Atomic Structure

Chlorine Diatomic Mass spec Cl2

These lines would be in addition to the lines at 35 and 37. The relative heights of the 70, 72 and 74 lines are in the ratio 9:6:1. However, you cannot make any predictions about the relative heights of the lines at 35/37 compared with those at 70/72/74 because those will depend on what proportion of the molecular ions break up into fragments.

Isotopes

Variations of the same element that differ in their neutron number. This means that isotopes of an element all have the same number of protons and therefore the same Atomic number, so they occupy the same position in the periodic table

What three subatomic particles make up atoms?

protons, neutrons, electrons

Electron Facts

Electrons carry a relative negative charge. Electrons have very little (almost no) mass but each is approximately 1840 times lighter than either a proton or a neutron. In a neutral atom, the number of electrons will equal the number of protons in the nucleus. An atom with a differing amount of protons and electrons - due to the loss or gain of one or more electrons - will carry a net charge and is known as an ion.

TOF mass spectrometry Examples

For example, the analysis of a geological sample for detection and/or measurement of relative abundance of materials; a biopsy sample from a patient for diagnosis or medical research; or a crime scene sample to help solve crimes; to name just a few. Historically and most topically, it was used during the Second World War to isolate and enrich uranium isotopes for Project Manhattan, the top-secret war effort by the Allies, aimed at developing the first nuclear weapons.u

The energy needed to remove an electron from an atom is called

Ionisation energy

An atom of an element has a net charge of 1+. Which subatomic particle of the atom could have changed in number?

Minus an electron

Proton Facts

Protons carry a relative positive charge. A proton has a relative mass of 1. The number of protons in the nucleus of an atom determines its position on the periodic table. The number of protons gives an element its 'Atomic Number'.

Mass number also known as...

Relative Atomic Mass, Atomic Mass Number, nucleon number, or simply the weight

Define first ionisation energy

The energy required to remove one mole electrons from one mole of isolated gaseous atoms to form one mole of gaseous ions

First ionisation energy trends (dips)

The graph above shows that the increase in ionisation energy across both periods is not smooth. There are 2 dips in each: Beryllium to Boron in period 2 Nitrogen to Oxygen in period 2 Magnesium to Aluminium in period 3 Phosphorus to Sulphur in period 3 These observations can be explained by looking at the electron configurations.

Mass Spectrometry Step 2 Acceleration

The positive ions are then guided by an intense electric field through an array of electrode plates with a small opening in their centre: >Positive ions, passing the first accelerator electrode in a beam, gain stronger momentum >This momentum accelerates the beam further to the intermediate plate, gaining even more momentum >The high momentum accelerates the positive ion beam further to the final plate where the uniformly high-velocity beam will move preferentially towards the deflecting magnet

What is the electron configuration of sodium?

Sodium has 11 electrons. Following the chart, we can write this in the configuration: Na = 1s22s22p63s1 So, we can see there are 3 energy levels and 4 orbitals containing the 11 electrons.

Electron configuration of carbon

The first shell s can hold 2 electrons maximum, so this is 1s2. This leaves 4 more electrons to fill. Shell number 2 can have 2 subshells 2s and 2p. 2s can only hold 2 electrons, so this is 2s2. This leaves only 2 more electrons to fill. Next to be filled will be subshell 2p. Although it can hold up to 6 electrons, there are only 2 electrons left, so this will be 2p2. The full electron configuration for Carbon, therefore, is 1s2 2s2 2p2 as shown below (you can check this against the interactive periodic table).

Shielding effect in the periodic table

The more shells and subshells of electrons there are shielding the valence electron from the attraction pull of nucleus, the less the 1st IE needed to remove an electron. Across a period, there is no change in the number of shells that electrons occupy; the first period (hydrogen and helium) confines its electrons to one shell, the second period to two shells, and the third period to three shells, as shown above. Consequently, the shielding effect is constant across a period, but increases going down a group. Therefore, shielding does not contribute to the increasing 1 st IE trend seen across periods 2 and 3, but it does contribute to the decreasing 1st IE seen down the group between period 2 and 3.

The only difference between isotopes is

Their mass number. Chemically all isotopes react the same way

Mass Spectrometry Step 1 Ionisation

We know that atoms can have a net charge by either losing or gaining an electron; these are called ions. Mass spectrometers generate positively charged ions by: > Having a high vacuum environment, ensuring an empty chamber and no air molecule interference > Creating an electron beam by heating a metal coil inside the ionising chamber, directing the beam towards an electron trap > Passing vaporised samples perpendicular to the electron beam, allowing fragmentation of the samples, as well as knocking off one or more electrons to create positively charged ions (1+ or 2+, respectively) >Using the slightly positively charged ion-repeller plate to direct the generated positive ions out of the ionising chamber

1st ionisation energy as an exemplified equation

Where g means gaseous. The energy needed to remove electrons from one mole of X is in kilo-joule per mole (kJ/mol) unit. The bigger the energy value, the harder it is to ionise that element to its 1+ ion form, for example, in elements that normally form negatively charged ions or those that do not normally form ions, such as the inert or noble gases.

Several factors affect the first ionisation energy

>The charge on the nucleus - the more protons there are, the stronger the attraction to its orbiting electrons and the higher the 1st IE needed to remove an electron. >The atomic radius - the further away the electron is from the nucleus (bigger radius), the less the 1 st IE needed to remove an electron. >The number of electrons (in shells and sub-shells) between the electron and the nucleus - this is called the shielding effect ; the more shells and subshells of electrons there are shielding the valence electron from the attraction pull of nucleus, the less the 1 st IE needed to remove an electron.

Electron shell rules to remember

>The further away a shell is from the nucleus (the higher shell number), the higher the energy level. >The lower number shells, or the lower energy shells, are always filled up first. >Therefore, if there is an incomplete shell, it will be the outermost one, called the valence shell. >The subshells are annotated s, p, d, f, and g and they can hold 2, 6, 10, 14, and 18 electrons in them respectively. >The first shell, closest to the nucleus, can only have one subshell s. >The second shell away from the nucleus can have 2 subshells in them: s and p. >The third subshell can have 3 subshells: s, p and d . >The fourth can have 4 subshells: s, p, d and f. >The fifth can have 5 subshells: s, p, d , f and g. >For our purposes, knowledge up to 4 levels is sufficient

Basic principles of mass spectrometry

A mass spectrometer performs the following four steps (in a vacuum): Creates positive ions from a neutral sample Accelerates the ions to a high kinetic energy Separates the ions according to their mass/charge ratio by magnetic field deflection Detects and measures the relative abundances of ions and their relative masses; the information being represented as a mass spectrum.

Mass

A measure of the amount of matter in an object

Time of Flight Mass Spectrometry

A time-of-flight (TOF) Mass spectrometry is an analytical technique that can identify and measure chemical components of a sample, based on mass-to-charge ratio and how long a flight an ion takes through the spectrometer. It is used in many different fields where the systematic breakdown of pure or complex samples for identification or measurement purposes are needed.

How many orbitals and electrons do these contain: a. 1s b. 2p c. 3s d. 4d e. 5f

A. 1 orbital, 2 electrons B. 3 orbitals, 6 electrons C. 1 orbital, 2 electrons D. 5 orbitals, 10 electrons E. 7 orbitals, 14 electrons

Which has the highest energy: a. 3p, 3d or 4s? b. 3d, 4d or 5s? c. 4f, 5p or 6s?

A. 3d B.4d C.4f

How do these factors affect the value of the first ionisation energy of a substance: a. Nuclear charge b. Atomic radius c. Shielding

A. As the nuclear charge increases, the ionisation energy increases. This is because attraction for the outermost electron increases and more energy is requires to remove an electron B. As atomic size increases, ionisation energy decreases. This is because the attraction of the positive nucleus for the negative electron decreases and less energy is required to remove an electron C. As shielding increases, ionisation energy decreases. This is because the outermost electron is screened from the attraction of the nucleus by the repelling effect of the inner electron.

Write equations for: a. The first ionisation energy of Li b. The third ionisation energy of Al

A. Li -> Li+ + e- B. Al2+ -> Al3+ + e-

What is a. A shell? b. A subshell? c. An orbital? d. Electron spin? e. A valence shell?

A. The area around the nucleus where electrons orbit B. Shells which are divided into smaller shells; s,p,d and f, and are made up of orbits. These hold up to 2, 6. 10 and 14 electrons respectively C. Regions around the nucleus that can hold up to two electrons. These make up sub shells D. A quantum mechanical property of electrons, in an electron pair inside an orbital, if one spins upwards the other must spin downwards. They cannot spin in the same direction E. The outermost shell of an atom which contains the valence electrons - the electrons which contribute to chemical reactions

How does different neutron count affect isotopes?

Although isotopes have the same chemical properties, the different neutron count gives them different physical properties. Some isotopes can be radioactive, referred to as radioisotopes. For example, Carbon-14 is radioactive, whereas Carbon-12 and Carbon-13 are both stable isotopes.

Example of an isotope is the element carbon

An example is the element Carbon, which has 6 protons and a mass number of 12.011. This mass number is an average of all the mass numbers of Carbon isotopes.

Effective nuclear charge trend in the periodic table

As we go across a period from left to right, there is a trend of increasing nuclear charges as the atomic number increases. This contributes to the trend of increasing 1 st IE across the period and as we saw in the bar graph, as the more protons there are in a nucleus, the stronger the positive pull to the orbiting electrons and therefore, the more energy needed to remove them.

Atomic radius effect trend in the periodic table

Atomic radii decrease as we go across a period and up a group. This means that as we traverse across the period from left to right, or bottom to top, the tighter the atom is and therefore, the stronger the nuclear pull to the electrons. This leads to higher IE requirements, as seen in the bar graph above. This also contributes to the observation in the bar graph that the 1st IE needed in period 2 is generally higher than that in period 3; because as we go up a group, the radii get smaller and therefore, the attraction between the nucleus and the electrons is stronger, leading to an increase in IE needed to remove an electron.

Say you are given an unknown atom, but you know it carries 8 protons and is a neutral atom. What is the identity of the element? And how many electrons does it carry?

Carbon 8 electrons

Standard relative element to which all elements are compared ?

Carbon 12

What is carbon 12 known for?

Carbon-12 is the most abundant of the 3 and has a special function. It has 6 protons, 6 neutrons and 6 electrons, and a mass number of exactly 12. It is used as the standard relative against which the atomic weights of all elements are measured.

the mass spec of chlorine (a diatomic element)

Chlorine is taken as typical of elements with more than one atom per molecule. We'll look at its mass spectrum to show the sort of problems involved. Chlorine has two isotopes, 35Cl and 37Cl, in the approximate ratio of 3 atoms of 35Cl to 1 atom of 37Cl. You might suppose that the mass spectrum would look like this: However, it is not really that simple. The problem is that chlorine consists of molecules, not individual atoms. When chlorine is passed into the ionisation chamber, an electron is knocked off the molecule to give a molecular ion, Cl 2+. These ions will not be particularly stable and some will fall apart to give a chlorine atom and a Cl + ion. The term for this is fragmentation. Cl2+ -> Cl + Cl+ If the Cl atom formed is not then ionised in the ionisation chamber, it simply gets lost in the machine - it is neither accelerated nor deflected. The Cl+ ions will pass through the machine and will give lines at 35 and 37, depending on the isotope; producing exactly the same pattern as seen in the last diagram. The problem is that you will also record lines for the un-fragmented Cl 2+ ions. Think about the possible combinations of chlorine-35 and chlorine-37 atoms in a Cl2+ ion. Both atoms could be 35Cl, or both atoms could be 37Cl, or you could have one of each sort. That would give you total masses of the Cl2+ ions of: 35 + 35 = 70 35 + 37 = 72 37 + 37 = 74

Valence electrons

Electrons on the outermost energy level of an atom

Electrospray ionisation

Electrospray ionization (ESI) is a technique to generate ions for mass spectrometry using electrospray by applying a high voltage to a liquid to produce an aerosol. High energy electrons from a cathode are fired at the sample which knocks one electron off the sample, ionising it.

Atomic number facts

Every element in the periodic table has a different Atomic Number The Atomic Number (Z) equals the number of protons in an element The Atomic Number (Z) also equals the number of electrons in a neutral element The Atomic Mass Number (A) is the total number of protons AND neutrons in an element The Atomic Mass Number (A) is always bigger than the Atomic Number (Z), except for Hydrogen, which has the value 1 for both.

Following the order of subshell filling, write out the electron configuration for the element titanium (atomic number 22).

If we follow the order as shown by the arrows on the electron sub shell configuration diagram, the electron configuration of titanium will be as follows: 1s2 2s2 3s2 3p6 4s2 3d2 Note how 4s2 was filled before 3d. Even though 4s should be the outermost in this case, it actually has lower energy when compared to 3d so gets filled first.

Why does potassium follow the electron configuration trend but chromium and copper do not?

In potassium the electrons fill up in the way you might expect; following the logical order 1,2,3,4 and s,p,d,f but in chromium and copper the 4s shell has only one electron in it before filling the 3d shell

Mass Spectrometry Step 3 Deflection

In the deflection chamber, the poles of an electromagnet are placed astride the bent tube: >This produces a high-intensity, variable magnetic field within the tube. >The field causes a change in the direction of movement of these positive ions. >The magnitude of this deflection depends entirely on the mass-to-charge ratio of the ion. Lighter ions and ions with 2 or more positive charges are deflected more. >Thus the ion beam is split up into a series of separate beams, each of which has particles of one specific mass/charge ratio (m/z or m/e). Ions with the smallest mass/charge ratio will deflect the most.

Mass Spectrometry Step 4 Detection

In the diagram above, only stream B of ions, with intermediate mass/charge ratios move towards the ion detector - a metal box that is connected to an amplifier wire. So what happens here? >When a positive ion B hits the metal box, the ion is neutralised by gaining an electron from the metal, leaving an empty space for an electron in the metal. >The connecting wire then donates an electron to fill in the empty space in the metal. >This flow of electrons can be detected as an electric current that can be amplified and recorded. >The more positive ions hit the metal box, the more electron flow occurs and the higher the reading recorded. This is a measurement of the number of ions arriving at the detector. >The mass is measured on the Carbon-12 scale, i.e. it has an atomic mass of exactly 12 units. However, the mass, and therefore mass/charge (m/z) calculated also factors in the size of the magnetic field used to bring the ion into the detector. >A mass spectrum results from the electrical current reading calibrated to the m/z settings directly. >This produces a 'plot' of intensities of successive ion beams, where the height of each peak is proportional to the number of ions of a given mass/charge ratio reaching the collector plate in unit time. Obviously, the mass spec analysis will not be very comprehensive unless the coverage is improved to include ions from streams A and C. This can be done by varying the strength of the magnetic field used.

TURN OVER There are three things we can glean from this chart: 1. The number of isotopes - the two peaks in the mass spectrum shows that there are 2 isotopes of boron - with relative isotopic masses of 10 and 11 on the Carbon-12 (12C) scale. 2. The abundance of the isotopes - the relative sizes of the peaks gives a direct measure of the relative abundances of the isotopes. The tallest peak is often given an arbitrary height of 100, but it can be any value. You can find the relative abundances by measuring the lines on the diagram. In this case, the two isotopes (with their relative abundances) are: boron-10 (10B) ~ 23 boron-11 (11B) ~ 100 3. The relative atomic mass of boron - Suppose you had 123 typical atoms of boron. 23 of these would be 10B and 100 would be 11B. The total mass of these would be (23 x 10) + (100 x 11) = 1330 The average mass of these 123 atoms would be 1330 / 123 = 10.8 (to 3 significant figures). 10.8 is the relative atomic mass of boron.

Intererporate this mass spectra

TURN OVER In the diagram, the vertical axis labelled "relative abundance" is related to the current received by the chart recorder - and therefore to the number of ions arriving at the detector: the greater the current, the more abundant the ions. The relative abundance can be in unit numbers or percentages. >In the mass spectrogram above, produced by injecting a molybdenum sample: >The most abundant ion has a mass/charge ratio of 98. >The original sample is not perfectly pure and contains ions with varying mass/charge ratios of 92, 94, 95, 96, 97 and 100. >This result shows the molybdenum sample to consist of 7 different isotopes. >Assuming that the ions all have a charge of 1+, this means that the masses of the 7 isotopes on the carbon-12 scale are 92, 94, 95, 96, 97, 98 and 100. >The average mass of molybdenum from this experiment would be 96, close to the published 95.65 relative mass (check your periodic table). >How do you know if there were also 2+ ions present? In this case: >Every one or some of the lines in the chart diagram would also have another line at exactly half its m/z value (for example, 98/2 = 49). >Those lines would have shorter peaks than the 1+ ion lines because the chances of forming 2+ ions are much less than forming 1+ ions.

Interpolating Mass Spectrograms

TURN OVER 1. The number of isotopes - the 5 peaks in the mass spectrum show that there are 5 isotopes of zirconium - with relative isotopic masses of 90, 91, 92, 94 and 96 on the 12C scale 2. The abundance of the isotopes - this time, the relative abundances are given as percentages. Again you can find these relative abundances by measuring the lines on the diagram. In this case, the 5 isotopes (with their relative percentage abundances) are: zirconium-90 (90Zr) 51.5 zirconium-91 (91Zr) 11.2 zirconium-92 (92Zr) 17.1 zirconium-94 (94Zr) 17.4 zirconium-96 (96Zr) 2.8 3. The relative atomic mass of zirconium - Suppose you had 1000 typical atoms of zirconium. 515 of these would be 90Zr, 112 would be 91Zr and so on. The total mass of these 1000 typical atoms would be The average mass of these 100 atoms would be 91318 / 1000 = 91.3 (to 3 significant figures). 91.3 is the relative atomic mass of zirconium.

Interporate this mass spec

Describe and explain, with reference to electron configurations, the trend of first ionisation energy across period 2

Ionisation energy increases across period 2 due to increasing nuclear charge and decreasing atomic radius. Since the electrons all go into the same shell, the shielding if the ionising electron is about the same. Therefore, the outer electrons are increasingly more attracted to the positive nucleus so more energy is needed to remove an electron

Neutron Facts

Neutrons carry no charge. A neutron has a relative mass of 1. Neutrons make up the rest of the mass of the nucleus (along with protons). The number of neutrons that the nucleus contains can vary within different atoms of the same element. This gives rise to the isotopes of an element.

Neutron number is equal to ?

Number of Neutrons = Atomic Mass Number - Atomic Number

Full electron configuration for every element in periods 2 and 3

Remember that: The first shell 1s does not contain sub-shells. The second shell can contain the subshells 2s 2p. The third shell can contain the subshells 3s 3p 3d.