Atomic Structure

History of the atom Robert Boyle (1st to define an element) believed all matter was made up of tiny particles 150 years later John Dalton proposed his atomic theory based on experiments he carried out

Dalton’s Atomic Theory: All matter consists of tiny particles called atoms All atoms are indivisble so cannot be split into simpler particles Atoms cannot be created or destroyed

Discovery of the electron: William Crookes worked with discharge tubes studying if gases conduct electricity

When electricity passes through gases they form invisible rays that travel from the negative electrode (cathode) to the positive electrode (anode) Crookes called these rays cathode rays

John Thomson studied cathode rays He found they were rays of negative particles that were lighter than a hydrogen atom (the smallest particle known then) so particles smaller than atoms must exist He also discovered the e/m (charge to mass) ratio of the electron Cathode rays are streams of negatively charged particles called electrons

Milikan’s Oil Drop experiment: Robert Milikan carried out an experiment to measure the negative charge of an electron He sprayed a fine mist of oil between oppositely charged plates He passed electricity through the oil droplets to ionise them (make them charged)

He focused on one oil drop and adjusted the charge on the plates until the droplet stayed suspended in mid air The electrical force on the plate was then equal to the gravitational force on the droplet He could then calculate the charge on the electron

Thomson’s Plum Pudding Model: He knew atoms were neutral He knew they contained negative particles called electrons He thought there must be a positive charge to balance the negative one

Thomson thought the atom looked like a plum pudding A large sphere of positive charge with negative electrons stuck in it Like fruit in a plum pudding

Discovery of the nucleus: Ernest Rutherford carried out the Gold Foil experiment He pointed alpha particles at sheets of gold foil He used a Geiger-Muller tube to detect them on the other side

Based on Thomson’s plum pudding model Rutherford made the following assumptions: 1) most particles would pass straight through because the atom is full of empty space 2) some particles would be slightly deflected by electrons

What actually happened (his results): 1) most particles passed straight through 2) some were deflected at small angles 3) some were deflected at very large angles 4) some were deflected back on their paths

Expected: Actual:

Rutherford’s conclusions were: 1) most of the atom is empty space 2) the atom does contain electrons 3) the atom has a large positive mass at its centre – called the nucleus

Discovery of the proton: Rutherford continued his experiment using lighter elements like N and F He found that these elements released positive particles when hit with alpha particles He called these positive particles protons

Discovery of the neutron: Chadwick (one of Rutherford’s researchers) carried out the same experiment with Be He found that it released neutral particles when hit with alpha particles He called these neutral particles neutrons

Sub-atomic particles: Mass: Charge: Location: Electron 1/1836 a.m.u. -1 Outside nucleus Proton 1 a.m.u. +1 Nucleus Neutron

Atomic Structure. Atoms consist of a nucleus at the centre, containing protons and neutrons, with electrons orbiting the nucleus in shells. Atomic number = Number of protons in the nucleus. Mass number = Number of protons plus number of neutrons in nucleus.

No. electrons = no. protons No. neutrons = mass no. – atomic no. No. protons = atomic no. No. electrons = no. protons No. neutrons = mass no. – atomic no. Element: No. electrons: No. protons: No. neutrons H 1 He 2 Li 3 4 ----

Isotopes Isotopes are atoms of the same element that have different mass numbers due to the number of neutrons in the nucleus, e.g., H-1, H-2 and H-3 or C-12 and C-14 Relative atomic mass (Ar) is the average mass number of the isotopes of the element compared to 1/12th the mass of the carbon-12 isotope. Mass spectrometer is used to measure Ar.

Mass Spectrometer Principle – Positive ions are separated according to their mass in a magnetic field The lighter the ions the more they will be deflected by the magnetic field 1) Vaporisation: Sample is vaporised and injected into sample inlet

2) Ionisation: 3) Acceleration: Sample passes through the electron gun where it is bombarded with electrons causing electrons in the sample atoms to fall off so they become positively charged 3) Acceleration: Ions pass through oppositely charged plates where they are attracted to the negative plate and are accelerated to high speeds

4) Separation: Ions pass into a magnetic field where they are separated based on their mass. Lighter particles are easy to deflect compared to heavier ones. The magnetic field can be adjusted so that all isotopes present can be measured. 5) Detection: The detector responds to the number of ions that hit it which are then registered on a chart called a mass spectrum – relative abundance peaks are shown for the different relative atomic masses

Mass Spectrum for chlorine

Mass spectrum for aspirin

Uses of the mass spectrometer: Identify presence of isotopes Measure relative abundance of isotopes Measure relative atomic & relative molecular mass - Identify unknown compounds

Calculations: Q: A mass spectrometer is used to measure a sample of boron. It is found that the sample contains 18.7% B-10 and 81.3% B-11. Calculate the relative atomic mass of boron. A: 18.7 x 10 = 187 81.3 x 11 = 894.3 100 atoms = 1081.3 1 atom = 10.813 Ar = 10.813

Q: A sample of chlorine consists of 75% Cl- 35 and 25% Cl-37. Calculate the average mass of an atom of chlorine. A: 75 x 35 = 2625 25 x 37 = 925 100 atoms = 3550 1 atom = 35.50 Ar = 35.5