chlorine atoms, but different physical properties such as density, because atoms of one isotope are heavier than atoms of the other.

Question 20

How is mass spectrometry done?

Answer 20

the sample is injected into the instrument and vaporised. the atoms within the sample are then bombarded with high energy electrons. As a result, the atoms lose some of their electrons to form positively charged ions, known as cations.

the resulting ions are then accelerated by an electric field and deflected by a magnetic field. the degree of deflection depends on the mass to charge ratio. particles with no charge are not affected by the magnetic field and will therefore never reach the detector. the species with the lowest mass and highest charge will be deflected the most.

Question 21

Emission spectra and absorption spectrum.

Answer 21

a pure gaseous element subjected to a high voltage under the reduced pressure will glow- in other words, it will emit light. when this light passes through a prism, it produces a series of lines against a dark background. this is known as an emission spectrum. in contrast when a cold gas is placed between the prism and a source of visible light of all wavelengths, a series of dark lines within a continuous spectrum will appear. this is known as the absorption spectrum.

Question 22

how do you carry out a flame test?

Answer 22

• clean the end of the potassium or nichrome wire by dipping it into a small sample of dilute hydrochloric acid and placing it in a non luminous bunsen burner flame. repeat until no flame colour is observed
• dip the end of the flame test wire into one of the salt samples , and place it in the edgre of the non luminous bunsen burner flame, noting down the identity of the metal in the salt and the colour observed.

Question 23

Flame test colours (zinc, potassium, strontium, sodium and copper)

Answer 23

Zinc- Yellow

Potassium- Lilac

Strontium- Red

Sodium- Orange

Copper- Blue/green

Question 24

what is the EM spectrum in order of decreasing wavelength, increasing frequency and increasing energy

Answer 24

R–>M–>I–>V–>U–>X–>G

Question 25

what Is the speed of light

Answer 25

3.00x10^8

Question 26

how do you calculate ionisation energy?

Answer 26

E=hf

where E is the specific energy possessed by the photon, expressed in joules, J

h= plancks constant, 6.63x10^-34 Js

f= frequency of the radiation, expressed in hertz, Hz, or inverse seconds, s^-1

Question 27

What postulations did the Bohr model lead to in regard to electron position?

Answer 27

1. the electron can only exist in certain stationary orbits around the nucleus. these orbits are associated with discrete energy levels.
2. when an electron in the orbit with the lowest energy level absorbs a photon of the right amount of energy, it moves to a higher energy level and remains at that level for a short time (excited state)
3. when the electron returns to a lower energy level it emits a photon of light. this photon represents the energy difference between the two levels.

Question 28

where is the most stable states of hydrogen?

Answer 28

state at n=1 (n is the principle quantum number), where the electron has the lowest possible energy. this energy level is known as the ground state of the atom. In contrast, the energy levels with n=2,3… are called excited states. atoms in excited states are unstable and spontaneously return to the ground state by emitting photons of specific wavelengths.

Question 29

what energy is released at n=1?

Answer 29

Electron transitions to the ground state, n=1, release higher energy, shorter wavelength ultraviolet light

Question 30

What energy is released at n=2

Answer 30

visible light

Question 31

what energy is released at n=3

Answer 31

while electrons returning to n=3 produce lines in the infrared region of the electromagnetic spectrum.

Question 32

Limitations of the Bohr model

Answer 32

1. the model could not predict the emission spectra of elements containing more than one electron. it was only successful with the hydrogen atom.
2. it assumed the electron was a subatomic particle in a fixed orbit around the nucleus
3. it could not account for the effect of electric and magnetic fields on the spectral lines of atoms and ions
4. it could not explain molecular bonding and geometry

Question 33

what is Heisenbergs uncertainty principle?

Answer 33

it is impossible to precisely know the location and momentum of an electron simultaneously. as opposed to bohrs model which stated that electrons exhibited fixed momentum in specific circular orbits

Question 34

what are schrodingers wave functions?

Answer 34

electrons in atoms in terms of their probability density, using heisenbergs idea that the momentum and position of electrons are uncertain. instead of saying that electrons will follow a defined travel path, this theory gives the probability that an electron will be found at a specific region of space where there is a high probability of finding an electron.

Question 35

what are atomic orbitals?

Answer 35

there are several types of atomic orbitals, and each orbital can hold a maximum of two electrons. each orbital has a characteristic shape and energy. the first four atomic orbitals, in order of increasing energy are labelled s,p (x,y,z) ,d and f. subsequent orbitals are theoretical, and these are labelled alphabetically (g,h,i,k and so on)

Question 36

What are the main energy levels in an atom (as opposed to sublevels).

Answer 36

the principle quantum number (n), introduced by the bohr model represents the main energy levels. these energy levels are split into sublevels comprised of atomic orbitals. for example, for n=1,2 and 3, the s atomic orbitals are 1s, 2s, 3s… As n increases, the s orbitals are further distanced from the nucleus

Question 37

Orbital diagrams

Answer 37

the electrons are shown in pairs or singular, and shown by arrows, one up and one down (in a paired electron) the arrows represent the direction of the spin

Question 38

What are Atomic Orbitals?

Answer 38

Atomic orbitals are regions of space where there is a high probability of finding electrons. electrons are charged negatively, and like charges repel each other, so two electrons should not be able to occupy the same region of space. a pair of electrons with opposite spins behave like magnets facing in opposite directions.

Question 39

what is the Pauli exclusion principle?

Answer 39

only two electrons can occupy the same atomic orbital and those electrons must have opposite spins.

Question 40

what is Hunds rule?

Answer 40

every degenerate orbital in a sublevel is singly occupied before any orbital is doubly occupied and that all electrons in singly occupied orbitals have the same spin. this means that the three p orbitals must have one electron with the same spin in each of them before any orbital can become doubly occupied with an electron of opposite spin.

Question 41

What are exceptions to the Aufbau principle?

Answer 41

The Aufbau principle correctly predicts the order of filling atomic orbitals for most elements. However, when atoms loose electrons to form ions the electrons in the sublevel with the highest principle quantum number (n) are lost first.
The ground state for copper and chromium are also different for those predicted by the Aufbau principle. The predicted electron configuration of copper is [Ar] 4s^2 3d^9, as the Aufbau principle suggests that the lower energy 4s orbital should be filled first. However the observed ground state electron configuration for copper is [Ar] 4s^1 3d^10. For chromium the predicted configuration is [Ar] 4s^2 3d^4 and the observed is [Ar] 4s^1 3d^5. In each case, promoting a 4s electron to a 3d level leads to a more stable electron configuration. In the case of copper, this gives a full 3d sublevel, and in the case of chromium, there are no paired electrons but rather six half occupied orbitals, each containing an electron with the same spin

Question 41

What is the Aufbau principle?

Answer 41

as electrons are added to atoms, the lowest available energy orbitals fill before higher energy orbitals do.
when electron shells fill, generally the following order is observed.
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p

Question 42

Condensed electron configurations.

Answer 42

as the atomic number of an element increases, the full electron configuration gets longer and it can be time consuming to write. the chemistry of atoms and ions is mostly determined by their valence electrons, that is, the outermost electrons, rather than the inner core electrons. a more convenient way to write the electron configurations is to highlight the valence electrons and represent the inner core electrons as the previous group 18 (noble gases) element in the period table.

Question 43

what is ionisation energy?

Answer 43

the minimum energy required to eject an electron out of a neutral atom or molecule in its ground state.

Question 44

how does ionisation energy vary going down the groups and across the periods?

Answer 44

First ionisation energy generally decreases going down the groups of the periodic table and increases across the periods.

Going down a group, the number of sublevels increases. the outermost electrons are shielded from the pull of the nucleus by the electrons in the lower energy sublevels. the more sublevels the greater the shielding and therefore less energy is required to remove electrons from the outermost sublevel.

Going across a period, the number of protons in the nucleus increases, so the outermost electrons are held closer to the nucleus by the increased nuclear charge. At the same time, the shielding effect remains nearly constant because the number of inner electrons does not change. therefore more energy is required to remove the outermost electron, so IE increases along the period

Question 45

what is the discontinuity of ionisation energy across the period of group 2 and group 3 elements? (dip in the graph)

Answer 45

The valence electron configuration of beryllium is 2s^2 while for boron it is 2s^2 2p^1. the paired 2s^2 electrons shield the single 2p electron in boron from the nucleus, making the electron slightly easier to remove.

The same trend can be observed in comparing group 2 to group 3 elements in any period. For example, the 3s62 electrons shield the lone 3p^1 electron in aluminium, so the first ionisation of aluminium is lower than that of magnesium.

Electrons are always removed from the highest occupied energy level, and from the highest energy sublevel within that level.

Question 46

what is the discontinuity of ionisation energy across the period between the group 15 and 16 elements? (dip in the graph)

Answer 46

From group 15 to 16 there is also a drop in ionisation energy. The electron configuration of nitrogen is 1s^2 2s^2 2p^3 while for oxygen it is 1s^2 2s^2 2p^4.

Nitrogen has a more stable electron configuration than oxygen as it has a half filled p sublevel and therefore more energy is required to remove an electron from nitrogen. This is because the paired electrons in oxygen occupy the same region of space and have increased repulsion. However, in nitrogen the three electrons in the 2p orbitals do not come into close proximity.

Question 47

How do you calculate ionisation energy from spectral data?

Answer 47

As the principle quantum number of energy levels increases, the distance between the levels converges to a continuum. This can be observed by the convergence of spectral lines in the hydrogen emission spectrum.

E= hc/λ

ionisation energy in kJ mol^-1= ((energy needed to remove one electron from an atom)*6.022x10^23)/λ

Question 48

Why is it harder to remove electrons in successive ionisation energies/

Answer 48

It requires more energy to remove the second and successive electrons from an atom because the number of protons exceeds the number of remaining electrons while the electron-electron repulsion decreases (removing from an increasingly positive ion) As a result electron clouds are pulled closer to the nucleus and held tighter by increased electrostatic attraction. Once all the valence electrons are removed so that only the stable noble gas configuration remains, the energy required to remove the next electron increases sharply.

Question 49

Why do certain gaseous vapours only emit certain frequencies of visible light?

Answer 49

• Electrons can only occupy energy levels
• Which are fixed characteristic energies/ distances from the nucleus
• So when an excited electron falls to a lower energy level it emits a known, discrete energy corresponding to the energy groups
• E= hv, so only frequency observed